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Chem Exam 2

TermDefinition
Buffer Resists pH change when strong acid/base is added
Components of a buffer Conjugate acid-base pair
pH range 0-14
pH less than 7 Acidic
pH more than 7 Basic
pH equal to 7 Neutral
Henderson-Hasselbach equation Used to find pH/pOH
Buffer capacity Ability of buffer to resist change in pH
Highest buffer capacity pH = pKa or ratio is 1:1
Valid buffer ratio Between 10:1 and 1:10
Three ways a buffer can be created 1. More weak acid than strong base 2. More weak base than strong acid 3. Weak acid and weak base conjugate pair
Acid-base titration Neutralization reaction
Equivalence point When moles of acid and base are equal
pH of equivalence point of strong acid and strong base 7
pH of equivalence point of weak acid and strong base More than 7
pH of equivalence point of strong acid and weak base Less than 7
Weak polyprotic acids Have more than 1 equivalence point
Half equivalence point pH = pKa and acid = conjugate base
Buffer region Flat part of graph
After equivalence point Only conjugate base/acid left
Strong acid added to buffer Weak base decreases, conjugate acid increases
Strong base added to buffer Weak acid decreases, conjugate base increases
pH indicator is usually Organic compound, weak acid or base
Active zone of indicator + or - 1 of pH
Semisoluable salts Primarily solid in water
s Molar solubility
Q Reaction quotient, used when not at equilibrium
Ksp Solubility product constant, used at equilibrium
Larger s More soluble, less stable
How to determine higher solubility Compare Ksp if number of ions is the same, s if different.
Factors that increase solubility Moves reaction forward
Factors that decrease solubility Moves reaction backwards
Common ion effect Semisoluable salt is placed into a solution that contains components of the equilibrum, soluility is decreased compared to pure water
Precipitation reactions Forms a solid from aqueous ions
Q=K Saturated solution at equilibrium, no precipatate
Q>K Precipitate will form until remaining solution is saturated
Q<K Unsaturated, no precipitate
Fractional precipitation Separating substances by means of their gradual precipitation
Complex ion Metal cation and ligand with an overall charge
Complex formation Increases solubility of semisoluble component
Amphoteric solids React as acid or base
Thermodynamics Science of heat and work
First law of thermodynamics Energy cannot be created or destroyed (conservation of energy)
SI unit of energy J
Internal energy Sum of potential and kinetic energy
Kinetic energy Energy of motion
Potential energy Amount of energy that can turn into kinetic
Enthalpy ΔH°, idicates endothermic/exothermic
Exothermic processes -ΔH°, favorable
Endothermic processes +ΔH°, unfavorable
Phase change effect on ΔH° +ΔH° as T increases, -ΔH° as T decreases
Bonds broken +ΔH°
Bonds formed -ΔH°
Hess's Law Used to find ΔH°rxn from ΔH°f
Elemental state State of element found at standard conditions, ΔHf is 0
Standard conditions 1 mol, 1 atm, 298°K
Diatomic elements H2, N2, O2, F2, I2, Cl2, Br2
Polyatomic elements S8, P4
Spontaneous change Does not require additional energy
Usually spontaneous ΔH° < 0
Usually nonspontaneous ΔH° > 0
Entropy S°, measure of chaos, natural tendency to increase
Entropy equation S° = k ln W
3 types of motion 1. Translational 2. Vibrational 3. Rotational
Microstate Quantitized state of the system of molecules
2 ways to find ΔS°sys 1. Number of microstates 2. Change in heat
ΔS° increases as T Increases
ΔS° increases as phase changes From solid to liquid to aqueous to gas
ΔS° increases as number of bonds Increases
ΔS° increases as size Increases
Calculating entropy is similar to Hess's Law
+ΔS° More disorder, favorable
-ΔS° Less disorder, unfavorable
Solid or liquid dissolved in solution +ΔS°
Gas dissolved in solution -ΔS°
Gas dissolved in gas +ΔS°
Increase in moles +ΔS°
Decrease in moles -ΔS°
Order for entropy 1. Highest phase moles 2. Phase/temperature 3. # of elements 4. Size 5. Intermolecular forces
Second law of thermodynamics Spontaneous processes lead to increase in entropy of the universe
ΔS°univ ΔS°sys + ΔS°surr
ΔS°surr - ΔH°rxn / T
Phase transitions effect on ΔS°univ ΔS°univ = 0
Third law of thermodynamics Entropy of a pure, perfectly formed crystelline substance at 0°K is absolute zero
Sign of ΔS°surr is aways opposite of ΔH°sys
ΔG° Gibb's free energy
K > 1 Product favored
K < 1 Reactant favored
K = 1 At Equillibrium
At standard conditions, a spontaneous, product favored reaction -ΔG°
At standard conditions, a nonspontaneous, reactant favored reaction +ΔG°
How do you increase buffer capacity? Increase concentration of the buffer
At equilibrium, the ΔS°univ is 0
When moving to equilibrium, the ΔS°univ is ΔS°univ > 0
When a reaction is spontaneous at all T, the reverse reaction is Nonspontaneous at all T
Crossover temperature The T that the reaction turns from spontaneous to nonspontaneous
Crossover temperature equation T = ΔH°rxn/ΔS°rxn
Hess-type equation Used to find ΔG°rxn from products and reactants
Solid heated below melting point is Nonsponspontaneous
Solid heated above melting point is Spotaneous
Solid heated at melting point At equilibrium
How can you make an unfavorable reaction favorable? Couple it to a favorable reaction
n Number of mols
When does Q = K? At equilibrium
What does ΔG (without a circle) mean? Not at standard conditions
ΔG < 0 (nonstandard) Can release free enegy until ΔG = 0, which is maximum work obtainable
ΔG > 0 (nonstandard) Can gain free enegy until ΔG = 0, which is minimum work necessary
R Gas constant, 8.314 J/mol*K, on formula sheet
T Temperature in °K
°K Kelvin, °C + 273
Oxidation number Follows electron charge
Oxidation state Hypothetical charge that an atom would have if electrons were transferred completely and not shared
Elemental state oxidation number 0
Monatomic ion oxidation number Charge of ion
Neutral compound sum of oxidation numbers 0
Polyatomic sum of oxidation numbers Charge of ion
Group 1A oxidation number +1
Group 2A oxidation number +2
Group 7A oxidation number Usually -1
Fluorine -1
Hydrogen with nonmetals +1
Hydrogen with metals and boron -1
Oxygen Usually -2
Oxygen with peroxides -1
Redox reaction Reduction-oxidation reaction, involves transfer of electrons from one reactant to another
OIL RIG Oxidation is loss (of electrons), reduction is gain (of electrons)
Redox reactions occur in both Ionic and covalent compounds
Higher positive oxidation number as product than as reactant Reactant is oxidized
Lower positive oxidation number as product than as reactant Reactant is reduced
Reducing agent Species that is oxidized
Oxidizing agent Species that is reduced
Half reaction Only contains the oxidizing or reduction half of the equation
Material balancing Add H2O and H+ after original materials are balanced
Charge balancing Add e- to balance charge after material balancing
In acidic solutions Add H2O to balance oxegen, then add H+ to balance hydrogen
In basic solutions Add same amount of OH- to both sides to convert remaining H+
When coupling redox reactions, make sure to Balance e-
Electrochemical cells Cells that gain energy from chemical reactions
Parts of an electrochemical cell Catgode, anode, salt bridge, path for electron transfer
Salt bridge Transfers cations to cathode solution and anodes to anion solution
Voltaic cell (aka galvanic cell) Spontaneous, ΔG° < 0, K > 1, +E°
In a voltaic cell, which side is positively charged? Cathode
In a voltaic cell, which side is negatively charged? Anode
Electrolytic cell Nonspontaneous, ΔG° > 0, K < 1, -E°
In an electrolytic cell, which side is positively charged? Anode
In an electroytic cell, which side is negatively charged? Cathode
In all electrochemical cells, electrons always flow from Anode to cathode
In all electrochemical cells, reduction takes place at the Cathode (red cat)
In all electrochemical cells, oxidation takes place at the Anode (an ox)
In all electrochemical cells, cations flow to the Cathode
In all electrochemical cells, anions flow to the Anode
In all electrochemical cells, mass is gained at the Cathode
In all electrochemical cells, mass is lost at the Anode
E° Electrical potential
E°cell E°cathode - E°anode
Sign of E°cell for voltaic cells +E°
Sign of E°cell for electrolytic cells -E°
In a voltaic cell, which side has the higher E° value? Cathode
In a voltaic cell, which side has the lower E° value? Anode
In an electrolytic cell, which side has the higher E° value? Anode
In an electrolytic cell, which side has the lower E° value? Cathode
Standard cell notation Anode | oxidized product | reduced ion | cathode
What must be at both ends of standard cell notation? A solid
What are the two inactive electrodes? Pt (s) and C (gr)
In standard cell notation, different phases are separated with a Straight line
In standard cell notation, similar phases are separated with a Comma
Reactant species of a reduction half reaction with a high E° value can be very good Oxidizing agents
Product species of a oxidation half reaction with a high E° value can be a very good Reducing agents
To see a distinct color in a mixture of two colors, you need one color to have 10 times the intensity of the other
Auto ionization of water 2(H2O) = (OH-) + (H3O+)
Equillibrium expression for auto ionization of water Kw = [H3O+][OH-]
Kw Autoionization constant of water, 1E-14
Special cases in solubility Sulfide and carbonates
Created by: lprocopio
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