Review Chapters 6 -10 :
General, Organic, & Biological
Chemistry
Janice Gorzynski Smith
Chapter 6 & 7 Concepts
 Energy
 conversions, conservation of energy
 Breaking bonds requires E, forming bonds releases E
 Endothermic & Exothermic Reactions
 Energy diagrams, Activation Energy, heat absorbed or released
 Factors affecting rates of reactions
 Concentration, temperature, catalysts
 Equilibrium
 Equilibrium constant expressions
 Le Chatlier Principle
 States of matter: g, l, s & their properties
 Effect of intermolecular forces on behavior
 Gas laws: combined, ideal, & dalton’s law partial pressure
 Intermolecular forces
 London-Dispersion, Dipole-Dipole, Hydrogen-Bonding
 Relative strength, importance in g, l, s behavior
 Phase Changes
 Navigate a heating/cooling curve
 Enthalpy of phase changes
2
Equations & Conversions
P1V1
=
T1
P2V2
K =
T2
[C]c [D]d
[A]a [B]b
1 cal = 4.184 J
PV = nRT
L • atm
R = 0.0821 mol • K
L • mm Hg
R = 62.4
mol • K
Ptotal = PA
[products]
[reactants] =
+
PB
1,000 cal = 1 kcal
1,000 J = 1 kJ
+
PC
1 kcal = 4.184 kJ
Equations to memorize in orange
Energy of Reactions
ENDOTHERMIC
Transition State
E
Energy
required to
break
bonds
Energy
released
as bonds
form
Ea
Ea
ENDOTHERMIC
ΔH
Reactants
Heat + A + B  C + D
Products have weaker bonds and a higher
energy then Reactants.
Heat is absorbed by the system.
ΔE +
ΔH +
Heat absorbed
ΔH
Heat released
Products
EXOTHERMIC
EXOTHERMIC
A + B  C + D + heat
Products have stronger bonds and a lower
energy then Reactants.
Heat is released by the system.
ΔE -
ΔH -
Rates of Reactions
Increase the Rate of a Reaction
Increase Temperature
Increase Average KE of particles, so more
likely to collide with enough energy to
overcome Ea
Increase Concentration Reactants
Increase the number of collisions per second
Add a Catalyst
Decrease Ea
Greater likelyhood
that particles will
have enough KE to
react
Same likelyhood rxn
will happen when
particles collide, but
more collisions
Equilibrium & Le Chatlier’s Principle
aA + bB
equilibrium
constant = K =
cC + dD
[products]
[reactants] =
[C]c [D]d
[A]a [B]b
K > 1 products favored K < 1 reactants favored K = 1 equilibrium
A+B
C + D + heat
A + B + heat
C+D
product
reactant
Eq Shift
reactant
Eq Shift
increase

increase

decrease

decrease

product

increase

increase

decrease

decrease

T increase

T increase

T decrease

T decrease
Intermolecular Forces
London Dispersion Forces
Weakest
Dipole-Dipole Forces
Hydrogen Bonds
Ion-Dipole Forces
Strongest
Forces experienced by states of matter
Gas < Liquids < Solids
Increasing Average Kinetic Energy
Physical Properties
Property of s, l, g
Increases
Decreases
Example
Water has a high boiling point because it has H-bonding,
dipole, and dispersion forces. It is close to heptane
(C7H16), a heavier molecule that only experiences
dispersion forces .
The melting point of ionic solids is extremely high
compared to water which experiences all other
intermolecular forces, but not ion-dipole forces. (NaCl is
1074 K and water is 273 K)
Boiling Point
increasing total
intermolecular forces
decreasing total
intermolecular forces
Melting Point
increasing total
intermolecular forces
decreasing total
intermolecular forces
Retention of V &
Shape
Decreasing
Increasing intermolecular intermolecular forces,
forces and decreasing T & and increasing kinetic
P
energy of particles or T &
P
Gases will fill the volume and shape of the container that
holds them, while solids will retain their own shape and
volume regardless of the container.
Surface Tension
with increasing
intermolecular forces
The molecules on the surface have less neighbors (and
therefore less stabilizing intermolecular forces) and so have a
higher potential energy, which the material will try to reduce
with its shape (sphere): water beading.
Viscosity
Vapor Pressure
with decreasing
intermolecular forces
increasing intermolecular decreasing intermolecular Not just a property of liquids, also gases and solids.
Amorphous solids change shape over time because of their
forces and decreasing
forces and increasing
viscosity.
temperature
temperature
Decreasing intermolecular Increasing intermolecular Ether has weaker intermolecular forces than water and a
higher vapor pressure, so it evaporates much faster then
forces and increasing
forces and decreasing
water.
temperature
temperature
Gas Behavior
Non Rigid Container:
Piston
balloon
P1V1
T1
=
P2V2
P constant
V increase w/ T or
# of moles
T2
PV = nRT
Ptotal = PA + PB + PC
Rigid Container:
Closed Flask
V constant
P increase w/ T or
# of moles
Phase Changes
SOLID
fusion
LIQUID
freezing
evaporation
GAS
condensation
deposition
sublimation
endothermic System absorbs energy from surrounds in the form of heat
o Requires the addition of heat
exothermic System releases energy into surrounds in the form of heat or lig
o Requires heat to be decreased
Phase Changes
gas
TEMPERATURE
l <--> g
liquid
evaporation
or vaporization
ΔHvap
endothermic
s <--> l
solid
fusion
ΔHfus
endothermic
HEAT ADDED
Chapter 8 & 9 Concepts
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Identify the solvent and solute in a solution
Like dissolves like, predict which molecules will form solutions
Predict the effect of temperature or pressure on a solution
Perform concentration calculations & conversions
Perform dilution calculations
Predict relative changes in colligative properties between multiple
solutions
Understand osmotic pressure & how your kidney’s work.
Identify an acid/base reaction, the acid, base, conjugate acid/base
Caculate Ka, Kb
Use Kw to determine concentration of H3O+ or OHDiscuss how water acts as both an acid and a base
Perform titration calculations
Communicate how a buffer prevents large pH changes
12
CH 8 Equations & Conversions
Molarity = moles of solute (mol)
V of solution (L)
M1V1
=
M2V2
CH 9 Equations & Conversions
[H3O+][ A −]
[HA]
Ka =
-
Kb
=
Kw
=
[OH ][BH+]
[ B]
[H3O+][OH−] = 1.0 x 10−14
pH = -log[H3O+]
•Acidic solution:
pH < 7  [H3O+] > 1 x 10−7
•Neutral solution:
pH = 7  [H3O+] = 1 x 10−7
•Basic solution:
pH > 7  [H3O+] < 1 x 10−7
Solutions, Solubility, & Concentration
1. The solute is the substance present in a lesser amount.
2. The solvent is the substance present in a larger amount.
Solubility is the amount of solute that dissolves in a given
amount of solvent. REMEMBER: LIKE DISSOLVES LIKE.
 In aqueous or liquid phase solutions solubility increases with
increasing temperature
 Gases dissolved in liquids increase solubility with decreasing
temperature and increasing pressure
Communicate how much of a solute is dissolved in a solvent using
concentration:
 % w/v
Dilution: Adding more solvent to the initial solution.
 % v/v
The number of moles solute DOES NOT CHANGE.
 % mass / mass
M1V1
=
M2V2
 ppm
initial values
final values
 Molarity
Colligative Properties
Colligative properties are properties of a solution that depend on the
concentration of the solute but not its identity.
 One mole of any nonvolatile solute raises the boiling point
of 1 kg of H2O the same amount, 0.51 oC.
 One mole of any nonvolatile solute lowers the freezing point
of 1 kg of H2O by the same amount,1.86 oC.
Reverse Osmosis
Apply pressure to
reverse osmosis.
This is how our
kidneys filter blood
Acids / Bases
•A Brønsted–Lowry acid is a proton (H+) donor.
Strong:
•A Brønsted–Lowry base is a proton (H+) acceptor.
Weak:
gain of H+
H A
acid
+
A − +
H B+
conjugate conjugate
acid
base
B
base
loss of H+
−
H
O
Conjugate
base
Kw
remove H+
=
H
[H3O+][OH−]
O
H
H2O as a base
H2O as an acid
+
H
add H+
H
O
H
conjugate acid
Acid / Base Equilibrium & pH
H3O+(aq) + A
HA(g) + H2O(l)
acid dissociation
constant
Ka =
[H3O+][ A −]
[HA]
pH = -log[H3O+]
-
=
Low pH (0 ~ 7)
[H3O+] high
Acidic Conditions
OH- (aq) + BH+ (aq)
B (g) + H2O(l)
Base dissociation K
b
constant
− (aq)
[OH ][BH+]
[ B]
High pH (7 ~ 14)
[H3O+] low
Basic Conditions
Common Acid / Base Reactions
Neutralization reaction: An acid-base reaction that produces a salt and water.
H+(aq) + OH− (aq)
H—OH(l)
A bicarbonate base, HCO3−, reacts with one H+ to form carbonic acid, H2CO3.
H+(aq) + HCO3−(aq)
H2CO3(aq)
H2O(l) + CO2(g)
A carbonate base, CO32–, reacts with two H+ to form carbonic acid, H2CO3.
2 H+(aq) + CO32–(aq)
H2CO3(aq)
H2O(l) + CO2(g)
Titration
AH + B  A- + BH+
Acid + Base  Conjugate Base + Conjugate Acid
mole–mole
conversion
factor
M (mol/L)
conversion
factor
Moles of
base
[1]
Volume of
base
[2]
Moles of
acid
[3]
Volume of
acid
M (mol/L)
conversion
factor
Buffers
pH of buffer = -log[H3O+] where
[H3
O +]
=
Ka
x
[HA]
[ A −]
Chapter 10 Concepts
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Interpret Atomic number and mass number
Know radioactive particles: alpha, beta, positron, gamma
Write & solve radioactive decay equations
Determine the number of half lives that pass in a given
amount of time.
 Familiar with measurements of the amount of radioactivity
 Familiar with measurements of radiation absorbed
 Understand how radioisotopes are used in medicine
22
Atomic Symbols & Nuclear Particles
atomic number (Z) =
alpha particle: a
the number of protons
or
mass number (A) =
the number of protons
+
the number of neutrons
beta particle:
β or
mass number (A)
positron: β+
or
atomic number (Z)
12
6
number of protons 6
number of neutrons 12 – 6 = 6
gamma ray:
0
e
−1
0
e
+1
C
g
4
He
2
Nuclear Equations & Half Life
original
nucleus
radiation
emitted
new
nucleus
=
4
He
2
0
e
−1
+
0
e
+1
radiation
emitted
g
The half-life (t1/2) of a radioactive isotope is the time it
takes for one-half of the sample to decay.
Radioactivity
amount of radioactivity
1 Ci = 3.7 x 1010 Bq.
radiation absorbed
The rad—radiation
absorbed dose
The rem—radiation
equivalent for man
Radioisotopes can be
injected or ingested to
determine if an organ is
functioning properly or to
detect the presence of a
tumor.