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Glossary - Chemistry (Intro)

Chemistry: The science that studies the properties of substances and how substances react with one another.
Chemistry in contrast to physics is involved with the change of the outer electron-layers only, whereas
physics involves the state of the nucleus as well.
C. Energy: Energy stored within the structural units of chemical substances [J], see thermochemistry.
C. Equation: An equation that uses chemical symbols to show what happens during a chemical reaction.
C. Equilibrium: A state in which the rates of the forward and reverse reactions are equal.
C. Formula: An expression showing the chemical composition of a compound in terms of the symbols for
the atoms of the elements involved; see stoichiometry.
C. Kinetics: The area of chemistry concerned with the speeds, or rates, at which chemical reactions occur;
see stoichiometry.
C. Property: Any property of a substance that cannot be studies without converting the substance into
some other substances.
Electro C.: The branch of chemistry that deals with the use of chemical reactions to produce electricity,
the relative strengths of oxidizing and reducing agents, and the use of electricity to produce chemical
change.
Inorganic C.: The branch of chemistry that deals with compounds other than organic compounds.
Organic C.: The branch of chemistry that deals with carbon and usually hydrogen compounds, excluding
carbohydrates.
Nuclear C.: The study of the structure of nuclei, of the changes this structure undergoes, and of the
consequences of those changes for chemistry.
Thermo C.: The study of heat changes in chemical reactions.
Closed System: A system that allows the exchange of energy (usually in the form of heat) but not mass with its
surrounding environment.
Cohesion: The intermolecular attraction between like molecules (see physics - matter).
Compound: 1) A specific combination of elements that can be separated into elements by using chemical
techniques. 2) A substance consisting of atoms of two or more elements in a defined ratio.
Distillation: The separation of a mixture by making use of the different volatilities of its compounds.
Fractional D.: Separation of the components of a liquid mixture by repeated distillation, by making use of
their differing volatilities.
Element: A substance that cannot be separated into simpler substances by chemical means; e.g.: C, N, Fe, Na,
etc.
Notation of E.: Elements of the periodic table are assigned with a mass- and atomic number to quantify its
number of protons (Z) and number of protons and neutrons (A); see chemistry atom.
Representative E.: Elements in groups 1A through 7A, all of which have incompletely filled s or p
subshell of highest principal quantum number.
Experiment: A test carried out under carefully controlled conditions.
Group: The elements in a vertical column of the periodic table - see period.
Isomere: (Gk. isos, equal; meros, part) One of a group of compounds identical in atomic composition but
differing in structural arrangement; e.g., glucose and fructose
Matter: Anything that occupies space and possesses mass.
Properties of M.: 1) A physical property can be measured and observed without changing the
composition or identity of a substance. 2) in order to observe it, a chemical change has to be carried out;
e.g.: hydrogen burns with oxygen to form water - see chemistry-solid -liquid, -gas.
• Extensive P.o.M.: A physical property of a substance that depends of the size of the sample; e.g.:
mass, internal energy, enthalpy, entropy, etc.
• Intensive P.o.M.: A physical property of a substance that is dependent of the size of the sample; e.g.:
density, molar volume, temperature, etc.
Mineral: A naturally occurring substance with a range of chemical composition.
Mixture: A type of matter that consists of more than one substance and may be separated into components by
making use of the different physical properties.
Heterogeneous M.: A mixture in which the individual components, although mixed together, lie in
distinct regions, even on a microscopic scale; e.g.: a mixture of sand and sugar, ect.
Homogenous M.: A mixture in which the individual components are uniformly mixed, even on an
atomic scale; e.g.: air, solutions, etc.
Ore: The natural mineral source of a metal; e.g.: Fe2O3, hematite; etc.

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Period: The horizontal row of the periodic table; the number of the period is equal to he principal quantum
number of the valence shell of the atoms; periods from 1 to 7.
P. Table: A chart in which the elements are arranged in order of atomic number and divided into groups
and periods in a manner that shows the relationships between the properties of the elements.
Group: The vertical column of the P.T.; the number of the group is equal to the number of electrons in the
valence shell of the atoms; e.g.: groups from I to VIII.
Groupings within the P.T.:
Actinid Series: Elements that have incompletely filled 5f subshells or readily give rise to cations that have
incompletely filled5f subshells.
Alkali Metal: The group 1A elements like Li, Na, K, Rb, Cs, and Fr.
Alkali Earth Metal: A group of grayish-white , malleable metals easily oxidized in air, comprising Be,
Mg, Ca, Sr, Ba, and Ra.
Halogen: The nonmetallic elements in Group 7A, F, Cl, Br, I, and At.
Lanthanide Series (are earth series): Elements that have incompletely filled 4f subshells or readily give
rise to cations that have incompletely filled 4f subshells.
Metalloid: An element with properties intermediate between those of metals and nonmetals; i.e.: it has the
physical appearance and properties of a metal but behaves chemically like a nonmetal; e.g.: arsenic,
polonium, etc.
Metal: 1) A substance that conducts electricity as well as heat, has a metallic luster, is malleable and
ductile, forms cations and has basic oxides. 2) A metal consists of cations held together by a sea of
electrons (have the tendency to form positive ions in ionic compounds); i.e.: iron, copper, uranium, etc.
Noble Gas: Nonmetallic elements in group 8A; He, Ne, Ar, Kr, Xe, and Rn.
Transition Metal.: Elements that have incompletely filled d subshells or readily give rise to cations that
have incompletely filled d subshells; i.e.: it belongs to the central part of the periodic table, between
Groups II and III; a member of the d-block of the periodic table; i.e.: vanadium, iron, gold, etc.
ppm - Parts per Million: 1) The ratio of the mass of a solute to the mass of the solvent, multiplied by 106; 2) The
mass percentage composition multiplied by 104;
Phase: A particular state of matter; a homogeneous part of a system in contact with other parts of the system but
separated from them by a well defined boundary; a substance may exist in solid, liquid, and gas phases
and in certain cases, in more than one solid phase; e.g.: white and gray tin are two solid phases of tin; ice,
liquid, and vapor are three phases of water; etc.
P. Change: Transformation from one phase to another.
P. Diagram: A map divided into regions that tells us which phase is the most stable under corresponding
conditions of temperature and pressure. The lines separating the areas in the diagram are called phase
boundaries. The points on a phase boundary show the conditions under which two phases coexist in
dynamic equilibrium.
Distinct Points in a P.D.:
KP, Critical point: The liquid-vapor boundary terminates at this point; for water at 101,3[kPa] = 100[°C].
TP, Triple point: The pint where 3 phase boundaries meet; for water it occurs at 0,61[kPa] and 0,01[°C];
under these conditions all three phases (ice, water, vapor) coexist in dynamic equilibrium.
BP, Boiling point: The temperature at which a liquid boils when the atmospheric pressure is 1[atm] or
101,3[kPa].
FP, Freezing point: The temperature at which it freezes (melts) at 1[atm]; i.e.: liquid and solid state are in
dynamic equilibrium.
Significant Figures: The number of meaningful digits in a measured or calculated quantity.
STP - Standard Temperature and Pressure: 0[°C] = 273.15[K] and 1[atm] = 101.325[kPa].
Substance: A form of matter that has a definite or constant composition (the number and type of basic units
present) and distinct properties.

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SI - Base Units: (F. système international): Modern system of definitions and metric notation, now spreading
throughout the academic, industrial, and commercial community; these are: ampere, area, joule, kelvin,
kilogram, meter, mole, newton, rad, second, volume.
n Amount of Substance [mol]
The amount of substance that contains as many elementary entities (atoms, molecules, or other particles)
as there are atoms in exactly 12 grams of 12C isotope.
I Electric Current [ampere, = A]
the flow of 1 coulomb (1C = 6.25x1018 electrons) of charge/s.
F Force [kg⋅m/s2] [newton, = N]
the force that will give an object of 1 kg an acceleration of 1 m/s2 = [kg m/s2].
l Length [meter, = m]
the length of the path traveled by light in vacuum during a time of 1/299792458 of a second. Area: [m2].
Volume: [m3] Quantity of space an object occupies.
L Light Intensity [N⋅m/(s⋅sr)] = [J/(s⋅sr)] = [W/sr] = [candela, Cd]
light intensity of a monochromatic radiation with a frequency of 540x1012 oscillations /s [Hz] with a
power in the direction equal to 1/683 [Js or W/sterarian].
m,M Mass [kilogram, = kg]
one kilogram is the amount of mass in 1 liter of water at 4°C.
rad Radian: The radian is the 2D plane angle between two radii of a circle which cut off on the circumference
an arc equal in length to the radius: 1 rad = 57.3°; π rad = 180°;
sr - steradian: Is the solid 3D angle which, having its vertex in the center of a sphere, cuts off an area equal
to that of a flat square with sides of length equal to the radius of the sphere.
T Thermodynamic Temperature [kelvin, = K]
defined to be 1/273.15 the thermodynamic temperature of the triple point of water; ice melts therefore at
273.15 K and water boils at 373,15 K (both at atmospheric pressure).
t Time: Second [s]; the time taken by a 133Cs-atom to make 9 192 631 770 vibrations.
W Work [N⋅m] [joule, = J]
the specific heat of work at 15°C is given as 4185.5 J/kg⋅C done by a force of 1 newton acting over a
distance of 1 meter.

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SI-Derived Units:
α dissociation constant of acid / base [-]
β mass concentration (solute over solvent) [g/l]
ρ density [kg/m3] [g/l]
η coefficient of efficiency [-]
λ wavelength [m]
π circle’s constant = 3.14159 [-]
π osmotic pressure [Pa]
τ level of titration [-]
µ dipole moment [C⋅m] [D]
ℜ,R gas constant = 8.314 510 [J/(mol⋅K)]
A area, cross-sectional area [m2]
b molality (solute over solvent) [mol/g]
c speed of light 2.99792458⋅108 ≈ 3⋅108 [m/s]
c specific heat capacity [N⋅m/(kg⋅K)] [J/(kg⋅K)]
c molar concentration [mol/l]
d distance [m]
D debye = 3.336⋅10-30 [C⋅m]
e charge of an electron = 1.602 177 3349⋅10-19 [C]
E energy [kg⋅m2/s2] = [N⋅m] = [Pa⋅m3] [joule, J] 1eV = 1.6022⋅10-19 [J]
F Faraday’s constant = 96485 [C/mole] 1[cal] = 4.178⋅103 [J]
-34
h plank’s constant = 6.626 075 540⋅10 [J⋅s]
H ∆H enthalpy [J/mol]
k Henry’s law variable [mol/(l⋅Pa)]
k kinetic rate law [l/(mol⋅s)] or [1/s] or [mol/(l⋅s)] varies, depending upon the rate order.
kB Boltzman’s constant = 1,380 658 12⋅10-23 [J/K]
kF Coulomb’s force constant = 8.987 551 79⋅109 [N⋅m2/C2] ≈ 9x109 = 1/(4⋅π⋅ε0)
KA,B dissociation constant of acids and bases [mol/l]*
KC equilibrium constant for molar concentrations [mol/l]*
Kf complex formation constant [mol/l]*
KP equilibrium constant for partial pressures [Pa/l]*
Ksp solubility product [mol/l]*
-14
Kw water auto-dissociation constant = 1⋅10 [mol2/l2]
l length [m]
me mass of an electron = 9.109 389 754⋅10-31 [kg]
mp mass of a proton = 1.672 623 110⋅10-27 [kg]
mn mass of a neutron = 1.674 928 610⋅10-27 [kg]
M, Mr molecular mass [g/mol], [amu]
n molar amount [mol]
NA Avogadro’s constant = 6,022 136 736⋅1023 [1/mol]
p impulse [kg⋅m/s]
p pressure [N/m2] 1 bar = 105 [Pa] [pascal, Pa] 1atm = 1.01325[bar] = 101.325[kPa]
q electric charge of an electron (see e) [A⋅s] [coulomb, C]
Q heat capacity [J]
QC reaction quotient [mol/l]*
r radius [m]
s specific heat capacity [J/(kg⋅K)]
S entropy [kg⋅m/(s2⋅K)] = [N⋅m/K] [J/K]
U internal heat [J]
v velocity [m/s]
v(Ax) kinetic reaction rate [mol/(l⋅s)]
V electric potential [J/C] [volt, V]
V volume [m3]
w percent by mass [%]
x molar fraction [-]
SI-Prefixes: 1018 exa- E; 1015 peta- P; 1012 tera- T; 109 giga- G; 106 mega- M; 103 kilo- k; 10 deka- d;
10-3 milli- m; 10-6 micro- µ; 10-9 nano- n; 10-12 pico- p; 10-15 femto- f; 10-18 atto- a;

*) varies, depending upon stoichiometric coefficient of products and reactants; e.g.: [l/mol], [mol2/l2], etc.

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Periodic Table of Elements

Group
Pieriod
I II III IV V VI VII VII
I
1 H He
2 Li Be B C N O F Ne
3 Na Mg Al Si P S Cl Ar
4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
6 Cs Ba La *) Hf Ta W Re Os Ir Pt Au Hg Ti Pb Bi Po At Rn
7 Fr Ra Ac **)
Block s d f d p

*) Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
**) Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr

Element amu Z A Element amu Z A Element amu Z A

Ac Actinium (227) 89 H Hydrogen 1.008 1 Pm Promethium (147) 61
Ag Silver 107.9 47 He Helium 4.003 2 Po Polonium (209) 84
26.98 178.5 Praseodymium 140.9
Al Aluminium 13 Hf Hafnium 72 Pr 59
Am Americum (243) 95 Hg Mercury 200.6 80 Pt Platinium 195.1 78
Ar Argon 39.95 18 Ho Holmium 165.0 67 Pu Plutonium (209) 94
As Arsenic 74.92 33 Ga Gallium 69.72 31 Ra Radium 226.0 88
At Astatine (210) 85 Gd Gadolinium 157.3 64 Rb Rubidium 85.47 37
Au Gold 197.0 79 Ge Germanium 72.59 32 Re Rhenium 186.2 75
B Boron 10.81 5 I Iodine 126.9 53 Rh Rhodium 102.9 45
Ba Barium 137.3 56 In Indium 114.8 49 Rn Radon (222) 86
Be Beryllium 9.012 4 Ir Iridium 192.2 77 Ru Ruthenium 101.1 44
Bi Bismuth 209.0 83 K Potassium 39.10 19 S Sulfur 32.07 16
Bk Berkilium (247) 97 Kr Krypton 83.80 36 Sc Scandium 44.96 21
Bo Boron 10.81 5 La Lanthanum 138.9 57 Se Selenium 78.96 34
Br Bromine 79.90 35 Li Lithium 6.941 3 Si Silicon 28.09 14
C Carbon 12.01 6 Lr Lawrencium (260) 103 Sb Antimony 121.8 51
Ca Calcium 40.08 20 Lu Lutetium 175.0 71 Sm Samarium 150.4 62
Cd Cadmium 112.4 48 Md Mendelevium (258) 101 Sn Tin 118.7 50
Ce Cerium 140.1 58 Mg Magnesium 24.30 12 Sr Strontium 87.62 38
Cf Californium (249) 98 Mn Manganese 54.94 25 Ta Tantalum 180.9 73
Cl Chlorine 34.45 17 Mo Molybdenum 95.94 42 Tb Terbium 158.9 65
Cm Curium (247) 96 N Nitrogen 14.01 7 Tc Technetium (98) 43
Cs Cesium 132.9 55 Na Sodium 22.99 11 Te Tellurium 127.6 52
Cr Chromium 52.00 34 Nb Noibium 92.91 41 Th Thorium 232.0 90
Co Cobalt 58.93 27 Ne Neon 20.18 10 Ti Titanium 47.88 22
Cu Copper 63.55 29 Nd Neodymium 144.2 60 Tl Thallium 168.9 69
Dy Dysprosium 162.5 66 Ni Nickel 58.69 28 Tm Thulium 168.9 69
Er Erbium 167.3 68 No Nobelium (259) 102 U Uranium 238.0 92
Es Einsteinium (252) 99 Np Neptunium 237.0 93 V Vanadium 50.94 23
Eu Europium 152.0 63 O Oxygen 16.00 8 W Tungsten 183.8 75
F Flourine 19.00 9 Os Osmium 190.2 76 Xe Xenon 131.3 54
Fe Iron 55.85 26 P Phosphorus 30.97 15 Zn Zinc 65.39 30
Fm Fermium (257) 100 Pa Protactinium 231.0 91 Zr Zirconium 91.22 40
Fr Francium (223) 87 Pb Lead 207.2 82 Y Yittrium 88.91 39
Pd Palladium 106.4 46 Yb Ytterbium 173.0 70

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Glossary - Chemistry (Atom)

Atom: The basic unit of an element that can enter into chemical combination (see physics - nuclear).
Z - Atomic Number: The number of protons in the nucleus of an atom; it determines the identity of the
element and the number of electron in the neutral atom.
A. Orbital: see orbital.
A. Radius: ½ the distance between the two nuclei in two adjacent atoms of the same element in a metal.
For elements that exist as diatomic units, the atomic radius in ½ the distance between the nuclei of the
two atoms in a particular molecule; The A.R. decreases from left to right across the period of the periodic
table, whereas it increases going down a group of the periodic table.
A. Weight: The weight of a representative atom of an element relative to the weight of an atom of 12C,
which has been assigned the value 12 - see stoichiometry - mass.
Average Atomic Mass: The average mass of the naturally occurring mixture of isotopes as listed in the
periodic table of elements - see chemistry introduction - periodic table.
Parts of the atom:
Electron: A subatomic particle that has a very low mass and me = 9.1095⋅10-28 [g]
carries a single negative electric charge (see physics - introduction). qe = -1.6022⋅10-19 [C]
Neutron: A subatomic particle that bears no electric charge. mn = 1.67495⋅10-24 [g]
Its mass is slightly greater than a proton’s.
Proton: A subatomic particle having a single positive electric mp = 1.67252⋅10-24 [g]
charge. The mass of a proton is about 1840 times that of an electron. qp = +1.6022⋅10-19 [C]
Aufbau Principle: As protons are added one by one to the nucleus to build
up the elements, electrons similarly are added to the atomic orbitals (see table below).
Bohr’s Atomic Model: Energy of the electron orbiting the hydrogen R, Rydberg c. = 2.18⋅10-18 [J]
nucleus is quantized into fixed values; i.e. electron can only „move“ n, principal quantum number
along fixed steps of energy (no values in-between) c, speed of light 3⋅108 [m/s]
2
E = R⋅h⋅c/n [N⋅m] = [J] h, Planck c. = 6.63⋅10-34 [J⋅s]
Dalton’s Atomic Theory: Hypotheses about the nature of matter.
• Elements are composed of atoms. All atoms of a given element are identical (except for isotopes),
having the same size, mass, and chemical properties. Atoms of one element are different from atoms of
all other elements.
• Compounds are composed of atoms of more than one element. In any compound, the ratio of numbers
of atoms is either an integer or a simple fraction.
• Chemical reactions involve only the separation, combination, or rearrangement of atoms; it does not
result in their creation or destruction (except nuclear reaction).
De-Broglie’s Hypothesis: The proposal that every particle has h, Planck c. = 6.63⋅10-34 [J⋅s]
wavelike properties; hence it is impossible to determine both m, mass [kg]
position and velocity simultaneously: c, speed of light = 3⋅108 [m/s]
λ = h/(m⋅c) = h/p [m] p, impulse [kg⋅m/s]
Electron: see atom - parts of.
E. Affinity: The energy change when an electron is accepted by an atom (or ion) in the gaseous state.
E. Configuration: The distribution of electrons among the various orbitals in an atom or molecule.
E. Density: The probability that an electron will be found at a particular region in an atomic orbital.
E. per Shell: The principles outlined by both the atomic Bohr model and the rules of the quantum number,
only a certain amount of e can occupy each orbital: eN, number of e/shell [-]
eN = 2⋅n2 n, principal shell number [-]
Exited State: A state that has higher energy that the ground state.
Exclusion Principle: see pauli.
FC - Formal Charge: 1) The electric charge of an atom assigned on the assumption that there is perfect covalent
bonding. 2) FC = number of valence electrons in the free atom - number of lone-pair electrons - ½ x
number of shared electrons; see chemistry molecule.
Ground State: The lowest energy state of a system.
Heisenberg uncertainty Principle: It is impossible to know x, position [m]
simultaneously both momentum and the position of a particle with m, mass [kg]
certainty (see physics - nuclear). c, speed of light = 3⋅108 [m/s]
∆x⋅∆p = h/(4⋅π) = ∆x⋅∆(m⋅c)⋅h/(4⋅π) [kg⋅m2/s] h, Planck c. = 6.63⋅10-34 [J⋅s]

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Hund’s Rule: The most stable arrangement of electrons in subshells is the one with greatest number of parallel
spins; e.g.: C with 6 electrons, where the 2p2 of the 1s2 2s2 are parallel - compare chemistry molecule -
hybridization.
Inert Pair Effect: The two relatively stable and unreactive outer s electrons.
Ion: An atom or molecule that has lost or gained one or more electrons, and thus becomes positively or
negatively charged; i.e.: Al3+ (mono-atomic ion), SO4- (poly-atomic ion); see also chemistry-molecule &
thermochem.
Anion: An ion with a net negative charge; characteristic of non-metallic elements, e.g.: F-, SO4-, etc.
Cation: An ion with a net positive charge; characteristic of metallic elements, e.g.: Na+, NH4+, Al3+, etc.
Ionic Radius: The radius of a cation or an anion as measured in an ionic compound (see table below); it
decreases from left to right across the a period and increases down a group of the periodic table; in the
same period, anions generally have larger radii than cations.
Ionization: Conversion to cations by the removal of electrons; see chemistry-thermochemistry ∆H.
e.g.: K(g) → K+(g) + e-(g)
Isoelectronic Species: Species with the same number of atoms and the same number of valence electrons;
e.g.: SO2 and O3; Cl- and Ar; Na+ and Ne, NO3- and CO32- (Ar, Ne are noble gases).
Isotope: One of several possible forms of a chemical element that differ from others in the number of neutrons in
the atomic nucleus, but not in chemical properties e.g.: 21Na, 22Na, 23Na.
I. Dating: The determination of the ages of objects by measuring the activity of the radioactive isotopes
they contain; particularly, radiocarbon dating by using 14C.
Magnetism: A materials ability to store magnetic energy after being magnetized describes 3 distinct classes:
Diamag.: Repelled by a magnet; a diamagnetic substance contains only paired electrons, resulting in a
magnetic flux which is lower inside than outside (µr<1, slightly repelling when exposed to a magnetic
field; temperature independent) - H2, Cu, Hg....
Ferromag.: Strongly attracted by a magnet; the unpaired spins in a ferromagnetic substance are aligned in
a common direction; the magnetic flux inside is extremely high compared to outside; (µr>>1, the higher
the temperature the less ferromagnetic will be the substance) - Fe, Ni. Co...
Paramagnet.: Attracted by a magnet. A paramagnetic substance contains one or more unpaired electrons;
resulting in a magnetic flux which is higher inside than outside (µr≥1, slightly attracting, the lower the
temperature the more ferromagnetic the substance will become) - O2, Cr, Pt....
Notation of Elements: Elements of the periodic table are assigned with a mass- and atomic number to quantify
its number of protons (Z) and number of protons and neutrons (A).
Orbital: The idealized representation of electron moving around its nucleus.
Atomic O.: The wave function (ψ) of an electron in an atom (see quantum number).
Degenerate O.: Orbitals that have the same energy.
Electron Configuration: The occupancy of orbitals in an atom or molecule; i.e.: N 1s2 2s2 2p3,
N2 σ2 σ*2 π4 σ2 (valence electrons only); see chemistry molecule.
Partial Charge: A charge arising from small shifts in the distribution of electrons.
Pauli Exclusion Principle: No two electrons in an atom can have the same four quantum numbers.
Photon: A particle of light (see physics-optics); The energy contained in a photon can be put as follows:
E = h⋅f = h⋅c/λ [J]
Quantum: The smallest quantity of energy that can be emitted, h, Planck c. 6.63⋅10-34 [J⋅s]
or absorbed in the form of electromagnetic radiation c, speed of light 3⋅108 [m/s]
(see physics-optics). λ, wavelength [m]
c = λ⋅f [m/s] f, frequency [1/s] [Hz]
Quantum Number: The quantum numbers of quantum mechanics which describe the distribution of
electrons, labels the state of the electron and specifies the value of a property in atoms.
Paired Electrons: Two electrons with opposed spins (↑↓).
Parallel Electrons: Electrons with spins in the same direction (↑↑).
1. n - Principal QN.(shell number): The average distance of the electron from the nucleus in a particular
orbital; can have integral values of 1, 2, 3, and so forth (higher values ≈ greater average distance) e.g.:
1 = 1st period, 7 = 7th period;
2. l - Angular Momentum QN.(shubshell of one shell): Its value reflects the orbital shape; it correlates
with n; (l = n-1); which reveals 0 for the s-, 1 for p-, 2 for d- 3 for f-, 4 for g- 5 for the h-orbital.
3. ml - Magnetic QN.: Describes the orientation of the orbital in space; depends upon the value of l; (ml
= 2⋅l+1), e.g.: ml = 1 a sphere; ml = 3 gives -1/0/+1 (x,y,z-orientation); ml = 5 gives -2/-1/0/1/2 etc.
4. ms - Electron Spin QN.: According to the electromagnetic theory, spinning electrons posses a
magnetic orientation; (ms = n) ms can either be -1/2 (↓)or +1/2(↑); with ms =3 giving 3 magnetic spins: -
1
/2 / +1/2 / -1/2.

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∆ψ - Schrödinger Wave Equation: The fundamental equation of quantum mechanics; it interprets the wave
nature of material particles in terms of probability wave amplitudes.
It specifies the possible energy state the electron can occupy ∆ψ, LaPlace operator [-1/m]
and identifies the corresponding wave function (ψ). ψ, wave function [m]
These energy states and wave functions are characterized ψ2, probability of density [-]
by the quantum numbers; The concept of electron density PE, potential energy [J]
gives the probability that an electron will be found in a E, total energy [J]
particular region of an atom: m, mass [kg]
∆ψ+(8⋅π2⋅m/h2)⋅(E-PE)⋅ψ = 0 h, Planck c. = 6.63⋅10-34 [J⋅s]
Series: A family of spectral lines arising from transitions that π, circle’s constant = 3.142 [-]
have one state in common (see spectrum):
Bracket S.: Energy spectrum located in the infrared (ni = 4)
Paschen S.: Energy spectrum located in the infrared (ni = 3)
Balmer S.: A family of spectral lines (some of which lie in the visible region, others in the UV band), in
the spectrum of atomic hydrogen; (ni = 2)
Lyman S.: Energy spectrum located entirely in the UV (ni = 1)
Shell: All the orbitals of a given principal quantum number >n<; i.e.: the single 2s and three 2p orbitals of the
shell with n = 2;
S.-Nomenclature: s - sharp; p - principal; d - diffuse; f - fundamental
Subshell: All the atomic orbitals of a given shell of an atom with the same value of the quantum number
>l<; i.e.: the five 3d orbitals of an atom.
Spectroscope: An optical device that separates light into its constituent frequencies in form of spectral lines.
Spectrum: 1) The frequencies or wavelengths of electromagnetic radiation emitted or absorbed by substances
2) The splitting of white light into its distinct components.
Emission S. of Hydrogen: A line spectrum of radiation emitted R, Rydberg constant 2.18⋅10-18 [J]
by the excited hydrogen electron. (integers of ni > nf) - see series ni, initial principal QN
∆E = R⋅h⋅c(1/ni2 - 1/ nf2) [J] nf, final principal QN
S. Line: Radiation of a single wavelength emitted or absorbed h, Planck’s c. = 6.63⋅10-34 [J⋅s]
by an atom or molecule. c, speed of light = 3⋅108 [m/s]
Line Spectra: Spectra produced when radiation is absorbed or
emitted by substances only at some wavelengths.
Spin: The intrinsic angular momentum of an electron; the spin cannot be eliminated and may occur in only two
senses, denoted as ↑ and ↓ - see quantum number.

Atomic and ionic radii of common elements before and after ionization [pm]:

I II III IV V VI VII VIII

H He (30)
1 He- (154)
Li (142) Be (112) B (88) C (77) N (70) O (66) F (64) Ne
2 Li+ (60) Be+ (31) O2- (140) F- (138)
Na (186) Mg (160) Al (143) Si (117) P (110) S (104) Cl (99) Ar
3 Na+ (95) Mg2+ (65) Al3+ (50) S2- (184) Cl- (181)
K (231) Ca (197) Sc (160) Cu (128) Zn (133) Ga (122) Ge (122) As (121) Se (117) Br (114) Kr
4 K+ (133) Ca2+ (97) Sc3+ (81) Cu+ (96) Zn2+ (74) Ga3+ (62) Ge4+ (53) Se2- (198) Br- (195)
Rb (244) Sr (215) Y (180) Ag (144) Cd (149) In (162) Sn (140) Sb (141) Te (137) I (133) Xe
5 Rb+ (148) Sr2+ (113) Y3+ (93) Ag+ (126) Cd2+ (97) In3+ (81) Sn4+ (71) Te2- (221) I- (216)
Cs (262) Ba (217) La (188) Au (144) Hg (150) Ti (171) Pb (175) Bi (146) Po (140) At (140) Rn
6 Cs+ (169) Ba2+ (135) La3+ (115) Au+ (137) Hg2+ (110) Tii3+ (95) Pbi4+ (84)

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The order in which atomic orbitals are filled according to the Aufbau principle:

Cor I II III IV V VI VII VIII

e 1s → →
2s → → 2p
He 3s → → 3p
Ne 4s 3d → → 3d 4p
Ar 5s 4d → → d4 5p
Kr 6s 5d 4f 5d 6p
Xe 7s 6d 5f 6d
Rn 2 s-b g. 1 14 f-block grp. (1+) 9 d-block groups 6 p-block groups

The ground-state electron configuration of the elements:

Element Electron Element Electron config. Element Electron

config. config.
Ac Actinium [Rn] 7s26d1 H Hydrogen [ ] 1s 1
Pm Promethium [Xe] 6s24f5
Ag Silver [Kr] 5s14d10 He Helium [ ] 1s2 Po Polonium [Xe] 6s24f145d106p4

Al Aluminium [Ne] 3s23p1 Hf Hafnium [Xe] 6s24f145d2 Pr Praseodymium [Xe] 6s24f3

Am Americum [Rn] 7s25f7 Hg Mercury [Xe] 6s24f145d10 Pt Platinium [Xe] 6s14f145d9
Ar Argon [Ne] 3s23p6 Ho Holmium [Xe] 6s24f11 Pu Plutonium [Rn] 7s25f6
As Arsenic [Ar] 4s23d14p3Ga Gallium [Ar] 4s23d14p1 Ra Radium [Rn] 7s2
2 7 1
At Astatine
2 14 10
[Xe] 6s 4f 5d 6p 5
Gd Gadolinium [Xe] 6s 4f 5d Rb Rubidium [Kr] 5s1
1 14 10 2 1 2
Au Gold [Xe] 6s 4f 5d Ge Germanium [Ar] 4s 3d 4p Re Rhenium [Xe] 6s24f145d5
B Boron [He] 2s22p1 I Iodine [Kr] 5s24d105p5 Rh Rhodium [Kr] 5s14d8
2
Ba Barium [Xe] 6s In Indium [Kr] 5s24d105p1 Rn Radon [Xe] 6s24f145d106p6
2
Be Beryllium [He] 2s Ir Iridium [Xe] 6s24f145d7 Ru Ruthenium [Kr] 5s14d7
Bi Bismuth
2 14 10
[Xe] 6s 4f 5d 6p 3
K Potassium [Ar] 4s1 S Sulfur [Ne] 3s23p4
2 9
Bk Berkilium [Rn] 7s 5f Kr Krypton [Ar] 4s23d14p6 Se Selenium [Ar] 4s23d14p4
2 1 5 2 1
Br Bromine [Ar] 4s 3d 4p La Lanthanum [Xe] 6s 5d Si Silicon [Ne] 3s23p2
2 2
C Carbon [He] 2s 2p Li Lithium [He] 2s1 Sb Antimony [Kr] 5s24d105p3
2
Ca Calcium [Ar] 4s Lr Lawrencium [Rn] 7s25f146d1 Sc Scandium [Ar] 4s23d1
2 10
Cd Cadmium [Kr] 5s 4d Lu Lutetium [Xe] 6s24f145d1 Sm Samarium [Xe] 6s24f6
2 1 1
Ce Cerium [Xe] 6s 4f 5d Md Mendelevium [Rn] 7s25f13 Sn Tin [Kr] 5s24d105p2
Cf Californium [Rn] 7s25f10 Mg Magnesium [Ne] 3s
2
Sr Strontium [Kr] 5s2
Cl Chlorine [Ne] 3s23p5 Mn Manganese [Ar] 4s 2 5
3d Ta Tantalum [Xe] 6s24f145d3
Cm Curium [Rn] 7s25f76d1 Mo Molibydenum [Kr] 5s 1 5
4d Tb Terbium [Xe] 6s24f9
Cs Cesium [Xe] 6s1 N Nitrogen [He] 2s2 3
2p Tc Technetium [Kr] 5s24d5
Cr Chromium [Ar] 4s13d5 Na Sodium [Ne] 3s1
Te Tellurium [Kr] 5s24d105p4
Co Cobalt [Ar] 4s23d7 Nb Noibium [Kr] 5s14d4 Th Thorium [Rn] 7s26d2
1 10
Cu Copper [Ar] 4s 3d Ne Neon [He] 2s22p6 Ti Titanium [Ar] 4s23d2
2 10 2 4 2 14 10 1
Dy Dysprosium [Xe] 6s 4f Nd Neodymium [Xe] 6s 4f Tl Thallium [Xe] 6s 4f 5d 6p
2 12 2 8 2 13
Er Erbium [Xe] 6s 4f Ni Nickel [Ar] 4s 3d Tm Thulium [Xe] 6s 4f
Es Einsteinium [Rn] 7s25f11 No Nobelium [Rn] 7s25f14 U
2 3 1
Uranium [Rn] 7s 5f 6d
2 7 2 4 1
Eu Europium [Xe] 6s 4f Np Neptunium [Rn] 7s 5f 6d V Vanadium [Ar] 4s23d3
2 5 2 4
F Flourine [He] 2s 2p O Oxygen [He] 2s 2p W Tungsten [Xe] 6s24f145d4
Fe Iron [Ar] 4s23d6 Os Osmium [Xe] 6s24f145d6 Xe Xenon [Kr] 5s24d105p6
2 12
Fm Fermium [Rn] 7s 5f P Phosphorus [Ne] 3s23p3 Zn Zinc [Ar] 4s23d10
1
Fr Francium [Rn] 7s Pa Protactinium [Rn] 7s25f26d1 Zr Zirconium [Kr] 5s24d2
Pb Lead [Xe] 6s24f145d106p2 Y Yittrium [Kr] 5s24d1
Pd Palladium [Kr] 4d10 Yb Ytterbium [Xe] 6s24f14
2
the symbol [He] is called the helium-core and represents 1s
[Ne]-core: 1s22s22p6
[Ar]-core: 1s22s22p63s23p6
[Kr]-core: 1s22s22p63s23p64s23d104p6
[Xe]-core: 1s22s22p63s23p64s23d104p65s24d105p6
[Rn]-core: 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p6

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Orbitals of the 1st four elements

perio Shell Sub- Orbital Subshell Number of

d (n) shell (m) nomenclature orbitals per
(l) subshell
1 1 0 0 1s 1
2 2 0 0 2s 1
1 -1, 0, +1 2p 3
3 3 0 0 3s 1
1 -1, 0, +1 3p 3
2 -2, -1, 0, +1, +2 3d 5
4 4 0 0 4s 1
1 -1, 0, +1 4p 3
2 -2, -1, 0, +1, +2 4d 5
3 -3, -2,-1, 0, +1, +2, 4f 7
+3

Electron-configuration of the 1st ten elements

Element Orbital Diagram Configuration

1s  2s  2px 2py 2pz
1H ↑ 1s1
2He ↑↓ 1s2
3Li ↑↓ ↑ 1s2 2s1
4Be ↑↓ ↑↓ 1s2 2s2
5B ↑↓ ↑↓ ↑ 1s2 2s2 2p1
6C ↑↓ ↑↓ ↑ ↑ 1s2 2s2 2p2
7N ↑↓ ↑↓ ↑ ↑ ↑ 1s2 2s2 2p3
8O ↑↓ ↑↓ ↑↓ ↑ ↑ 1s2 2s2 2p4
9F ↑↓ ↑↓ ↑↓ ↑↓ ↑ 1s2 2s2 2p5
10Ne ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 1s2 2s2 2p6

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Glossary - Chemistry (Molecule)

Bond: A link between bonds.
B. Energy (Bond Dissociation Energy): The enthalpy change required to break a bond in a mole of
gaseous molecules - see thermochemistry.
B. Length: The distance between the centers of two bonded atoms in a molecule.
B. Order: see orbital-molecular.
Covalent B.: A bond in which two electrons are shared by two atoms. In a coordinate covalent bond the
pair of electrons is supplied by one of the two bonded atoms;
e.g.: H-Cl , by sharing each others electrons, both obtain noble gas configuration.
• π - Pi B.: A covalent bond formed by sideways overlapping orbitals; its electron density is
concentrated above and below the plane of the nuclei of the bonding atoms.
• σ - Sigma B.: A covalent bond formed by orbitals overlapping end-to-end; it has its electron density
concentrated between the nuclei of the bonding atoms.
• Coordinative (Dative) B.: The lone pair of one compound is occupied by a cation;
e.g.: NH3 + H+ → NH4+
Double B.: Two electron pairs shared by neighboring atoms; i.e.: one σ-bond and one π-bond.
Hydrogen B.: see forces.
Ionic B.: Electrostatic force that holds ions together in an ionic compound;
e.g.: Na + e- + Cl ↔ Na+(cation) + Cl-(anion)
Metallic B.: Ions (cations) in fixed in their proper lattice-position are held together by a sea of electrons
(have the tendency to form positive ions in ionic compounds); atoms of metals give off electrons into the
common pool of the conducting-band to obtain noble gas configuration.
Multiple B.: Bonds formed when two atoms share two or more pairs of electrons.
Nonbonding Electrons: Valence electrons that are not involved in covalent bond formation.
Peptide B.: A type of covalent bond joining amino acids in a polypeptide.
Polypetide B.: Several amino acids linked together by peptide bonds.
Triple B.: Three electron pairs shared by two neighboring atoms; i.e.: one σ-bond and two π-bonds.
Catenation: The ability of atoms to form bonds with one another.
Compound: A substance composed of atoms of two or more elements chemically united in fixed properties.
Covalent Radius: The contribution of an atom to the length of a covalent bond - see table below;
e.g.: Cl-Cl ; r = 198 [pm]
q - Charge: Is a whole number multiple of one electron and cannot be quantized below it (1e or 1p-net charge);
no net charge has ever been found - it can’t be created or destroyed, only transformed (compare 1st law of
thermodynamics) e = 1.602 18⋅10-19 [As] = [C]. Like charges repel, opposite charges attract;
electric charges can be isolated, magnet poles cannot.
FC - Formal Charge: 1) The electric charge of an atom assigned on the assumption that there is perfect
covalent bonding. 2) FC = number of valence electrons in the free atom - number of lone-pair electrons -
½ the number of shared electrons;
e.g.: overall e- availability of NO3- given by the octet rule would be: eV’ = 4 elements⋅8 = 32;
N accounts for only 5, and O for (3⋅6) 24, plus one extra (e-): eV’’ = gives a total of 24;
the difference ½⋅(32-24) = 8 gives the maximum of possible bonds between these two atoms; in the case
of NO3- ½⋅8 = 4 bonds, with N in the middle (a planar Y-shaped molecular structure).
Partial C.: A charge arising from small shifts in the distribution of electrons, resulting in an electric
dipole, see there.
Coloumb’s Law: The potential energy between two ions is directly proportional to the product of their charges
and inversly proportional to the distance of separation between them (see physics - electromagnetism).
The relationship among electrical force, charge, and distance:
F = kF⋅q1⋅q2/d2 [Nm2/C 2 ⋅ C2/m2] = [N] kF, coulomb c. ≈ 9⋅109 [N⋅m2/C2]
If the charges are alike in sign, the force is repelling; q, charge [A⋅s] 1.602⋅10-19 [C]
If the charges are unlike, the force is attractive. d, distance [m]
Cyanides: Compounds containing the CN- ion.

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Dipole: A molecule with a separate regions or net negative and net positive charge, so that one end acts as a
positive pole and the other as a negative pole.
D. Moment: The electrostatic force required to align a dipolar molecule parallel to the electrostatic field;
the force required increases as the separation of the molecular charges decreases; The product of charge
and the distance between the charges in a molecule: q, charge [C]
µ = q⋅d [C⋅m] = [D]; 1Debey = 3.336⋅10-30 [C⋅m] d, distance [m]
The overall DM is obtained by adding the individual vector-amounts of the involved atoms.
D. Forces: Forces that act between polar molecules; see force.
Electric D.: A positive charge next to an equal but opposite negative charge.
Induced D.: The separation of positive and negative charges in a neutral atom (or non-polar molecule)
caused by the proximity of an ion or polar molecule.
Electronegativity: The measure of the ability of an atom to attract electron to itself when it is part of a
compound; it is formed by the arithmetic sum of the actual bonding energy (A-B) and the hypothetical
unpolar bond between A-B; the difference reveals the polar character of a bond; it increases from left to
right across the period of the periodic table and decreases down a group; see table below.
Esters: Compounds that have the general formula R’COOR; R’ can be H or an alkyl group or an aryl group.
Ether: An organic compound containing the R-O-R’ linkage, where R and R’ are alkyl and/or aryl groups.
F - Force: An influence that changes the state of motion of an object; i.e.: electrostatic force from an electric
charge; a mechanical force from an impact, etc.; m, mass [kg]
FL = m⋅a [kg⋅m/s2] = [N] a, acceleration [m/s2]
Intermolecular F. (Van der Waals): The close-ranging, relatively
weak attraction exhibited between atoms and molecules. Attractive forces between molecules that hold
atoms together; the boiling point of an element or compound reflects the strength of the bonds involved:
• Dipole-Dipole F.: Forces acting between polar molecules; alike charges repel, opposite attract.
Hydrogen Bond.: A special, weak chemical bond (dipole-dipole interaction) between a hydrogen atom
bonded to an atom of high electronegative elements (F, N, or O) and another atom of one of the three
electronegative elements, accounting for the high boiling point of water HF, and NH3.
• Ion-Dipole F.: The force of attraction between an ion and the opposite partial charge of the electric
dipole of a polar molecule; e.g.: HCl, NaCl, AgCl, ect.
• London F. (Dispersion): The force of attraction that arises from the interaction b/w electric dipoles on
neighboring polar / nonpolar molecules; it increases w/ the size of the molecule (molecular mass); e.g.:
hydrocarbon with <4 central C-atoms, gaseous; >5-17<, liquid, >18 waxy solid.
Hybridization: The process of mixing the atomic orbitals in an atom (usually the central atom) to generate a set
of new atomic orbitals - hybrid orbitals in a molecule (see orbital);
Hydrates: Compounds that have a specific number of water molecules attached to them; e.g.: CuSO4⋅5H2O;
Hydride: A binary compound of a metal or metalloid with hydrogen; the term is often extended to include all
binary compounds of hydrogen.
Bonds w/ C,N,O: Generally those compounds wher a hydrogen atom is bonded to; hydrocarbons (C),
ammonia (N), phoshine (P), water (O), hydrogen sulfide (S), hydrogen fluoride (F), etc.
Complex H.: Elements of the 3rd group linked to hydrogen in a tetrahedral bond; e.g.: NaBH4, LiAlH4.
Covalent H.: Hydrogen bonds are bonded covalently with metallic elements of group 3 as well with Be,
and Mg; e.g.: BeH2, GaH3.
Metallic H.: A compound of certain d-block metals and hydrogen; in which H fills the gap in-between the
metallic atoms; e.g.: TiHx (x<2).
Molecular H.: A compound of hydrogen and a nonmetal; e.g.: CH4 (methane), SiH4 (silane).
Saline H.: A compound of hydrogen and a strongly electropositive metal; e.g.: Li+, Na+, K+, Cs+, Fr+, Ca+,
Sr+, Ba+ in which hydrogen is present as H-.
Hydrocarbons: (Gk. hydro, water; L. carbo, charcoal) Organic compound consisting of H and C atoms only.
Aliphatic HC: Hydrocarbons that do not contain the benzene group or the benzene ring.
Aromatic HC: A hydrocarbon that contains one or more benzene rings.
Unsaturated HC: Hydrocarbons that contain carbon-carbon double bonds or carbon-carbon triple bonds.
Ion: An atom or molecule that has lost or gained one or more electrons, and thus becomes positively or
negatively charged; i.e.: Al3+ (mono-atomic ion), SO4- (poly-atomic ion); see also chemistry-atom.
Anion: An ion with a net negative charge, i.e.: F-, SO4-, etc; see table below.
Cation: An ion with a net positive charge, i.e.: Na+, NH4+, Al3+, etc; see table below.
Ionic Compound: Any neutral compound containing cations and anions.
Ion Pair: A species made up of at least 1 cation and at least 1 anion held together by electrostatic forces.
Polarized I.: The distorted electron cloud of an ion (or atom); see also chemistry atom; e.g.: Anion
(>cloud; >charge): I- (r = 220pm) easier distortable than F- (r = 133pm); S2- easier polarizable than Cl-.
Cation (<cloud; <charge): Li2+ more covalent than Cs+; Be2+ more covalent then cations of 4th period.

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Lewis:
L. Formula: (for ionic compounds) A representation of the structure of an ionic compound, showing the
formula unit of ions in terms of their Lewis diagram.
L. Diagram: (for atoms and ions) The chemical symbol of an element, with a dot for each valence
electron.
L. Dot Symbol: The symbol of an element with one or more dots that represent the number of valence
electrons in an atom of the element.
L. Resonance: A blending of Lewis structures into a single composite, hybrid structure; see resonance;
e.g.: O-S=O ↔ O=S-O
L. Structure: A representation of covalent bonding using Lewis dot symbols. Shared electron pairs are
shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on
individual atoms.
Ligand: A molecule or an ion that id bonded to the metal ion in a complex ion.
Lone Pairs: Valence electrons that are not involved in covalent bond formation. LP-electrons occupy more space
and experience a greater repulsion from neighboring lone pairs than bonding pairs; see valence - VSEPR.
Mesomerism: see resonance.
Molecule: 1) Smallest possible unit of a compound, that possesses the chemical properties of the compound. 2)
A definite and distinct, electrically neutral group of bonded atoms; i.e.: H2, NH3, CH3COOH, etc.
M. Equation: Equations in which the formulas of the compounds are written as though all species existed
as molecules or whole units; e.g.: C6H12O6, C2H3CO2H, ect.
M. Orbital: see orbital.
M. Weight: The relative weight of a molecule when the weight of the most frequent kind of carbon atom
(its isomers) is taken as 12; the sum of the relative weights of the atoms in a molecule;
e.g.: MH2O = 2⋅H + O = 18.016 [g/mol]
Diatomic M.: A molecule that consists of two atoms; i.e.: H2, CO, etc.
Homonuclear Diatomic M.: Diatomic molecules containing atoms of the same element; e.g.: O2, H2, N2,
etc.
Polar M.: A molecule that possess a dipole moment; see there.
Octet Rule: An atom other than hydrogen and helium that tends to form bonds until it is surrounded by eight
valence electrons by sharing or transferring them - see also chemistry - stoichiometry.
BX = ½⋅(ΣVe + ΣV8) B, number of bonds
e.g.: N2 (2N) has the equivalent of 10e-, 16e- are possible to Ve, number of valence electrons
obtain noble gas configuration:16-10 = 6 ≅ 3 covalent bonds V8, number of octet-electrons
(:N≡N:); elements from the 2nd period can’t form more than 4 covalent bonds.
Exception to OR.: 1) Rule valid for covalent bonds only; 2) there are irregularities from the 3rd period of
the periodic table onwards, since d-orbitals provide extra electron, seemingly interfering with the octet
rule.
Orbital: The idealized representation of electron moving around its nucleus.
Hybrid O.: Atomic orbitals obtained when two or more nonequivalent orbitals of the same atom combine;
a mixed orbital is obtained by blending together atomic orbitals on the same atom;
e.g.: sp3 hybrid orbital of carbon in CH4: to make 4 covalent bonds for H possible the outer bonding shell
in the C atom has to be altered in the following way:
C = 2s22p2 = 2s(↑↓) + 2p(↑_, ↑_, _ _) → 2sp3 (↑_, ↑_, ↑_, ↑_) with H having only a 1s orbiting electron it
connects via a σ-bond to the hybridized C-atom to give rise to a tetrahedrally shaped molecule.
Molecular O.: An orbital that results from interaction of the atomic orbitals of the covalently bonding
atoms; it spreads over all the atoms in a molecule. The formation of a bonding molecular orbital caused by
the orientation of electron spin within the orbital; e.g.: H-H atom: 1sσ* = antibonding (repelling, opposite
spin); 1sσ = bonding (attraction, same spin);
i) every molecular orbital can accept two electrons;
i) bonding molecular orbitals are energetically far more stable than antibonding MO’s;
e.g.: E1sσ < E1sσ*; or E2sσ < E2sσ* < E2pxσ < E2pyπ; E2pzπ < E2pyπ*; E2pyπ* < E2pxσ*
i) orbitals are filled with electrons after the Aufbau principle, see atom;
i) the overall number of molecular orbitals consists of the sum of participating atomic orbitals.

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Delocalized Molecular O.: Molecular orbitals that are not confined between two adjacent bonding atoms
but actually extend over three or more atoms, or even spread over the entire molecule; e.g.: π-bonding.
• Antibonding MO.: A molecular orbital that is of higher energy and lower stability than the atomic
orbital from which it was formed; energy has to be expended to build this molecule; probability of
finding an electron within the boundary surface of the two atomic orbitals is low - cancellation in the
internuclear region due to opposite spin -1/2 (↓) & +1/2(↑);
• Bonding MO.: A molecular orbital that is of lower energy and greater stability than the atomic orbitals
from which it was formed (energy will be released), hence more stable; (probability of finding an
electron within the boundary surface of the two atomic orbitals is high - addition in the internuclear
region occurs due to same spin -1/2 (↓) & -1/2 (↓) or +1/2(↑) & +1/2(↑).
Covalent Bond.: A bond in which two electrons are shared by two atoms; see bond π-, σ - bond.
Bond Order: The difference between the numbers of electrons in bonding molecular orbitals and
antibonding molecular orbitals, divided by two;
e.g.: H2 = ½⋅(2-0) = 1; bonding dominate over antibonding electrons, therefore H2 is a valid molecule.
e.g.: He2 = ½(2-2) = 0; bonding and antibonding electrons cancel each other, He2 is inexistant molecule.
Resonance (mesomerism): The use of two or more Lewis structures to represent a particular molecule (deducted
from its formal charge - see there).
R. Structure: One of two or more alternative Lewis structures for a single molecule that cannot be
described fully with a single Lewis structure; see also Lewis resonance.
Stereoisomerism: The occurrence of two or more compounds with the same types and numbers of atoms and the
same chemical bonds but different spatial arrangements; e.g. isomers like glucose and fructose = C6H12O6
Valence: The number of bonds that an atom can form.
V. Band: The full band of orbitals in a solid.
V. Bond Theory: The description of bond formation in terms of the pairing of spins in the atomic orbitals
of neighboring atoms.
V. Electron: The outer e of an atom, involved in chemical bonding; i.e.: n = 2 shell of period 2 atoms.
V. Shell: The outermost electron-occupied shell of an atom, usually needed in bonding.
V. Shell-Electron-Pair Repulsion Model (VSEPR): A theory for predicting the shapes of molecules,
using the fact that electron pairs repel one another.
• electron pairs repel another, hence orient themselves in a manner with the least obstruction;
Lone pair versus lone pair repulsion > LP versus bonding pair repulsion > BP versus BP repulsion
• the molecular shape is determined by the position of the central atoms only
CO2 or C2H2: 2 pointed → linear;
BF3 or SO3 or CO3- 3 pointed → trigonal-planar,
CH4 or BF3 or H3PO3: 4 pointed → tetrahedral;
PF5 or SO3: 5 pointed → trigonal-bipyramidal;
SF6 or SeF6: 6 pointed → octahedral;
IF7: 7 pointed → pentagonal-bipyramidal;
V. Shell Expansion: The use of d orbitals in addition to s and p orbitals to form a covalent bond.

Intermolecular attraction (p = 1atm): *) Hydrogen bonding

Molecule Dipole Moment Attracting force [kJ/mol] Melting point Boiling point
[C⋅m] = [D] Dipol-dipol F. London [K] [K]
F.
CO 0.4 0.0004 8.74 74 82
HI 1.3 0.025 27.9 222 238
HBr 2.6 0.69 21.9 185 206
HCl 3.4 3.31 16.8 158 188
NH3 5.0 13.3 14.7 195* 240*
H2O 6.1 36.4 9.0 273* 373*

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Common ionic compounds:

Substance Notation Substance Notation Substance Notation
Azide N3- Hydroxide OH- Nitrate NO3-
Ammonium NH4- Hydrogen Carbonate HCO3- Nitride N3-
Bromide Br- Hydrogen Phosphate HPO4- Nitrite NO2-
Carbonate CO32- Hydronium H3O+ Oxide O2-
Chlorate ClO3- DiHydrogen Phosphate H2PO4- Oxonium OH3+
PerChlorate* ClO4- Hydrogen Sulfate HSO3- PerOxide* O22-
Chloride Cl- Iodide I- Phosphate PO43-
Chlorite ClO2- Iron-II ion Fe2+ Phosphide P3-
HypoClorite ClO- Iron-III ion Fe3+ Phosphonium PH4+
Cobalt-II ion Co2+ Lead-II ion Pb2+ Sulfate SO42-
Cobalt-III ion Co3+ Lead-IV ion Pb4+ Sulfide S2-
Copper-I ion Cu+ Manganese-II ion Mn2+ Sulfite SO32-
Copper-II ion Cu2+ Manganese-III ion Mn3+ Tin-II ion Sn2+
Cyanide CN- Mercury-I ion Hg22+ Tin-IV ion Sn4+
Floride F- Mercury-II ion Hg2+

(*) per (L. allover) the elements max. ability to combine with O; i.e.: highest O-content

Relative electronegativity of elements from the main groups [V]

H 2.2 He -
Li 1.0 Be 1.6 B 2.0 C 2.6 N 3.0 O 3.4 F 4.0 Ne -
Na 0.9 Mg 1.3 Al 1.6 Si 1.9 P 2.2 S 2.6 Cl 3.2 Ar -
K 0.8 Ca 1.0 Ga 1.8 Ge 2.0 As 2.2 Se 2.6 Br 3.0 Kr -
Rb 0.8 Sr 0.9 In 1.8 Sn 2.0 Sb 2.1 Te 2.1 I 2.7 Xe -
Cs 0.8 Ba 0.9 Ti 2.0 Pb 2.3 Bi 2.0 Po 2.0 At 2.2 Rn -

5
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Common molecular compounds

Compound Notation Compound Notation Compound Notation

Acetaldehyde CH3CHO CycloHexane C6H12 Octane C6H18
Acetic Acid CH3COOH Hexane Oxalic Acid (COOH)2
Acetylene C2H2 Oxygen O2
Acetone CH3COCH3 Hydrazine N2H4 Ozone O3
Aniline C6H5NH2 Hydrobromic Acid HBr Phenol C6H5OH
Aluminium Dioxide Al2O3 Hydrobromous Acid HBrO Phosphine PH3
Ammonia NH3 Phosphoric Acid H3PO4
Ammonium Chlorate NH4ClO3 Hydrochloric Acid HCL Phosphourous Acid H3PO3
Ascorbic Acid C6H8O6 Hydrochlorous Acid HClO Phosphorous Pentachloride PCl5
Barium Chloride Dihydrate BaCl⋅2H2O Phosphorous Trichloride PCl3
TetraPhosphorous Decaoxide
Barium Hydroxide Ba(OH)3 Hydrobromic Acid HBr P4O10
Barium Peroxide BaO2 Hydrocyanic Acid HCN CycloPropane C3H6
Benzene C6H6 Hydroflouric Acid HF Propane C3H8
Benzoic Acid CH5CO2H Hydrogen Bromide HBr Propylene C3H6
Butane C4H10 Hydrogen Peroxide H2O2 Potassium Dichromate K2Cr2O7
Potass. Dihydrogenphosphate
DiBorane B2H6 KH2PO4
Borone Triflouride BF3 Hydrogen Sulfide H2S Potassium KOH
Hydroxide
PerBromic Acid HBrO4 Hydroiodic Acid HI Potassium Permanganate KMnO4
Calcium Carbonate CaCO3 Hydrosulfuric Acid H2S
Calcium Oxide CaO Silane SiH4
Calcium Phosphate Ca3(PO4)2 Iodic Acid HIO3 Silicon Carbide SiC
Iodine Heptaflouride IF7 Silicon Tetrachloride SiCl4
Carbon Dioxide CO2 DiIodine Pentaoxide I2O5 DiSilicon Hexabromide Si2Br2
Carbon Disulfide CS2 PerIodic Acid HIO4 Sodiumchloride NaCl
Carbon Monoxide CO Sodiumfloride NaF
Isopropanol C3H8O Sodium Hydroxide NaOH
Cesium Sulfide Cs2S Lead-II Oxide PbO Sodium Peroxide Na2O2

Strontium Nitrate Tetrahydrate Sr(NO3)⋅7H2O

Chloric Acid HClO3 Lithiumcarbonate Li2CO3
Chlorine Triflouride ClF3 Lithium Chloride Monohydrate LiCl⋅H2O
DiChloride Heptaoxide Cl2O7 Lithiumsulfide Li2SO3 Sucrose C12H22O12
PerChloric Acid HClO4
Chloroform CHCl3 Magnesium Sulfate Heptahydrate MgSO4⋅7H2O Sulfuric Acid H2SO4
Chlorous Acid HClO2 Sulfur Dichloride SCl2
Copper-II Bromide CuBr2 Mercury-I Nitrite Hg2(NO2) Sulfur Dioxide SO2
2
Copper-II Nitrite Cu(NO2)2 Sulfur Hexaflouride SF6
Copper Sulfate Pentahydrate CuSO4⋅5H2 Methane CH4 Sulfur Trioxide SO3
O
Ethanol C2H5OH Methanol CH3OH TetraSulfur Tetranitride S4N4
Ethane C2H6 Methylamine CH3NH2 Sulfurous Acid H2SO3
Ethylene C2H4
Formaldehyde HCHO Nitric Acid HNO3 Toluene C7H8
Formic Acid HCO2H Nitrogen Dioxide NO2 Urea CO(NH2)2
Fructose C6H12O6 Nitrogen Trichloride NCl3 Water H2O
Glucose C6H12O6 DiNitrogen N2O5
Pentaoxide
Glycine NH2CH2CO2H DiNitrogen N2O4
Tetraoxide

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Glossary - Chemistry (Stoichiometry)

Adsorption: The process of binding a substance to a surface by releasing energy; see also catalyst.
Chemical A.: Molecules are bound chemically via lone pairs (σ, bonds, covalent-like forces) to the
surface; e.g.: catalytic hydration of olefins.
Physical A.: Molecules are bound via dispersion forces to the surface (releases less energy then in
chemically adsorbed processes; e.g.: charcoal filter.
Alcohol: An organic compound containing the hydroxyl group -OH.
Aldehydes: Compounds with a carbonyl functional group and the general formula RCHO, where R is an H-atom,
an alkyl, or an aryl group.
Allotrops: Two or more forms of the same element that differ significantly in chemical and physical properties.
Alloy: A solid solution composed of two or more metals, or of a metal or metals with one or two nonmetals.
Amalgam: An alloy of mercury with another metal or metals.
Amines: Organic bases that have the functional group -NR2, where R may be H, an alkyl group, or an aryl group.
Arrhenius Behavior: see kinetics.
Aryl Group: A group of atoms equivalent to a benzene ring or a set of fused benzene rings, less one H-atom.
NA - Avogadro’s Number: The number of particles in a mole: 6.02252⋅1023 [particles/mole]
Catalyst: A substance that increases the rate of a reaction without being consumed in the reaction; activation
energy EA is lowered significantly (increases rate of reaction) once catalyst is used in process without
changing the overall energy of reaction (∆H);
e.g.: N2O(g) → (Au) → N2O(docked onto Au-cat)
N2O(docked onto Au-cat) → N2(g) + O(docked onto Au-cat)
2O(docked onto Au-cat) → O2(g) and so forth
Homogeneous C.: A catalyst is homogeneous if it is present in the same phase as the reactants;
e.g.: Br2(aq) for the decomposition of H2O2(aq);
Heterogeneous C.: A catalyst is heterogeneous if it is in a different phase from the reactant;
e.g.: Pt/Rh(s) for the decomposition of NO and CO (Ostwald process)
Chain Reaction: Is a reaction in which an intermediate reacts to produce another intermediate - see radicals, or
physics nuclear decay.
Chemical Equation: see equation.
Compound: A substance composed of atoms of two or more elements chemically united in fixed properties.
Concentration Units: The following equations are commonly used when
operating with liquid mixtures.
β - Mass Concentration: The mass of solute per liter of solution: m(x), mass [g]
β(x) = m(x) / V(Sln) [g/l] V(Sln), volume of solution [l]
w - Mass Percent Composition (weight percentage): The percent
by mass of one solvent over the total mass of solvent and solution:
e.g.: 30% NaCl in 70% water
w(A) = 100⋅m(A) / (m(A) + m(B) + etc.) [%]]
b - Molatity: The number of moles of solute dissolved in one m(A), mass of solute A [kg
kilogram of solvent; b is not affected by temperature change: n(x), molar amount [mol]
b(x) = n(x) / m(Svt) [mol/g] m(Svt), mass of solvent [g]
Remember: A solution with a given molality always maintains the
same molar with whatever solute, as long as the solvent has not changed!
n - Molar Amount: The amount of an element per molar mass:
n(x) = m(x) / M(x) [mol] m(x), mass [g]
c - Molar Concentration (M - molarity): The number of moles M(x), molar mass [g/mol]
of solute in one liter of solution (c increases with temperature): n(x), molar amount [mol]
c(x) = n(x) / V(Sln) [mol/l] V(Sln), volume of solution [l]
x - Molar Fraction: The ratio by mole of one solvent over the
sum of moles of all components:
x(A) = n(A) / (n(A) + n(B) + n(C) + etc.) [-]
Molar Solubility: The number of moles of solute in one liter of a saturated solution (mol/L) = c.

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Equation: An expression showing the chemical formulas of the reactants and products (both in symbols).
Ionic EQ.: An equation that shows dissolved ionic compounds in terms of their free ions;
i.e.: Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) → AgCl(s) + Na+(aq) + NO3-(aq)
Net Ionic EQ.: The equation showing the net change, obtained by canceling the spectator ions in an ionic
equation; i.e.: Ag+(aq) + Cl- (aq) → AgCl (s)
Writing ionic and net ionic EQ: 1) Write a balanced molecular EQ for the reaction; 2) Rewrite the
equation to indicate which substance are in ionic form in solution (all electrolyte in solution dissociate into
anions and cations; group-I elements); 3) Identify and cancel spectator ions (appear on both sides of the
EQ) to arrive at the net ionic EQ;
Balanced EQ.: A chemical equation in which the same number of atoms of each element appear on both
sides of the equation; i.e.: 2H2 + O2 → 2H2O
Skeletal EQ.: An unbalanced equation that summarizes the qualitative information about the reaction;
i.e.: H2 + O2, → H2O,
Product: A substance formed in a chemical equation.
Reactant: A starting material in a chemical reaction; a reagent taking part in a specified reaction.
Symbol: One- or two-letter abbreviation of an element’s name.
Equilibrium: A state in which there are no observable changes as times goes by.
KC - E. Constant: A number equal to the ratio of the equilibrium concentration of gaseous products to the
equilibrium concentrations of gaseous reactants, each raised to the power of its stoichiometric coefficient
(ignored when in their solid or liquid state); by convention, numerator stands for products, and
denominator stands for reactants; KC is temperature dependent;
any solid that precipitate or compound that liquefies, is left out. kF, pF, kR, pR, rate c. of forward /
e.g.: aA +bB ↔ cC +dD reverse reactions [mol/(l⋅s)]
KC = k F/k R = cc(C)⋅cd(D) / (ca(A)⋅cb(B)) [var] KC > 103: favors products strongly
KP = pF/pR = pc(C)⋅pd(D) / (pa(A)⋅pb(B)) [var] KC > 10-3-103: reactants & products at equilibrium.
∆n
where KP = KC⋅(R⋅T) KC < 10-3: favors reactants strongly
∆n, moles of gaseous products - moles of gaseous reactants KC, EC for concentrations
QC - Reaction Quotient: The ratio of the product of the KP, EC for partial pressures
concentrations of the products to that of the reactants (defined c, molar concentration [mol/l]
like the equilibrium constant) at an arbitrary stage of reaction; p, partial pressure [N/m2]
e.g.: N2(g) + 3H2(g) → 2NH3(g) ∆n = (c+d) - (a+b) [mol]
QC < KC: reverse-reaction towards products (shift to right) KC = c2(NH3)/(c(N2)⋅c3(H2))
QC = KC: system in equilibrium QC = c02(NH3)/(c0(N2)⋅c03(H2))
QC > KC: forward-reaction towards reactants (shift to left) c0(X), initial arbitrary c. [mol/l]
E. Vapor Pressure: The vapor pressure measured under
dynamic equilibrium of condensation and evaporation; see chemistry liquid and physics - heat.
Dynamic E.: The condition in which a foreword process and its reverse are occurring simultaneously at
equal rates; e.g.: vaporizing and condensing; chemical reactions at equilibrium, etc.; e.g.:
H2(g) + I2(g) ↔ 2HI(g); vF, vR, rate of forward/reverse
vF = k F⋅c(H2)⋅c(I2); vR = k R⋅c2(HI); k F, k R, rate constants [1/s]
vF = vR in equilibrium: c2(HI)/(c(H2)⋅c(I2)) = k F/k R = K c(x), molar concentration [mol/l]
Heterogeneous E.: An equilibrium state in which the reacting species are not all in the same phase.
e.g.: CaCO3(s) ↔ CaO(s) + CO2(g) KC = c(CaO)⋅c(CO2) / c(Ca2CO3)
solid substance has to be converted from density [g/cm3] to molar concentration [mol/l].
Homogeneous E.: An equilibrium state in which all reacting species are in the same phase.
e.g.: N2O4(g) ↔ 2NO2(g) KP = p2(NO2) / p(N2O4)
Multiple E.: If a reaction can be expressed as the sum of two or more reactions, the equilibrium constant for
the overall reaction is given by the product of the equilibrium constants of the individual reactions: e.g.:
H2CO3(aq) ↔ H+(aq) + HCO3-(aq) K1 = c(H+)⋅c(HCO3-) / c(H2CO3)
HCO3 (aq) ↔ H (aq) + CO3 (aq)
- + 2-
K2 = c(H+)⋅c(CO32-) / c(HCO3-)
H2CO3(aq) ↔ 2H (aq) + CO3 (aq)
+ 2-
KC = K1⋅K2 KC = c2(H+)⋅cCO32-) / c(H2CO3)
Law of Mass Action: For an equilibrium of the form aA + bB ↔ cC + dD, the reaction quotient
QC = cc(C)⋅ cd(D) / ca(A)⋅ cb(B); evaluated by using the equilibrium molar concentrations of the reactants and
products, is equal to a constant KC which has a specific value for a given reaction and temperature.
Van’t Hoff Equation: It shows how the equilibrium constant KC
varies with temperature. T, temperature [K]
ln(KC2/ KC1) = -(1/ T2 - 1/ T1)⋅∆H°/R KC, equilibrium constant [var]
Functional Group: That part of a molecule characterized by a special R, gas constant 8.314 [kJ/mol]
arrangement of atoms that is largely responsible for the chemical ∆H°; std enthalpy of react.
[kJ/mol]
behavior of the parent molecule.
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Formula (Chemical F.): A collection of chemical symbols and subscripts showing the composition of a
substance.
Empirical .F.: A chemical formula that shows the relative numbers of atoms of each element in a
compound, i.e.: CH for acetylene (C2H2)or benzene (C6H6), etc.
Molecular F.: A combination of chemical symbols and subscripts showing the actual numbers of atoms of
each element present in a molecule; i.e.: H2O, SF6, C6H12O6; etc.
Structural F.: A chemical formula that shows the groupings of atoms in a compound.
Lewis Dot Symbol: The symbol of an element with one or more dots that represent the number of valence
electrons in an atom of the element.
Lewis Structure: A diagram showing how electron pairs are shared between atoms in a molecule;
e.g.: H-Cl::: or ::O=C=O::
Haber Bosch Process: The catalyzed synthesis (Pt) of ammonia at high pressure and high temperature:
N2(g) + 3H2(g) ↔ 2NH3
Half Life: see kinetics.
Hydrogenation: The addition of hydrogen, especially to compounds with double and triple-carbon bonds.
Ion: (Gk, to go) An atom or molecule that has lost or gained one or more electrons, and thus becomes positively
or negatively charged; i.e.: Al3+ (mono-atomic ion), SO4- (poly-atomic ion); see chemistry atom &
molecule.
Complex I.: An anion containing a central metal cation bonded to one or more molecules or ions; e.g.:
Ag(NH3)2+, CdCl42-, Cu(NH3)42+, Cu(OH2)42+, Fe(CN)63-, Fe(CN)64+, Zn(OH)42-;
Spectator I.: Ions that are present in the reaction but remain unchanged (compare net ionic w/ ionic eq.)
Kinetics: The speed, or rate, at which a chemical reaction occurs. Ea, activation energy [J/mol]
Arrhenius Behavior: A reaction shows Arrhenius behavior k, rate constants [1/s]
if a plot of ln(k) against 1/T is a straight line. A, collision frequency [-]
e.g.: k = A⋅e-Ea/(R⋅T); ln(k) = ln(A)-(Ea/(R⋅T) e, eulers number 2.7183 [-]
for a given reaction at 2 different T’s: R, gas constant 8.314 [J/mol]
ln(k2/k1) = EA⋅(T2-T1)/(R⋅T1⋅T2) T, temperature [K]
K. Theory (kinetic molecular theory); A theory of properties of an ideal gas in which point-like molecules
are in continuous random motion and there are no interactions between them; see chemistry gas.
t1/2 - Half Life: The time required for the concentration of a reactant to decrease to half of its initial
concentration; see O, 1st, 2nd order reactions.
Kinetic Reactions:
k - KR. Constant: The constant of proportionality in a rate law; k is dependent upon the temperature and
can only be determined experimentally; e.g.: c(A) at t1 - c(A) at t2/(t2 - t1)
k = decrease of c(reactants)/t [l/(mol⋅s)] or [1/s] or [mol/(l⋅s)] depending upon the rate order.
Rate Law: An EQ expressing the instantaneous reaction rate in terms of the concentrations, at that instant,
of the substances taking part in the reaction; it can only be determined experimentally; it cannot be written
down from the stoichiometry of the chemical equation for the reaction;
e.g.: NO2(g) + CO(g) →NO(g) + CO2(g) k[NO2]2
KR. Order: The power to which the concentration of a single substance is raised in a rate law;
e.g.: k[SO2]⋅[SO3]-1/2, than the reaction is 1st order in SO2 and of order -½ in SO3 and +½ overall.
• Pseudo 1st Order KR.: A reaction with a rate law that is effectively first order because one substance
has a virtually constant concentration.
• 0 Order KR.: A reaction with a rate that is independent of the concentration of the reactant; i.e.: the
catalyzed decomposition of ammonia; plot of cA is linear over t:
e.g.: v(Ax) = -k [mol/(l⋅s)] c(A) = -k⋅t + c(A0) [mol/l] t1/2 = c(A0) / (2⋅k) [s]
• 1st Order KR.: A reaction whose rate depends on reactant concentration raised to the first power;
(directly proportional to the concentration); plot of ln(cA) is linear over t:
e.g.: v(Ax) = -k⋅c(Ax); [1/s]⋅[mol/l] c(A) = c(A0)⋅e-k⋅t [mol/l] t1/2 = ln(2) /k [s]
• 2 Order KR.: A reaction whose rate depends on reactant concentration raised to the second power or
nd

on the concentration of two different reactants, each raised to the first power (directly proportional to
the square of the concentration); plot of 1/cA is linear over t:
e.g.: v(Ax) = k⋅c2(Ax) [l/(mol⋅s)]⋅[mol2/l2] 1/c(A) = k⋅t + 1/c(A0) [mol/l] t1/2 = 1 / (k⋅c(A0)) [s]
v(Ax) - KR. Rate: The change in concentration of a substance divided by the time it takes for the change to
occur: an increase in product automatically implies a decrease in reactant and vice versa;
e.g.: v(Ax) = dc(Ax)/dt [mol/(l⋅s)] with v(Ax) positive = increase; v(Ax) negative = decrease.
KR. Sequence: A series of reactions in which products of one reaction take part as reactants in the next;
e.g.: 2C(s) + O2(g) → 2CO(g), followed by 2CO(g) + O2(g) → 2CO2(g)
• Mono-Molecular KRS.: A single molecule forms the product; v = k⋅c(A); e.g.: radioactive decay
• Di-Molecular KRS.: Involves two molecules; v = k⋅c(A)⋅c(X); with A and X 1st order, with the overall
reaction itself as 2nd order.
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Law of Definite Proportions: Different samples of the same compound always contain its constituent elements
in the same proportions by mass.
Law of Mass Action: see equilibrium.
Law of Multiple Proportions: If two elements can combine to form more than one compound, the masses of
one element that combine with a fixed mass of the other element are in ratios of small whole numbers.
Le Chatelier’s Principle: If an external stress is applied to a system at equilibrium, the system will adjust itself
in such a way as to partially offset the stress; i.e.: a reaction at equilibrium tends t o proceed in the
endothermic reaction when the temperature is raised.
• Change of concentration: When the concentration of any of the reactants or products at equilibrium are
changed, the position of the equilibrium shifts so as to reduce the change in concentration that was
made, with the net effect that KC remains unchanged.
• Change of pressure and volume: i) In gaseous substances with ∆n≠0 changes of pressure and volume
do influence the equilibrium; e.g.: The amount of product (NH3) at equilibrium increases with rising
pressure; N2 + 3H2 ↔ 2NH3 + heat since the formation of two molecules of NH3 causes the
disappearance of four molecules of reactants (N2 + 3H2), thus decreasing the total number of
molecules.
i) In reactions involving liquids, moderate pressure changes do not affect the equilibrium;
i) In reactions involving gaseous substance with ∆n=0, changes of pressure will not affect the equilibrium;
• Change of Temperature: In gaseous substances a temperature increase favors an endothermic reaction,
and a temperature decrease favors exothermic reactions:
T-increase (T2 < T1): exothermic reaction: ∆H° < 0; K2/K1<1 or K2<K1
endothermic reaction: ∆H° > 0; K2/K1>1 or K2>K1; equilibrium concentration of products
increases; position of equilibrium shifts to right; K increases;
T-decrease (T2 < T1): endothermic reaction: ∆H° > 0; K2/K1<1 or K2<K1;
exothermic reaction: ∆H° < 0; K2/K1>1 or K2>K1; equilibrium concentration of products decreases;
position of equilibrium shifts to left; K decreases;
Lewis: Lewis dot symbol and Lewis structure see formula.
Limiting Reagent: The reactant used up in a reaction - see yield.
Lysis: (Gk. lysis, loosening) A process of disintegration or destruction.
Glycolys.: Splitting of glucose molecules; (C6H12O6) into two molecules of pyruvate, resulting in the
release of energy (exothermic) in the form of two ATP molecules; does not require O2.
Hydrolys.: The splitting apart of two covalently bound subunits by the insertion of a molecule of water;
the OH-group being incorporated into one fragment and the H-atom in the other; reverse process of
synthesis occurs in digestion.
Protolys.: Dissociation of protons in an acid HA ↔ H+ + A-.
Mass: The quantity of matter in a object, or concentrated energy (see SI-units, weight).
M. Defect: The difference between the mass of an atom and the sum of the masses of its protons,
neutrons, and electrons; since part of the mass is needed to hold the atom together (potential energy).
A - Mass Number: The total number of neutrons and protons present in the nucleus of an atom; i.e.: 146C
has 14 nucleons (6 protons and 8 neutrons).
Atomic Mass Unit (amu): A mass exactly equal to one 12th of the mass of one carbon-12 atom (see
molecular mass): 1 amu = m12C / 12 = 1.6605⋅10-24 [g]
Average Atomic M.: see chemistry atom.
Law of Conservation of M.: Matter can be neither destroyed nor created.
Mass Units: The following equations are commonly used in dealing with masses in chemical equations:
β - Mass Concentration: The mass of solute per liter of solution: m(x), mass [g]
β(x) = m(x) / V(Sln) [g/l] V(Sln), volume of solution [l]
ρ - Density: The mass of a substance divided by its volume m(Sln), mass [g]
ρ(Sln) = m(Sln) / V(Sln) [g/l] V(Sln), volume of solution [l]
Mr - Molecular M.: The sum of the atomic masses (in amu) present in the molecule;
i.e.: Mr(H2O) = 2⋅1.008 (atomic mass of H) + 16.00 (atomic mass) of O = 18.02 [amu].
n - Molar Amount: The amount of an element per molar mass: m(x), mass [g]
n(x) = m(x) / M(x) [mol] M(x), molar mass [g/mol]
M - Molar Mass: The relative mass (g, or kg) per mole of atoms
(amu) , molecules, or other particles - see also atom-mass: mav, average mass [amu]
M(x) = mav⋅NA [g/mol] NA, Avogadro’s c.6.022⋅1023 [atoms/mol]
i.e.: M(H2O) = 2⋅1.008 [g/mol]of H + 16.00 [g/mol] of O = 18.02 [g/mol];
Molar Mass Fraction: Ratio of the number of moles of one
component of a mixture to the total number of moles of all components in the mixture; i.e.: the respective
mass of H and O in a given sample of water is obtained by:
mH = mH2O ⋅ 2⋅MH/MH2O [g]; mO = mH2O ⋅ MO / MH2O [g];
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Molar: The quantity per mole; i.e.: molar mass (the mass per mole), molar volume (the volume per mole), etc.
b - Molality: see chemistry-liquids.
c - Molar Concentration: see chemistry-liquid.
Molar Solubility: see chemistry-liquid.
c - Molarity: see chemistry-liquid.
Mole: (L, massive, heap) The SI base unit for the amount of substances that contains as many elementary entities
(atoms, molecules, or other particles) as there are atoms in exactly 12 grams of the carbon-12 isotope;
always equal to Avogadro’s number = 6.02205⋅1023.
M. Method: The approach of determining the amount of product formed in a reaction, based on the fact
that the stoichiometric coefficients in a chemical EQ can be interpreted as the number of moles of each
substance; the MM consists of the following steps:
1. Write correct formula for all reactants and products and balance the resulting EQ;
2. Convert the quantities of some or all given or known substances (usually reactants) into moles;
3. Use the coefficients in the balanced EQ to calculate the number of moles of the sought or unknown
quantities (usually products);
4. Using the calculated numbers of moles and the molar masses, convert the unknown quantities to
whatever units are required (usually grams);
5. Check that the answer is reasonable in physical terms
Octet Rule: An atom other than hydrogen and helium that tends to form bonds until it is surrounded by eight
valence electrons by sharing or transferring them - see chemistry solid.
Percent Composition: The percent by mass of each element in a compound.
Products & Reactant: see equation.
Radical: An atom, molecule, or ion with at least one unpaired electron; e.g.: NO, •:O:•, •CH3;
R. Reaction Mechanism: The pathway that is proposed for an overall reactions and accounts for the
experimental rate law; e.g.: H2 + Cl2 → 2HCl splits into 4 sub-reactions:
• Initiation: Cl2 →(blue light) → 2Cl• blue contains more energy than red light
• Chain R.: Cl• + H2 → 2HCl + H•; H• + Cl2 → HCl + Cl•...in a closed loop
• Back R.: H• + HCl → H2 + Cl•
• Termination: Cl• + Cl• → Cl2 + E; H• + H• → H2 H• + Cl• → HCl
once the light is switched off....
Reaction: A chemical change in which one substance responds to the presence of another, to a change of
temperature, or to some other influence; i.e.: the sequence of elementary steps that leads to product
formation.
QC - R. Quotient: see equilibrium.
Elementary R.: An individual reaction step in a reaction mechanism; e.g.: H• + Cl• → HCl.
Reversible R.: A reaction that can occur in both directions until dynamic equilibrium is reached, see
there;
e.g.: H2(g) + I2(g) ↔ 2HI(g)
Classification of R.: first 5 based on the composition of products and reactants.
1. Synthesis: Formation of compounds from simpler starting materials; e.g.: 2H2(g) + O2(g) → 2H2O(l)
2. Decomposition: Formation of simpler substances from more complex starting materials;
e.g.: CaCO3(s) → CaO(s) + CO2(g)
3. Replacement R.: Exchange of partners; e.g.: 2NaCl(aq) + Pb(NO3)2(aq) → 2NaNO3(aq) + PbCl2(s)
4. Combustion: Reaction with oxygen to form CO2, H2O, N2, and oxides of any other elements present;
e.g.: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
5. Corrosion: Reaction of a metal with oxygen to form the metal oxide; e.g.: 4Fe(s) + 3O2(g) →
2Fe2O3(s)
Gasevolution: Formation of gas (driving force: escape of gas);
e.g.: CaCO3(s) + 2HCl → CaCl2(s) + H2O(l) + CO2(g)
Precipitation: Formation of precipitate when one solution is added to another (driving force: formation of
insoluble precipitate); e.g.: 3CaCl2(aq) + 2Na3PO4(aq) → Ca3(PO4)2(s) + 6NaCl(aq)
Neutralization: Reaction between any acid and a base (driving force: formation of solvent)
e.g.: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
RedOx: Transfer of electrons from one species to another (accompanied by atoms in many cases - driving
force: e-transfer to achieve greater stability); see chemistry redox-reaction;
e.g.: 2Mg0(s) + O20(g) → 2Mg+1O-1(s)

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Stoichiometry: The mass relationships among reactants and products in chemical reactions.
S. Amount: The exact molar amount of reactants and products that appear in a balanced chemical EQ.
S. Coefficient: The number of moles of each substance in a chemical equation;
i.e.: 1 and 2 in: H2 + Br2 → 2HBr
S. Point: The stage in a titration when exactly the right volume of solution needed to complete the
reaction has been added (see chemistry - acid and base).
S. Proportions: Reactants in the same proportions as their coefficients in the chemical equation; i.e.:
equal amounts of H2 and Br2 in the reaction mentioned above.
S. Relation: An expression that equates the relative amounts of reactants and products that participate in a
reaction; i.e.: 1 mol H2 = 2 mol HBr.
Reaction S.: The quantitative relation between the amounts of reactants consumed and products formed in
chemical reactions as expressed by the balanced chemical equation for the reaction.
Titration: The analysis of composition by measuring the volume of one solution (the titrant) needed to react with
a given volume of another solution (the anylate).
Anylate: The solution of unknown concentration in a titration.
Titrant: The solution of known concentration added from a buret in a titration.
Van’t Hoff Equation: see equilibrium.
Yield: The outcome of a chemical reaction, expressed in grams, mole, liters, etc.
Y. of Reaction (actual yield): The quantity of product obtained from the reaction.
Percentage Y.: The percentage of the theoretical yield of a product achieved in practice
Y% = 100⋅YA/YT [-]: YA, achieved yield [g]
Theoretical Y.: The amount of product predicted by the YT, theoretical yield [g]
balanced equation when all of the limiting reagent has reacted.
Limiting Reagent: The reactant that governs the theoretical yield of product in a given reaction.

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Glossary - Chemistry (Gas)

Avogadro’s Law: Equal volumes of different gases at the same temperature and pressure contain equal numbers
of molecules. 1 [mol] (6.022⋅1023 molecules) of gas at 0[°C] and 1[bar] occupies 22.414 [l] = mole-
volume.
Boyle’s Law: At a given temperature, the product of pressure and volume of a given mass of gas is constant; the
volume of a fixed amount of gas maintained at constant temperature is inversely proportional to the gas
pressure: p1⋅V1 = constant = p2⋅V2 [Pa⋅m3]
Colligative Properties: see chemistry - liquid.
Collision Theory: The theory of elementary gas-phase bimolecular reactions in which it is assumed that
molecules react only if they collide with at least enough kinetic energy for bonds to be broken.
Activated Complex T.: A theory of reaction rates in which it is postulated that the reactants form an
activated complex; i.e.: compounds temporarily formed by the reactant molecules as a result of collision.
Activated Complex: A combination of the two reactant molecules that can either go on to form products
or fall apart into the unchanged reactants;
e.g.: A2 + X2 ↔ A2X2 → 2AX; where A2X2 is very short-living.
EA - Activation Energy: The minimum energy needed for reaction; the height of the activation barrier;
is usually endothermic, until activated complex-level is reached, followed by a more or less exothermic
reaction; resulting in a net exo-/endothermic reaction: EAU, Uphill-EA [J]
e.g.: ∆U = EAU - EAD [J] EAD, Downhill-EA [J]
Condensation: The phenomenon of going from gaseous state to the liquid state (see physics - matter).
C. Reaction: A reaction in which two smaller molecules combine to form a larger one. Water is invariably
one of the products of such a reaction.
Dalton’s Law of Partial Pressures: see pressure.
ρ - Density: Mass of a substance per unit volume; m, mass [kg]
ρ = m/V [kg/m3] V, volume [m3]
e.g. Flour: ρ(F2) = MF2/Vm = 38/22.4 [g⋅mol/(l⋅mol)] = [g/l] R, gas constant 8314 [kJ/mol]
Diffusion: The spreading of one substance through another one; T, temperature [K]
the speed of diffusion can be expressed as: M, molar mass [g/mol]
v = √(3⋅R⋅T/M)
Effusion: The process by which a gas under pressure escaped from one compartment of a container to another by
passing through a small opening.
Graham’s Law of E.: The rate of effusion of a gas is inversely t, time [s]
proportional to the square root of its molar mass: M, molar mass [g/mol]
tA/tB = √( MA/ MB)
Gas: A fluid form of matter that fills the container it occupies and can easily be compressed into a much smaller
volume (gas is a substance at higher temperature than its critical temperature, vapor is a gaseous form of
matter at a temperature below its critical temperature).
1[mol] (6.022⋅1023 molecules) of an ideal gas at 0[°C] and 1[bar] occupies 22.414 [l] = mole-volume.
Ideal G.: A hypothetical gas whose pressure-volume-temperature behavior
can be completely accounted for by the ideal gas equation n, molar amount [mol]
• Ideal G. EQ.: It expresses the relationships among . R, gas c. 8,314 [J/(K⋅mol)]
p, V, T, and amount of V, volume [l] [m3]
p⋅V = n⋅R⋅T [kg⋅m] gas T, Temperature [K]
Real G.: Effects of intermolecular forces and the fact that the molecules or atoms involved possess small
but distinct volume deviate (n = P⋅V/R⋅T) versus p significantly; pO, observed pressure [Pa]
• Van der Waals EQ: An equation that describes a, proportionality c. [-]
p, V, and T of a non-ideal gas. n, number of molecules [mole]
pIdeal = pO + a⋅n2/V2 [N/m2] = [Pa] V, volume [m3]
n⋅R⋅T = (pO + a⋅n2/V2)⋅(V-n⋅VX) [kg⋅m] VX, molecular volume [m3]
Kinetic G. Theory: A theory of properties of an ideal gas in which
point-like molecules with no interactions between them, are in continuos motion at temperature > 0 [K]:
1. A gas is composed of atoms or molecules, which are far apart from each other, empty space in-b/w.
2. Gas molecules are in constant random (brownian) motion along a straight line until collision with the
container wall or other gas-molecules; the higher the temperature, the faster they move;
3. The molecules exert no (repelling/attracting) force on each other or on the container; collisions are
elastic - their total energy of the molecules remains constant; hence pressure is directly related upon
the speed and number of collisions of molecules against the container walls;
4. The average kinetic energy (KE) of the molecules of a gas is proportional to the absolute temperature.
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vD - Maxwell-Boltzmann Speed Distribution: Displays the NA, Avogadro c. 6,022⋅1023 [1/mol]

most probable spectrum of molecular speeds available to the v, speed [m/s]
system at a particular temperature; e.g.: molecular oxygen m, mass [g]
has an average speed of 200[m/s] at 73[K], whereas it n, molar amount [mol]
increases to about 400[m/s] at 273[K]: R, gas c. 8,314
[J/(K⋅mol)]
KE = ½⋅m⋅vD2 [kg⋅m2/s2] T, temperature [K]
p⋅V = n⋅NA⋅KE⋅2/3 = n⋅R⋅T [m3⋅N/m2] M, molar mass [g/mol]
= NA⋅KE⋅2/3 = R⋅T [kg⋅m/mol]
= NA⋅m⋅vD2⋅2/6 = M⋅v2⋅1/3 → vD = √(3⋅R⋅T/M) see diffusion.
Charles and Gay Lussac’s Law: The volume of a fixed amount of gas maintained at constant pressure is
directly proportional to the absolute temperature of a gas; i.e.:
the volumes of the reactants and the gaseous products V, volume [m3] or [l]
are always in the ratio of small whole numbers; T, temperature [K]
V1/T1 = constant = V2/T2 vu, volumetric unit 22.41 [l]
e.g.: 2⋅vu of CO(g) + 1⋅vu of O2(g) → 2⋅vu of CO2(g)
Henry’s Law: The solubility of a gas in a liquid is proportional to the pressure of the gas over the solution; see
chemistry liquids.
Kinetic Theory (kinetic molecular theory); A theory of properties of an ideal gas in which point-like molecules
are in continuous random motion and there are no interactions between them.
Maxwell-Boltzmann Speed Distribution: see gas -kinetic gas theory.
X - Mole Fraction: see pressure, or chemistry stoichiometry - concentration units
Phase: A particular state of matter; a homogeneous part of a system in contact with other parts of the system but
separated from them by a well defined boundary; a substance may exist in solid, liquid, and gas phases
and in certain cases, in more than one solid phase; e.g.: white and gray tin are two solid phases of tin; ice,
liquid, and vapor are three phases of water; etc. see chemistry introduction.
P. Change: Transformation from one phase to another.
P. Diagram: A summary in graphical form of the conditions of temperature and pressure at which the
various phases of a substance exist; the cross point of vaporization-, and fusion conduct to the triple point.
Pressure: Force applied per unit area. p, pressure [N/m2] [Pa]
Atmospheric P.: The pressure exerted by Earth’s atmosphere. 1 bar = 105 [Pa]
Critical P.: The minimum pressure necessary to bring about 1 atm = 1.01325 [bar]
liquefaction at the critical temperature. 1 atm = 101.325 [kPa]
px - Partial P.: The pressure of one component in a mixture of
gases it would exert if it alone occupied the container: px, partial pressure [N/m2] [Pa]
pΣ = pA + pB + pC + pD + etc.
Dalton’s Law of partial P.: The partial pressure of a gas in a mixture is independent of other gases
present; the total pressure is the sum of the partial pressure of all gasses present:
pA = xA⋅pT [Pa] xA, mole fraction [-]
Standard Atmospheric P.: The pressure that supports a pT, total pressure [Pa]
column of mercury exactly 76 cm high at 0°[C] at sea level = 101[kPa].
Standard Temperature and P. (STP): 0°[C] and 101[kPa].
X- Mole Fraction: A dimensionless quantity expressing the ratio of moles of one component to the
number of moles of all components present; see stoichiometry - concentration units; e.g.: 2-gas mixture A
& B:
XA + X B = 1
XA = nA/(nA + nB) [-] n, molar amount [mol]
Radical: see stoichiometry.
Sublimation: The process in which molecules go directly from the solid into the vapor phase.
Temperature: How hot or cold a sample is; the intensive property that determines the direction in which heat
will flow between two objects in contact.
Critical T.: The temperature above which a substance cannot exist as a liquid i.e.: gaseous.
Triple Point: The point at which the vapor, liquid, and solid states of a substance meet in a phase diagram and
are in dynamic equilibrium; e.g.: TP of H2O = 0.0099[°C] and 0.61[kPa]
Van der Waals Equation: see gas - real gases.
Vapor: The gaseous phase of a substance (specifically, of a substance that is a liquid or a solid at the temperature
in question); see chemistry - liquid.
V. Pressure: The pressure exerted by the vapor of a liquid (or solid) when the vapor and the liquid (or
solid) are in dynamic equilibrium; the higher the intermolecular forces the lower the vapor pressure;
e.g.: H2O(g) ↔ H2O(l)
• VP Lowering: see colligative properties.

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Glossary - Chemistry (Liquid)

Boiling Point: The temperature at which the vapor pressure of a liquid is equal to the external atmospheric
pressure (KELiquid = KE of surrounding environment); the higher the intermolecular forces the lower the
vapor pressure the higher is also the boiling point
BP Elevation: see colligative properties.
Colligative Properties: A property that depends only on the number of solute particles present in a solution and
not on their chemical composition;
Boiling Point Elevation: The vapor pressure at any ∆TS, change in temperature [°C]
temperature of a given solution is lower than that of the pure bSle, molality of solute [mol/g]
solvent; the liquid- vapor curve for the solution lies below that kb-Slt, boiling point coefficient of
of the pure solvent intersecting the boiling point at a higher solvent [°C/(g⋅mol)]
temperature: ∆Tb = kSlt⋅bSle [°C]
Freezing Point Depression: The lowering of the freezingpoint of a solution caused by the pressure of a
solute; when the solution begins to freeze, only the solvent
solidifies and solute is left behind in the solution, raising the ∆TS, change in temperature [°C]
concentration of the solute as well as lowering the freezing kf-Slt, freezing point coefficient of
point of the solvent: solvent [°C/(g⋅mol)]
∆Tf = kSlt⋅bSle [°C]
Osmosis: (Gk. osmos, impulse or thrust) The diffusion of water, or any other solvent, across a
differentially permeable membrane; in the absence of other forces, the movement of water during osmosis
will always be from the region of greater potential to one of lesser water potential.
• O. Pressure: The pressure needed to balance the flow of c, molar concentration [mol/l]
solvent through a semipermeable membrane; R, gas constant 8,314 [J/K⋅mol]
i.e.: the pressure required to stop osmosis: T, absolute temperature [K]
π = c⋅R⋅T = n⋅R⋅T/V [mol⋅N⋅m⋅K/(K⋅mol⋅m3)] = [N/m2] = [Pa] n, molar amount [mol]
• Reverse O.: The passage of solvent out of a solution when V, volume [m3]
a pressure greater than the osmotic pressure is applied on the
solution side of a semipermeable membrane; e.g.: a method of desalination using high pressure to
force water through a semipermeable membrane from a more concentrated solution to a less
concentrated solution.
Vapor Pressure Lowering: The presence of a non-volatile solute raises the boiling point of the solvent,
because it can block the escape of solvent molecules but has no effect on the rate of return of the solvent
particles from the vapor to the solution.
Complex Formation: A metal ligand coordinate-covalent bond formation.
Complex Ion: Ions containing a central metal cation bonded to one or more molecules or ions like the
following sample show: Ag(NH3)2+, CdCl42- Cu(NH3)4+, Fe(CN)63-, Fe(H2O)63+, PtCl4-, etc.
e.g.: AgCl(s) + 2NH3(aq)↔ Ag(NH3)2+(aq) + Cl-(aq)
according to Le Chattelier’s principle, the removal of Ag+ ions from the solution to form Ag(NH3)2+ ions
will cause more AgCl to dissolve, whereas if NH3 is absent, more AgCl would precipitate.
Ligand: A group attached to a central metal ion in a complex, i.e.: molecules or ions that surround the
metal atom and the ligands can be thought of a Lewis acid-base reaction; every ligand has at least one
unshared pair of valence electrons (lone pairs), hence is a Lewis base; e.g.: H2O, NH3, CO, Cl-, etc.
Kf CF. Constant: A measure of the tendency of a metal to form a particular complex ion; the larger Kf,
the more stable the complex ion;
Concentration Units: see chemistry stoichiometry.
Condensation: The phenomenon of going from gaseous state to the liquid state (see physics - matter).
C. Reaction: A reaction in which two smaller molecules combine to form a larger one. Water is invariably
one of the products of such a reaction.
Dalton’s Law of Partial Pressures: The partial pressure of a gas in a mixture is independent of other gases
present; the total vapor pressure is the sum of the partial pressure of all gasses present (see there):
Deposition: The process in which the molecules go directly from the vapor into the solid phase.
Desalination: Purification of seawater by the removal of dissolved salts; see osmosis.
Diffusion: (L. diffundere, to pour out) The net movement of suspended or dissolved particles from a more
concentrated region to a less concentrated region by virtue of their kinetic energy (result of the random
movement, brownian motion - see physics-matter) of individual molecules; the process tends to distribute
such particles uniformly throughout the medium; see also chemistry - gas.
Dilution: see solution.

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Distillation: The separation of a mixture by making use of the different volatilities of its components;
Fractional D.: A procedure for separating liquid components of a solution that is based on their different
boiling points; based on Dalton’s law of partial pressures;
see chemistry-gas: xA, mole fraction [-]
pA = xA⋅p [Pa] pT total pressure [Pa]
Vacuum D.: Boiling off humidity by lowering the temperature below freezing and subsequently
decreasing the pressure; e.g.: freeze-dried coffee.
Electrolyte: 1) An ionically conducting medium. 2) A substance that, when dissolved in water, results in a
solution that can conduct electricity; see table below.
E. Rule: For a net potential of zero, the positive and negative charges must add up to zero; a solution must
contain essentially as many anionic as cationic charges.
Non-E.: Is a solution in which no proportion of the solute molecules are ionized, hence does not conduct
electricity; e.g.: C6H12O6(aq); see table below
Strong E.: Is a solution in which a large proportion of the solute molecules are ionized (complete
dissociation into ions); 1mol K2SO4 → 3mol ions
e.g.: NaCl(s) → Na+(aq) + Cl-(aq) 1mol NaCl → 2mol ions
Weak E.: Is a solution in which only a small proportion of the solute molecules are ionized ionized (partly
dissociation into ions); e.g.: CH3COOH(aq); see table below
Equilibrium: The state of final balance of a multi-compound homogenous mixture;
for KC, QC, Dynamic E. etc., see chemistry-stoichiometry equilibrium.
Evaporation: The escape of molecules from the surface of a liquid; also called vaporization.
Freezing Point: The temperature at which a liquid freezes (crystallization); the normal freezing point is the
freezing temperature under a pressure of 1atm, i.e.: liquid and solid phase are in equilibrium.
FP Depression: see colligative properties.
Henry’s Law: The solubility of a gas in a liquid is proportional to the pressure of the gas above the solution;
solubility ∝ c; applicable only if: solution is low in concentration;
low pressures; solvent does not react with gas. k, Henry’s variable [mol/(l⋅Pa)]
c = k⋅p [mol/l] p, partial pressure [1/atm] [Pa]
Hydration: A process in which an ion or a molecule is surrounded by water molecules arranged in a specific
manner; e.g.: water - H2O molecules attach to a central ion (ion-dipole interaction).
Hydrated Anion: Hydrogen bonds form between the H of water and the central anion; e.g.: SO42-.
Hydrated Cation: Ion-dipole forces between the O of water and the central ion; e.g.: Be2+.
H. Crystals: Hydrated ions remain intact even in a solidified structure;
e.g.: [Fe(OH2)6]3+ Cl33- actual structure notation: Fe2Cl3⋅6H2O
or [Cu(OH2)4]2+ [SO4(H2O)]2- notation: CuSO4⋅5H2O
Enthalpy of H.: see chemistry - thermochemistry.
Hydrophillic: Water-liking.
Hydrophobic: Water-fearing.
Intermolecular Forces: see chemistry molecule.
Law of Mass Action: see chemistry stoichiometry - equilibrium
Ligand: see complex formation.
Liquid: A fluid form of matter that takes the shape of the part of a container it occupies and is almost
incompressible.
Melting Point: The temperature at which solid and liquid phases coexist in equilibrium.
Miscible: Two liquids that are completely soluble in each other in all proportions are said to be miscible.
Mixture: A combination of two or more substances in which the substances retain their identity.
Heterogeneous M.: The individual components of a mixture remain physically separated and can be seen
as separate components.
Homogeneous M.: The composition of the mixture, after sufficient stirring, is the same throughout the
solution.
Osmosis: see colligative properties.
Precipitation Reaction: see solution-properties.
Phase: see chemistry introduction.

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Raoult’s Law: The vapor pressure of an ideal solution of a nonvolatile solute is directly proportional to the mole
fraction of the solvent in the solution; i.e.: The partial pressure of the solvent over a solution is given by
the product of the vapor pressure of the pure solvent pSlt, partial pressure of solvent [Pa]
and the mole fraction of the solvent in the solution. pSle, partial pressure of solute [Pa]
pSlt = xSlt⋅p°Slt [Pa] xSlt, mole fraction of solvent [-]
p(Σ) = pSlt + pSle + etc. p°Slt, vapor pressure of pure solvent [Pa]
Deviation of RL.: Non-ideal (real) solutions deviate from the ideal pattern given by RL:
• Negative DoRL.: Total- and partial pressure are less than predicted by RL; attractive forces between
solute/solvent is greater than with like ones; the heat of solution in negative (exothermic).
• Positive DoRL.: Total- and partial pressure are greater than predicted by RL; attractive forces between
solute/solvent is lesser than with like ones; the heat of solution is positive (endothermic).
Semipermeable Membrane: A membrane that allows solvent molecules to pass through, but blocks the
movement of solute molecules; see osmosis.
Solubility: The maximum amount of solute that can be dissolved in a given quantity of a specific solvent at a
specific temperature (for gases: at a specific pressure); the concentration of a saturated solution of a
substance; e.g.: how much of a salt can be dissolved in a solvent.
KS - S. Constant: see solubility product;
Ksp - S. Product: The product of relative ionic molar concen- c, ionic molar conctr. [mol/l]
-trations of the constituent ions in a saturated solution, each raised A-, C+, anion, cation [-]
to the power of its stoichiometric coefficient in the equilibrium EQ; KC, equilibrium constant[-]
Ksp = KC⋅c(HA) = c(A-)⋅c(C+) [mol2/l2]
e.g.: Hg2Cl2(s) ↔ Hg22+(aq) + 2Cl- (aq); Ksp = c(Hg2++)⋅c2(Cl-) [mol3/l3]
Qsp - S. Quotient: The molar analogue of the solubility product, but with the molar concentrations not
necessarily those at equilibrium.
Qsp ≥ Ksp precipitate will form, whereas if Qsp < Ksp, still more salt can be added and will dissolve.
S. Rules: Solubility pattern of a range of common compounds in water - see table below.
• unpolar and polar substances are not miscible; e.g.: oil and water.
• like dissolves like; e.g.: ionic bonded element dissolve well in polar solvents, NaCl in H2O.
Pressure and Solubility: The process always implies a reduction of volume:
S. in Liquids: Since liquids are almost incompressible there is little influence in a change of volume.
S. in Gases: With increasing pressure more gas becomes dissolved in the solution; e.g.: soft-drink bottle
effect, diving (see Henry’s law).
Properties of aqueous solutions:
Acid-Base Reaction: Is a reaction in which protons are transferred; it is a proton transfer reactions; see
chemistry acid-base.
Gas-Forming Reaction: e.g.:
Ammonium salt: NH4+ + OH- → NH3(g) + H2O
Carbonate: CO3- + 2H+ → H2CO3 → CO2(g) + H2O
Sulfide: S2- + 2H+ → H2S(g)
Sulfate: SO3- + 2H+ → H2SO3 → SO2(g) + H2O
Oxidation-Reduction Reaction (RedOx): A reaction in which electrons are transferred; it is an electron
transfer reaction; see chemistry - stoichiometry-redox.
Precipitation Reaction: A reaction in which an insoluble solid product (which separates from the
solution) is formed when two solutions are mixed; i.e.: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
Temperature and Solubility: Temperature always influences solubility in a negative or positive way:
Endothermic S.: With increasing temperature, solubility increases, the endothermic (energy consuming)
process can proceed; follows the principle of Le Chatelier (see chemistry stoichiometry)
Exothermic S.: Solubility decreases with lower temperatures; gases are exothermic (energy releasing)
even though solubility decreases with raising temperatures.

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Sln - Solution: A homogeneous mixture of two or more substances.

Aqueous S.: A solution in which the solvent is water.
Enthalpy of S.: see chemistry - thermochemistry.
S. Concentration: The amount of solute present in a given quantity of solution.
Ideal S.: Any solution that obeys Raoult’s law at any concentrations. Real solutions resembles ideal
solutions more closely the lower the concentration; below 0.1[mol/kg] for non electrolytic solutions and
0.01[mol/kg] for electrolytic solutions.
Dilution: A procedure for preparing a less concentrated solution from a more concentrated solution;
Reminder: n(concentrades solvent) = n(diluted solvent):
If a solute is added in very small quantities compared to the solvent, than the vapor pressure of the liquid
can be said to be equal to that of the pure solvent; any diluted liquid given as a %-value usually refers to
mass-% in a 100g of solution;
being so diluted Raoult’s law can be implemented: ∆p, change of vapor pressure of solution
molar fraction of solute B = x(B) = ∆p/p°(A) [-] p°(A), vapor pressure of pure solvent [Pa]
Saturated S.: At a given temperature, the solution that results when the maximum amount of a substance
has dissolved in a solvent; dissolved and undissolved solute are in dynamic equilibrium.
• Oversaturated S.: (supersaturated) A solution that contains more solute than it has the capacity to
dissolve (unstable).
• Unsaturated S.: A solution that contains less solute than it has the capacity to dissolve.
Standard S.: A solution of accurately known concentration, used for acid-base tritations to calculate the
dosage of molar amounts out of a used volume; see chemistry-acid-base.
Slt - Solvent: The substance (usually one, or more) present in larger amount in a solution.
Sle - Solute: The substance present in smaller amount in a solution.
γ - Surface Tension: A measure of the force that must be applied to surface molecules so that they experience the
same force as molecules in the interior of the liquid; the amount of energy required to stretch or increase
the surface by unit area; decreases with increasing temperature; work has to be expended to bring
molecules to the surface.
Titration: see chemistry acid-base.
Triple Point: The point at which the vapor, liquid, and solid states of a substance are in equilibrium.
Vapor: The gaseous phase of a substance (specifically, of a substance that is a liquid or a solid at the temperature
in question).
V. Pressure: The pressure exerted by the vapor of a liquid (or solid) when the vapor and the liquid (or
solid) are in dynamic equilibrium; the higher the intermolecular forces the lower the vapor pressure
a liquid’s vapor pressure in an open beaker cannot increase since surrounding atmospheric pressure is
more or less constant, therefore temperature (at boiling point) remains constant until the entire liquid has
evaporated;
e.g.: H2O(g) ↔ H2O(l)
• VP Lowering: see colligative properties.
Enthalpy of V.: see thermochemistry.
Vaporization (evaporation): The escape of molecules from the surface of a liquid; a certain number of
molecules in a liquid (at any temperature) possess sufficient energy to escape from the surface; see
chemistry gas - Maxwell-Boltzman speed distribution.
Molar Heat of Vaporization: see chemistry - thermochemistry;
Viscosity: A measure of a fluid’s resistance to flow; decreases with increasing temperature.
Volatility: The readiness with which a substance vaporizes: A substance is typically regarded as volatile if it its
boiling point is below 100°C.
Volatile: as a measurable vapor pressure.
Water: Liquid substance of H-O-H molecules, in which intermolecular forces (H-bonds) are responsible for the
high boiling point; water molecule is bent due to the lone pair of the O and possess a dipole moment, since
the O-atom (more electronegative) deprives the H-atoms off their electron clouds.
Hard W.: Water that contains Ca2+ and Mg2+ ions.
Soft W.: Water that mostly free of Ca2+ and Mg2+ ions.

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Molal boiling point elevation and freezing point depression constants of several common compounds:

Solvent freezing point kf-Slt boiling point kb-Slt

[°C] [°C⋅kg/mol] [°C] [°C⋅kg/mol]
Acetic acid 16.6 -3.90 117.9 2.93
Benzene 5.5 -5.12 80.1 2.53
Carbon -22.8 -29.8 76.8 5.02
Tetrachloride
Chloroform -63.5 -4.68 61.2 3.63
Cyclohexan 6.6 -20.0 80.7 2.79
Ethanol -117.3 -1.99 78.4 1.22
Naphtalin 80.2 -6,8 - -
Water 0 -1.86 100 0.52

Classification of solutes in aqueous solution

Strong Electrolyte Weak Electrolyte Non- Electrolyte

HCL CH3COOH (NH2)2CO (urea)
HNO3 HF CH3OH (methanol)
HCLO4 HNO2 C2H5OH (glucose)
H2SO4 NH3 C6H12O6 (glucose)
NaOH H2O C12H22O11 (sucrose)
Ba(OH)2
ionic compounds

Solubility Rule (based on Group-I &-II &-III elements, NH4+ Al3+ Pb2+ Sn2+:

soluble compounds semi-soluble insoluble compounds

Bromides (Br-) all* Pb2+, Hg2+ Ag+, Hg22+, TI+
Carbonates (CO32-) Group-I elements, NH4+ all remaining*
Chlorides (Cl-) all* Pb2+ Ag+, Hg22+, TI+
Chlorates (ClO3-) all*
Florides (F-) single charged cations multiple charged cations
Hydroxides (OH-) Group-I elements Ba2+, Ca2+, Sr2+ all remaining
Ba2+, NH4+ Mg2+
Nitrates (NO3-) all*
Sulfates (SO42-), all* Ca2+, Ag+ Sr2+ Ba2+, Pb2+, Hg22+
Sulfides (S2-) Group-I & -II elements, all remaining;
NH4+, (Fe23+, Al3+, Cr3+ form
hydroxides)
Sulfites (SO42-) Group-I elements, NH4+, all remaining*
Oxides (O2-) all Group-I elements, most hydroxides (OH-)
NH4+ nitrate (NO3-),
chlorate (ClO3-),
perchlorate (ClO4-)
3-
Phosphates (PO4 ) Group-I elements, NH4+ all remaining*
*) elements taken in consideration: Group-I & -II elements, NH4+, Ag+, TI+, NH4+, Mn2+, Fe2+, Co2+,
Ni2+, Cu2+, Zn2+, Cd2+, Hg22+, Sn2+, Pb2+, Al3+, Cr3+, Fe3+
Compounds of group-I elements: H+, Li+, Na+, K+, etc.
Compounds of group-II elements: Be2+, Mg2+, Ca2+, Sr2+, etc.
Compounds of group-III elements: transition metals like Fe, V, Au, etc.

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Glossary - Chemistry (Solid)

Amorphous solid: A solid that lacks a regular 3-dimensioal arrangement of atoms or molecules;
i.e.: a liquid with an extremely high viscosity index e.g.: glass, tar, etc.
Bond: see chemistry molecule.
Born-Haber Cycle: The cycle that relates lattice energies of ionic compounds to ionization energies, electron
affinities, heats of sublimation and formation, and bond dissociation energies; see thermochemistry.
Closest Packing: The most efficient arrangements for packing atoms, molecules, or ions in a crystal.
Cohesion: The intermolecular attraction between like molecules (see physics - matter).
Colligative Properties: see chemistry - liquid.
Conductor: Substance capable of conducting electric current (see physics - electromagnetism).
Coordination Number: see lattice.
Crystal: An regular arrangement of atoms, ions, and molecules of periodically repeated, identically constituted,
congruent lattice consisting of unit cells; see table below.
Crystalline Solid: A solid that possesses rigid and long-range order; its atoms, molecules, or ions occupy
specific positions; e.g.: NaCl, diamond, graphite, etc.
Fractional C.: The separation of a mixture of substances into pure components on the basis of their
differing solubility.
Crystallization: The process in which dissolved solute comes out of solution and forms crystals.
Recrystallization: Purification by repeated dissolving and crystallization.
Types of crystals: see table below.
Covalent C.: The structure is primarily determined by the geometry of the bonds formed by each atom;
the lattice ranges from face-centered cubic (diamond) to cubic close-packed (tetrahedral of SiO4 form
SiO2) and to triangular pyramidal (As4).
Ionic C.: Are composed of charged species with anions and cations different in size; the lattice is face-
centered (NaCl and most others cN = 6) or simple cubic (CsCl, CsBr, CsI, cN = 8) where every anion tends
to gather as many cations and vice versa (1:1 structure).
Metallic C.: Generally body-centered, face-centered, or hexagonal close-packed with the bonding
electrons delocalized over the entire crystal; the cohesive of these e-forces are responsible for the metals
strength;
All metals, except Mn, Hg, Ga, In, Ge, Sn have one of the 3 types of structures:
• Monoatomic Body-Centered: The lattice is body-centered cubic, one atom is situated at each lattice
point, giving 2 atoms per unit cell (1 + 8⋅1/8); cN = 8; e.g.: Li, Na, K, V, Cr, Fe, Rb, Nb, Mo, Cs, Ba,
Ta, W
• Cubic Close-Packed: The lattice is face-centered cubic, one atom situated at each lattice point, giving
4 atoms per unit cell (6⋅½ + 8⋅1/8); cN = 4; e.g.: Al, Ca, Ni, Cu, Sr, Rh, Pd, Ag, IR, Pt, Au, Pb
• Hexagonal close-packed: The close-packed layers are perpendicular to the vertical axis or
perpendicular to the body’s diagonal of the unit cell; e.g.: Be, Mg, Sc, Ti, Co, Zn, Y, Zr, Tc, Ru, Cd,
La, Hf, Re, Os, Ti
Molecular C.: Structure is determined by the size and the shape of the atoms involved and range from
face-centered cubic (CO2) to face-centered orthorhombic (Cl2, Br2, I2)
Deposition: The process in which the molecules go directly from the vapor into the solid phase.
Freezing Point: The temperature at which a liquid freezes (crystallization); the normal freezing point is the
freezing temperature under a pressure of 1atm, i.e.: liquid and solid phase are in equilibrium;
Lattice: A regular periodic framework throughout an area or space, made of interwoven strips, as in crystals.
L. Energy: The energy required to completely separate one mole of a solid ionic compound into gaseous
state; LE depends upon the charge of ions (the higher, the stronger), and the size of ions (the smaller, the
stronger the force of attraction).e.g.: Na+(g) + Cl-(g) → NaCl(s) ∆HL = -788 [kJ/mol];
i.e.: the energy needed for vaporizing this solid +788 [kJ/mol]; see thermochemistry.
L. Enthalpy: The standard enthalpy change for the conversion of an ionic solid to a gas of ions.
cN - L. Coordination Number: The number of nearest neighbors of an atom (ion), in a solid (of opposite
charge); for complex ions (see chemistry soichiometry) it is the number of atoms directly attached to the
central metal ion; the higher cN, the stronger the intermolecular bond; see unit cell.
L. Points: the positions occupied by atoms, molecules, or ions that define the geometry of a unit cell.
Melting Point: The temperature at which solid and liquid phases coexist in equilibrium.
Metals: 1) A substance that conducts electricity as well as heat, has a metallic luster, is malleable and ductile,
forms cations and has basic oxides. 2) A metal consists of cations held together by a sea of electrons (have
the tendency to form positive ions in ionic compounds); i.e.: iron, copper, uranium, etc.

1
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Phase: see chemistry - introduction.

Rays: (see physics - nuclear).
α Alpha R.: Helium ions with a positive charge of +2; also known as alpha particles.
β Beta R.: Electrons.
γ Gamma R.: High energy radiation.
X-Ray: Electromagnetic radiation with wavelengths ranging from 10[pm] to about 1000[pm]; see physics
- optics.
XR-Diffraction: The analysis of crystal structures by studying the θ, diffraction angle [°]
interference pattern in a beam of x-rays: λ wavelength [m]
To determine the bond length: n, integer, 1, 2, 3, [-]
2⋅d⋅sinθ = n⋅λ
Salt: An ionic compound made up of a cation other than H+ and an anion other than OH- or O2-.
S. Hydrolysis: The reaction of the anion or cation, or both, of a salt with water.
Solid: A rigid form of matter that maintains the same shape whatever the shape of the container.
Sublimation: The process in which molecules go directly from the solid into the vapor phase; can be measured
easily in a vacuum chamber where at a given temperature a distinct vapor pressure can be detected (solid
and gas phases are in equilibrium).
S. Vapor Pressure: The pressure exerted by the vapor of a solid when the vapor and or solid are in
dynamic equilibrium; the higher the intermolecular forces the lower the vapor pressure;
see chemistry - liquid.
Triple Point: The point at which the vapor, liquid, and solid states of a substance are in equilibrium; see
chemistry - gas.
Unit Cell: The basic repeating unit of the arrangement of atoms, molecules, or ions in a crystalline solid; when
stacked together can reproduce the entire crystal; see table below.
Scc - Simple Cubic (primitive cubic): The unit cell of a space lattice is obtained by joining 8 points to
give a solid with pairs of parallel faces (cube); cN = 6.
Bcc - Body-Centered Cubic: A cube with an additional point in the center in the middle of the cube; cN =
8.
Fcc - Face-Centered Cubic: A cube with a n additional point in the center of each face; cN = 12.
Monoclinic: a ≠ b ≠ c; α = γ = 90° (β ≠ 90°)
Hexagonal: a = b ≠ c; α = β = 90° (γ = 120°)
Orthorhombic: a ≠ b ≠ c; α = β = γ = 90°
Rhombohedral: a = b = c; α = β = γ ≠ 90°
Tetragonal: a = b ≠ c; α = β = γ = 90°
Triclinic: a ≠ b ≠ c; α ≠ β ≠ γ ≠ 90°

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Common lattice energies

Type of compound Compound Ions Σ of ionic radius Lattice E. [kJ/mol]

1+, 1- NaCl Na+, Cl- 95 + 181 = 276 - 788
CsCl Cs+, Cl- 169 + 181 = 350 - 669
1+, 2- Na2O 2Na+, O2- 95 + 140 = 235 -2570
Cs2O 2Cs+, O2- 169 + 140 = 309 -2090
2+, 1- MgCl2 Mg2+, 2Cl- 65 + 181 = 246 -2525
2+, 2- MgO Mg2+, O2- 65 + 140 = 205 -3890

The cubic unit cell

Type of # of points # of points in # of points in total cN

cell at corners faces center of cube
scc 8 x 1/8 0 0 1 6
bcc 8 x 1/8 0 1 2 8
fcc 8 x 1/8 6x½ 0 4 12

Types of crystals and general properties

Type of Particles Force(s) holding General properties Examples

crystal involved the unit together
Ionic positively and electrostatic hard, brittle, high melting point, poor NaCl,
negatively attraction conductor of heat and electricity LiF,
charged ions MgO,
CaCO3
Covalent atoms covalent bond hard, high melting point, poor conductor C
(spatial) of heat and electricity (diamond)
3-dimensional
, SiO2
(quartz)
Covalent atoms covalent bond + sift, high melting point, good conductor of C
(layers) dispersion electricity (graphite)
2-dimensional
forces
atoms and ions covalent bond + high melting point, poor conductor of mica, clay
electrostatic electricity
attract.
Covalent atoms covalent bond + fibrous, partly meltable to a viscose SiS2
(chain) dispersion or liquid,
1-dimensional
dipole-dipole
forces
atoms and ions covalent bonds fibrous, poor conductor of heat and asbestos
+ electrostatic electricity
attraction
Molecular polar molecules dispersion-, soft, low melting point, poor conductor of Ar, CO2,
dipole-dipole- heat and electricity I2, H2O,
(polar) forces, hydrogen C12H22O11
bonds (sucrose)
(nonpolar) nonpolar dispersion H2, Cl2,
molecules forces only CH4
Metallic positive ions metallic bond soft to hard, low to high melting point, all
(cations) and good conductor of heat and electricity, metallic
highly motile easily malleable elements;
electrons Na, Mg,
Fe, Cu,
etc.

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Glossary - Chemistry (Thermochemistry)

Adiabatic Process: A process in which no heat exchange occurs between the system and its surrounding
environment (see physics - heat).
Bond: A link between bonds - see chemistry - molecule.
Double B.: Two electron pairs shared by neighboring atoms.
Multiple B.: Bonds formed when two atoms share two or more pairs of electrons.
Triple B.: Three electron pairs shared by two neighboring atoms.
Born-Haber Cycle: The cycle that relates lattice 1. Na(s) + ½ Cl2(g) ∆HSublimation
energies of ionic compounds to ionization, 2. Na(g) + ½ Cl2(g) ½ ∆HDissociation
energies electron affinities, heats of sublimation, 3. Na(g) + Cl(g) ∆HIon
and formation, and bond dissociation energies; 4. Na+(g) + e- + Cl(g) ∆HEa
e.g.: ∆Hf of NaCl is obtained out of the energy 5. Na+(g) + Cl-(g) ∆HLattice
differences of steps 1 to 5 Σ: NaCl ∆Hf of NaCl
Calorimeter: A device to measure the heat released/absorbed by a process under constant volume (isochorous)
i.e.: W = 0 → ∆U = Q; the internal energy of a process will be converted into a change of heat.
Bomb C.: A combustion chamber w/n an isolated, sealed tank with stirrer, igniter and thermometer;
Calorimetry: The use of a calorimeter to measure the thermochemical properties of reactions.
Clausius Clapeyron’s Equation: The quantitative relationship between
the vapor pressure pV of a liquid and the absolute temperature T: pv, vapor pressure [N/m2] [Pa]
ln(pv) = C - ∆HV/(R⋅T) C, constant [1/K]
Conversion of p1, T1 to p2, T2: R, gas constant 8314 [J/mol]
log(p1/p2) = ∆HV/R ⋅ {T2-T1/(T2⋅T1)} T, temperature [K]
Collision Theory: The theory of elementary gas-phase bimolecular reactions in which it is assumed that
molecules react only if they collide with at least enough kinetic energy for bonds to be broken.
Activated Complex T.: A theory of reaction rates in which it is postulated that the reactants form an
activated complex; i.e.: compounds temporarily formed by the reactant molecules as a result of collision.
Activated Complex: A combination of the two reactant molecules that can either go on to form products
or fall apart into the unchanged reactants;
e.g.: A2 + X2 ↔ A2X2 → 2AX; where A2X2 is very short-living.
E - Energy: The capacity to do work or supply heat; KE + PE = constant! Units in: [kg⋅m2/s2] = [N⋅m] = [J]
EA - Activation Energy: The minimum energy needed for reaction; i.e.: the height of the activation
barrier; usually endothermic, followed by a more or less exothermic reaction; resulting in a net exo-
/endothermic reaction (see stoichiometry - kinetics Arrhenius behavior): EAU,
Uphill-EA [J/mol]
e.g.: ∆H = EAU - EAD [J/mol] EAD, Downhill-EA [J/mol]
Bond E.: The energy needed in diatomic molecules to dissociate 1mol of molecules into gaseous atoms;
e.g.: H2(g) → 2H(g, atomic) = 436[kJ/mol]; i.e.: H = 218.0[kJ/mol]
Dissociation E.: The amount of energy needed to break a bond - see table below.
Free E.: Energy available to do work.
U - Internal E. (energy of reaction): The total energy of a system; the sum of KE and PE of all particles;
absolute amount of U is not detectable, but relative changes can be expressed as follows:
∆U = Q + W (energy supplied as heat and work) [J]
∆U < 0: internal energy decreases by giving of energy to its surrounding environment;
∆U > 0: internal energy increases by taking up energy from the surrounding environment;
∆U = 0: no change in internal energy (W = 0; → ∆U = QV); QV = heat supplied at constant volume;
Ionization E.: Energy required to remove of electrons from an element; see ∆H - electron enthalpy.
Kinetic E.: Energy available because of the motion of an object; see physics - mechanics: KEL = ½⋅m⋅v2
[J]
Lattice E.: The energy required to completely separate one mole of a solid ionic compound into gaseous
state; i.e.: energy of + energy is needed to vaporized this solid substance. LE depends upon charge of ions
(the higher, the stronger) and the size of ions (the smaller, the stronger the force of attraction):
e.g.: Na+(g) + Cl-(g) → NaCl(s) ∆HL = -788 [kJ/mol];
Law of Conservation of E.: see 1st law of thermodynamics.
Potential E.: Energy available by virtue of an object’s position; see physics - mechanics: PE = m⋅g⋅y [J]
Standard Free E. of Change(∆G°): The free-energy change
when reactants in their standard states are converted to products in their standard states.
Standard Free E. of Formation (∆Gf°): The free-energy change when 1 mole of a compound is
synthesized from its elements in their standard states.
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Thermal E.: Energy associated with the random motion of atoms and molecules.
H - Enthalpy: A thermodynamic quantity used to describe heat changes taking place at constant pressure
(isobar); i.e.: reservoir of energy that can be obtained as heat;
∆H = ∆U + W = Hfinal - Hinitial [kJ/mol] ∆U, internal energy [J]
∆H = ∆U + p⋅∆V = ∆U + R⋅∆n⋅T [kJ/mol] W, work [N⋅m] [J]
∆H < 0: heat released (exothermic reaction); p, pressure [N/m2] [Pa
∆H > 0: heat absorbed (endothermic reaction); ∆V, change in volume [m3]
Reminder: when using ∆H, don’t forget to add R, gas c. 8.314 [J/(K⋅mol)]
the reactants- and products phase! ∆n, molar amount [mol]
E. of Chemical Change: Processes involved in chemical changes; T, temperature [K]
• E. of Formation: see Born Haber cycle.
• E. of Hydration: Energy which is released from ions(g) to dissolved hydrated ions(aq) in a
hypothetical process; ∆HH is always negative.
e.g.: K+(g) + Cl-(g) → (H2O) → K+(aq) + Cl-(aq) ∆HH = -684.1[kJ/mol]
• E. of Reaction: The difference between the enthalpies of the products and the enthalpies of the
reactants; measured in [J/mole] (compare physics - heat).
Endothermic R.: Processes that absorb heat from the surrounding environment, ∆H > 0;
i.e.: ½ H2(g) + ½ I2(s) → HI(g) ∆H = +25.9 [kJ]
Exothermic R.: Processes that give off heat to the surroundings, ∆H < 0;
i.e.: H2(g) + ½ O2(g) → H2O(g) ∆H = -241.8 [kJ]
• E. of Solution: The heat generated or absorbed when a certain amount of solute is dissolved in a
certain amount of solvent; i.e.: the sum of energy required to break the lattice structure and energy set
free during solvatization; e.g.: ∆H = HSln - HComponents
1. KCl(s) → (H2O) → K+(g) + Cl-(g)....lattice energy (positive) ∆HS = +701.2 [kJ/mol]
2. K (g) + Cl (g) → (H2O) → K (aq) + Cl (aq)....hydration energy
+ - + -
∆HS = -684.1 [kJ/mol]
Σ KCl(s) → (H2O) → K+(aq) + Cl-(aq)....endothermic, heat take-up ∆HS = +17.1 [kJ/mol]
i) most non-ionic bonds are endothermic (lattice energy smaller than in ionic bonds);
i) most gases are exothermic (no lattice present to be broken up);
E. of Physical Change: Processes involved in physical changes;
• ∆Hfrez E. of Freeezing: The negative of the enthalpy of melting.
• ∆Hvap E. of Vaporization: The difference in enthalpy per mole between the vapor and liquid states of
a substance; e.g.: ∆Hvap = Hvapor - Hliquid water: +40.7[kJ/mol]
• ∆Hmel E. of Melting: The difference in enthalpy per mole between the solid and liquid states of a
substance; e.g.: ∆Hmel = Hliquid - Hsolid water: +6.01[kJ/mol]
• ∆Hsub E. of Sublimation: The enthalpy change per mole of molecules when a solid changes into vapor;
e.g.: ∆Hsub = Hvap - Hsolid always endothermic (positive)
Electron E.: Uptake or release of energy by giving off, or accepting an electron.
• ∆HEa - Electron Affinity: The energy released when an electron is added to a gas-phase atom or ion of
the elements forming an anion; the negative of the electron-gain enthalpy (see table below); e.g.:
∆H > 0: Ne(g) + e- → Ne- (g) energy is needed to add an e ∆HEa = +29 [kJ/mol] or [eV/atom]
∆H < 0: Fe(g) + e- → Fe- (g) energy is released ∆HEa = -328 [kJ/mol]
with the highest value in the upper right and the lowest values at the lower left of the periodic table.
• ∆HIon - Ionization Energy: The minimum energy required to remove an electron from the ground state
of a gaseous atom, molecule , or ion. The second ionization energy is the ionization energy for
removal of a second electron and has to be higher than the first; (∆H always positive)
e.g.: 1st IE: Na → Na+ + e- ∆HIon = +496 [kJ/mol] = [eV/atom]
2 IE: Ag → Ag + e
nd + 2+ -
∆HIon = +4563 [kJ/mol]
I.E. increases from left to right across the period and decreases down a group in the periodic table.
Electron Gain E.: see electron enthalpy.
Lattice E.: The standard enthalpy change for the conversion of an ionic solid to a gas of ions.
∆Hf° Standard E. of Formation: The heat change that results when 1 mole of compound is formed from
its elements in their standard states, i.e.: 101.3[kPa] and 298.15[K] = 25[°C];
e.g.: ∆H° = Σn∆Hf°(products) - Σn∆Hf°(reactants)
∆H° Standard E. of Reaction: Enthalpy change of the reaction carried out under standard-state
conditions.
∆HC° Standard E. of Combustion: The change in enthalpy per mole of substance when it burns (reacts
with Oxygen) completely under standard-state conditions;
e.g.: C converts to CO2, H to H2O, N to N2, S to SO2;

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Entropy: A direct measure of the randomness or disorder of a system. ∆H, enthalpy [J/mol]
Q - Heat: The amount of energy in form of heat; n, molar amount [mol]
Q = ∆H⋅n = C⋅∆T [J] C, heat capacity [J/K]
C - Heat Capacity: The amount of heat required to raise the T, temperature [K]
temperature of a given quantity of the substance by 1°C; m, mass [g]
i.e.: to raise 1g of water from 14.5°C to 15.5°C s, specific heat [J/(g⋅K)]
the energy of 4.184[J] (=1cal) is required: C = m⋅s [J/K]
H. of Dilution: The heat change associated with the hydration process - see chemistry liquid.
H. of Hydration: The heat change associated with the hydration process.
H. of Solution: see enthalpy of solution. Q, heat [N⋅m] [J]
Specific H. Capacity: The heat capacity per gram. m, mass [kg
s = Q/(m⋅∆T) [N⋅m/(kg⋅K)] = [J/(kg⋅K)] T, temperature [K]
Hess’s Law: When reactants are converts to products, the change in enthalpy is the same whether the reaction
takes place in one step or in a series of steps.
Molarity: The number of moles of solute in one liter of solution.
M. Heat of Fusion: The energy [kJ] required to melt one mole of solid;
e.g. 1[mol] at 101[kPa]: H2O(s) → H2O(l) ∆HFus = +6.01 [kJ/mol]
M. Heat of Condensation: Equal to molar heat of vaporization but with opposite sign.
M. Heat of Crystallization: Amount of heat withdrawn from one mole of liquid at crystallization; equal
to molar heat of fusion but with opposite sign.
M. Heat of Sublimation: The energy [kJ] required to sublime one mole of solid;
∆HSub = ∆HFus + ∆HVap [kJ/mol]
M. Heat of Vaporization: The energy [kJ] required to vaporize one mole of liquid; i.e.: the energy
needed to liberate single molecules from the liquid pool (D=volumsarbeit); see Claudius-Clapeyron’s
equation;
e.g. 1 mol of at 25°C at 1atm: H2O(l) → H2O(g) ∆HVap = +43.8 [kJ/mol]
• the higher the intermolecular forces the lower the vapor pressure;
• the lower the temperature the smaller ∆HV (Tcrit: ∆HV = 0!).
Temperature: 1) How hot or cold a sample is. 2) The intensive property that determines the direction in which
heat will flow between two objects.
Absolute T.: Temperature scale with the lowest possible temperature: T = 0K (K = °C + 273.15)
Critical T.: The temperature above which a gas will not liquefy.
Thermochemistry: The study of heat changes in chemical reactions.
T. Equation: An equation that shows both the species involved and enthalpy relations;
e.g.: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ∆H = -890 [kJ/mol]
Thermodynamics: The study of the transformation of energy from one form to another.
1st Law: Energy can be converted from one form to the another, but cannot be created or destroyed.
2nd Law: The entropy of the universe increases in a spontaneous process and remains unchanged in an
equilibrium process.
3rd Law: The entropy of a perfect crystalline substance is zero at the absolute zero of temperature.
Van’t Hoff Equation: see stoichiometry - equilibrium.
W -Work: The energy expended during the at of moving an object against an opposing force; one joule of work
is done when a force of 1 newton is exerted over a distance of 1 meter (compare power):
W = F⋅d [N⋅m] = [J] F, force [N]
Endergonic: (Gk. endo, within; ergon, work) Describing a d, distance [m]
chemical reaction that requires energy to proceed; (compare exergonic).
Exergonic: (L. ex, out; ergon, work) Energy-yielding, as in a chemical reaction; applied to a “downhill“
process; compare endergonic).

Molar Heat of Vaporization and Fusion of selected substances [kJ/mol]

Substance Boiling Point ∆HVap Melting Point ∆HFus

[°C] [kJ/mol] [°C] [kJ/mol]
Argon (Ar) -186 6.3 -190 1.3
Methane (CH4) -164 9.2 -183 0.84
Ethyl ether (C2H5OC2H5) 34.6 26.0 -116.2 6.90
Ethanol (C2H5OH) 78.3 39.3 -117.3 7.61
Benzene (C6H6) 80.1 30.8 5.5 10,9
Water (H2O) 25 43.8 0 6.01
Water (H2O) 100 40.79
Mercury (Hg) 357 59 -39 23.4
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Standard bond energies of formation for common substances (25°C and 101.3kPa) [kJ/mol]

Compound ∆Hf° Compound ∆Hf° Compound ∆Hf°

C-C +347 Br-Br +190 N-N +159
C=C +619 Cl-Cl +243 N=N +418
C≡C +812 F-F +155 N≡N +941
C-H +413 H-Br +364 O-Cl +207
C-Cl +326 H-Cl +431 O-F +184
C-F +485 H-F +565 O-H +463
C-N +293 H-H +435 O-O +143
C=N +616 H-I +297 O=O +494
C≡N +879 I-I +149 P-Cl +326
C-O +335 P-H +318
C=O +707 N-Cl +201 S-Cl +276
C≡O +1070 S-H +364

Standard enthalpies of formation of common substances (25°C and 101.3kPa) [kJ/mol]

Compound ∆Hf° Compound ∆Hf° Compound ∆Hf°

AgCl(s) -127.0 Cl-(aq) -167.4 I2(g) 0
Al2O3(s) -1669.8 Cl2(g) 0 KCl(s) -436.0
B5H9(s) +73.2 CO(g) -110.5 KClO3(s) -391.2
B2O3(s) -1273.5 CO2(g) -393.5 KClO4(s) -435.1
BaCO3(s) -1218 COCl2(g) -223 LiAlH4(s) -100
BaO(s) -588.1 CS2(l) +87.86 MgO(s) -601.83
C (atomic) +716.7 Cu2+(aq) +64.4 MgCO3(s) -1112.9
C +1.9 CuO(g) -155.2 NaCl(s) -411.0
(diamond)
CaCO3(s) -1206.9 F2(g) 0 NaOH(s) -425.6
CaO(s) -635.5 Fe2+(aq) -87.9 N2(g) 0
Ca(OH)2(s) -986.59 Fe2O3(s) -822.2 NF3(g) -113
CF4(g) -913.4 H+(aq) 0.0 NH3(g) -46.19
CH4(g) -74.85 H2(g) 0 NH3(l) -66.9
C2H2(g) +226.7 HBr(g) -36.2 NH4NO3(s) -365.1
C2H4(g) +52.3 HCl(g) -92.3 NO(g) +90.37
C2H6(g) -84.68 HCN(g) +130.5 NO2(g) +33.8
C3H8(g) -103.8 HF(g) -269 N2O(g) +81.6
C4H10(g) -126.1 HgBr2(s) -169 N2O4(g) +9.16
C5H12(g) -146.4 HI(g) +25.9 OH-(aq) -229.95
C6H6(l) +49.04 HNO3(l) -173.2 O2(g) 0
C6H12O6 (s) -1268 H2O(g) -241.8 O3(g) 142.2
CH3Cl(l) -132 H2O(l) -285.9 PH3(g) +9.25
CH3NH2(g) -28 H2O2(l) -187.6 SO2(g) -296.9
CH3OH(g) -201.2 H2S(g) -20.2 SO3(g) -395.2
CH3OH(l) -238.6 Zn2+(aq) -152.3
C2H5OH(l) -277.6 ZNO(s) -348.0
all other solid elements, like Ag(s), Al(s), Hg(l), etc. have values of ∆Hf° = 0.

Electron affinity of common elements (theoretical values) [kJ/mol]

I II III IV V VI VII VIII

H -73 He (+21)
1
Li -60 Be (+240) B -27 C -122 N0 O -141 F -328 Ne (+29)
2
Na -53 Mg (+230) Al -43 Si -134 P -72 S -200 Cl -349 Ar (+35)
3
K -48 Ca (+156) Ga -29 Ge -116 As -77 Se -195 Br -325 Kr (+39)
4
Rb -47 Sr (+168) In -29 Sn -121 Sb -101 Te -190 I -295 Xe (+41)
5
Cs -45 Ba (+52) Ti -29 Pb -35 Bi -91 Po -183 At -270 Rn (+41)
6

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Glossary - Chemistry (ElechtroChemical Series)

Battery: An electrochemical cell, or often several electrochemical cells connected in series, that can be used as a
source of direct electric current at a constant voltage.
Conductor: Any material through which charge easily flows (carries electric current) when subject an external
electrical force; are also good heat conductors as well; electrons in the outer atomic shell are loose
(valence shell).
C. Band: An incompletely filled band of orbitals in a solid.
Conductivity: The intrinsic property of a substance to conduct electric current; reciprocal to resistivity.
Metallic C.: An externally imposed potential difference forces the electrons in the conduction band,
which are the carrier of the electric current, to move according to the gradient imposed.
Electrolytic C.: Carrier of the electric current are ions, results in a net transport of mass.
Corrosion: The deterioration of metals by an electrochemical process.
Daniell Cell: Consists of Cu- and Zn electrodes dipping into solutions of CuII-sulfate and Zn-sulfate; the 2
solutions make contact through a porous cup, which allows ions to pass through, to complete the circuit.
Electrochemistry: The branch of chemistry that deals with the use of chemical reactions to produce electricity,
the relative strengths of oxidizing and reducing agents, and the use of electricity to produce chemical
change.
Electrochemical Cell (Voltaic Cell): A system consisting of two electrodes in contact with an electrolyte.
Galvanic.C. An electrochemical cell used to produce electricity by means of a spontaneous redox
reaction.
Here anode (-), cathode (+) to collect, give off the electrons of the spontaneously occurring reaction; e.g.:
redox-reaction at the electrodes resulting in two half-cell reactions:
Zn(s) → Zn2+(aq) + 2e- 2e- + Cu2+(aq) → Cu(s)
overall reaction: Zn(s) + Cu (aq) → Zn2+(aq) + Cu(s)
2+

Electric C.: An electrochemical cell in which an electric current is used to cause chemical change;
Here anode (+), cathode (-) to start a reaction by imposing an external force during electrolysis.
e.g.: Zn electrode (anode) + Cu electrode (cathode) produces an electric current if emerged in an
electrolyte; resulting in two half reactions:
Zn(s) → Zn2+(aq) + 2e- 2e- + Cu2+(aq) → Cu(s)
overall reaction: Zn(s) + Cu (aq) → Zn2+(aq) + Cu(s)
2+

Electrode: A metallic conductor that makes contact with an electrolyte in an electrochemical cell - see there.
Anode: (Gk: an, up) The electrode at which oxidation occurs; attracts anions; e.g.: Cl-.
Cathode: (Gk: cat, down) The electrode at which reduction occurs; attracts cations; e.g.: Na+.
SHE - Standard Hydrogen E.: A H-electrode that is in its standard state (H+ ions at concentration
1[mol/l] and H-pressure 101[kPa]) and is defined as having E° = 0:
H2 → 2H+ + 2e-
2H+(aq, 1molar) + 2e- → H2(g, 1atm) E = 0
Electrolysis: 1) A process in which a chemical change is produced by passing an electric current through a
liquid. 2) The process of driving a reaction in a non-spontaneous direction by passing an electric current
through a solution: e.g.: electrolysis in aqueous solution of Na2SO4:
cathode-reaction: anode reaction:
H2O ↔ H+ + OH- H2O ↔ H+ + OH-
2H+ + 2e- → H2(g) 2OH- → ½O2(g) + H2O + 2e-
2H2O + 2e ↔ H2(g) + 2OH
- -
H2O → ½ O2(g) + 2H++ 2e-
Net result of these two half reactions: H2O → ½ O2(g) + H2(g)
• electrolysis in aqueous solution of NaCl yields: 2H2O + 2Na+ + 2Cl- ↔ H2(g) + Cl2(g) + 2Na+ + 2OH-
• electrolysis in aqueous solution of CuSO4 yields; H2O + Cu2+ + SO42- ↔ Cu(s) + H2(g) + 2H+ + SO42-

1
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Electrolyte: 1) An ionically conducting medium. 2) A substance that, when dissolved in water, results in a
solution that can conduct electricity; see table below.
E. Conduction: see conductor.
E. Rule: For a net potential of zero, the positive and negative charges must add up to zero; a solution must
contain essentially as many anionic as cationic charges; see chemistry - liquid.
Non-E.: Is a solution in which no proportion of the solute molecules are ionized, hence does not conduct
electricity; e.g.: C6H12O6(aq)
Strong E.: Is a solution in which a large proportion of the solute molecules are ionized (complete
dissociation into ions); e.g.: NaCl(s) → Na+(aq) + Cl-(aq) 1mol K2SO4 → 3mol ions
Weak E.: Is a solution in which only a small proportion of the solute 1mol NaCl → 2mol ions
molecules are ionized (partly dissociation into ions); e.g.: CH3COOH(aq)
Electronegativity: The ability of an atom to attract electron toward itself in a chemical bond.
Faraday’s Law of Electrolysis: The amount (in moles) of product formed by an electric current is chemically
equivalent to the amount (in moles) of electrons supplied:
q = FEQ⋅z⋅n = FEQ⋅z⋅m/M [C] q, disposing charge [C]
F. Constant: The charge per mole of electrons; FEQ, equivalent charge [C/mol]
F = 6.022⋅1023 [1/mol] x 1.6022⋅10-19 [C] = 96485 [C/mol] z, amount of charge/ion [1/mol]
Half Cell Reaction: Oxidation and reduction reactions at the n, deposited molar amount [mol]
electrodes - see electrochemical cell. m, deposited mass [g]
Redox Couple: Consists of the oxidized and reduced species M, molar mass [g/mol]
taking part in the half reaction; e.g.: Red (Zn) / Ox (Zn2+) F, Faraday c. 96485 [C/mol]
Nernst Equation: The EQ expressing the cell potential in terms of the KC, equilibrium constant[var]
concentrations of the reagents taking part in the cell reaction; E°, stand. reduction pot. [V]
E = E° - ln(q)⋅R⋅T/(NOX⋅F) = E° - log(KC)⋅0.05916/NOX [V] R, gas constant 8314 [J/mol]
e.g.: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu (s) T, temperature [K]
= 1.1-log(c(Zn2+)/c(Cu2+))⋅0.0592/2 NOX, oxidation number [-]
Redoxpotential of any half reaction of a metal Mn+, molar concentration on metal
n+
E = E° + log(M )⋅0.05916/NOX [V] ion in solution [mol/l]
V - Potential (voltage): The electric (pressure) potential energy per PE, potential energy [J]
amount of charge, measured in volts, see physics electromagnetics. q, charge [A⋅s] [C]
V = PE/q [J/(A⋅s)] = [J/C] = [V]
E° - Standard Cell P.: The cell potential when the concentration of each type of ion taking part in the cell
reaction id 1[mol/l] and all the gases are at 1[atm] pressure. E° is the sum of its two standard electrode
potentials: E° = E°(cathode) + E°(anode) [V]; see table below.
E°(ox/red) - Standard Electrode P.: The contribution of an electrode to the standard cell potential see
redox couple (half cell reaction).
Standard Reduction P.: The voltage measured as a redox reaction occurs at the electrode when all
solutes are 1[mol/l] and all gases are at 101[kPa].
Rules for SRP: i) all E° values of the half reactions are read from left to right
i) the more positive the greater the tendency to be reduced
i) half reactions are reversible; any electrode can act either as an anode or as a cathode
i) under STP conditions any species on the left of a given halfreaction will react
spontaneously with a species that appears on the right above it (diagonal rule)
i) a change in stoichiometric coefficient does not effect the value of E°
i) E° changes sign whenever a halfreaction is reversed

Classification of solutes in aqueous solution

Strong e-lyte Weak e-lyte Non-e-lyte

HCL CH3COOH (NH2)2CO (urea)
HNO3 HF CH3OH (methanol)
HCLO4 HNO2 C2H5OH (glucose)
H2SO4 NH3 C6H12O6 (glucose)
NaOH H2O C12H22O11 (sucrose)
Ba(OH)2
ionic
compounds

2
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Standard reduction potentials at 25[°C] ECS

Red oxidizing agent reducing agent Ox E° [V]

(removes e - reduces itself) (supplies e - oxidizes itself)
weak F2(g) + 2e- → 2F-(aq) strong +2.87
O3(g) + 2H+(aq) + 2e → O2(g) + H2O(l) +2.07
Co3+(aq) + e- → Co2+(aq) +1.82
2H2O2(aq) + 2H+(aq) + 2e- → 2H2O +1.77
PbO2(s) + 4H+(aq) +SO42-(aq) + 2e- → PbSO4(s) + 2H2O(l) +1.70
Ce4+(aq) + e- → Ce3+(aq) +1.61
MnO4 (aq) + 8H+(aq) + 5e- →
-
Mn2+(aq) + 4H2O +1.51
Au3+(aq) + 3e- → Au(s) +1.50
Cl2(g) + 2e- → 2Cl-(aq) +1.36
CrO2 (s) + 14H+(aq) + 6e- →
2-
2Cr3+(aq) + 7H2O(l) +1.33
MnO2(s) + 4H+(aq) + 2e- → Mn2+(aq) + 2H2O(l) +1.23
O2(g) + 4H+(aq) + 4e- → 2H2O(l) +1.23
Br2(l) + 2e- → 2Br-(aq) +1.07
NO3 (aq) + 4H+(aq) + 3e- →
-
NO(g) + 2H2O(l) +0.96
2Hg2+(aq) + 2e- → 2Hg22+(aq) +0.92
Hg22+(aq) + 2e- → 2Hg(l) +0.85
Ag+(aq) + e- → Ag(s) +0.80
Fe3+(aq) + e- → Fe2+(aq) +0.77
O2(g) + 2H+(aq) + 2e- → H2O2(aq) +0.68
MnO4(aq) + 2H2O + 3e- → MnO2(s) + 4OH-(aq) +0.59
I2(s) + 2e- → 2I-(aq) +0.53
O2(g) + 2H2O + 4e- → 4OH-(aq) +0.40
Cu2+(aq) + 2e- → Cu(s) +0.34
AgCl(s) +e- → Ag(s) + Cl-(aq) +0.22
SO4 (aq) + 4H+ + 2e- →
2-
SOs(g) + 2H2O(g) +0.20
Cu2+(aq) + e- → Cu+(aq) +0.15
Sn4+(aq) + 2e- → Sn2+(aq) +0.13
2H+(aq) + 2e- → H2(g) 0.00
Pb2+(aq) + 2e- → Pb(s) -0.13
Sn2+(aq) + 2e- → Sn(s) -0.14
Ni2+(aq) + 2e- → Ni(s) -0.25
Co2+(aq) + 2e- → Co(s) -0.28
PbSO4(s) + 2e- → 2Pb(s) + SO42-(aq) -0.31
Cd2+(aq) + 2e- → Cd(s) -0.40
Fe2+(aq) + 2e- → Fe(s) -0.44
Cr3+(aq) + 3e- → Cr(s) -0.74
Zn2+(aq) + 2e- → Zn(s) -0.76
2H2O(l) + 2e- → 2H2(g) + 2OH-(aq) -0.83
Mn2+(aq) + 2e- → Mn(s) -1.18
Al3+(aq) + 3e- → Al(s) -1.66
Be2+(aq) + 2e- → Be(s) -1.85
Mg2+(aq) + 2e- → Mg(s) -2.37
Na+(aq) + e- → Na(s) -2.71
Ca2+(aq) + 2e- → Ca(s) -2.87
Sr2+(aq) + 2e- → Sr(s) -2.89
Ba2+(aq) + 2e- → Ba(s) -2.90
K+(aq) + e- → K(s) -2.93
strong Li+(aq) + e- → Li(s) weak -3.05

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Glossary - Chemistry (Solutions / RedOx)

Disproportionation Reaction.: see Redox R.
Oxidation: 1) Combination with oxygen. 2) A reaction in which an atom, ion, or molecule loses an electron;
e.g.: Ca(s) → Ca2+(s) + 2e- (represents the Ca in CaO).
Oxidation corresponds to an increase of oxidation number.
O. Agent: A substance that can accept electrons from another substance or increase the oxidation number
in another substance (being oxidized); where the substance itself is reduced;
e.g.: O2, O3, MnO4-, Fe3+
Nox - O. Number: The effective charge on an atom in a compound, calculated according to a set of rules.
An increase in ON. corresponds to oxidation, and a decrease to reduction.
O. Reaction: The half-reaction that involves the loss of electrons;
O. Rules: Rules to assign ON; see also table below.
1. In free elements, each atom has an ON. of (0); e.g.: H2, He, Br2, Na, Be, K, O2, P4, S8, etc.
2. Single-atomic ions; ON. is equal to the charge on the ion; e.g. Na+ (+1), Mg2+ (+2), Fe3+ (+3), etc.
3. In neutral molecules, the sum of the ON. of all the atoms must be 0; in polyatomic ions, the sum of
ON. of all atoms in the ion must be equal to the net charge of the ion; e.g.: NH4+ (-3) + 4(+1) = (+1).
4. The most negative element F has an ON. of (-1); the 2nd-most negative, O usually (-2).
5. H has an ON. of (+1), e.g. HCl; when bonded to metals, the O.N. is (-1), e.g.: LiH, NaH, CaH2.
6. Compounds consisting of non-metallic atoms, O.N. is equal to the individual ionic net charge;
e.g.: PCl3 (+3) + 3(-1) = (0)
RedOx Reaction: A reaction in which there is either a transfer of electrons or a change in the oxidation numbers
of the substances taking part in the reaction. Oxidation and reduction takes place simultaneously, because
an electron that is lost by one atom is accepted by another. Oxidation-reduction reactions are important
means of energy transfer in living systems. e.g.: Ca(s) + ½O2(g) → CaO(s);
Balancing RedOx R.: Balancing Redox-Reactions is done with the ion-electron method:
1. Write unbalanced EQ for the reaction in ionic form;
2. Separate the equation into two-half reactions;
3. Balance the atoms other than O and H in each half reaction separately;
4. For reactions in acidic, basic media, add H2O to balance the O atoms and H+ to balance H atoms;
5. Determine NOX, add electrons to one side of each half-reaction to balance the charges; if necessary
balance electrons by multiplying one or both half reactions by an appropriate coefficient;
6. Add the two half reactions together and balance the final EQ by inspection (e cancel on both sides);
7. Verify that the EQ contains the same types, masses, numbers of atoms, and charges on both sides;
Types of Redox R.:
Combination R.: Any combustion reaction that involves elemental oxygen or other e.g.: A + B → C
Decomposition R.: A breakdown reaction of a compound into two or more components;
e.g.: C → A + B
Displacement R.: An ion, or atom, in a compound is replaced by an ion, or atom of another one;
e.g.: A + BC → AC +B
• Hydrogen D.: All group-I and some group-II elements (Ca, Sr, Ba) will displace H from water;
• Metal D.: Metals ranking higher in their period will displace the lower ranked elements.
• Halogen D.: The power of elements in group-7 decrease as we move down from F to I; molecular F
can replace Cl, Br, and I etc.
Disproportion R.: A RedOx reaction in which a single element is simultaneously oxidized and reduced;
e.g.: 2Cu+(aq) → Cu(s) + Cu2+(aq).
Komproportion R.: The opposite reaction of disporportion; in which two compounds react to form a
product; e.g.: 2MnO4-(s) + 3Mn2+(aq) → 5MnO2(aq)
Reduction: (L. reductio, bringing back) 1) The removal of oxygen form (bringing back a metal from its oxide) or
the addition of hydrogen to a compound. 2) A reaction in which an atom, an ion, or a molecule gains an
electron; reduction takes place simultaneously with oxidation;
e.g.: ½O2(g) + 2e- → O2- (represents the O in CaO).
Reduction corresponds to a decrease in oxidation number.
R. Agent: A substance that can donate electrons to another substance or decrease the oxidation numbers
in another substance (being reduced), where the substance itself is oxidized; e.g.: H2; H2S, SO32-
R. Reaction: The half-reaction that involves the gain of electrons; e.g.: S(s) + 3Fe2(g) → SF6

1
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Oxidation Number of elements in their compounds:

I II III IV V VI VII VIII

1 H He
+1
-1
2 Li Be B C N O F Ne
+1 +2 +3 +4 +5 +2 -1
+2 +4 -1/2
-4 +3 -1
+2 -2
+1
-3
3 Na Mg Al Si P S Cl Ar
+1 +2 +3 +4 +5 +6 +7
+-4 +3 +4 +6
-3 +3 +5
-2 +4
+3
+1
-1
4 K Ca Sc Ti V +5 Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
+1 +2 +3 +4 +4 +6 +7 +3 +3 +2 +2 +2 +3 +4 +5 +6 +5 +4
+3 +3 +5 +6 +2 +2 +1 -4 +3 +4 +3 +2
+2 +2 +4 +4 -3 -2 +1
+3 +3 -1
+2 +2
5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
+1 +2 +3 +4 +5 +6 +7 +8 +4 +4 +1 +2 +3 +4 +5 +6 +7 +6
+4 +4 +6 +6 +3 +2 +2 +3 +4 +5 +4
+3 +4 +4 +2 -3 -2 +1 +2
+3 -1
6 Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Ti Pb Bi Po At Rn
+1 +2 +3 +4 +5 +6 +7 +8 +4 +4 +3 +2 +3 +4 +5 +2 -1
+4 +6 +4 +3 +2 +1 +1 +1 +2 +3
+4
7 Fr Ra
+1 +2

• The sum of the oxidation numbers of all the atoms in the species is equal to its total charge.
• For the atoms in element form it is 0.
• For elements of group I = +1
II = +2
III = +3 (except B)
IV = +4 (except Si and C)
• For H +1 in combination with nonmetals -1 in combination with metals
• For F is always -1.
• For O -2 unless combined with F -1 in peroxides O22-
-½ in superoxides O2-
-1/3 in ozonides O3-

2
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Glossary - Chemistry (Solutions / Acid and Bases)

Acid-Base Concepts: Acids are molecules or ions that can donate protons to other molecules; in aqueous
solutions generate cations by proton transfer (H+) to water molecules (H3O+); bases are molecules or ions
that can accept protons from other molecules; in aqueous solutions generate anions (OH-) by accepting
protons from water molecules; In terms of ionic charge, the acid in general is considered more positive
than the base.
A.B. Indicator: see indicator.
A.B. reaction: A reaction between an acid and a base; reaction between conjugated acid-base pairs.
e.g.: CH3COOH(aq) + H2O(l) ↔ H3O+(aq) + CH3COO-(aq)
Arrhenius: Acid-base concept based on aqueous solutions only.
• A. Acid: A compound that contains hydrogen and releases hydrogen ions (H+) in water; e.g.: HCL,
CH3COOH; but not CH4.
• A. Base: A compound that produces hydroxide ions (OH-) in water; e.g.: NaOH, NH3, but not Na.
Bronsted: Acids and bases description based on their proton acceptability (H+) and donation (H-).
• B. Acid: A substance capable of donating a proton; a source of H-ions, H+;
e.g.: HCL, CH3COOH; HCO3-, NH3
• B. Base: A substance capable of accepting a proton; a species to which hydrogen ions H+ can bond;
e.g.: OH-, Cl-, CH3CO2-, HCO3-, NH3;
Conjugate A.B. Pair: A bronsted acid and its conjugated base; i.e.: two substances related by gain or loss
of an H+ ion (proton); e.g.: HCl and Cl-, NH4 and NH3.
• C. A.: The bronsted acid formed when a bronsted base has accepted a proton;
e.g.: NH4 is the conjugate acid of NH3.
• C. B.: The bronsted base formed when a bronsted acid has donated a proton;
e.g.: NH3 is the conjugate acid of NH4.
Lewis A.B.: His theory is based upon the donation of electron pairs from one compound to another; hence
does not necessarily exclude compounds which are not made partly of oxygen and hydrogen and is not
limited to water as a solvent, instead any net ionic equation contains L-acid and L-bases.
• L. Acid: A substance that can accept a pair of electrons; e.g.: H+, Fe3+, Cu2+, BF3, etc.
• L. Base: A substance that can donate a pair of electrons; e.g.: OH-, H2O, NH3, CN, F-, etc.
Polyprotic A.B.: A bronsted acid or base that can donate or accept more than one proton; see table.
• Monoprotic A.: A bronsted acid with one acidic hydrogen atom; each of the acid yields an H-atom;
e.g.: HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq).
• Diprotic A.: Each unit of the acid yields two H-atoms; e.g.: KA1 > KA2
H2SO4(aq) + H2O(l) → H3O+(aq) + HSO4-(aq) KA1 = completely dissociated
HSO4-(aq) + H2O(l) ↔ H3O+(aq) + SO42-(aq) KA2 = 1.3⋅10-2
• Triprotic A.: These yield three H ions; e.g.: KA1 > KA2 > KA3
+

H3PO4(aq) + H2O(l) → H3O+(aq) + H2PO4-(aq) KA2 = 7.6⋅10-3

H2PO4 (aq) + H2O(l) → H3O (aq) + HPO4 (aq)
- + 2-
KA2 = 6.2⋅10-8 ∆ = 105!
HPO4 (aq) + H2O(l) → H3O (aq) + PO4 (aq)
2- + 3-
KA2 = 1.1⋅10 -12
∆ = 104!
Weak A.B.: The stronger electronegativity, and the larger the element, the more likely an H-atom can be
drawn away from it: 1) Acids and bases that are only partially ionized in aqueous solutions at normal
concentrations. 2) Acids and bases for which Ka and Kb, respectively, are small compared to 1;
e.g.: HF, CH3COOH (weak acids); NH3, CH3NH2 (weak bases).

1
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Acid: A substance that dissociates in water by yielding hydrogen ions (H+), with subsequent hydration;
e.g.: HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq).
KA - Ionization constant: The equilibrium constant for the proton transfer of water(acid concentration);
e.g.: HA + H2O ↔ H3O+ + A- KA = KC⋅c(H2O) = c(H3O+)⋅c(A-) / c(HA)
Binary A.: An acid that contains only two elements.
Carboxylic A.: An acid that contains the carboxyl group -COOH.
Strength of A.: H-A = Hydrogen Acid bond.
• Acid - the more polar the H-A bond (due to a larger difference of electronegativity), the stronger the
acid of the same period; e.g.: H2O < HF <HCl < HBr < HI.
• The weaker the H-A bond (due to greater atomic radii), the stronger the acid; this effect is dominant
for acids of the same group; e.g.: HF < HCl < HBr < HI
• OxoAcid - the greater the number of O-atoms attached to the central atom (the greater its oxidation
number (NOX), the stronger the acid; e.g.: HClO < HClO2 < HClOO3 < HClO4
• If the same number of O-atoms are attached to the central atom (= greater EN of the central atom), the
stronger the acid; e.g.: HClO < HBrO < HIO
Amphoteric: The ability to react with both acids and bases; e.g.: H2O; HCO3-, NH3, HSO4-, see oxides.
Anylate: see titration.
Arrhenius: see acid-base.
Base: A substance that dissociates in water by yielding hydroxide ions (OH-) with subsequent hydration;
e.g.: NaOH(s) → H2O → Na+(aq) + OH-(aq)
KB - B. Ionisation Constant: Equilibrium constant for the proton transfer from water(base ionization);
e.g.: B + H2O ↔ BH+ + OH- KB = c(BH+)⋅c(OH-) / c(B)
Strength of A.B.: The stronger the base the weaker its conjugate acid and vice versa.
Buffer: A solution that resist any change in pH when small amounts of acid or base are added; i.e. a weak acid or
base with its conjugated counterpart in relative high concentrations; see Henderson-Hasselbalch.
Acid B.: Stabilizes solutions at pH < 7; e.g.: a solution containing both CH3CO2H and CH3 CO2-.
Base B.: Stabilizes solutions at pH > 7; e.g.: a solution containing NH3 and NH4+.
B. Capacity: The amount of hydronium ions (H3O+) or hydroxide ions (OH-) that can be added to a buffer
solution without changing its pH by more than 1 unit.
B. Range: The pH range over which a buffer solution is most effective.
B. Solution: A solution of a weak acid or base and its salt; both components must be present. The solution
has the ability to resist changes in pH upon the addition of small amounts of either acid or base.
α - % Dissociation: Overall concentration of a weak acid. c(0), initial concentration [mol/l]
α = 100⋅c(A-) / c(0) = 100⋅c(A-) / (c(HA) + c(A-)) [-] c(A-), conc. conj. base [mol/l]
KA = α2⋅c(0) / (1-α) [mol/l] c(HA), concentr. of acid [mol/l]
Equivalence Point: The point at which the acid has completely reacted (neutralized) with the base - see titration.
Henderson-Hasselbalch Equation: An approximate equation for
estimating the pH of a solution containing a weak conjugate acid KA, ioniz. const. [mol/l]
and base, and specifically that of a buffer solution. c(A-), conc. of conj. base [mol/l]
pH = log(KA) + log(c(A-)) - log(c(HA)) [-] c(HA), concentr. of acid [mol/l]
pH = pKA + log(c(A-) / c(HA)) [-]
Hydroxyl: see ion.
Indicator: A weak organic acid or base that changes color when it shift from acid to its base form; i.e. acid-base
neutralization (acid-base indicator) or from its oxidized to its reduced form (a redox indicator); e.g.:
HInd(aq) + H2O(l) ↔ H3O+(aq) + Ind-(aq) where c(H3O+) is the H+ concentration (pH) of the anylate.
K(Ind) / c(H3O+) = c(HInd) / c(Ind-) [mol/l]
if = c(HInd) / c(Ind-) > 1, color of acid (Hind) predominates
if = c(HInd) / c(In-) < 1, color of conjugate base (Ind-) predominates
Preconditions are: both 1) HIn and In- have to be water soluble, 2) HIn and In- have to separate colors, 3)
Indicator concentration c(HIn) has to be low, 4) c(H3O+) of indicator should measure only pH of solution.
Ion: An atom / molecule that has lost / gained one or more electrons, and thus becomes positively or negatively
charged; i.e.: Al3+ (mono-atomic ion), SO4- (poly-atomic ion); see chemistry-atom and molecule.
Common I. Effect: The shift on equilibrium caused by the addition of a compound having an ion in
common with the dissolved substances e.g.:
acid and electrolyte of the acetate ion - CH3COOH and CH3COONa;
Hydration of I.: The attachment of water molecules to a central ion.
Hydronium I.: The ion H3O+; H5O3+; H7O5+;
Hydroxyl Group: An OH- group; a negatively charged ion formed by the disassociation of a water
molecule-compound that consists only of hydrogen and carbon atoms.
Oxoanion: An anion derived from an oxoacid; e.g.: HCO3-, CO32-.

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Ionization: 1) Conversion to cations by the removal of electrons. 2) The donation of a proton from a neutral acid
molecule to a base with the formation of the conjugate base (an anion, in this instance) of the acid
(dissociation into ions of a compound in solution);
e.g.: K(g) → K+(g) + e-(g)
or: CH3COOH(aq) + H2O(l) → H3O+(aq) + CH3CO2-(aq)
Lewis: see acid-base.
Neutralizing Reaction: A reaction between an acid and a base resulting in the formation of solute and a salt,
where the acid provides the anions and the base the cations: e.g.:
HCl(aq) + NaOH(aq) → H2O + Na+(aq) + Cl-(aq)
Oxoacid: An acid that contains O, H and another central element; e.g.: H2CO3, HNO3, HNO2, HClO;
• Metallic OA.: H-O-Z (Z stands for the metallic compound), showing lesser electronegativity than O,
i.e.: H-O bond is more covalent, leading to hydroxide (OH-) separation when dissolved in water;
• Non-Metallic-OA.: H-O-Z (Z stands for the non-metallic compound), here the O-Z bond is more
covalent, leading to hydrogen (H+) separation when dissolved in water; e.g.: HOI, HOBr, HOCl, ect.
Multiple OA.: Acidity increases with extra O-atoms in the n = 0, weak acidity
compound (relocate electron-density), which facilitates H+ n = 1, intermediate acidity
dissociation; (HO)mZOn n = 2, strong acid
e.g.: HOCl < HOClO < HOClO2 < HOClO3. n = 3, very strong acid
Non-OA.: Acids in which the H-atom is not bounded to an O-atom; e.g.: HF, HCl, H2S, etc. in which
acidity increases moving down a group: NH3 < H2O < HF or PH3 < H2S < HCl
Oxide: see table below
Acidic O.: An oxide of nonmetallic elements (p-block) that reacts with water to give an acid;
• Soluble: React with water by forming acids; like CO2; N2O5, SO2, SO3
e.g.: CO2 + 2H2O → H3O(aq) + HCO3-.
• Insoluble: Acid oxides can react with basic oxides; CaCO3(s) → T → CaO(s) + CO2(g)
e.g.: CaO(s) + SiO2(s) → CaSiO3(s)
Amphoteric: The ability to react with both acids and bases; e.g.: H2O; HCO3-, NH3, HSO4-,
• A. Oxide: An oxide that exhibits both acidic and basic properties.
Basic O.: The oxides (Bronsted base) of metallic elements (s-block) are generally basic e.g.: MgO, Na2O
etc. except BeO.
• Soluble: Group-I and II elements; CaO, SrO, BaO posses ionic structure already in the solid phase;
e.g.: CaO + H2O → Ca2+ + O2- + H2O → Ca2+ + 2OH-.
• Insoluble: Become soluble only during a neutralization reaction;
e.g.: Fe2O3(s) + 6H2O+(aq) → 2Fe3+(aq) + 9H2O(l)
pH (potential hydrogen): The negative logarithm of the hydronium ion pH = -log(c(H3O+))
concentration in a solution, at 25°C; KC = c(H3O+)⋅c(OH-) / c2(H2O)
e.g.: auto-dissociation of pure water 2H2O(l) ↔ H3O+(aq) + OH-(aq)
pH<7 indicates an acidic solution, and pH>7 a basic solution; KW = 10-14[mol2/l2] = 10-7[mol/l]
pH for weak acids/bases: simplification is applicable see autoionization of water for KW
acid: pH = ½⋅(pKA - logc0) KA, acid ionization const. [-]
base: pH = ½⋅(pKB - logc0) KB, base ionization const. [-]
pH for intermediate acids/base: no simplification applicable here: c0, initial concentration [mol/l]
acid: x = c(H3O+) = -½⋅KA + √(1/4⋅KA2 + KA⋅c0) pH = ½⋅pKW + ½⋅(KB + logc0)
base: x = c(OH-) = -½⋅KB + √(1/4⋅KB2 + KB⋅c0)
pH-Curve: The graph of the pH of a reaction mixture
against volume of titrant added in an acid-base titration.
pOH (potential hydroxide) The negative logarithm of the pOH = -log(c(OH-))
hydroxide ion concentration in a solution, at 25°C; KW = 10-14[mol2/l2]
ionic product of pH and pOH reflects pKW pH + pOH = 14
Salt: The product (other than water) of the reaction between an acid and a base; see chemistry liquid - solubility
equilibria e.g.: NaCl, K2SO4, etc.
• S. from strong bases and acids: pH of solution remains unaffected and oscillates around pH = 7;
e.g. NaCl, KNO3, etc.
• S. from strong bases and weak acids: pH shifts to levels >7; e.g.: KNO2, NaCN, Ca(CH3CO2)2, etc.
• S. from weak bases and strong acids: pH shifts to levels <7; e.g.: NH4NO3, FeBr2, AlCl3, ect.
• S. from weak acids and bases: both acidic or basic solutions are possible, depending upon the acidic-
or basic-character of the anions.
Substitution Reaction: 1) A reaction in which an atom (or a group of atoms) replaces an atom in the original
molecule. 2) In complexes, a reaction in which on Lewis base expels another and takes its place;
e.g.: [Fe(H2O)6]3+(aq) + 6CN-(aq) → [Fe(CN)6]3-(aq) + 6H2O(l)

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Titration: The analysis of composition by measuring the volume of one solution (titrant) needed to react with a
given volume of another solution (the anylate).
T. involving a strong acid and base: Equivalence point oscillates around pH 7 (steep rise at EP).
T. i.a. weak acid and strong base: Equivalence point oscillates around levels pH 7 - 14.
T. i.a. weak base and strong acid: Equivalence point oscillates around levels pH 1 - 7.
T. i.a. a weak acid and base: Equivalence point oscillates around pH 7, with a flat EP, hard to determine.
T. Curve: The sigmoidal profile of pH (anylate) over the volume of added titrant or level of titrant τ.
Anylate: The solution of unknown concentration in a titration.
Equivalence Point (stoichiometric point): The stage in a titration
when exactly the right volume of solution needed to complete n, molar amount [mol]
the reaction has been added; n(An)An = n(Ti)Ti V, volume of probes [l]
cAn⋅VAn = cTi⋅VTi c, molar concentration [mol/l]
Titrant: The solution of known concentration added from a buret in a titration
KW - Water Autionization Constant: The equilibrium constant for the auto-ionization of water,
2H2O(l) ↔ H3O+(aq) + OH-(aq) KW = KA⋅KB
KW = c(H3O+)⋅c(OH-) = KC⋅c2(H2O) = 10-14[mol2/l2] pKW = pKA⋅pKB

Representative elements and their highest oxidation states

I II III IV V VI VII VIII

Li2O BeO B2O3 CO2 N2O5 OF2

Na2O MgO Al2O3 SiO2 P4O10 SO3 Cl2O7
K2O CaO Ga2O3 GeO2 As2O5 SeO3 Br2O7
Rb2O SrO In2O3 SnO2 Sb2O5 TeO3 I2O7
Cs2O BaO Ti2O3 PbO2 Bi2O5 PoO3 At2O7

s-block p-block

basic oxide acidic oxide amphoteric oxide

Common acids and bases in aqueous solution:

acid/base conjugate a/b KA ** KB pKA * pKB

a PerChloric Acid HClO4+ H2O ↔ H3O+ + ClO4- ≈-
10
a Hydroiodic Acid HI + H2O ↔ H3O+ + I- ≈-
10
a Hydrobromic Acid HBr + H2O ↔ H3O+ + Br- ≈ -9
a Hydrochloric Acid HCl + H2O ↔ H3O+ + Cl- ≈ -6
a Sulfuric Acid H2SO4 + H2O ↔ H3O+ + HSO4- ≈ -3
Hydrogen sulfate ion HSO4- + H2O ↔ H3O+ + SO42- 1.3⋅10-2 1.89
a Nitric Acid HNO3 + H2O ↔ H3O + + NO3- -1.3
a Hydronium H3O+ + H2O ↔ H3O+ + H2O 1.0⋅100 -1.7
b Urea NH2CONH2 + H2O ↔ OH- + NH2CONH3+ 1.3⋅10-14 13.90

a Trichloroacetic Acid CCl3CO2H + H2O ↔ H3O+ + CCl3CO2- 3.0⋅10-1 0.52

a Benzenesulfonic Acid C6H5SO3H + H2O ↔ H3O+ + C6H5SO3- 2.0⋅10-1 0.70
a Iodic Acid HIO3 + H2O ↔ H3O+ + IO3- 1.7⋅10-1 0.77
a Sulfurous Acid H2SO3 + H2O ↔ H3O+ + HSO3- 1.6⋅10-2 1.81
HSO3- + H2O ↔ H3O+ + SO32- 5.6⋅10-8 7.25
a Phosphorous Acid H3PO3 + H2O ↔ H3O+ + H2PO3- 1.6⋅10-2 1.79
H2PO3- + H2O ↔ H3O+ + HPO32- 7.0⋅10-7 6.15
a Oxalix Acid H2C2O4 + H2O ↔ H3O+ + HC2O4- 5.9⋅10-2 1.22
HC2O4- + H2O ↔ H3O+ + C2O42- 5.4⋅10-5 4.27
a Chlorous Acid HClO2 + H2O ↔ H3O+ + ClO2- 1.0⋅10-2 2.00
a Dichloric Acid Cl2HC2O2H + H2O ↔ H3O+ + Cl2HC2O2- 3.3⋅10-2 1.48

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a Phosphoric Acid H3PO4 + H2O ↔ H3O+ + H2PO4- 7.6⋅10-3 2.12
Dihydrogen phosphate ion H2PO4- + H2O ↔ H3O+ + HPO42- 6.2⋅10-8 7.21
Hydrogen Phosphate ion HPO42- + H2O ↔ H3O+ + PO43- 1.0⋅10-12 12.0
a Chloroacetic Acid CH2ClCO2H + H2O ↔ H3O+ + CH2ClCO2- 1.4⋅10-3 2.85
a Arsenic Acid H3AsO4 + H2O ↔ H3O+ + H2AsO4- 2.5⋅10-4 3.60
H2AsO4- + H2O ↔ H3O+ + HAsO42- 5.6⋅10-8 7.25
HAsO4 + H2O ↔ H3O+ + AsO43- 3⋅10-13 12.5
a Lactic Acid CH3CH(OH)CO2H + H2O ↔ H3O+ + CH3CH(OH)CO2- 8.4⋅10-4 3.08
a Nitrous Acid HNO2 + H2O ↔ H3O+ + NO2- 4.3⋅10-4 3.37
a Hydroflouric Acid HF + H2O ↔ H3O+ + F- 3.5⋅10-4 3.45
a Formic Acid HCO2H + H2O ↔ H3O+ + HCO2- 1.8⋅10-4 3.75
a Cyanic Acid HOCN + H2O ↔ H3O+ + OCN- 1.2⋅10-4 3.9
a Benzoic Acid C6H5CO2H + H2O ↔ H3O+ + C6H5CO2- 6.0⋅10-5 4.19
b Aniline C6H5NH2+ H2O ↔ OH- + C6H5NH3+ 4.3⋅10-10 9.37
a Stickstoffwasserstoff Acid HN3 + H2O ↔ H3O+ + N3- 1.9⋅10-5 4.7
a Acetic Acid (HAc) CH3CO2H + H2O ↔ H3O+ + CH3CO2- 1.8⋅10-5 4.74
b Nicotinic Acid C5H4NCO2H+ H2O ↔ H3O+ + C5H4NCO2- 1.4⋅10-5 4.85
b Pyridine C5H5N + H2O ↔ OH- + C5H6O+ 1.8⋅10-9 8.75
b Hydroxalamine NH2OH + H2O ↔ OH- + NH2OH2+ 1.1⋅10-8 7.97
a Propane Acid C2H5CO2H + H2O ↔ H3O+ + C2H5CO2- 1.3⋅10-6
a Carbonic Acid H2CO3 + H2O ↔ H3O+ + HCO3- 4.3⋅10-7 6.37
Hydrogen Carbonate ion HCO3- + H2O ↔ H3O+ + CO32- 4.8⋅10-11 10.3
a Hydrosulfuric Acid H2S + H2O ↔ H3O+ + HS- 1.1⋅10-7 7.0
Hydrogen Sulfide ion HS- + H2O ↔ H3O+ + S2- 1.0⋅10-14 14.0
Water H2O + H2O ↔ H3O+ + OH- 1.0⋅10-7 7.0 7.0
a Hypochlorous Acid HClO + H2O ↔ H3O+ + OCl- 3.0⋅10-8 7.53
b Hydrazine N2H4 + H2O ↔ N2H3+ + OH- 9.8⋅10-7 6.0
b Nicotine C10H14N2 + H2O ↔ OH- + C10H15N2+ 1.0⋅10-6 5.98
b Morphine C17H19O3N + H2O ↔ OH- + C17H20O3N+ 1.6⋅10-6 5.79
b Hydrazine NH2NH2 + H2O ↔ OH- + NH2NH3+ 1.7⋅10-6 5.77
a Hyprobromous Acid HBrO + H2O ↔ H3O+ + OBr- 2.0⋅10-9 8.69
b Pyridine C6H5N + H2O ↔ OH- + C6H5NH+ 1.5⋅10-9 5.2
a Bromous Acid H3BO3 + H2O ↔ H3O+ + H2BO3- 9.1
a Boric Acid B(OH)3 + 2H2O ↔ H3O+ + B(OH)4- 7.2⋅10-10 9.14
a Hydrocyanic Acid HCN + H2O ↔ H3O+ + CN- 4.9⋅10-10 9.31
b Ammonia NH3 + H2O ↔ OH- + NH4+ 1.8⋅10-5 4.74
a Phenol C6H5OH + H2O ↔ H3O+ + C6H5O- 1.3⋅10-10 9.89
a Hypoiodous Acid HIO + H2O ↔ H3O+ + IO- 2.3⋅10 -11 10.64

b Trimethylamine (CH3)3N + H2O ↔ OH- + (CH3)3NH+ 6.5⋅10 -5

4.19
b Methylamine (CH3)3NH2 + H2O ↔ OH- + (CH3)3NH3+ 3.6⋅10-4 3.44
b Dimethylamine (CH3)2NH + H2O ↔ OH- + (CH3)2NH2+ 5.4⋅10-4 3.27
b Ethylamine C2H5NH2 + H2O ↔ OH- + C2H5NH3+ 6.5⋅10-4 3.19
b Triethylamine (C2H5)3N + H2O ↔ OH- + (C2H5)3NH+ 1.0⋅10-5 2.99
a H2O2 + H2O ↔ H3O+ + HO2- 11.6
b H2BO3- + H2O ↔ OH- + H+ 1.3
b HBO32- + H2O ↔ OH- + H+ 0.2
b Hydroxide ion OH- + H2O ↔ OH- + H2O 1.0⋅100 0.0
b Hydride ion H- + H2O ↔ OH- + H2 -1.7
b Amide ion NH2- + H2O ↔ OH- + NH3 ≈ -9
b O2- + H2O ↔ 2OH- ≈-
10
**) KA⋅KB = KW = 1.0⋅10-14 1.3⋅10-14 *) pKA+pKB = pKW = 14