FINAL EXAM Flashcards
Define solubility
measure of max amount of solute in a solution
Define saturated, unsaturated and supersaturated solution.
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
Describe the three types of interactions that determine the extent to which a solute is
dissolved in solution.
solute to solute interactions
solvent to solvent interactions
solute to solvent interactions (likes dissolve likes –> miscible)
List and describe the factors that affect the solubility of a solute.
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
Use van’t Hoff factors to determine the colligative properties of electrolyte solutions.
non-electrolytic solutions: i =1
electrolytic solutions: i = # of ions the solute will dissociate into
Use colligative properties to determine the percent dissociation (or percent ionization) of
an electrolyte in solution.
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
Describe the main factors that can increase the rate of a reaction.
increase in concentration of reactant
increase in temperature
adding a catalyst
Describe the collision theory of chemical kinetics.
reacting molecules must collide
molecules must have correct orientation
activation energy must be exceeded
Define effective collision and activation energy.
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
Predict the units of the rate constant k for a reaction.
look at the reaction order
Define reaction mechanism.
what we think actually happens during a reaction, a series of elementary steps that lead to an overall reaction
Determine the rate law of a reaction given its rate determining step.
based off the species in the rate determining step
Understand the criteria that must be met for a proposed mechanism to be plausible.
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
Define elementary reaction.
a reaction that occurs in a single step within a reaction mechanism
Define catalyst and intermediate.
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
Distinguish between a spontaneous and non-spontaneous process and provide examples of each
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
Define entropy.
measurement of randomness or disorder in a system
Describe the conditions for standard entropy.
the entropy of 1 mol of a substance in its standard state at 1 atm
List key trends in standard entropy of atoms and molecules.
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
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.
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
Give in your own words the second law of thermodynamics.
the state of the entire universe’s entropy is always increasing
Determine whether a process is spontaneous given ΔSsurr and ΔS sys.
ΔS universe = ΔS surr + ΔS sys
as long as ΔS universe is positive, a process is spontaneous
Give in your own words the third law of thermodynamics.
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
Define Gibbs free energy.
max useful work that can be done by a system on its surroundings in a spontaneous process at constant T and P
Define standard free energy of formation.
free energy change when reactants and products are in their standard states (pure material, 25 C, 1 atm, 1 M of solution)
Define reversible process.
a process in which the system and surroundings are returned to their initial state
Define equilibrium.
forward and reverse reactions of a process occur at equal rates (ratio of reactants to products is constant)
Differentiate between equilibrium constant and reaction quotient. How is the reaction quotient connected to the
equilibrium achievement?
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
Use the equilibrium constant to predict the relative amounts of reactants to products at equilibrium.
K –> infinity, reaction goes to completion
K –> 0, no reaction occurs
K > 1, more products than reactants
K < 1, more reactants than products
Differentiate between heterogeneous and homogeneous equilibria.
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)
Predict the direction of a reaction given initial concentrations of reactants and products and the value of the equilibrium constant.
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
Give in your own words the meaning of Le Châtelier’s principle
when stress is applied to a system at equilibrium, the system will shift to reduce applied stress and re-establish equilibrium
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.
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
how do catalysts and inhibitors effect equilibrium
neither have an effect on equilibrium when added to a reaction, just effect the speed at which equilibrium is met
manipulating chemical equations
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)
Define conjugate acid and conjugate base and cite examples of conjugate pairs.
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)
Give in your own words definitions of acids/bases: Brønsted and Lewis.
Bronsted Acid: a proton (H+) donor
Bronsted Base: a proton (H+) acceptor
Lewis Acid: an electron pair acceptor
Lewis Base: an electron pair donor
Rank species of each of the following acids based upon relative strength: hydrohalic,
oxoacids, and carboxylic.
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
Use the pH scale to classify a solution as being acidic, basic or neutral.
pH < 7 –> acidic
pH = 7 –> neutral
pH > 7 –> basic
Identify an acid or base as being strong or weak.
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
Describe the common ion function in buffers
a common ion prevents the weak acid/base from ionizing as much as it usually does and keeps the pH from changing
Describe how a buffer works when absorbing external acids or bases.
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
Predict the pH of a titration given the type of acid and base (strong or weak)
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
Use a titration curve to identify what type of titration was employed (e.g.
weak acid plus strong base)
look at pH where curve starts and steepness and what pH equivalence point is at
Distinguish between endpoint and equivalence point
equivalence point: point at which moles of acid = moles of base
endpoint = point in titration where the indicator changes color
Select an indicator given the acid and base species used in a titration
look at pH range of where indicator works in comparison to pH of equivalence point
Define solubility product constant (Ksp)
measures the extent to which a substance will dissolve in water (salt and dissolved ions reach equilibrium)
Cite examples of how the following can affect solubility: common ions and pH
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
Define key electrochemical terms: galvanic cell, anode, cathode and salt bridge
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
Describe the conditions used to calculate the standard reduction potential for a reaction
all solutes are 1 M and all gases are at 1 atm, temperature is 25 C
Explain what an electrochemical cell is and how it works
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
Define concentration cell
cells where the cell potential is generated entirely by a difference in concentration at the two electrodes (same ions in anode and cathode)
