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Significant Figures
Name
methane
ethane
propane
butane
pentane
hexane
heptane
octane
nonane
decane
….ene
…yne
Ruler A – 4.40 cm
Ruler B – 4.4 cm
Add 1 more decimal than the smallest measurement can
measure.
Type of 0
placeholder
placeholder
trapped
precise
0.00003
3,0000
3,0003
3.0000
Significant?
no
no
yes
yes
Formula
CH4
C2H6
C3H8
C4H10
C5H12
C6H14
C7H16
C8H18
C9H20
C10H22
CnH2n
CnH2n-2
Covalent prefixes
Number
Prefix
Number
Prefix
Adding sig figs – do not count the # of sig figs! Count the #
of decimal places.
1
mono-
6
hexa-
2
di-
7
hepta-
% error
3
tri-
8
octa-
4
tetra-
9
nona-
5
penta-
10
deca-
= I valueacc - valueexpI
valueacc
x 100%
106 m = 1 m
109 nm = 1 m
1012pm = 1 m
micro

nano
n
pico
p
Substance
Liq. Water
Gas at STP
Density
1.00 g/mL
mass of 1 mol/22.4L
Average atomic mass = % mass + % mass +…..
(isotope 1) (isotope 2)
x
x
1 mole = 6.022 x 1023 particles. (Avogadro's number)
1 12C atom = 12.00 amu (atomic mass)
1 mole 12C atom = 12.00 g (molar mass)
MnO4- permanganate
S2O3-2 thiosulfate
CrO4-2 chromate
Cr2O7-2 dichromate
O2-2
peroxide
ClO-
hypochlorite
ClO2ClO3ClO4C2O4-2
chlorite
chlorate
perchlorate
oxalate
Fe(CN)6-4 hexacyanoferrate(II)/ ferrocyanide
A = mass number
Z = atomic number
A = number of protons + neutrons
Z = number of protons
A – Z = # of neutrons
Isotopes= atoms with same number of protons but
different numbers of neutrons
Ions= electrons vary
Isomer – different arrangements if atoms within a
molecules
PROTONS, Atomic # define the element!
Fe(CN)6-3 hexacyanoferrate(III)/ ferricyanide
BO3-3
borate
Hg2+2
mercury (I) ion
Molecular formula = the way a molecule really appears in
nature
Empirical formula = smallest whole number ratio of
elements
ex:
= whole number always!
=
Glucose has a molecular formula (MF) of C6H12O6
Ba(NO3)2  Ba(NO2)2 + O2
8.
Glucose’s empirical formula (EF) is CH2O
Glucose’s molar mass (MM) is 180 g/mol
Glucose’s empirical mass (EM) is 30 g/mol
Diatomic molecules - BrNClHOF
Other Molecular Elements
o Phosphorus--one form P4
o
Sulfur—often S8
o
Carbon--4 allotropes, diamond and graphite,
buckyballs, graphene, all covalent networks
Other compounds to memorize
1. Hydrogen Peroxide – H2O2
2. Ammonia – NH3
REACTIONS TO MEMORIZE:
1.
Synthesis reactions of metal oxides with water: .
ex:
Na2O + H2O  2NaOH
CaO +H2O  Ca(OH)2
2.
Synthesis reactions of nonmetal oxides with water:
ex:
SO3 + H2O  H2SO4
N2O5 + H2O  2HNO3
3.
4.
Decomposition of Binary Compounds
ex:
2H2O  2H2 + O2
MgS  Mg + S
H2O2  H2O + O2
Decomposition of Metal Carbonates:
ex:
CaCO3  CaO + CO2
Na2CO3  Na2O + CO2
5.
Decomposition of Metal Hydroxides:
ex:
Ca(OH)2  CaO + H2O
2NaOH  Na2O + H2O
6.
Decomposition of Metal Chlorates:
ex:
2KClO3  2KCl + 3O2
Mg(ClO3)2  MgCl2 + 3O2
7.
Decomposition of Metal Nitrates:
Decomposition of acids: certain acids can decompose to
form nonmetal oxides and water.
ex:
H2CO3  CO2 + H2O
H2SO3  SO2 + H2O
9.
EF X 6 = MF; EM X 6 = MM for glucose!
Solving EF problems:
1. If given percentage, drop % and use g. (assume 100 g)
2. if given mass, just leave it!
3. Convert g of each element to moles to get mole ratio.
4. Divide by smallest to get lowest whole number ratio.
5. if dividing by smallest does not give whole number ratio, then
multiply all by a whole number:
.2,.4,.6,.8 x 5
.25, .75 x 4
.33, .67 x 3
2NaNO3  2NaNO2 + O2
10.
11.
Combustion: elements in a compound separate and
combine with oxygen to make most reasonable cmpds.
ex:
ABC + O2  AO + BO + CO2
C2H5OH + 3 O2  2 CO2 + 3 H2O
CS2 + 3 O2  CO2 + 2SO2
Single Replacement: A metal replaces a metal or a
nonmetal replaces a nonmetal IF the element doing the
replacing is higher on the activity series. Hydrogen
replaces/is replaced by metals, though.
Ex:
MgCl2 + 2Na  2 NaCl + Mg
Cl2 + 2 NaI  2 NaCl + I2
2 HCl + Mg  MgCl2 + H2
Double replacement: Only works if 1 of following is
produced: 1. Precipitate 2. Gas 3. Water
If the following is made
NH4OH
H2SO3
H2CO3
It decomposes into…
H2O + NH3
H2O + SO2
H2O + CO2
All compounds containing group 1 metals, acetates, nitrates, or
ammonium are soluble.
BaSO4, PbI2, and AgCl are common insoluble substances!
There is no such thing as an oxide ion in solution! It will quickly
become a hydroxide ion!
Preparing Solutions:
1. Always use volumetric flask.
2. Add solute or concentrated solution.
3. Fill to the line with deionized water.
4. Shake.
McVc = MdVd
Mc = molarity of concentrated solution (used to make new soln)
Vc = molarity of concentrated solution (used to make new soln)
Md = concentration of solution produced
Vd = volume of solution produced = vol. of concentrated soln +
volume of water!
Leo says Ger!!!!
Losing
Electrons is
Oxidation
Gaining
Electrons is
Reduction!
EX: 2 AgNO3 + Zn  Zn(NO3)2 + 2 Ag
2 Ag+ + Zn  Zn+2 + 2 Ag
Silver is reduced; zinc is oxidized
2 moles of electrons are transferred (not 4!)
Half reactions:
2 Ag+ + 2 e-  2 Ag
Zn  Zn+2 + 2e-
Amount of energy to change temp of 1oC of H2O = 1 cal
4.184 J = 1 cal
1000 cal = 1 kcal = 1 Cal
Half reactions are often easy!
If not single replacement, use….
All Oranges Have Citrus Energy!
200 g water at 70oC
1. Separate into half reactions. Leave out spectators.
2. A = balance atoms (except H and O)
3. O = balance oxygens
(add H2O for acid rxns; add 2 x OH- for bases)
H = balance hydrogens (add H+ for acid rxns; add H2O for bases)
5. C = balance charge (by adding electrons)
6. E = balance electrons (by multiplying half reactions)
4.
Assigning Oxidation Numbers:
1.
2.
In free elements, each atom has an oxidation number of 0.
For ions consisting of a single atom, the oxidation number is
equal to the charge on the ion.
3. For binary ionic compounds, assign oxidation numbers
(charges) as you always have!
For covalent molecules, or polyatomic ions, assign oxidation
numbers in the order shown below.
F-O-H, closest to F, farthest from F
4.
5.
6.
7.
Fluorine is always assigned first an oxidation number of –1
when in a compound.
Next, assign oxygen. The oxidation number of oxygen is -2 in
most compounds.
Exceptions: Oxygen in peroxides is -1 Ex. H2O2 Na2O2
Then, assign hydrogen. The oxidation number of hydrogen is
usually +1, but may be –1 when combined with a metal.
For example: H in hydrides is -1 Ex. NaH
For compounds in which both atoms cannot have the
oxidation number which is equal to the charge the element
commonly has, one closest to fluorine “wins”.
HEAT
Total amount of
kinetic energy of a
sample
A measure of
average kinetic
energy of a sample
q
Joules
T
o
SPECIFIC HEAT
The amount of
energy required to
change the temp of
1 g of substance by
1oC
C
J/goC
cal/goC
J/moloC
HEAT
CAPACITY
The amount of
energy required to
change the temp
of a given
substance by 1oC
none
J/oC
TEMPERATURE
1 mL H2O = 1 g H2O
Enthalpy = heat
C, K
cal/oC
100 g water at 70oC
Same T
Same KEavg of particles
Same speed of particles
Same specific heat of water
More heat energy (q)
Less heat energy (q)
Heat of formation
∆Hf
Heat of fusion
∆Hfus
Heat of
vaporization
∆Hvap
Heat of reaction
∆Hrxn
Heat of
combustion
∆Hcomb
Ionization energy
IE
Lattice Energy
U
Standard entropy
∆S
Gibbs Free Energy
of formation
∆Gf
Gibbs free energy
of reaction
∆Grxn
Energy required to form an
element from its standard
state
Energy required to melt a
substance or RELEASED when
frozen
Energy required to boil a
substance or RELEASED when
condensed
Energy required/released
during a reaction (per mole)
Energy required/released
during a comb. reaction (per
mole)
Energy required to remove 1
electron from an atom
Energy required to break 1
mole of ions in a crystal
lattice into gaseous ions
Measure of disorder
compared to that of a solid
crystal at 0K
Measure of spontaneity;
ability to do work on
surroundings as compared to
elements in standard state
Measure of spontaneity;
ability to do work on
surroundings
Ways to Calculate the Enthalpy of reaction, ∆Hrxn:
1. ∆Hf (products) - ∆Hf (reactants)
2. bonds broken – bonds made
3. sum of ∆Hrxn for reactions that add up to new reaction
4. measure amount of energy gained/lost by water. Divide
by moles of 1 reactant reacted.
Person
Dalton
Experiment
-
Mendeleev
Thomson
Cathode rays
Millikan
Rutherford
Einstein
Planck
Bohr
Oil drop experiment
Gold foil experiment
Photoelectric Effect
Line emission spectrum of
(excited) hydrogen gas
De Broglie
Schrodinger Heisenberg
-
Discovered/Proposed
Atomic Theory – 5 postulates, incorrect about all atoms of a given element
identical and atoms are indivisible
Designed Periodic Table; left holes for 3 missing elements and predicted their
properties
All matter contains electrons (and therefore protons); determined mass:charge
of electron; Proposed plum pudding model of atom (p+/e- spread throughout)
Determined exact mass and charge of electron
Discovered positively charged, very dense nucleus
photons
Calculated “size” of photon – Planck’s constant
Existence of energy levels/ quantized energy states of electrons
De Broglie equation; wavelength of any moving object
Schrodinger’s equation; calculates probability of finding electron in a given
region (orbital!) within an atom by treating electron as probability wave function
Uncertainty Principle; the greater the precision in measuring a small object’s
location, the greater the uncertainty in measuring its velocity and vice versa
(can’t know an electron’s location and velocity simultaneously)
Excited electrons have gained energy and jumped to a
higher energy level. They possess more energy.
They fall back down to a lower energy state and must
release energy in the form of 1 quantum/photon, E = hν
Only certain sized photons (lines of
frequency/wavelength) are emitted so each element
has its own distinct line emission spectrum. This is due
to the existence of quantized energy states in atoms.
Pauli Exclusion Principle
Aufbau Rule
Hund’s Rule (Bus Rule!)
Sublevel
No 2 electrons in the same atom can have the same 4 quantum numbers; thus no 2
electrons can be in the same energy level/sublevel/orbital AND have same spin; 2
electrons in same orbital must have opposite spins
Electrons fill up orbitals from lowest energy to highest energy (this may not be in
numerical order! See aufbau box below) 2 is higher than 1; d is higher than p, etc….
If 2 equal energy orbitals are available, electrons each go to separate orbitals (with
same spin) before pairing up 2p: ____ _____ _____
orbitals within
Shape of orbitals
Picture
s
s
Spherical
# of
electrons in
an orbital
2
p
px, py, pz
Dumbbell shaped
2
6
d
dxy, dxz, dyz, dz2, dx2 –y2
4 lobed
2
10
f
g
7 different f’s
9 different g’s
8 lobed/too complex
!!!!!
2
2
14
18
Don’t even try
AAAAH!
# of electrons in
sublevel
2
Quantum numbers describe where a given electron is in
an atom.
1st quantum #
Energy level
1, 2, 3, 4….
nd
2 quantum #
sublevel
s, p, d, f
3rd quantum #
Orbital/orientation px, py, pz,for example
4th quantum #
spin

Aufbau box: (add arrows)
Which one has higher
Coulombic attraction?
Why?
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
5g
6s
6p
6d
6f
6g
6h
7s
7p
7d
7f
7g
7h
8s
8p
8d…….
7i
Example of orbital notation:
C: ____ ____ _____ _____ ______
1s
2s 2p
Examples of electron configuration:
Mg: (you do it!)
K: 1s22s22p63s23p64s1
Mn: 1s22s22p63s23p64s23d5
Ag: 1s22s22p63s23p64s23d104p65s14d10
Pb: (you do it!)
# unpaired electrons = #electrons alone in orbital
C has 2 unpaired electrons
K has 1, Ag has 1, Pb has 2, Mn has 5, Mg has 0
Unpaired electrons make element paramagnetic.
Drawing Lewis Structures for COVALENT compounds:
1. The atom with lowest EN is the central atom.
2. Count the total number of valence electrons.
MAGIC NUMBER! (adjust for ions)
3. Place one sigma bond (a pair of electrons)
between each pair of bonded atoms.
4. Subtract the total number of valence electrons
used for bonds from the MAGIC NUMBER.
5. Place lone pairs about each terminal atom
except for hydrogen atoms. Subtract the
number of lone pairs from the sum.
6.
7.
If e- are still left at this point, assign them to the
central atom. If the central atom is from the
third or a higher period, it can accommodate
more than four pairs of electrons.
If the central atom is not yet surrounded by
four electron pairs, convert one or more
terminal atom lone pairs to pi bond pairs.
IONIC
Metal + nonmetal
Metal +polyatomic ion
Poly ion + nonmetal
Poly ion + poly ion
Strongest type of bond
Electrostatic attraction
between oppositely
charged ions
Electrons transferred,
creating ions
Forms crystal lattice
And thus
crystals
Formula units
Higher MP, Higher BP
(must break bond to
change phase)
Lattice Energy
Higher charge, smaller
size= stronger bond
Large EN difference
If soluble, forms
electrolytes when
dissolved
Electrolytic when
molten, not solid
High % ionic character
COVALENT
Nonmetal + nonmetal
Within poly ion
Weaker bond
Electrostatic attraction
between nuclei and
shared electrons
Electrons shared – NO
IONS!
A few form covalent
networks (SiO2, , most
form individual molecules
molecules
Lower MP, lower BP
(only break IMFS NOT
bonds to change phase)
Bond energy
Higher bond order, smaller
atom = stronger bond
Forms polar and nonpolar
bonds
Medium (polar) or Small
(nonpolar) EN difference
Nonelectrolytes, may
dissolve but won’t ionize!
Nonelectrolyte
Low % ionic character
Show resonance!
Total
number of
sigma
bonds and
lone pairs
on central
atom
Shape
Atoms
bonded to
central
atom
(# of
sigma
bonds)
Lone
pairs
on
central
atom
2
Linear
2
0
sp
yes
3
Triangular
Planar
3
0
sp2
yes
3
Bent
2
1
sp2
no
4
Tetrahedral
4
0
sp3
yes
4
Triangular
pyramidal
3
1
sp3
no
4
Bent
2
2
sp3
no
5
Trigonal
bipyramidal
5
0
dsp3
or
sp3d
yes
5
Unsymmetrical
tetrahedron
4
1
dsp3
sp3d
no
Molecular Shape
Picture
Hybridization
Symmetrical in
3d space?
And bond angles
(see saw)
5
T-shaped
3
2
dsp3
sp3d
no
5
Linear
2
3
dsp3
sp3d
yes
6
Octahedral
6
0
d2sp3
sp3d2
yes
6
Square
pyramidal
5
1
d2sp3
sp3d2
no
6
Square planar
4
2
d2sp3
sp3d2
yes
Example of a
Lewis structure
that is this
shape. (Yes,
draw the Lewis
structure!)
TYPE
USES
TO DETERMINE/Produce
HOW
Mass Spectroscopy
Magnets
Photoelectron
Spectroscopy
Beam of Light
on Metal
Distinguishes isotopes (separates
by mass), can be used to determine
% abundance of each isotope
Binding Energy of electrons in
different sublevels within a given
atom
Atomic absorption
spectroscopy or Atomic
Emission Spectroscopy
Electrical or
Heat Energy
added to excite
electrons
The magnet bends heavier
isotopes more than lighter
ones
“Light” beam (could be other
form of ER) gives energy to
electrons and removes them
from the atom- this energy is
measured…takes more
energy to remove innermost
electrons than outermost
electrons
Energy excites electrons and
they absorb a specific
wavelength, then when they
fall back down, they release a
specific wavelength
Nucleus vibrational states
altered by radio waves, gives
“signature” of element
Line Emission spectrum (or line
absorption spectrum) which
identifies the element
NMR Spectroscopy
(nuclear magnetic
resonance)
Radio waves
Identity of a substance (types of
(LOW Energy
atoms present)
right?) added to
“excite” nucleus
Infrared or Visible or UV
IR or Visible
Identity of a substance/molecule,
Electrons in bonds also have
spectroscopy (example:
light or UV light (types of bonds present) OR
quantized energies and these
Using spectroscopes)
to excite
concentration of substance (when
energies can be used to
electrons in
we used spectrophotometer)
identify a molecule
bonds
Mass Spectrometry Data: …each bar represents the mass and abundance of a given isotope. What element? _______
PES data: (photoelectron spectroscopy)…each bar represents the abundance AND energy to remove each of the
electrons from an atom (binding energy!). The outermost electrons take the least amount of energy to remove.
1s
3s
2p
2s
When bond is formed, 2 atoms rest at a distance of
minimum potential energy. This gives the typical
bond length of the molecule.
Boiling point – temp at which v.p equals atmospheric
pressure.
Normal b.p. –temp at which v.p equals 1 atm.
---------------------------------------------------------∆E = q + w
∆E = internal energy
q = heat added to the system
w = work done on the system
So, for example if heat is added to the system AND work
is done on the system, then the internal energy of the
system increases and has a positive sign.
If a reaction is exothermic and it expands in volume,
work is done on the surroundings and so work is
negative, and heat is negative, and the internal energy
decreases and change in energy has a negative sign.
Top of hump = activation complex/transition state/highenergy intermediate
---------------------------------------------------------
Taller the “hump”, slower the reaction
Activation energy is used to break bonds, create an
“effective” collision
Effective collision:
1. enough energy to react/break bonds
2. correct orientation
A 2 step mechanism might result in double hump.
During phase change, added or removed heat doesn’t
change kinetic energy of particles and thus, does not
change the temp. Energy is used to break IMFs or
released when formed.
Slow step only determines rate of reaction. (slow step
has higher hump)
Activation energy
with catalyst
How is rate law determined experimentally?
1. Method of Initial Rates: Start with different
concentrations of reactants and measure initial
rate.
a. If tripling A causes rate to increase 9fold, exponent of A is 2.
2. Measure the concentration over time. Make a
graph to see what order.
1 mol gas at STP = 22.4 L
2 ways a gas may not be ideal:
1. Large size
2. Polar
5 factors that affect rate of reaction:
1. nature of reactants
2. concentration: increases # of collisions
3. phase/ surface area: increases # of collisions
4. temperature: increases # of collisions AND # of
effective collisions
5. catalyst: Increases # of effective collisions by
lowering activation energy
Consider this reaction: A+ B   D+ E
Rate = k [A]x[B]y[C]z NOT to be confused with
equilibrium expression:
An unideal gas:
1. Tends to have inelastic collisions, and/or
2. Have a size that DOES matter
3. AND doesn’t obey the gas laws perfectly.
Any gas is more real (unideal) at
1. High pressure (size of particles does matter)
2. Low temperature (particles more likely to have
inelastic collisions)
Any gas is more ideal at
1. Low pressure
2. High temperature
2 balloons each filled with different gases at the same
volume in the same room.
K=
1. Exponents (in rate law) are NOT determined by
balanced equation.
2. Exponents can match slow step of mechanism.
3. Exponents can only be determined
experimentally.
4. k depends on T.
5. A catalyst can be included (C perhaps)
6. Products are (almost) never included – no
denominator (in most cases)
Rate law
Rate depends on [A]
Integrated Rate law
Plot needed to give a
straight line
Relationship of rate
constant to slope of
straight line
Half life
Half life depends on [A]
He
CO2
Same V, T, and P
Same KE.
Same number of gaseous particles
Higher mass, higher density
Lower mass, lower density
Lower speed of particles
Higher speed of particles
Zero Order
Rate = k
NO!
[A] = -kt + [A]o
1st order
Rate = k[A]
ln [A] = -kt + ln [A]o
[A] vs. t
ln [A] vs. t
2nd order
Rate = k[A]2
Yes
1 = kt + 1
[A]
[A]o
1/[A] vs. t
Slope = -k
Slope = -k
Slope = k
t1/2 = [A]o/2k
yes
t1/2 = 0.693/k
NO!
t1/2 = 1/k[A]o
yes
Yes
What are IMFs?
1. NOT bonds
2. Attractions between neighboring species
3. MUCH Weaker than bonds usually
Increasing m.p., b.p , Hfus, Hvap,
critical point, specific heat
Increasing vapor pressure,
volatility, likelihood of being gas
at room temp
The 5 types of intermolecular forces (IMFS) listed from
strongest to weakest are:
1. Ion-ion (NOT really an IMF)
Example: Salt bonds
2. Ion-dipole
Example: Salt dissolving in water
3. Dipole-dipole (Strongest type is hydrogen
bonding)
Example: Water condensing
4. Dipole-induced dipole
Example: CO2 dissolving in water
5. Induced dipole-induced dipole
AKA London dispersion forces
AKA Van der Waals forces
Example: freezing nitrogen
Ways to increase solubility of a
Solid
Increase temp
Increase surface area
Agitation, stirring
gas
Decrease temp
Increase pressure above
Don’t agitate
ΔHsoln (heat of solution): SUM OF
-
Energy required to break the water-water IMFs
Energy required to break IMFS in solute
Energy released when IMFS form between
solute and solvent.
Hydrogen bond: the ATTRACTION between a H of one
molecule (that is bonded to a very EN element such as
F, O, N) and a very EN element (O, F, N) of another
molecule. NOT A BOND!!!!!!!!!
Ways to measure the concentration of a solution:
The stronger the IMFS, the lower v.p. at given T.
The stronger the IMFs, the higher the normal b.p.
Types of crystalline solids:
Ionic –formed from ionic bonds, very strong, brittle, high MP,
think NaCl
Molecular – formed from the regular arrangement of IMFS
between covalent molecules, low MP, not strong, think
snowflakes
Network – covalent bonds throughout, high MP, hard,
nonconducting (usually), think SiO2 (sand/glass), C (diamond)
Graphite is a unique network solid made only of carbon
atoms. They are bonded together in sheets. The sheets are
attracted to each other by weak IMFS. (rubs off easily) Within
the sheets, there is resonance throughout, allowing graphite
(a nonmetal) to conduct electricity!
Molarity (M) = mol solute/L solution
Molality (m) = mol solute/kg solvent
Mole fraction (X) = mol solute/ mol solution
% by mass = g solute/g solution
% by volume = mL solute/mL solution
m/v % = g solute/mL solution
Colligative property: any property of a solution that
changes based SOLELY on the NUMBER of solute
particles dissolved. (not mass or size or IMF formed!)
-freezing point depression
-Vapor pressure lowering
-Boiling point elevation (as a result of v.p.
lowering!)
WHY? Solute particles interfere with the process of
evaporating, boiling, and freezing.
Alloys are metal solutions. (homogeneous mixtures)
Substitutional alloy: formed by 2 metals of similar radii;
mp or ability to resist corrosion or conductivity might
be affected. Hardness is NOT.
Strong acid:
Weak acid:
Strong base:
Weak base:
Soluble ionic compound:
Insoluble ionic compound:
Soluble covalent compound:
Insoluble covalent compound:
Interstitial alloy: formed by 2 metals of different radii;
adding in smaller metals tend to make the alloy harder.
For both alloys, adding metals with more valence
electrons increases conductivity.
DO NOT MIX CONCEPTS! Periodic trends do NOT explain boiling point for example.
WHY …?
Does X have
a larger radius than Y?
a higher ionization energy than Y?
a higher electronegativity than Y?
a higher reactivity than Y?
Does Y make a given shape? Have a given bond angle?
ONLY ReasonsForce
you will
of use
attraction
Tell how many protons and electrons
each species has
Then, use
proton-electron attraction (Coulomb’s LAW)
OR electron-electron repulsion (Coulomb’s LAW)
OR shielding (electrons don’t feel protons pulling on them
due to intervening energy levels)
VSEPR THEORY!!! Bonding pairs of electrons experience
repulsion with each other so they spread out in 3d space
as far as possible; Lone pair electrons induce greater
repulsion and cause angles between bonding electrons to
be smaller than expected.
Is Y polar?
Does Y
dissolve in Z?
Have a high freezing pt, low volatility, low vp, high
critical pt, high melting pt, high Hfus, high Hvap?
Does SOLUTION A
Have a high b.p., high conductivity, low f.p., low v.p?
Electronegativity difference (polar or nonpolar bonds?)
AND Symmetrical or Unsymmetrical shape
Identify intermolecular force involved
Compare strengths of intermolecular force
Compare [ ] of solutions
Ksp problem types: always AB  A+ + B1. Calculate solubility using Ksp: Do RICE solve for x
2. Calculate Ksp using solubility: Do RICE, solve for K
3. Calculate solubility within a solution of common ion: Do
RICE, solve for X (should not have all 0’s on I row)
4. Calculate if precipitate forms when mix 2 solutions:
- MV = MV to determine new concentration
- Plug into equil expression to get Q
- Compare to K
5. Calculate how much of a given substance is required to
form precipitate, (or compare which precip forms 1st)
- Solve equil expression for missing value!
Le Chatelier’s principle:
If a reaction is stressed, the reaction responds in such a way as
to remove that stress.
Equilibrium can be disturbed by:
1. Adding products (shifts left)
2. Adding reactants (shifts right)
3. Removing reactants (shifts left)
4. Removing products (shifts right)
5. Adding volume or decreasing partial pressure (shifts
towards side with fewer gaseous or aqueous particles)
6. Increasing temperature (shifts to get to new K)
a. Shifts right if endothermic
b. Shifts left if exothermic
Buffer: combination of (weak) conj. acid/base pair that resists
change in pH
A good buffer:
1. desired pH = pKa of acid
2. [conj. acid] = [conj. base]
3. [conj.acid] and [conj. base] high (increases buffering
capacity)
How to approach ANY titration calculation:
1. Molarity x volume = moles for both substances.
Subtract to see what remains. Does a conjugate
acid or base get produced? (it will in any weak
titration.) Determine its moles.
2. Divide by volume to get molarity for substance
remaining.
3. THINK….how do I get pH of this/these
substances?
a. –log [strong]
b. If weak, RICE
c. If weak acid and weak base are present,
do H-Hasselbach.
Equilibrium canNOT be disturbed by:
1. Adding inert gas (no effect)
2. Adding solid (no effect)
The ONLY way to change K is to change the temperature
K >1 products preferred at equilibrium
At equivalence point of ANY titration: MaVa = MbVb,
moles acid – moles base = 0
Midpoint = half moles neutralized (AND [weak
acid/base] = [conj. acid/base]); pH = pKa
Titrations
Starting point
Midpoint
Equivalence point
Strong acid/strong
base
high
Not relevant
Exactly 7
Vertical region
Very long
Strong acid/weak
base
not as high
pH= pKa
Below 7
(conjugate acid
present)
Not as long
Strong base/
Strong acid
low
Not relevant
Exactly 7
Very long
Always pick an indicator whose pKa = pH at the equivalence point of the titration.
Strong base/weak
acid
not as low
pH = pKa
Above 7
(conjugate base
present)
Not as long
Various Ways to Describe Acid Strength
Property
Strong Acid
Weak Acid
Ka Value
Too large to measure
<< 1
Position of the dissociation
equilibrium
Equilibrium concentration of
H+ compared to the original
concentration of HA
Equilibrium concentration of
HA compared to the original
concentration of HA
Percent Dissociation
right
left
[H+ ] = [HA]o
[H+ ] << [HA]o
[HA] ~ 0 M
[HA] << [HA]o
[HA] ~ [HA]o
100%
< 5%
Strength of conjugate base
Conjugate base is nonexistent
How to represent in a net
ionic equation
H+ + A-
Conjugate base is also weak, but
inversely proportional in strength to
the acid (Kb = Kw/Ka).
HA
Sign
Of  H0
1. Exothermic
2. Exothermic
3. Endothermic
4. Endothermic
+
+
System
Entropy
Change
Increasing
Decreasing
Increasing
decreasing
Galvanic Cell/
Voltaic Cell/
Electrochemical Cell
Spontaneous
Electrolytic Cell
Nonspontaneous
Ex: electrolysis
water
Sign
of  S0
Sign of
 G0
Sign of E0
(if a redox)
K
+
+
-
-
+
>1
Spontaneous?
(Thermodynamically
favorable?)
always
+ or -
- or +
Don’t know
Low T
+ or -
- or +
Don’t know
High T
+
-
<1
never
Chemical energy  electrical energy
E=+
ΔG = -
anode=
oxidation
Cathode =
reduction
Electrical energy  chemical energy
E=-
ΔG = +
Anode=
oxidation
Cathode=
reduction
Ex: battery!
As reaction progresses, K doesn’t
change BUT
When rxn is shifting
right
When rxn is shifting left
When a reaction is at
equilibrium
Q approaches K
Q<K
Q>K
Q=K
G approaches 0
G=-
G=+
G=0
E approaches 0
E = +
E=-
E=0
No ions
formed
usually
+1
+1/-1
Increasing nuclear charge explains everything: increasing ability to attract
electrons, decreasing AR, increasing EN, increasing IE, which both explain
+2
-4
-3
-2
-1
Most
reactive
nonmetal,
highest
EN
S block
decreasing reactivity across metals, increasing reactivity among nonmetals
+3
Alkali ne earth metals
Alkali metals
p block
d block (1 behind!)
5770
Zn
+2
Ag Cd
+1 +2
(+2)
+4
Sn
Pb
Transition metals
*
(+3)
+5
+6
numeral
Variable positive charge = roman numeral
89102
+7
**
Increasing # of energy levels
(increased shielding) explains everything:
Decreasing ability to attract electrons,
Increasing AR, Decreasing IE, Decreasing EN
Which explain increasing reactivity among
metals, decreasing reactivity among
nonmetals
Sb
roman numeral
Roman
Most
reactive
metal,
lowest IE,
lowest EN
*
Lanthanide series, 4f
Variable positive charge = roman numeral
Rare earth metals
**
Actinide series, 5f
Highest IE
No measurable EN (too low), high IE, INERT!
H
5770
*
89102
**
*
**
Page 16: Look on the internet for an interesting molecule that has many atoms. Draw its Lewis structure below and 1. Give its molar mass, 2. Give its mass per
molecule. Then pick 2 central atoms that have different hybridizations and label their 3. Hybridization, 4. Bond angles, 5. Shape about that atom
A good website to try might be http://www.chm.bris.ac.uk/motm/motm.htm (Molecule of the Month).
Which atom?
1.
2.
3.
4.
5.
4.
5.
Which atom?
3.