FINAL EXAM

Question 1

Define solubility

Answer 1

measure of max amount of solute in a solution

Question 2

Define saturated, unsaturated and supersaturated solution.

Answer 2

unsaturated solution: contains less solute than in saturated solution

saturated solution: max amount of solute that stays in solution, extra solute with precipitate

supersaturated solution: contains more solute than in saturated solution, extremely unstable

Question 3

Describe the three types of interactions that determine the extent to which a solute is
dissolved in solution.

Answer 3

solute to solute interactions
solvent to solvent interactions
solute to solvent interactions (likes dissolve likes –> miscible)

Question 4

List and describe the factors that affect the solubility of a solute.

Answer 4

temperature: most solids, as temperature increases solubility increases, gases are less soluble as temperature increases because particles move faster and can escape the liquid more easily

pressure: as pressure increases, solubility increases for gases

polarity: likes dissolve likes

molecular size

Question 5

Use van’t Hoff factors to determine the colligative properties of electrolyte solutions.

Answer 5

non-electrolytic solutions: i =1
electrolytic solutions: i = # of ions the solute will dissociate into

Question 6

Use colligative properties to determine the percent dissociation (or percent ionization) of
an electrolyte in solution.

Answer 6

physical properties of solutions that depend on # of solute particles but NOT on the identity of the solute
i.e. boiling point, freezing point, and osmotic pressure

Question 7

Describe the main factors that can increase the rate of a reaction.

Answer 7

increase in concentration of reactant
increase in temperature
adding a catalyst

Question 8

Describe the collision theory of chemical kinetics.

Answer 8

reacting molecules must collide
molecules must have correct orientation
activation energy must be exceeded

Question 9

Define effective collision and activation energy.

Answer 9

effective collision is collision which follows the requirements of the collision theory and results in a reaction
activation energy is the minimum energy required for a reaction to proceed

Question 10

Predict the units of the rate constant k for a reaction.

Answer 10

look at the reaction order

Question 11

Define reaction mechanism.

Answer 11

what we think actually happens during a reaction, a series of elementary steps that lead to an overall reaction

Question 12

Determine the rate law of a reaction given its rate determining step.

Answer 12

based off the species in the rate determining step

Question 13

Understand the criteria that must be met for a proposed mechanism to be plausible.

Answer 13

the sum of each single-step reaction (elementary) must yield an overall equation and the rate law of the rate-determining step must agree with the experimentally determined rate law

Question 14

Define elementary reaction.

Answer 14

a reaction that occurs in a single step within a reaction mechanism

Question 15

Define catalyst and intermediate.

Answer 15

catalyst: a substance that alters the rate of a chemical reaction by lowering the activation energy for that reaction, can either be heterogeneous (in different phase as reactants) or homogeneous (in same phase as reactants)

intermediate: species that appear in a reaction mechanism but not in the overall reaction, formed in an earlier step and used in a later step

Question 16

Distinguish between a spontaneous and non-spontaneous process and provide examples of each

Answer 16

a spontaneous process is one that occurs without any external input to the system, a non-spontaneous process is one where external input is needed for the process to occur

at a temp. greater than 0 C, ice melting is spontaneous; at a temp. lower than 0 C, ice melting is non-spontaneous

Question 17

Define entropy.

Answer 17

measurement of randomness or disorder in a system

Question 18

Describe the conditions for standard entropy.

Answer 18

the entropy of 1 mol of a substance in its standard state at 1 atm

Question 19

List key trends in standard entropy of atoms and molecules.

Answer 19

standard S of gas > standard S of liquid > standard S of solid

for similar molecules, standard S tends to increase with: increasing molar mass, increased # of atoms in formula

Question 20

Predict the sign of ΔS of a process and use the sign to indicate whether the system has undergone an increase or
decrease in entropy.

Answer 20

entropy is increased when:
moles of products > moles reactants
more complex molecules are broken into smaller, simple molecules
phase change to a more disordered phase (gas > liquid > solid)
generally, dissolving a solute in a solvent

Question 21

Give in your own words the second law of thermodynamics.

Answer 21

the state of the entire universe’s entropy is always increasing

Question 22

Determine whether a process is spontaneous given ΔSsurr and ΔS sys.

Answer 22

ΔS universe = ΔS surr + ΔS sys
as long as ΔS universe is positive, a process is spontaneous

Question 23

Give in your own words the third law of thermodynamics.

Answer 23

the entropy of a pure crystal is 0 at 0 K

the entropy of a system approaches a constant value when its temperature approaches absolute zero

Question 24

Define Gibbs free energy.

Answer 24

max useful work that can be done by a system on its surroundings in a spontaneous process at constant T and P

Question 25

Define standard free energy of formation.

Answer 25

free energy change when reactants and products are in their standard states (pure material, 25 C, 1 atm, 1 M of solution)

Question 26

Define reversible process.

Answer 26

a process in which the system and surroundings are returned to their initial state

Question 27

Define equilibrium.

Answer 27

forward and reverse reactions of a process occur at equal rates (ratio of reactants to products is constant)

Question 28

Differentiate between equilibrium constant and reaction quotient. How is the reaction quotient connected to the
equilibrium achievement?

Answer 28

Q is K at non-equilibrium conditions and indicates the direction of reaction to reach equilibrium

Q < K, higher rate in forward direction
Q > K, higher rate in reverse direction
Q = K, forward rate = reverse rate, at equilibrium

Question 29

Use the equilibrium constant to predict the relative amounts of reactants to products at equilibrium.

Answer 29

K –> infinity, reaction goes to completion
K –> 0, no reaction occurs
K > 1, more products than reactants
K < 1, more reactants than products

Question 30

Differentiate between heterogeneous and homogeneous equilibria.

Answer 30

homogeneous –> all products/reactants are in the same phase
heterogeneous –> products/reactants in mixed phases, only include gaseous and aqueous materials (include phase in equilibrium expression)

Question 31

Predict the direction of a reaction given initial concentrations of reactants and products and the value of the equilibrium constant.

Answer 31

very large K –> product conc. very large or reactant conc. very small, reaction will go to completion, equilibrium favors products

very small K –> product conc. very small or reactant conc. very large, reaction proceeds very slowly or no reaction at all, equilibrium favors reactants

Question 32

Give in your own words the meaning of Le Châtelier’s principle

Answer 32

when stress is applied to a system at equilibrium, the system will shift to reduce applied stress and re-establish equilibrium

Question 33

Apply Le Châtelier’s principle toward determining the shift of a reaction at equilibrium given a change in one of the following: removal or addition of reactant or product, change in volume or pressure, and temperature change.

Answer 33

adding more reactant –> shift towards product
removing product –> shift towards product
adding more product –> shift towards reactant

pressure increased –> equilibrium shifts to produce smaller # moles of gas
pressure decreased –> equilibrium shifts to produce larger # moles of gas
same # moles of gas on both sides –> pressure doesn’t effect equilibrium

volume increased –> equilibrium shifts to produce larger # moles of gas
volume decreased –> equilibrium shifts to produce smaller # moles of gas
same # moles of gas on both sides –> volume doesn’t effect equilibrium

exothermic reaction:
increase T –> K decreases, favors reactants
decrease T –> K increases, favors products

endothermic reaction:
increase T –> K increases, favors products
decrease T –> K decreases, favors reactants

Question 34

how do catalysts and inhibitors effect equilibrium

Answer 34

neither have an effect on equilibrium when added to a reaction, just effect the speed at which equilibrium is met

Question 35

manipulating chemical equations

Answer 35

reversing an equation –> invert K (1/K)
multiplying coefficients by n –> raise K to nth power (K^n)
adding equations –> multiply K values (K1 * K2)

Question 36

Define conjugate acid and conjugate base and cite examples of conjugate pairs.

Answer 36

Conjugate acid: the substance formed from the base after the addition of an H+

Conjugate base: the substance formed from the acid after the loss of an H+

Examples:
HCL (acid) –> Cl- (conjugate base)
NH3 (base) –> NH4+ (conjugate acid)

Question 37

Give in your own words definitions of acids/bases: Brønsted and Lewis.

Answer 37

Bronsted Acid: a proton (H+) donor
Bronsted Base: a proton (H+) acceptor

Lewis Acid: an electron pair acceptor
Lewis Base: an electron pair donor

Question 38

Rank species of each of the following acids based upon relative strength: hydrohalic,
oxoacids, and carboxylic.

Answer 38

easier to lose H+, stronger acid
higher charge = stronger bond to H+ = weaker acid
shorter bonds are stronger so H+ harder to remove
less electronegative = stronger bond = weaker acid

Question 39

Use the pH scale to classify a solution as being acidic, basic or neutral.

Answer 39

pH < 7 –> acidic
pH = 7 –> neutral
pH > 7 –> basic

Question 40

Identify an acid or base as being strong or weak.

Answer 40

Strong acids: HCl, H2SO4, HNO3, HCLO4, HBR, HI

Strong bases: soluble hydroxides (NaOH, KOH, LiOH, Ba(OH)2)

Weak acids: H2CO3, H3PO4, CH3COOH, HF, HNO2, HCN, carboxylic acids

Weak bases: ammonia, amines, insoluble/slightly soluble hydroxides

Question 41

Describe the common ion function in buffers

Answer 41

a common ion prevents the weak acid/base from ionizing as much as it usually does and keeps the pH from changing

Question 42

Describe how a buffer works when absorbing external acids or bases.

Answer 42

when OH- is added, or H+ is reduced, the equilibrium will shift towards the conjugate base and the OH- will react with the H+ to form water

when H+ is added, the equilibrium will shift toward the undissociated acid

Question 43

Predict the pH of a titration given the type of acid and base (strong or weak)

Answer 43

Strong: from start of titration to near endpoint, pH increases slowly –> just before and after equivalence point pH increases rapidly –> at equivalence, moles acid = moles base and solution is neutral (salt)

Weak: initial pH is higher than strong –> pH changes near equivalence point are more subtle –> pH > 7 at equivalence point due to formation of basic salt

Question 44

Use a titration curve to identify what type of titration was employed (e.g.
weak acid plus strong base)

Answer 44

look at pH where curve starts and steepness and what pH equivalence point is at

Question 45

Distinguish between endpoint and equivalence point

Answer 45

equivalence point: point at which moles of acid = moles of base
endpoint = point in titration where the indicator changes color

Question 46

Select an indicator given the acid and base species used in a titration

Answer 46

look at pH range of where indicator works in comparison to pH of equivalence point

Question 47

Define solubility product constant (Ksp)

Answer 47

measures the extent to which a substance will dissolve in water (salt and dissolved ions reach equilibrium)

Question 48

Cite examples of how the following can affect solubility: common ions and pH

Answer 48

common ions: decreases the solubility of a weak electrolyte when present (from strong electrolyte)

pH: basic anion is more soluble in acidic solution, acidic cation is more soluble in basic solution

Question 49

Define key electrochemical terms: galvanic cell, anode, cathode and salt bridge

Answer 49

galvanic cell: an electrochemical cell where a spontaneous redox reaction occurs to produce electrical energy

anode: the electrode at which oxidation occurs

cathode: the electrode at which reduction occurs

salt bridge: keeps the two half-cells connected and allows ions to flow between them

Question 50

Describe the conditions used to calculate the standard reduction potential for a reaction

Answer 50

all solutes are 1 M and all gases are at 1 atm, temperature is 25 C

Question 51

Explain what an electrochemical cell is and how it works

Answer 51

electrons leave the anode and flow through the wire to the cathode

as electrons leave anode, the cations formed dissolve into the solution in the anode

as electrons reach the cathode, cations in the cathode are attracted to the now negative cathode

the electrons are taken by the cation, and the neutral metal is deposited on the cathode

salt bridge prevents charge imbalance –> anions move towards anode and cations move towards cathode

Question 52

Define concentration cell

Answer 52

cells where the cell potential is generated entirely by a difference in concentration at the two electrodes (same ions in anode and cathode)