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CHEM101 - Lecture notes All

General Chemistry I (Drexel University)

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CHAPTER 1:
Heterogenous mixture- visual distinction- chicken soup
Homogenous- air, coffee
Compound - hydrogen peroxide, ice
Elements cant be broken down
Properties of Matter
Intensive properties - independent of the amount of substance that is present
Density, boiling point, color
Extensive properties- dependent upon the amount of the substance present
Mass, volume, energy
SAME b4 and after chemical reaction:
-law of conservation of matter-
Sum of masses of all substances involved
# of atoms of each type involved
Law of definite proportions-compound always occur in fixed proportions
Law of multiple proportions= 1 g first element

Nuclear Atom
Nucleus
Contains protons and neutrons
Almost ALL atomic mass
Almost NONE volume
(+) charge
Electrons
Located outside nucleus
Almost NONE atomic mass
Almost NONE volume (electron cloud)
(-) charge
Atoms are neutral: if charge --> ion

Isotopes & Atomic Weight

Isotopes- atoms of same element with different mass

Mass#-identify isotope
Atomic#-identify element

CHAPTER 2:
1 mol = 6.022 "things"
Wave Nature of Light
Electromagnetic waves-movement of electrical charges
Movement produces fluctuations in electric/magnetic fields
Electromagnetic radiation

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Wavelength (λ) frequency (v) amplitude

C=λv, c=3 * m/s
Electromagnetic Spectrum
UV, X rays shorter wavelength, higher frequency radiation
Communications(Radio, microwave) longer wavelength, lower frequency
Visible light only tiny portion of spectrum
Quantum Theory
Planck: atoms absorb/emit electromagnetic energy only in discrete amounts
Smallest amount of energy, ____, given by
E= hv
H = 6.626*10^-34 Js
Photoelectric Effect:
Einstein consider electromagnetic energy to be bundled into little packets called Photon
Continuous Spectrum- white light- lights of all wavelength-rainbow
Line Spectrum-light from electrical discharge through gaseous element- bright lines at specific
wavelength

Bohr Model
Excitation- atom absorbs energy = difference btw energy levels
● Electron in higher e level
Emission - atom gives off energy as photon
● Electron drops to lower e level

E=0 when electron located infinitely far from nucleus

ΔE = Ef - Ei =
Nf > ni, E absorb
Nf < ni E releaese

Ground states and Excited states

Ground- electrons in lowest possible energy levels
Excited- electrons promoted to any level n>1
Promoted by absorbing energy (electric discharge, heat, lasers)
Eventually drop back to ground- relaxation

CHAPTER 3:
Orbitals
Region of probability of finding electron around nucleus
4 types (s, p, d, f)

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Max 2 electrons/orbital
Periodic table
Group- elements w similar properties
Period- atom list in increasing atomic number
Electron configuration- distribution of electrons among atomic orbital
1s^2
1: principal shell=energy level=n=1
S: orbital subshell
2: # of e in subshell
Orbital diagram-boxes represent orbitals within subshells, arrow represents electrons
Hand's Rule
1 electron in each orbital subshell before doubling up
Minimize repulsion
Degenerate- orbitals with same energy
For multi-electron atoms, energies of sublevels split
Charge interaction (coulomb's law)
Electrons attracted to nucleus
Electrons repelled by other electrons
Shielding
Electrons closer to nucleus reduce attraction of outer electrons
(effective nuclear charge ) --> highest closest to nucleus

Penetration
More electron penetration=stronger interaction
Lower sublevel = more electron penetration
E(s)< E(p) < E(d) < E(f)

Valence Shell
Outermost occupied principal shell
Controls reactivity
Electrons in valence shell= valence
Electrons in inner shell= core
Noble gas notation [Ar]4s^2 3d^6

Anions: gain electrons to complete val shell

Cations: lose electrons to attain complete val shell
Metals form cations
Less E to lose few electrons than add many e
Nonmetals form anions
Less E to lose few electrons than lose many e

Effective nuclear charge (

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# protons - # core electrons

Stays same going down a group

Atomic Radii (trend)

Increases down a column = bigger effect on radius
Val shell farther from nucleus
doesn’t change
Decreases left to right across row
Electrons added to same val shell
increases
Val shell held closer
Radius of atoms and ions
Cations smaller than corresponding atom Li > Li+
Remaining electrons feel higher effective nuclear charge
Anions larger than corresponding atom O < O2-
Extra electrons feel lower effective nuclear charge

Ionization Energy (trend)

Energy necessary to remove electron to form positive ion
Low value for metals
Electrons easy to remove
High value for non-metals
Electrons difficult to remove
First ionization energy
Energy to remove 1 electron from atom
Second ionization energy
Energy to remove 2nd electron from atom
Trend:
Harder to remove electrons from positive ion
Smaller radius/higher

CHAPTER 4:
Chemical bonds
Hold atoms together in molecules
Keep ions in place in solid ionic compounds
Ionic bond (metals and nonmetals)
Attractive force associated with oppositely charged ions due to transfer of electrons
Covalent bond (2+ nonmetals)
Attractive force associated with shared electron pairs between atoms

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Naming ions
Positive ions
Metal ion w only one charge state
= metal name + "ion"

Metal ion w multiple charge states

= metal name + (Roman numeral) to show charge

Negative ions
Monatomic ion
=add "-ide" to name stem + "ion"

Polyatomic ion
Familiar w most common ones
Formula Name
Carbonate ion

Sulfate ion

Phosphate

Ammonium

Ionic compounds
Held together by electrostatic forces
Electrically neutral
Form electrolytes when dissolved in water
Ions Compound Charges
(2+) + 2(1-) = 0

(2+) + (2-) = 0

Electrolyte=form ions in water

Percent composition
Ozone depletion effectiveness related to amount of chlorine/molecule
Chlorine-element that reacts w ozone
Higher percent chlorine=more damaging for same mass

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CHAPTER 5
Lewis Structures
Dots=valence electrons
Added clockwise-one on each side then pair
Show val e and bonding
1. Determine total val e
2. Draw structure, connecting atoms by single bond
3. Place lone pairs of e around terminal atoms to give each an octet
4. Assign any remaining e = lone pairs around central atom
5. If there are not enough e,
move lone pairs of e from terminal atom to form a multiple bond to central atom
Lewis Theory
Val electrons play fundamental role in chemical bonding
Atoms form chemical bonds to acquire electron configurations of noble gases
Bonding electron pair
Shared pair of elections in molecule
Lone pair electrons
Electron pairs that are not shared
Single bond
Covalent bond-one pair e shared
Multiple (covalent) bonds can also form
Double bond - 2 pairs e shared
Triple bond - 3 pairs e shared
EACH ATOM OBEYS OCTET RULE
Every atoms wants to hold 8 val e in outermost electron shell
Expectations to Octet Rule
Free radicals
Odd # of val e
Incomplete octets

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Compounds of Be, B, Al (very reactive)

Expanded valence shells
Central atom more than 8 electrons
Orbitals (d,f) - 3 period lower
Formal charge
Charge on atom if all shared electrons were equally shared
Most plausible lewis structure is one w no formal charge
Negative formal charges=most electronegative atoms
Charges on adjacent atoms should not have same sign

Sum of formal charges = actual charge

Resonance Structures
Molecule or ion represented by two or more plausible lewis structures
Structures differ only in electron distribution
Delocalized electrons
Spread out over several atoms
Part of resonance hybrid
True structure is single composite = resonance hybrid
Electronegativity (EN)
Ability of atom to attract its bonding electrons
Higher EN of atom = stronger attraction of electrons in bond
(trend)

Covalent Covalent Ionic

Nonpolar 0.0-0.4 Polar 0.4-2.0 2.0-3.3

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Symmetric Unsymmetric Metal & Nonmetal

Lone pair cancel out Lone pair don't cancel out
Can contain polar bond Don’t contain nonpolar bond
Molecular Polarity
Dipole moment=measure of polarity
Nonpolar - even distribution of color (warm)
Polar - cool to warm color
Ionic=largest dipole=most polar
Difference Molecule Moment (D)
Electronegativity 0
1.78
0.79

Geometry 1.85
0

Composition 0
1.92
1.04

Bond length
Distance between nuclei of 2 atoms joined by covalent bond
Function of parameter=atomic radius
Triple < Double < Single
Smaller radius = shorter bond
Bond order
Single vs double vs triple
Higher bond order = shorter bond length
Bond Strength
Single < Double < Triple
Stronger = shorter bond/smaller radius > longer

VSEPR (valence-shell electron-pair repulsion)

Method to determine geometry
Basis: pairs of val e in bonded atoms repel one another

Electron group
Val e in region around central atom

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Unshared pair of val e

Bond(single, double, triple)

Electron-group geometry
Orientation of e groups based on e repulsion
Based on # of e groups

Molecular geometry
VSEPR theory
Draw lewis e dot structure
Count # of bonding e groups around central atom
Double, triple bonds count as 1 e group
Count # of lone pairs of e
Match e group info to shapes

E groups E group geo Molecular geo

180˚ 2 Linear Linear

120˚ 3 Trigonal planar 0 Trigonal planar

1 Angular/Bent

109.5˚ 4 Tetrahedral 0 Tetrahedral

1 Trigonal pyramidal
2 Angular/Bent

120˚, 90˚ 5 Trigonal bipyramidal 0 Trigonal bipyramidal

1 Seesaw
2 T-shaped
3 Linear
90˚ 6 Octahedral 0 Octahedral
1 Square pyramidal
2 Square planar

CHAPTER 7
Chemical reactions
Chemical change- breaking and forming chemical bonds
Electrolysis

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Physical change- conversion between physical states/phases

Stoichiometric coefficients
Balance equation-change coefficient not subscripts/atoms
Mole (stoichiometric) ratios
Limiting reactant
Controls amount of product produced
Gets used up first
Smallest amount of product = limiting
Excess
Theoretical yield = amount of product predicted by stoichiometry
Actual yield = quantity of desired product actually formed

CHAPTER 8

Molarity (M) =

Mole Fraction (x) =

If solution diluted:
Volume increases
Moles stay the same
Concentration decreases

CHAPTER 9
Energy- capacity to do work
KE- due to motion
Thermal E- temperature
PE - due to position/composition
Chemical E- positions of electrons and nuclei

First Law of Thermodynamics

"energy cannot be created or destroyed"
Law of conservation of energy

Internal energy change of system is difference between its final and initial states
Energy transfer
Surroundings gain exact mount of energy lost by the system

Heat(q)

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Energy transfer --> thermal differences btw system and surroundings

"flows" automatically from higher T to lower T
"flows" end at thermal equilibrium
Properties related to temperature change
Type/characteristics of material
Amount of energy (heat q)
Amount of material (mass m)
Heat capacity of system
Quantity of heat required to change the temperature of system by 1˚C
Molar heat capacity
Heat capacity of 1 mol of a substance
Specific heat (c ) / Specific heat capacity
Heat capacity of 1 g of pure substance/homogenous mixture

Internal energy of a system is always increased by:

Adding heat to system
Have surroundings do work on system

Types of systems:
Open Closed Isolated

Coffee cup calorimeter

Determine heat of rxns in solution
-heat lost = heat gained

Bomb calorimeter
Determine heat of combustion
heat lost = heat gained

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Volume compression
Energy added by work physically
Energy released by temp decrease

Enthalpy (
Change in heat/energy of system

If heat added to system

(+) endothermic

If heat released to surroundings

(-) exothermic

Cyclic process

Extensive property
Depends on how much of substance is present

Energy and enthalpy=state functions

Property w unique value
Depends only on present state of system
Independent of path
Independent of history of system

CHAPTER 10
Properties of gases
Translucent
Constant, random motion
Compressible

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Exert pressure on whatever surrounds them

Expand into whatever volume is available
Easily diffuse into one another
Higher energy than solids/liquids
Kinetic-molecular theory
Molecules much smaller than intermolecular distance
Molecules in continuous, random, rapid motion
No fixed volume/shape
Collisions between molecules are elastic
Small intermolecular forces
Pressure= collision of gas molecules with wall of container
Temperature = related to average speed of gas molecules

Gas pressure
Pressure-force exerted per unit area by gas molecules as they collide w surfaces around them
Factors affect pressure
Number of gas particles
Volume of container
Average speed of gas particles
Measure pressure with barometer
Units:
1 atm = 760 torr
1 atm = 760 mm Hg
Molecular speeds
Higher molar mass, lower most-probable speed
Temperature increases, speed increases

Boyle's (const temp and mass)

Charles (const pressure and mass)

Avogadro (const pressure and temp)

Ideal Gas Law

P = atm

V=L
n= mol

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T=K
0˚C + 273 = 273 K

Dalton's law of partial pressure

Partial pressure=pressure a gas would exert if it were alone in container

Total pressure = sum of partial pressure of each gas

Mole fraction (X)

Ratio of moles of component to total moles

X value ranges from 0-1

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