UNIT 1 & 2 NOTES

ATOMIC STRUCTURE
The History of Atomic Structure
Democritus (400 B.C.)
A Greek philosopher, the first one to think about the existence of atoms.
Believed that matter was composed of tiny indivisible particles called atomos
(atoms).
No experimental evidence.

Lavoisier (1743-1794)
Worked on combustion reactions and established "Law of Consevation of Mass".
First one to realize the importance of combustion supporting gas, oxygen.

John Dalton (1766-1844)

A meteorologist and teacher.
Has supported evidence
Dalton's Theory:
1) All elements are composed of extremely small, indivisible particles called atoms.
2) Atoms of the same element have identical shape, size and chemical properties.
The atoms of any one element are different from those of another.
3) Atoms of different elements mix or combine in whole number ratios.
4) Chemical reactions occur when atoms seperate, join or rearrange. In a chemical
reaction, atoms of 1 element never change into another.

J. J. Thomson (1859-1940) - Cathode Ray Tube experiment to discover electrons

In the tube was an inert gas, a positive & a negative plates.
The gas particles were attacted to the postive plate
=> The particles must have a negative charge (electrons).
"Plum Pudding Model" = atom was a positive sphere with electrons shoved into the
sides of it.

Ernest Rutherford - Gold Foil Experiment to discover the nucleus

Shot high energy beam of positive alpha particles (= 2 protons + 2 neutrons) into
gold foil.
Expected results: the alpha particles would pass straight through the foil, but:
_Most alpha particles went through => the atom was mostly empty space
_Few particles were deflected at small angles => the atom contained something
small and positive (nucleus)
_Very rarely few particles were deflected at large angles => the atom contained a
small, very dense and positively charge centre (nucleus).
However, he could only account for half the mass of the atom.

James Chadwick (1891-1974) - Alpha particles bombardment of the metal

beryllium to discover neutrons
Bohr Model of Atom: Electrons, Flame Tests and AAS
_Bohr Model: electrons travel in orbits, or energy levels, around the nucleus. The
further the electron is from the nucleus; the more energy it has.

_Electromagnetic Spectrum: the classification of electromagnetic waves according to

their frequency.

_Photons: particles of light energy. Each wavelength of light has photons that have a different amount of
energy. The longer the wavelength, the lower the energy of the photons.
A photon of red light

A photon of blue light

Why an element’s line spectrum can be used to identify it?

 When atoms are energized, their ‘excited’ electrons jump to a higher energy level. Because of the
instability, the electrons will jump back down to its ground state by releasing the excess energy in the form
of light. The wavelengths of the emitted light depend on the difference in energy between the ground
state and the excited state. An element has its unique amount of energy in each energy level, so it only
emits specific wavelengths and light color.

_Flame Test: an analytical technique that relies on an element’s unique emission spectrum to identify its
presence in a mixture of compound, often used for metal ions.
1) A small sample of unknown compound is placed in a hot Bunsen burner flame.
2) Sample vaporizes and the heat of flame excites electrons.
3) Electrons return to ground state, emitting light with wavelength characteristic of element.
_Limitations: only used for small number of metal ions due to limited source of energy in the flame; unclear
results, which can be fixed by using a spectroscope to see line emission spectrum of a flame.
Metal Ions Flame Color
Boron Bright Green
Barium Pale Green
Calcium Red/ Orange
Copper Blue/ Green
Iron Gold
Lithium Red/ Crimson
Potassium Violet/ Lilac
Strontium Deep Red
Sodium Yellow
Atomic Absorption Spectroscopy – both quantitative and qualitative
1) Use lamp containing the same
element (gaseous state) as the
one being tested and emits
identical wavelengths of light to
be absorbed.
2) Sample is vaporized by the
flame and atoms are separated
so they can absorb light from
the hollow cathode lamp. Only
atoms of the element that we are looking for will absorb the emitted wavelengths from the HCL.
3) Light passes through a monochromator, which is set to select specific wavelength for analysis by
detector.
4) Detector measures the amount of light remaining without being absorbed = absorbance value, which
tells us the quantity of the element being tested.

Classifying of Matter

Homogeneous Heterogeneous

Pure substance Mixture

o Particles making up a substance are
all of the same kind o Contain two or more different kinds of substances
o Materials with distinct measurable o Can be composed of elements or compounds or both, which
properties (MP, BP, reactivity, can be separated out by physical properties
strength and density) o Material properties dependent on the identity and relative
amounts of the substances that make up the mixture
Element Compound
Cannot be Made from two or Solution Mixture that is not a solution
decomposed by more elements Ex. Air, sugar dissolved in Ex. Nuts mixture or cement
chemical reactions, chemically water solution and glass
either metal or combined
non-metal

Physical Properties Chemical Properties

Can be found by studying the substance itself rather Describe how it reacts chemically and its tendency to
than its chemical reactions form new substances
Ex. Solubility, state, MP, BP and conductivity Ex. Reactivity, ability to be stable or decompose with
heat and acidity
Separation of Mixture
Separation Magnetic Electrostatic Sieving Filtration
Method
Procedure Remove magnetic Separate materials Used for mixture of Separated insoluble
materials from a that have different solids or solid a liquid solids from a liquid or
mixture using a charge Unlikely to produce gas
magnet pure substance
Physical magnetism Electrical charge Particle size Particle size and
Properties solubility

Decantation Centrifugation Separating Funnel Gravity Evaporation

Separation
Separate solids A mixture is spun in Separates Heavier particles Heat causes a liquid to
from a liquid by a machine called immiscible liquids, fall to the bottom turn into a gas, leaving
pouring off the centrifuge. The high which form 2 layers when the behind any solids that
liquid to leave speed causes less due to difference in container is were dissolved in the
the solid behind dense substance to density shaken liquid, as well as
rise to the top impurities
Density Density Density and Force of gravity Boiling points
solubility

Crystallization (Fractional) Distillation

Purify impure solid substance Separate a solid from a liquid or a mixture
1) Dissolve both solid and impurities in of liquids.
solvent at high temperature. Hot 1) Solution is boiled, solvent changes to
filtration to remove any solid impurity vapor. Chromatography
2) Cool to room temp. Desired solid is 2) Vapor passes down a condenser,
less soluble at room temp, so it will where it is cooled and converted back
crystallize to liquid.
3) Crystals are filtered out of solution, 3) The liquid is collected as the distillate
washed with cold solvent and dried. in the flask.
Solubility Boiling point Intermolecular Forces

Chemical Bonding
 Giant Lattice Molecular
Metallic Ionic Covalent Network Molecular Covalent
Components Metals Metals and Non-metals Group IV Non-metals Non-metals
Examples Cu, Fe NaCl, CaO SiO2, C60 CO2, H2O, Cl2
Type of Positive ions Positive and negative
individual surrounded by ions Atoms Small individual
particles negative molecules
delocalized
electrons
Bonding Strong electrostatic Strong electrostatic Strong covalent Electronegativity
Force attraction attraction bonding due to Strong intramolecular
strong covalent bonding
electronegativity Weak intermolecular
forces
Directional    
Bonding
Delocalized    
Electrons Except for graphite
Model Mobile sea of Ionic lattice Continuous array Clusters of molecules
electrons (diamond)
Graphene (graphite)
Soccer ball shape
(fullerenes)
Lewis Dot  
Diagram

Ions with a transition metal = colored

Ions with an alkaline/ alkaline earth metal = white

 Properties of Types of Bonding

Properties Metallic Ionic Covalent

Molecular Network
Metallic Very high Moderate Low Low
characteristic
(reactivity/
tendency to lose
electrons)
M.P & B.P High High Low Very High
 Strong  Strong  Weak inter.  Strong intra.
electrostatic electrostatic Forces Molecular
attraction attraction covalent bonding
 The more between cations lattice
valence and anions
electrons, the
higher MP and
BP
Hardness Generally hard Hard Soft Very hard
 Strong bonding  Strong bonding  Weak inter.  Strong lattice
 Close packing Forces (except for
ions graphite)
Brittleness Not brittle Very brittle Soft Brittle
 Non- directional  Directional  Directional
bonding so bonding, shifted bonding, once
electrons can ions cause like disrupted the
slide over ions charged ions to lattice is brittle
repel
Malleability Malleable & ductile None, brittle None, soft None, brittle
(hammered into  Flexible
sheets) & delocalized
ductility (drawn electrons
into wires)
Solubility in Insoluble Soluble Insoluble Insoluble
water  Ion-dipole forces  No attraction
pull ions from its between water
lattice and the
molecules
Electrical Good Only conductive in Insulator Insulator
conductivity  Delocalized liquid and aqueous  Localized (except for graphite)
electrons phases electrons
 Charged ions
Thermal Good Quite good Poor Poor
conductivity  Delocalized  Vibrations of (except for graphite)
electrons and ions, but slow
vibration of ions
 Allotropes of Carbon
Allotropes:

 different forms of the same element;

 have similar chemical properties;
 but different physical properties due to different arrangement/ structure.

Structure
Diamond Fullerenes Graphite Carbon nanotubes
 Rigid 3-D tetrahedral  60+ carbon  Layers of 2-D  Walls made of 1
structure atoms hexagonal shaped single layer of
 Each atom is  Arranged in a rings. graphite rolled into
covalently bonded sphere or cage  6 rings (graphene)/ a cylinder shape
to 4 other atoms  Each atom is layer  Each atom is bonded
 No free electrons bonded to 3  Each atom is bonded to 3 other atoms
 Very strong covalent other atoms to 3 other atoms  1 delocalized
bonding  1 delocalized  1 delocalized electron/ electron/ atom
electron/ atom atom 
−9

 Weak inter.  Weak inter. 1+¿ dimension=1 nm=1× 10 m ¿

forces  Strong intra.
 Strong intra.
forces

Conductor of heat & electricity

Poor Semi Good Good
 No mobile electrons  Delocalized  Delocalized electrons
electrons can’t can move between
move between graphene sheets
molecules
Hard/ soft & brittle
Hard & brittle Hard & brittle Soft & tensile strength Same as graphite
 Continuous, strong  Graphene sheets can
array slip over each other
 Strong intra.
BP & MP
High High High High
 Chemistry Calculations
Relative Atomic Mass Ar: mass of an atom compared with one-twelfth of the mass of an atom of
carbon-1.
%a × Ar a +%b × Ar b
Ar =
100
Relative Molecular Mass Mr: mass of a molecule compared with one-twelfth of the mass of an atom of
carbon-12.
Ar =∑ of Ar of the component atoms

mass of a component
Percentage compos ition= ×100
mass of the total formula

Energy Changes in Chemical Systems

 Enthalpy/ Heat content (H): sum of potential and kinetic energy of a substance.
 Change in enthalpy: smaller in physical changes than chemical reactions

Δ H =H ( products ) – H ( reactants ) J mol−1 / kJ mol−1

Exothermic Reactions Endothermic Reactions

Release energy and heat up surroundings Absorb energy and cools the surrounding down
Negative ΔH: Enthalpy of product is reduced PositiveΔH: Enthalpy of product is increased
Energy Profile Diagram Energy Profile Diagram

2H2 (g) + O2 (g) 2H2O (l) ∆ H = -572 kJ CO2 (g) C (s) + O2 (g) ∆ H = +394 kJ
2H2 (g) + O2 (g) 2H2O (l) + 572 kJ CO2 (g) + 394 kJ C (s) + O2 (g)

Biofuels
 Biofuels: fuels made from organic materials & are considered “renewable”.

Biofuel Bioethanol Biodiesel

How it’s made Produced by fermentation of Made by processing feedstock:
starchy crops (sugar cane, corn, vegetable oil, soybean oil, animal
wheat) fats.
Advantages Few modifications are needed
 Carbon neutral, biodegradable and non-toxic
Suited to both small-scale home production or large scale industrial
Disadvantages Carbon dioxide is still produced due to transportation of feedstock,
biofuel production and distribution using fossil fuels
Deforestation to plant feedstock for resources

Organic Compounds
 An organic compound: made from carbon atoms.
 Structure & Properties:
o covalent bonds
o low MP & BP
o flammable
o soluble in non-polar solvents
o insoluble in water
 Hydrocarbons: organic compounds containing only carbon and hydrogen bonded covalently.
 4 classes of hydrocarbons: alkanes-saturated, alkenes, alkynes &aromatic compounds (benzene) –
unsaturated.
1. Alkanes: CnH2n+2

 Enprirical formula: ratio between C and H

Ex. C4H10 = C2H5
 Naming:
o prefix: meth-, eth-, prop-, but-, pent-, hex-, hept-, oct-, non-, dec-
o suffix: -ane
 State:
 1 – 4 carbon atoms: gas
 5 – 8 carbon atoms: liquid
 9 – 17 carbon atoms: thick liquid
 18+: solid
 Alkenes: CnH2n
 Isomers: same molecular formula but different atom arrangements; similar chemical properties and different
physical properties.
 Alkynes: CnH2n-2
 Benzene: C6H6
o Structure: Stable, flat hexagonal rings; Identical bonds,
which are intermediate in length; C atoms have double
bonds to either of its neighboring C atom.

Some benzene molecules can be carcinogen.

 Aromatic: Benzene based compounds but one or

more H atoms are replaced with other halides/
halogens or alkyl groups.
Ex. 1-bromo- 3- cloro- 5- methyl benzene
 Cyclic Hydrocarbons: alkanes, alkenes or alkynes that are arranged in rings; have 2 less hydrogen
atoms than non – clyclic forms; prefix: cyclo-

 IUPAC nomenclature for hydrocarbons:

1. Stem name is equivalent to the longest chain that contains the double/ triple bond.
2. Principal functional groups have to obtain the lowest number.
Priorities: double/ triple bond, halogens (F, Cl, Br, I), alkyl groups (methyl, ethyl, …)
3. Prefixes determined by the principal functional group.
4. Use 1 word to name the compound:
Name of each group is started with a number indicating its position.
Alphabetical order is applied when listing the groups.
Di, tri, tetra,… do not affect the order.
5. Hyphen seperates numbers and words.
Comma seperates numbers.
6. Number attached group from #2.
Rections of Hydrocarbons
 Addition Reactions:
o Occur with alkene or alkynes (unsaturated reactants).
o Faster than substitution.
o Reagents: H2(g), Cl2(aq) – clorine water, Br2(aq) – bromine water, HCl, HBr, HI.
o Observations: Yellow Cl2 water quickly turns colourless/ Red-brown Br2 water quickly turns
colourless.

 Substitution Reactions:
o Occur with saturated hydrocarbons (alkanes, benzene and aromatics).
o Reagents: clorine water, bromine water, catalyzed by UV, Pt, Ni and/or heat.
o Slower than addition, which is why it is used to distinguish between saturated and unsaturated
compounds.

 Combustion Reactions:
o Products: CO2(g) & H2O(g) & energy (complete combustion)/ CO(g)/ C(s) soot + H2O(g) & less energy
(incomplete combustion).
Ex. Complete combustion: 2C4H10(g) + 13O2(g) 8CO2(g) +10H2O(g) + 5754 kJ
Incomplete combustion: 2C4H10(g) + 9O2(g) 8CO(g) +10H2O(g) + 3490 kJ

Rates of Reaction
 Rate of reaction: rate at which reactants are used up or the rate at which products are formed;
measured in mol s-1, gs-1 or mL s-1.
 Collision Theory: conditions for reactions to occur:
1. Individual particles of the reacting substances must collide.
2. The collision energy must be equal to or greater than activation energy, Ea.
3. The reacting particles must collide with a suitable orientation.
 Increased chance of collisions, increased rate of reaction.

 Factors of ROR:
1) Nature of reactants: ionic reactions are faster than molecular ones because they have less bonds to be
broken and formed as the ions are held together by electrostatic forces of attraction.
2) Conc. of reactants: proportional relationship
3) State of subdivision of reactants: more surface area, faster ROR.
4) Temperature:
  Higher kinetic energy, faster velocity and therefore more collisions.
 More particles have sufficient activation energy, so more reactions can occur.
5) Gas pressure: decreased volume or adding more gas particles increase the pressure, only applied to gas
as liquid and solid usually can not be compressed.
6) Catalyst: not consumed in a reaction but provide an alternative easier pathway by lowering the E a.

 Activation Energy: energy that colliding particles must have to form an activated complex.
 Transition state: very unstable point of the reaction where bonds are breaking and forming.
 Energy Profile Diagram:

Transition State with

Activated Complex

 Examples of catalysts:
 Enzymes: not consumed, highly specific, lock &key model, specific temperature and pH.
 Transition metals and their compounds: Pt, Pd, Au, MnO2 & Rh.
 Nanoparticles: at least one dimension is 1 – 100nm, have a large surface area.
 Catalytic converters: reduce toxic gases using platinum, Rhodium, palladium with high SA,
increased ROR of CO & unburnt fuel with oxygen, so only carbon doixide and water are
produced; convert NO to N2 and O2.

VSPER Theory & Polarity

 Valence Shell Electron Pair Repulsion theory states that:
1) Pairs of outer shell electrons in atoms form charged clouds which are roughly spherical in shape.
2) These charged clouds repel each other and so are positioned as far apart as possible. This includes both
bonding and non-bonding pairs of electrons.
3) However, lone pairs have larger repulsive forces.

Shape Linear Triangular Planar Tetrahedral Pyramidal Bent (V-

(trigonal planar) shaped)
Bond Angle 180° 120° 109° 107° 104.5°
No. bond 2 3 4 4 4
pairs
No. lone - - - 1 2
pairs
 Shape
diagram

 Electronegativity:
 Attraction of an atom for shared electrons
 Increases from left to right
 Highest: flourine, lowest: lithium.
 Non-polar covalent bonds: only between non-metals, EN dif.= 0 to 0.4
 Polar covalent bonds: only between non-metals, EN dif.= 0.5 to 1.7
 Ionic bonds: between metals and non-metals, electrons are transferred, and
EN dif.= 1.8+

 Polar dipoles: exist in polar bonds/ molecules, negative end attracts more electrons.

Intermolecular Forces
 Dispersion Forces:
o Weakest intermolecular forces
o Exist between all molecules
o Temporary instantaneous dipoles
o Increases with increased molecular mass and in linear shaped molecules.
 Dipole-dipole Forces:
o Weak but permanent
o Between polar molecules only

 Hydrogen Bonding Forces:

o Strongest
o Between poalr molecules
o Between the hydrogen atom of a O-H, N-H or F-H bond to the N, O or F atom of another
molecule, which has a lone pair that attracts hyfrogen atoms.

Generally, network covalent > metallic > ionic > hydrogen bonding forces > dipole-dipole forces > disperson

Use intermolecular forces to compare & explain physical properties of covalent molecules
 M.P & B.P: increase with increased strength of intermolecular forces
  Vapour Pressure: tendency of a substance to evaporate, measured in kPa, Pa or atm; increases with
decreased intermolecular forces; when vapour pressure = atmospheric pressure => liquid boils.
 Solubility: when a new solution is formed, the new solute – solvent interactions are equal or stronger
in strength, so covalent network is insoluble due to the strong intramolecular forces; “like dissolves
like”.

Chromatography: paper chromatography, TLC, GC & HPLC

 Mechanism:
o Stationary phase: solid or liquid supported in a solid (paper)
o Mobile phase: liquid or a gas (water)
o If components move quickly = high retention factor => they are strongly attracted to the mobile
phase
o If components move slowly = low retention factor => they have strong interactions onto the
stationary phase.
distance of component
Retention Factor Rf = both ¿ the pen cil line
distance of solvent
 Paper chromatography:
o Stationary phase : polar paper (dipole dipole forces and water)
o Mobile phase: polar water (all intermolecular forces)
 Thin Layer Chromatography:
o Stationary phase: thin layer of silica gel SiO2, cellulose, starch or alumina Al2O3 coated on a
sheet of metal, plastic or glass.
o Mobile phase: non – polar liquid solvent or water
o Advantages: simple, rapid, inexpensive; small amount of test substance; non – destructive, only
involves physical seperation.
 Gas Chromatography:
o Sample is vaporised into gas molecules
o Stationary phase (coiled to increase SA): solid – gas chromatography (solid absorbant)
liquid – gas chromatography (liquid on an inert
support)
o Mobile phase (carrier gas): inert non-polar gas (He or N2)
o Polar and/or large molecules stay longer
o Non – polar volatile components elute (pass through) faster.
 High Performance Liquid Chromatography
o Mobile phase is pushed through stationary phase under high pressure.
o Stationary phase: solid particles of silica or polymers tightly packed.
o Mobile phase: liquid solvent.
o Normal phase HPLC: polar S.P, non-polar M.P; vice versa.
o Advantages:
 Can analyse compounds that decompose in gas chrom.
 Faster due to high pressure
 Small particles of adsorbent material on stationary phase create a large SA.
Properties of Gases and Kinetic Theory of Gases

Kinetic Theory Assumptions:

1) Gases consist of tiny particles moving in rapid, random, straight-line motion until they collide with
one another or with the container (Brownian motion).
2) Collisions between particles or witht the walls are perfectly elastic.
3) The size of the particles are negligible compared to the size of their container
 The particles have mass but no volume
 Distance between particles are larger than their size
4) Any attractive/repulsive forces between particles are negligible
5) Average kinetic energy increases as temperature increases

Therefore, gases can be compressed and diffuse.

Real Gases Ideal Gases

Have volume No volume
Particles attract and repel one another No forces exist between the particles
At temperature and pressure close to phase changes,
gase behavious is affected by intermolecular forces
Stronger intermolecular forces, less molar volume Molar volume = 22.71 L/mol

Properties of gases

 Pressure (P): force exerted by gas against the walls of the container, increases as temp and velocity
increase.
 Units: atm, mm Hg, torr, pascal Force
Pressure=
1 atm = 760 mm Hg = 760 torr = 100 kPa Area

 Volume (V): space occupied by gas, measured in L or mL.

 Amount (n): quantities of gas particles, in grams or moles.
 Temperature: average kinetic energy of all molecules
1 2
E k= m v
2

Gas Laws

1 P1 V 1=P2 V 2
 Boyle’s Law - P ∝
V

V1 V2
=
 Charles Laws – V ∝ T T 1 T2
 Gay – Lussac’s Law – P ∝T
P 1 P2
=
T1 T 2

 Avogrado’s Hypothesis: at the same temp and pressure, equal volumes of gases contain the same
amount of particles
V1 V2
=
n1 n2

 Molar Volume: volume of 1 mole, or 6.22 × 1023 particles of gas.

 Standard molar volume: volume of 1 mole of ideal gas at STP, or at 0°C and 100kPa = 22.71 L/mol.
V
n= at STP
22.71

 Ideal Gas Equation: when conditions are not at STP

Pressure in kPa n in moles
PV =nRT
Volume in L Temperature in Kelvin

 Combined Gas Law

P1V 1 P2V 2
=
T1 T2