Claire’s Semester 2, Honors Chemistry 1, Study Guide_

Claire’s Semester 2, Honors Chemistry 1, Study Guide:

Hi! I made this study guide for the final. I got all the info from the slideshows in moodle, my chem binder, and previous practice tests. I hope it's helpful! Feel free to share it with others!

Moles:

Mole - 6.02 x 10^23 representative particles of the substance and is the SI unit for measuring the amount of a substance
Avagadro’s number - 6.02x x 10^23
Molar mass - the mass of a mole of a specific element
Avagadro’s hypothesis - states that equal volumes of gasses at the same temperature and pressure contain equal numbers of particles
Percent composition - the relative amounts of the elements in a compound
Empirical formula - the lowest whole-number ratio of the atoms or molecules in a compound
Molecular formula - experimentally determined formula; a simple whole-number multiple of its empirical formula

7 elements exist as diatomic molecules: Br2, I2, N2, Cl2, H2,O2, F2
22.4 L = 1 mol at STP
D = M/V

Moles Sample Problems:

1) Calculate the percent composition of C3H8

mass of C in 1 mol C3H8 = 3 mol x 12.01 g/mol = 36.03 g
mass of H in 1 mol C3H8 = 8 mol x 1.01 g/mol = 8.08 g
molar mass of C3H8 = 36.05 g/mol + 8.08 g/mol = 44.13 g/mol

% C = mass of C in 1 mol C3H8/molar mass of
% C = 36.03 g/44.13 g x 100%
% C = 81.6%

% H = 100% - % C
% H = 18.4%

2) A compound is analyzed and found to contain 25.9% nitrogen and 74.1% oxygen. What is the empirical formula of the compound?

percent by mass of N = 25.9%, percent by mass of O = 74.1%

25.9 g N x 1mol N/14.01 g N = 1.85 mol N
74.1 g O x 1 mol O/16.00 g O = 4.63 mol O

Mole ratio of N to O = N1.85O4.63

1.85 mol N/1.85 = 1 mol N
4.63 mol O/1.85 = 2.50 mol O

Mole ratio of N to O = N1O2.50
Empirical formula = N2O5

3) How many moles is 5.69 g of NaOH?

5.69 g NaOH
I mol NaOH

40.00 g NaOH

0.142 moles of NaOH

4) What is the volume of 8.8 g of CH4 gas at STP?

8.8 g CH4
1 mol CH4
22.4 L CH4

16.05 g CH4
1 mol CH4

12 L CH4 *(2 sig figs)

5) What is the molar mass of a gas with a density of 1.964 g/L?

1.964 g
22.4 L
1 L
I mol

43.99 g/mol

6) Find the density of CO2 at STP

44.01 g CO2
1 mol CO2
1 mol CO2
22.4 L CO2

1.96 g/L

Chemical Reactions:

Reactants - the substances you start with
Products - the substances you end up with
Catalyst - a substance that speeds up a reaction, without being changed or used up by the reaction
Enzymes - biological or protein catalysts in your body
Precipitate - an insoluble salt formed in a reaction

Rules for balancing chemical equations:

  • Assemble correct chemical formulas for the reactants and products
  • Count the number of each type of atom appearing on each side of the arrow
  • Balance the elements one at a time by adding coefficients (the numbers in front)
  • Save balancing H and O for last!
  • Double-check to make sure everything is balanced
  • Never change a subscript to balance an equation
  • Never put a coefficient in the middle of a formula

Reading chemical equations:

  • “+” = “and”
  • (s) = solid
  • (l) = liquid
  • (g) = gas
  • (aq) = dissolved in water
  • (>>>) = “produces”

Sentence: copper reacts with chlorine to form copper (II) chloride
Word equation: copper + chlorine >>> copper (II) chlorine
Equation: Cu + Cl2 >>> CuCl2 chlorine is a diatomic molecules

5 major types of reactions:

  • Combination reactions (A + B >>> AB)
  • Some nonmetal oxides react with water to form a acid
  • Some metallic oxides react with water to form a base
    Ex. Ca + O2 >>> CaO
  • Decomposition reactions (AB >>> A + B)
  • Usually requires energy (heat, sunlight, electricity, etc.)
    Ex. NaCl >>> Na + Cl2
  • Single replacement reactions (A + BX >>> B + AX)
  • Reactants must be an element and a compound
  • Activity series can be used to determine if the reaction take place
  • Elements higher on the activity series replace those lower
    Ex. Na + KCl >>> K + NaCl
  • Double replacement reactions (AX + BY >>> AY + BX)
  • Reactants must be 2 ionic compounds in aqueous solutions
  • One product must be a precipitate, water, or a gas
    Ex. NaOH +FeCl3 >>> Fe(OH)3 + NaCl
  • Combustion reactions
  • If the combustion reaction is complete, the products will be CO2 and H2O
  • If the combustion reaction is incomplete, the product will be CO
    Ex. C4H10 + O2 >>> CO2 + H2O

Net Ionic Equations:

  • Fully balanced equation: AgNO3 + NaCl >>> AgCl + NaNO3
  • Ionic equation: Ag(1+) + NO3(1-) + Na(1+) + C(1-)l >>> AgCl + Na(1+) + NO3(1-)
    ***AgCl didn’t ionize because it is a precipitate
  • Net ionic equation: Ag(1+) + Cl(1-) >>> AgCl

Chemical Reactions Sample Problems:

1) Balance the following chemical equations

AgNO3 + Cu >>> Cu(NO3)2 + Ag
2AgNO3 + Cu >>> Cu(NO3)2 + 2Ag

P + O2 >>> P4O10
4P + 5O2 >>> P4O10

2) Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas

Fe2S3 (s) + HCl (g) >>> FeCl3 (s) + H2S (g)
Fe2S3 (s) + 6HCl (g) >>> 2FeCl3 (s) + 3H2S (g)

3) Complete the single replacement reactions

Fe + CuSO4 >>>
Cu + FeSO4
Pb + KCl >>>
K + PbCl2

4) Analyze the following: Copper (II) wire is placed in a aluminum nitrate solution

3Cu (s) + 2 Al(NO3)3 (aq) >>> 2Al (s) + 3Cu(NO3)2(aq)
no reaction will occur because aluminum is higher up on the activity series than copper

Stoichiometry:

Ways to interpret balanced chemical equations:

  • In terms of particles
  • Elements - made of atoms
  • Molecules - molecular compounds (made of only nonmetals)
  • Formula units - ionic compounds (metals and nonmetals)
  • In terms of moles
  • Coefficients tell us how many moles of each substance
  • A balanced equation is called a molar ratio
  • In terms of mass
  • Law of conservation of mass applies
  • In terms of volume
    mass and atoms are always conserved

How to calculate stoichiometry problems

  • Balance the equation
  • Convert mass in grams to moles
  • Set up mole ratios
  • Use mole ratio to calculate moles of desired chemical
  • Convert moles back into grams, if necessary

Limiting reagent - the reactant you run out of first (determines how much product is made)
Excess reagent - the one you have left over
Actual yield - what you get in the lab when the chemicals are mixed
Theoretical yield - what the balanced chemical equation says will be produced

Percent yield = actual/theoretical x 100
***tells us how “efficient” the reaction is

Stoichiometry Sample Problems:

1) 6.50 g of aluminum reacts with an excess of oxygen. How many grams of aluminum oxide are formed?

4Al + 3O2 >>> 2Al2O3

6.50 g Al
I mol Al
2 mol Al2O3
101.96 g Al2O3

26.98 g Al
4 mol Al
1 mol Al2O3

12.3 g Al2O3

2) How many liters of CH4 at STP are required to complete a reaction with 17.5 L of O2?

CH4 + 2O2 >>> CO2 + 2H2O

17.5 L O2
1 mol O2
1 mol CH4
22.4 L CH4

22.4 L O2
1 mol O2
1 mol CH4

8.75 L CH4

3) If 10.3 g of aluminum are reacted with 51.7 g of CuSO4, how much copper (grams) will be produced? Which reactant is limiting? How much of the excess reagent is remaining?

2Al + 3CuSO4 >>> 3Cu + Al2(SO4)3

10.3 g Al
1 mol Al
3 mol Cu
63.55 g Cu

26.98 g Al
2 mol Al
1 mol Cu

36.4 g Cu produced vs 20.6 g of Cu produced

51.7 g CuSO4
1 mol CuSO4
3 mol Cu
63.55 g Cu

159.61 g CuSO4
3 mol CuSO4
1 mol Cu

20.6 g of copper will be produced
CuSO4 is the limiting reagent

20.6 g Cu
1 mol Cu
2 mol Al
26.98 g Cu

63.55 g Cu
3 mol Cu
1 mol Al

10.3 - 5.83 = 4.5 g of aluminum remaining

States of Matter:

Kinetic energy - the energy an object has because of its motion
Potential energy - the stored energy in an object due to its position, properties, and forces acting on it
Kinetic theory - states that the tiny particles in all forms of matter are in constant motion
Gas pressure - defined as the force exerted by a gas per unit of surface area of an object
Vacuum - no particles present; no collisions and no pressure
Atmospheric pressure - the result of the collisions of air molecules with objects
Barometer - the measuring device used atmospheric pressure (depends on weather and altitude)
Pascal (Pa) - SI unit of pressure
STP - standard temperature and pressure; 0°C and 1 atm pressure (101.3 kPa)
Absolute zero - the temperature at which the motion of particles theoretically ceases (0 Kelvin)
Vaporization - the conversion from a liquid to a gas
Evaporation - when vaporization occurs at surface of a liquid that's not boiling
Vapor pressure - the pressure of a vapor in contact with its liquid or solid form
Condensation - the conversion of vapor/gas to a liquid
Dynamic equilibrium - rate of evaporation equals the rate of condensation
Boiling point - the temperature at which the vapor pressure of the liquid is just equal to the external pressure on the liquid
Normal boiling point - the boiling point of a liquid at a pressure of 101.3 kPa
Melting point - the temperature at which a solid turns to a liquid
Crystal-lattice - when the particles of a solid are arranged into a orderly, repeating, three-dimensional pattern
Allotropes - two or more different molecular forms of the same element in the same physical state
Sublimation - the change of a substance from a solid directly to a vapor, without passing through the liquid state
Effusion - gas escaping through a tiny hole in a container
Diffusion - the process of gasses mixing, the rate of diffusion is the rate of gas mixing

particles in any collection have a wide range of kinetic energy thus the term average kinetic energy is used (the higher the temperature, the wider the range of kinetic energies)
direct relationship between average kinetic energy and temperature
evaporation and boiling are cooling processes
once a liquid reaches its boiling point, the temperature doesn’t increase, the liquid just boils at a faster rate

The nature of liquids:

  • Liquid particles are in constant motion - often sliding past one another
  • Liquid particles are attracted to each other, unlike gas particles
  • Particles of a liquid spin and vibrate thus contributing to their average kinetic energy
  • Intermolecular forces decrease the amount of space between particles of liquids
  • Liquids are more dense than gasses
  • Increasing the pressure has hardly any effect on volume

The nature of solids:

  • Particles are relatively free to move
  • Solid particles tend to vibrate at fixed points
  • Particles are packed together in a highly organized manner
  • When a solid is heated, the particles vibrate more rapidly
  • Ionic solids have high melting points
  • Molecular compounds have relatively low melting points

3 basic assumptions of the kinetic theory as it applies to gasses:

  • Gas is composed of particles, usually molecules or atoms
  • Small, hard, spheres
  • Insignificant volume
  • No attraction or repulsion between particles
  • Particles in a gas move rapidly in constant random motion
  • Move in straight paths, only changing directions when they collide with one another or other objects
  • Collisions are perfectly elastic
  • Total kinetic energy remains constant

Ideal gasses don’t exist because…

  • Molecules DO take up space
  • There are attractive forces between particles
    Real and ideal gasses behave similarly at high temperatures and low pressures

4 variables that describe a gas:

  • Pressure (P) in kilopascals
  • Volume (V) in liters
  • As volume decreases, pressure increases
  • Temperature (T) in kelvin
  • As temperature increases, pressure increases
  • Amount (n) in moles
    Increasing the number of particles increase the pressure
    amount, volume, and temperature determine the pressure
    gas moves from areas of high pressure to area of low pressure

Boyle’s Law: P1V1 = P2V2 (T is constant)
Charles’s Law: V1/T1 = V2/T2 (P is constant)
Gay-Lussac’s Law: P1/T1 = P2/T2 (V is constant)
Combined Gas Law: P1V1/T1 = P2V2/T2
Ideal Gas Law: P x V = n x R x T
R = 8.31(L x kPa)/(mol x K)
Ideal Gas Law (#2): P x V = m x R x T / M
m: mass in grams, M: molar mass in g/mol
D = MP/RT
Dalton’s Law of Partial Pressure: Ptotal = P1 + P2 + P3 + …
Graham’s Law: RateA/RateB = √MassA/√MassB

Kelvin = °C + 273
Celsius = K - 273
temperature must be in KELVIN

States of Matter Sample Problems:

1) What happens when a substance is heated?
Particles absorb the energy; some of the energy is absorbed so potential energy increases and the remaining energy speeds up the particles so kinetic energy also increases - this increases the temperature.

2) Evaporation helps to keep our skin cooler on a hot day, why?
Liquids evaporate faster in higher temperatures because the added heat increases the average kinetic energy needed to overcome the attractive forces. Evaporation is a cooling process because as the liquid particles turn to water vapor, the particles left behind have a lower kinetic energy thus lowering the temperature.

3) A balloon contains 30.0 L of helium gas at 103 kPa. What is the volume of the helium when the balloon rises to an altitude where the pressure is only 25.0 kPa?

P1V1 = P2V2
V2 = P1V1/ P2
V2 = (103 kPa)(30.0 L)/(25.0 kPa)
V2 = 124 L

4) A balloon inflated in a room at 24°C had a volume of 4.00 L. The balloon is then heated to a temperature of 58°C. What is the new volume if the pressure remains constant?

T1 = 24°C + 273 = 297 K
T2 = 58°C + 273 = 331 K

V1/T1 = V2/T2
V2 = V1T2/ T1
V2 = (4.00 L)(331 K)/(297 K)
V2 = 4.46 L

5) If 2.50 moles of a gas occupy 5.60́ x 10^4 mL at STP, calculate the pressure of the gas if the amount of gas is halved and occupies a new volume of 13.5 L at -48ºC.

n1 = 2.50 moles
V1 = 5.60 x 10^4 mL
T1 = 273 K
P1 = 101.3 kPa
n2 = 1.25 moles
V2 = 13.5 L
T2 = 225 K
P2 = ???

P1V1/n1T1 = P2V2/n2T2
P2 = P1V1n2T2/V2T1n1
P2 = (101.3 kPa)(5.60 x 10^4 mL)(1.25 mol)(225 K)/(13.5 L)(273 K)(2.50 mol)
P2 = 173 kPa

Thermochemistry:

Thermochemistry - the study of energy changes that occur during chemical reactions and changes in state of matter
Chemical potential energy - energy stored in the chemical bonds of a substance
Heat - energy that transfers from one object to another because of a temperature difference between objects
Energy - capacity for doing work or supplying energy
Calorie - the quantity of heat needed to raise the temperature of 1 g of pure water 1ºC
System - the part of the universe as which you focus your attention
Surroundings - everything else in the universe
Law of conservation of energy - in chemical and physical processes, energy is neither created nor destroyed
Endothermic process - heat is absorbed from the surroundings
Exothermic process - heat is released into the surroundings
Specific heat - the amount of heat it takes to raise 1 g of the substance 1ºC
Calorimetry - the measurement of the heat flow into or out of a system for physical and chemical processes
Enthalpy (H) - accounts for the heat flow of a system at constant pressure
△H is negative for for an exothermic reaction and positive for a endothermic reaction
Heat of combustion - the heat of reaction for the complete burning of one mole of a substance
Molar heat of fusion(△Hfus) - the heat absorbed by one mole of a solid substance as it melts to a liquid at a constant temperature
Molar heat of solidification(△Hsolid) - the heat lost when one mole of a liquid substance solidifies at a constant temperature
Molar heat of vaporization(△Hvap) - the amount of heat required to vaporize one mole of a liquid at a constant temperature
Molar heat of condensation(△Hcond) - the amount of heat released when one mole of vapor condenses at its normal boiling point
Molar heat of solution(△Hsoln) - the enthalpy change caused by the dissolution of one mole of a substance
Hess’s law of heat summation - if you add two or more thermochemical equations to give a final equation, then you can also add the heats of reaction to give the final heat of reaction
Standard heat of formation(△Hfº) - the change of enthalpy that accompanies the formation of one mole of a compounds from its elements with all substances in their standard states
Heat of reaction - heat change for a equation
SC - standard conditions; 1 atm (101.3 kPa) and 25ºC

Specific heat: q/m x △T
Heat: m x c x △T
q: heat
q = mol x △Hfus
q = mol x △Hsolid
q = mole x △Hvap
q = mole x △Hcond
q = mole x △Hsoln
qsystem = △H = -qsurrpundings = -m x c x △T
△Hfus = -△Hsolid
△Hvap = -△Hcond
△fº = △Hfº (products) - △Hfº (reactants)

the standard heat of formation of a element in its standard state is “0”, this includes diatomic elements

Thermochemistry Sample Problems:

1) Find the specific heat given the following values

m = 95.4 g
△T = 23.0 ºC
q = 849 J

C = q/m x △T
C = (849 J)/(95.4 g)(23.0 ºC)
Specific heat = 0.387 J/(g x ºC)

2) The specific heat of ethanol is 2.4 J/(g x ºC). A sample of ethanol absorbs 676 J of heat, and the temperature rises from 22ºC to 64ºC. What is the mass of ethanol in a sample?

C = 2.4 J/(g x ºC)
△T = 64ºC - 22ºC = 42ºC
q = 676 J

C = q/m x △T
2.4 J/(g x ºC) = (676 J)/(m)(42 ºC)
6.7 g of ethanol

3) How many grams of ice at 0ºC will melt if 2.25 kJ of heat are added?

△Hfus = 6.01 kJ/mol
△H = 2.25 kJ

H2O (s) +6.01 kJ >>> H2O (l)

2.25 kJ
1 mol H2O (s)
18.0 g H2O (s)

6.01 kJ
1 mol H2O (s)

6.74 g H2O (s)

4) The temperature of 95.4 g piece of copper increases from 25.0ºC to 48.0ºC when the copper absorbs 849 J of heat. What is the specific heat of copper?

C = q/m x △T
C = (849 J)/(95.4 g)(48.0 - 25.0)
Specific heat of copper = 0.387 J/(g x ºC)

5) When 25.0 mL of water containing 0.025 mol HCl at 25.0ºC is added to a 25.0 mL cup of water containing 0.025 mol NaOH at 25.0ºC in a foam cup calorimeter, a reaction occurs. Calculate the enthalpy change (in kJ) during this reaction if the highest temperature observed is 32.0ºC. Assume the densities of the solutions are 1.00 g/mL.

Specific heat of water = 4.18 J/(g x ºC)
Mass = (50.0 mL) x (1.00 g/mL) = 50.0 g
△T = Tf - Ti = 32.0 - 25.0 = 7.0ºC

△H = -m x c x △T
△H = -(50.0 g)(4.18 J/(g x ºC)(7.0ºC)
△H = -1464 J
△H = -1.5 kJ

6) If 10.3 grams of CH4 are burned completely, how much heat will be produced? (See following equation)
CH4 (g) + 2O2 (g) >>> CO2 (g) + 2H2O (l) + 802.2 kJ

Note that this is a exothermic reaction because energy is a product, not a reactant

10.3 g CH4
1 mol CH4
802.2 kJ

16.05 g CH4
1 mol CH4

514 kJ of heat will released (△H = -514 kJ)

7) How many grams of ice at 0ºC will melt if 2.25 kJ of heat is added?

△Hfus = 6.01 kJ/mol
H2O (s) + 6.01 kJ >>> H2O (l)

2.25 kJ
1 mol ice
18.02 g ice

6.01 kJ
1 mol ice

6.74 g of ice will melt

8) How much heat (in kJ) is absorbed when 24.8 g H2O (l) at 100ºC and 101.3 kPa is converted to steam at 100ºC?

△Hvap = 40.7 kJ/mol
H2O (l) + 40.7 kJ/mol >>> H2O (g)

24.8 g H2O (l)
1 mol H2O (l)
40.7 kJ

18.02 g H2O (l)
1 mol H2O (l)

56.1 kJ of heat is absorbed

9) Calculate the standard heat of reaction
CH4 (g) + 202 (g) >>> CO2 (g) +2H2O (g)

△fº CH4 (g) = 74.86 kJ/mol
△fº 02 (g) = 0 kJ/mol (because it’s an element)
△fº CO2 (g) = -393.5 kJ/mol
△fº H2O (g) = -241.8 kJ/mol

△fº = △Hfº (products) - △Hfº (reactants)
△fº = [-393.5 + 2(-241.8)] - [74.86 + 2(0)]
△fº = -802.24 kJ (Exothermic reaction)

Solutions:

  • “Like dissolves like” (polarity)
  • The nature (polarity) of the solute determines whether it will dissolve and/or how much it will dissolve
  • Factors determining the rate of solutions include…
  • Stirring (agitation)
  • Surface area of the dissolving particles
  • Temperature
    higher temperature usually increases the amount that will dissolve
    as the temperature increases, gas solubility decreases
  • Solids dissolve best when heater, stirred, and crushed into smaller particles
  • Gasses dissolve best in low temperatures and high pressure
  • Solids will dissolve if the attractive force of the water molecules is stronger than the attractive force of the crystals

STP - standard temperature and pressure; 0°C and 1 atm pressure (101.3 kPa)
Solution - a homogeneous mixture that is mixed molecule by molecule
Solvent - the dissolving medium
Solute - the dissolving particles
Aqueous solution - a solution with water as the solvent
Saturated solution - contains the maximum amount of solute dissolved
Unsaturated solution - can still dissolve more solute
Supersaturated solution - a solution that is holding more solute than it theoretically can
Miscible - two liquids that can dissolve each other
Immiscible - two liquids that don’t mix or dissolve together
Concentrated solution - has a large amount of solute
Dilute solution - has a small amount of solute
Stock solution - pre-made solutions with known molarities
Colloid - a mixture that has particles ranging between 1 and 1000 nanometers in diameter, yet are still able to remain evenly distributed throughout the solution
Electrolytes - compounds that conduct an electric current in aqueous solutions
the strength of the electrolyte depends on the degree of ionization

Colligative properties of solutions…

  • Vapor pressure is lowered
  • Boiling point is elevated
  • Freezing point is lowered

Henry’s Law: S1/P1 = S2/P2
S - solubility, P - pressure
Dilution: MIVI = MFVF
M - molarity/concentration, V - volume, I - initial, F - final
Molarity (M): moles of solute/L of solution
Molality (m): moles of solute/kg of solvent
Percent by mass: mass of solute/mass of solution x 100
Percent by volume: volume of solute/volume of solution x 100

Water Molecules:

  • A simple tri-atomic molecule (H20)
  • Each O-H bond is highly polar because of oxygen’s high electronegativity
  • Bond angle: 105°
  • Due to the bent shape, the O-H bond polarities do not cancel so water is a polar molecule
  • Hydrogen bonding gives water a high surface tension and low vapor pressure
  • Ice is less dense than water because it has a open framework of water molecules

Solutions Sample Problems:

1) IV solutions are often administered to patients in the hospital. One saline solution contains 0.90 g NaCl in exactly 100 mL of solution. What is the molarity of the solution?

Known: 0.90 g of solute, 100 mL of solution, molar mass of NaCl (58.44 g/mol)
Unknown: molarity

0.90 g NaCl
1 mol NaCl
1000 mL
100 mL
58.44 g NaCl
1 L

Solution concentration = 0.15 M

2) A gas has a solubility in water of 16.9 g/L at 15°C and 505 kPa of pressure. What is its solubility in water at 15°C and 606 kPa of pressure?

S1/P1 = S2/P2

16.9 g/L/505 kPa = S/606 kPa
Solubility = 20.3 g/L

3) Calculate the molality of a solution prepared by dissolving 175 g of KNO3 in 1250 g of water

175 g KNO3
1 mol KNO3

101.10 g KNO3

1.73 moles of KNO3
1250 g solvent >>> 1.25 kg solvent

moles of solute/kg of solvent
1.73 mol KNO3/1.25 kg solvent
Molarity = 1.38 m

4) How many mL of alcohol are in 240 mL of a 95.0% (v/v) alcohol solution?

(240 mL)(0.95)
228 mL of alcohol

Kinetics, Reaction Rates, and Equilibrium:

Rate - a measure of the speed of any change that occurs within an interval of time
Activation energy - minimum energy needed to react
Activated complex - an unstable arrangement of atoms that forms momentarily (typically about 10^-13 seconds) at the peak of the activation energy barrier
the activation complex is also called the transition state
Catalyst - a substance that speeds up a reaction, without being consumed itself in the reaction
Enzyme - a large molecule (usually a protein) that catalyzes biological reactions
Inhibitors - interfere with the action of a catalyst; slow or even stop the reaction
Reversible reaction - a reaction in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously (double arrow is used to indicate this)
Le Chatelier’s principle - if stress is applied to a system in equilibrium, the system changes in a way that relieves the stress
Common ion - a ion that is found in both salts in a solution
Free energy - energy that is available to do work
Spontaneous reaction - occurs naturally, and favors the formation of products at the specified conditions
Non-spontaneous reaction - a reaction that does not favor the formation of products at the specified conditions
Entropy - a measure of disorder (J/mol x K)
Law of disorder - the natural tendency is for systems to move to the direction of maximum disorder, not vice-versa
Elementary reaction - a reaction in which the reactants are converted to products in a single step
Intermediate - the product of one of the steps in the reaction mechanism

Factors affecting rate:

  • Temperature
  • Increasing the temperature increases the rate of the reaction
  • Surface area
  • Increasing the surface area increases the rate of the reaction
  • Concentration
  • Increasing the concentration usually increases the rate of the reaction
  • Presence of a catalyst

What caused “stress” on a system (reaction)?

  • Concentration
  • Adding more reactant produces more product, and removing product as it forms will produce more product
  • Temperature
  • Increasing the temperature causes the equilibrium position to shift to the direction that absorbs heat
  • Pressure
  • Changes in pressure will only affect gaseous equilibria
  • Increasing the pressure will usually favor the side with fewer molecules

Equilibrium constant (Keq):

  • Chemists express the equilibrium in terms of a number
  • Only include aqueous and gas reactants and products
  • This value relates to the amounts (molarity) of reactants and products at equilibrium
  • aA + bB <> cC + dD
  • Keq is the ratio of product concentrations to reactant concentrations at equilibrium with each each concentration raised to the value of its balancing coefficient
  • Keq = [C]^c x [D]^d / [A]^a x [B]^b
  • Brackets []: molarity concentration
  • Keq > 1, products are favored at equilibrium
  • Keq < 1, reactants are favored at equilibrium
  • Keq expression doesn’t have units

How to find the solubility product constant (Ksp):

  • Write the balanced chemical equation
  • Split the chemicals into its ions
  • Write the equilibrium expression using the ions (don’t include liquids or solids)
  • Fill in the known values
  • Calculate answer
    see sample problem

Gibbs free energy equation: △G = △H - T△S
if △G is negative, the reaction is spontaneous; if △G is positive, the reaction is nonspontaneous
Rate (A >>> B): -△A/△t
Rate (aA + bB >>> cC + dD): k[A]^a[B]^b
k is the specific rate constant; if k is large than the products form quickly, if k is small than the products form slowly
the powers to which the reactants are raised to are called “orders”

3 things that affect rate:

  • Concentration (molarity) of the reactants
  • The power to which the reactants are raised
  • The value of k (which is different for every equation)

Kinetics, Reaction Rates, and Equilibrium Sample Problems:

1) Write an expression for equilibrium and calculate the equilibrium constant for this reaction.
A liter of a gas mixture at equilibrium at 10ºC contains 0.0045 mol of N2O4 and 0.030 mol of NO2.
N2O4 (g) <> 2NO2(g)

Keq = [NO2]^2/[N2O4]^1
Keq = [0.030 mol/L]^2/[0.0045 mol/L]^1
Keq = 0.20

2) Find the equilibrium expression and solubility product constant for this reaction.
AgCl (s) <> Ag+ (aq) + Cl- (aq)

Keq = [Ag+][Cl-]/[AgCl]
Ksp = [Ag+][Cl-]
find solubility values from the solubility table
Keq = 1.8 x 10^-10

3) If the exponent in the rate law of A is 2 and B is 3, then what is the order of A, order of 3, and overall order of the reaction?

The reaction is 2nd order in A, 3rd order in B, and 5th order overall.

Acids, Bases, and Salts:

Properties of acids:

  • Taste sour
  • Conduct electricity
  • React with metals to form H2 gas
  • Change the color of indicators (blue litmus paper turns red)
  • React with bases to form water and a salt
  • Have a pH of less than 7
  • React with carbonates and bicarbonates to produce a salt, water, and carbon dioxide gas
  • Examples: H2SO4, HNO3, HCl, H3PO4, HC2H3O

Neutralization reactions ALWAYS produce a salt and water

Properties of bases:

  • React with acids to form water and a salt
  • Taste bitter
  • Feel slippery
  • Can be strong or weak electrolytes in aqueous solutions
  • Change the color of indicators (red litmus paper turns blue)
  • Examples: NaOH, KOH, Mg(OH)2, Ca(OH)2

Acid-Base theories:
Arrhenius definition -1887

  • Acids produce hydrogen ions in aqueous solutions (H+)
  • Bases produce hydroxide ions when dissolved in water (OH-)
  • Limited to aqueous solutions
  • Only one kind of base (hydroxides)
    Bronsted-Lowry definition - 1923
  • Broader definition than Arrhenius
  • Acids are hydrogen-ion donors
  • Bases are hydrogen-ion acceptors
  • Acids and bases always come in pairs
  • Water is a base, makes hydronium ion
  • Conjugate base - the remainder of the original acid, after it donates its hydrogen ion
  • Conjugate acid - the particle formed when the original base gains a hydrogen ion
    Lewis definition
  • Acids are electron pair acceptors
  • Bases are electron pair donors
  • Most general of all 3 definitions

Polyprotic acids: monoprotic (HNO3), diprotic (H2SO4), triprotic (H3PO4)
Organic acids - acids that contain the carboxyl group (-COOH), organic acids are weak acids
Amphoteric - a substance that can act as both an acid and a base
Hydronium ion (H3O+) - the positive ion formed when a water molecule gains a hydrogen ion
pH - hydrogen power
Neutralization reaction - a reaction in which an acid and a base react in an aqueous solution to produce a salt and water
Titration - the process of adding a known amount of solution of known concentration to determine the concentration of another solution
Equivalence point - when the moles of hydrogen ions equals the moles of hydroxide ions
Indicators - used to show whether or not neutralization has occurred
Phenolphthalein is a common indicator; clear when acidic or neutral and pink when alkaline
Standard solution - solution of known concentration
End point - the point in the titration where the indicator changes color
Salt hydrolysis - a salt that reacts with water to produce an acid or base
Buffers - solutions in which the pH remains relatively constant, even when small amounts of acids or bases are added
Buffer capacity - the amount of acid or base that can be added before a significant change in pH

Ion product constant : Kw = [H+] x [OH-] = 1 x 10^-14 M^2
pH = -log[H+]
Acid dissociation constant (HA + H2O <> A- + H+): Ka = [H+][A-]/[HA]
Base dissociation constant (MOH + H2O <> OH- + M+): Kb = [M+][OH-]/[MOH]

Strengths of acids/bases:

  • Strong - ionize completely (more products)
  • Weak - ionize only slightly (less products)

Acids, Bases, and Salts Sample Problems:

1) Complete the following chemical equations:
HCl + Mg

MgCl2 + H2 (acids react with active metals to form a salt and hydrogen gas)
HCl + NaHCO3
NaCl + H2O + CO2
2HCl + Mg(OH)2
MgCl2 + H2O

2) Identify the Lewis acid and Lewis base in the following equation:
NH3 + BF3 >>> NH3BF3

Ammonia is the Lewis base because it donates a pair of electrons and boron trifluoride is the Lewis acid because it accepts a pair of electrons

3) Colas are slightly acidic. If the [H+] in a solution is 1.0 x 10^-5 M, is the solution acidic, basic, or normal? What is the [OH-] of this solution?

[H+] is 1.0 x 10^-5 M, which is greater than 1.0 x 10^-7 M, so the solution is acidic.
[OH-] = 1.0 x 10^-9 M

4) What is the pH of a solution with a hydrogen ion concentration of 4.2 x 10^-10 M?

pH = -log[H+]
pH = -log[4.2 x 10^-10]
pH = 9.38