Finals Flashcards
What is measured at constant pressure using a calorimeter?
delta H
Third Law of Thermodynamics
the entropy of a pure, perfectly crystalline substance at absolute zero (0K) is zero
Which are state functions and path functions?
state function: T, E, H, S, G, p, V, m, composition
path function: t,w and q
A reaction has a molar enthalpy (delta H) of formation of 50.63 kJ/mol. Under what conditions are the signs of delta E, w and q for this reaction?
*endothermic reaction
delta E = negative
w = negative
q = negative
How would you order the most stable to least stable to thermally decompose their elements under standard state conditions?
- Entropy (S)
1) most to least stable: solid, liquid, gas
2) bigger the molar mass, the less stable it is
First Law of Thermodynamics
Energy may be transferred as work or heat, but no energy can be lost, nor can heat or work be obtained from nothing
Second Law of Thermodynamics
There is an increase in entropy (disorder) as a spontaneous reaction occur in the universe; some energy lost as heat
The delta T is dependent on:
on the amount of heat transferred, q
on the direction of heat flow
inversely on the amount of material
on the identity of the material.
Heat capacity (C)
quantity of energy required to increase the temperature of a sample by 1°C
C is dependent on:
- mass of substance
- type of material
- state
- temperature
Specific heat capacity (Cs)
amount of heat energy required to raise the temperature of one g of a substance by 1°C
Jg-1K-1
Work (w)
energy used to move an object against an opposing force
w is dependent on:
- magnitude of the applied force
- displacement distance
As T increases, what happens to delta G
increases
Bond energy
energy required to break or make a bond
bond making= negative delta E
bond breaking= positive delta E
Exothermic vs Endothermic
Exothermic:if the reaction releases heat. Endothermic: if the reaction absorbs heat.
Enthalpy (delta H)
heat transferred into or out of a system at constant pressure.
Heat of vaporization
heat required to convert liquid to gas
Hess’s Law
A change in any state function is independent of path.
Thus, the energy change in a chemical reaction is independent of the manner in which the reaction takes place.
The enthalpy change for any overall process is equal to the sum of enthalpy changes for any set of steps that leads from the starting materials to the products.
Standard state
most stable form of a substance at T = 25 °C and p = 1 bar, and 1 M if it is in solution
delta H is dependent on:
temperature, concentration and pressure
Spontaneous direction
- preferred direction of reaction under specific conditions
- can only be reversed by the action of an outside force
What is the general tendency of heat?
Heat flows from high temperature to low temperature
Entropy (delta S)
the quantitative measure of dispersal resulting from an energy transfer
units: JK-1 or JK-1mol-1
As phases change from a solid to a liquid, and temperature increases, what happens to entropy?
entropy increases
delta S depends on:
- temperature
- phase changes (solid is most stable, gas is least satble)
- molar mass (larger the molecule, greater the entropy
- quantity of matter
can S ever be absolute zero?
no never at any temperature above 0K
Gibbs Free Energy (delta G)
Maximum useful work that can be obtained from a thermodynamic system at constant T, P.
Predicts in which direction a reaction is spontaneous
What does a negative and positive delta G indicate?
∆G > 0, process is reactant favoured
∆G < 0, process is product favoured
A reaction is spontaneous at all temperatures when ΔH°, ΔS° and ΔG° are?
ΔH° = -
ΔS° = +
ΔG° (high temperature)= -
ΔG° (low temperature)= -
A reaction is spontaneous at no temperatures when ΔH°, ΔS° and ΔG° are?
ΔH° = +
ΔS° = -
ΔG° (high temperature)= +
ΔG° (low temperature)= +
A reaction is spontaneous at high temperatures (entropy driven) when ΔH°, ΔS° and ΔG° are?
ΔH° = +
ΔS° = +
ΔG° (high temperature)= -
ΔG° (low temperature)= +
A reaction is spontaneous at low temperatures (enthalpy driven) when ΔH°, ΔS° and ΔG° are?
ΔH° = -
ΔS° = -
ΔG° (high temperature)= +
ΔG° (low temperature)= -
For a reaction to be thermodynamically favourable, what sign must be delta G?
negative delta G
The spontaneous direction of a phase change is dependent on:
pressure
Reaction mechanisms
the exact molecular pathway that starting materials follow on their way to becoming products
Reaction intermediates
species that are not part of the reaction stoichiometry
created in one step and consumed in a later step
Rate-determining step
- the slowest elementary step in a mechanism
- governs the rate of the overall chemical reaction because no net chemical reaction can go faster than its slowest step
- reaction that requires the most energy to proceed (highest activation energy)
rate is dependent on:
- concentration of a reactant
- temperature
- catalysts
units: molL-1s-1
Rate law
the effect of concentration on the rate of a particular chemical reaction
The rate law of an elementary reaction can be found:
written directly from the stoichiometry of the reactants
*****ONLY FOR ELEMENTARY REACTIONS
An overall reaction’s rate law can be found by:
the rate determining step since it is the slowest reaction
Half-life
the time required for the reactant concentration to drop to one-half its original value, t1/2
What are the units of the zero, first and second order?
zero=Ms-1
first=s-1
second=M-1s-1
Mechanism
- one or more elementary reactions describing how the chemical reaction occurs
- sum of the individual steps in the mechanism must give the overall balanced chemical equation
- reaction mechanism must be consistent with the experimental rate law
Activation energy
amount of energy to overcome an energy barrier
Activation energy is dependent on:
-temperature= as temperature increases, more reactants will have sufficient energy to overcome the activation barrier
What must happen in order for a reaction to be successful?
There must be sufficient collision energy Reactants must be the correctly orientated
The lower the probability of alignment, the lower the value of k and the slower the reaction.
catalysts
- substances that speed up the rate of a chemical reaction
- not consumed in a chemical reaction
- provide a different pathway with a lower activation energy.
- have no effect on equilibrium
Heterogeneous catalysts vs Homogeneous catalysts
Heterogeneous: catalyst is in a different phase than the reacting substance
Homogeneous: catalyst is in the same phase as the reacting substance
Equilibrium constant (Keq)
- related to the stoichiometry of the balanced net reaction.
- applies only at equilibrium.
- independent of initial conditions
What happens to Keq when
If you flip an equation, Keq (forward) = 1/Keq (reverse)
If you add equations together,
Keq (overall) = Keq,1Keq,2
If you multiply an equation by a coefficient n, then Keq,new = (old Keq)^n
What is the relationship between Q and Keq?
- If Q< Keq, the reaction goes to the right to make products.
- If Q= Keq, the reaction is at equilibrium and there is no net change.
- If Q> Keq, the reaction goes to the left to make reactants.
What is the relationship between Keq and temperature?
1) Keq of an exothermic reaction decreases with increasing temperature.
2) Keq of an endothermic reaction increases with increasing temperature.
Le Chatelier’s principle
- when a change is imposed on a system at equilibrium, the system will react in the direction that reduces the amount of change.
ex) addition or removal of reactants, products and heat; in gases, addition or removal of volume and pressure - *****does not apply to the addition of catalysts. Catalysts only change the rate of the reaction, not the equilibrium conditions.
Approximation method using ice tables
If the approximate value of x is less than 5% of the initial concentration, the approximation is valid!!
Acid
donates a proton
Base
accepts a proton
Spectator ions
ionic species that undergo no significant reactions
Monoprotic acids, Monoprotic bases, Polyprotic acids, Polyprotic bases
Monoprotic acids: can only donate one proton
Monoprotic bases: can only accept one proton
Polyprotic acids: can donate two or more protons
Polyprotic bases: can accept two or more protons
Amphiprotic
molecules or ions which can behave either way (either an acid or a base)
What happens to strong acids and baces
dissociate 100%
***Strong acids
HCl Hydrochloric acid HBr Hydrobromic acid HI Hydroiodic acid HNO3 Nitric acid HClO4 Perchloric acid H2SO4 Sulfur acid
***Strong bases
LiOH Lithium hydroxide
NaOH Sodium hydroxide
KOH Potassium hydroxide
Any soluble hydroxide (Group 1 and 2)
Hydronium ion
H3O+
Keq of water pure water at 25 degrees celsius and 1 bar
1*10^-14
Kw is dependent on:
equilibrium constant of pure water is dependent on temperature
A change of pH by one unit results in what change in [hyrdonium ion]?
tenfold change
All dissociations of weak acids and bases are determined by what, since it does not dissociate 100%?
equilibrium constants Ka and Kb
Oxoacid
an acid that contains an inner atom bonded to a number of oxygen atoms and acidic OH groups
- Anions that are conjugate bases of weak acids make a solution ______.
- Cations that are conjugate acids of weak bases make a solution _____.
- Anions that are conjugate bases of weak acids make a solution basic.
- Cations that are conjugate acids of weak bases make a solution acidic.
The stronger the acid, the ____ the conjugate base.
weaker
Rules on salts
Cations are potential acids, anions are potential bases
Anions of strong acids do not affect pH (Cl-, Br-, I-, NO3-, ClO4-)
+1 and +2 metal cations do not affect pH The anions of weak acids are weak bases and thus make the solution basic (RCOO-, CN-, F-, NO2-,…)
The cations of weak bases are weak acids and thus make the solution acidic (NH4+, RNH3+)
As acid strength increases, Ka value______ and the negative charge ______, polarity (electronegativity) _____, bond energy ______
increases
decreases
increases
weakens
How does the structure of a molecule make a solution more acidic?
- In order to donate a proton, a molecule must break a H–X bond.
- Electronegative atoms withdraw electron density from O–H and thus weaken this bond, and thus increase acidity
Indicator
an acid-base indicator changes colour based on pH and helps determine the endpoint of a reaction
Endpoint
reaction is complete indicated by the physical change in colour of the solution
Equivalence point / Stoichiometric endpoint
when the moles of acid = moles of base
How do you choose the best indicator for your reaction?
choose one that gives an endpoint very close to the equivalence point
As oxygen atoms are added to an acid, electronegativity ______ and acid strength ______
increases
increases
Half-equivalence point
half of the weak acid has been converted to conjugate base. [HA] = [A–]
In the titration of a polyprotic acid, what happens?
two equivalence points, the closer Ka1 and Ka2 are to each other, the less distinguishable the equivalence points
Which Ksp values indicate a soluble, slightly soluble and insoluble compounds?
- insoluble: Ksp «_space;1
- slightly soluble: (10^-5) < Ksp < (10^-2)
- soluble: Ksp > (10^-2)
Q vs Ksp
Q = Ksp: equilibrium solution is saturated
Q < Ksp not at equilibrium and not saturated Q > Ksp not at equilibrium and supersaturated
*If a solution is not at equilibrium, the reaction will shift either left or right until it is at equilibrium.
Oxidation
becomes more positive by losing electrons
Reduction
becomes more negative by gaining electrons
Oxidizing agent
the compound being reduced
Reducing agent
the compound being oxidized
Galvanic cell
each half reaction is separated into half-cells connected by a wire and salt bridge
Salt bridge
allows ion migration without free mixing of solutions
Anode
- negative
- oxidation happens here
- produces electrons
Cathode
- positive
- reduction happens here
- uses electrons
Shorthand cell notation
Anode | Oxidation Reagents || Reduction Reagents | Cathode
double line=salt bridge
single lines=phase boundaries
Standard electrical potential (Ecell)
-electrochemical cell potential under standard state conditions
-Electrical potential energy is converted to other forms of energy as the electrons flow through the external circuit
-The amount of potential energy depends on the two half reactions:
Potential = electrical potential energy difference
Positive reduction potentials
Species that are reduced more easily than H3O+have more positive reduction potentials
- Strong oxidizing agents
- occurs at cathode as a reduction
Standard Hydrogen electrode
2H30+ +2e = H2 + H2O
Standard Reduction Potential(Ecell)= 0V
Negative reduction potentials
- occurs at anode as oxidation
* strong reducing agent
Standard cell potential
the difference between two standard reduction potentials
Standard cell potential depends on:
- concentrations
- pressures
- *** DOES NOT DEPEND ON STOICHIOMETRY, when multiplied by an integer, it remains unchanged
Fuel cells
- batteries in which reactants are constantly being added
- Anode and Cathode both Pt coated metal
Electroplating
The process of depositing one metal on top of another, done usually for cosmetic reasons
Battery
a galvanic cell that generates electrical current to power a practical device
Alkaline Dry Cell Battery
oxidation of Zn under basic conditions with MnO2 and a passive graphite electrode
Lead storage batteries
provides power in automobiles, Pb anode and PbO2 cathode
Nickel-Cadmium Battery
Cd anode and NiO2 cathode
- Rechargeable because the nickel and cadmium hydroxides adhere tightly to the electrodes.
- Drawback=based on toxic heavy metal, also adds weight
Lithium Ion Battery
- Li ions are intercalated in both the anode and cathode materials
- Li ions move from the anode to the cathode during discharge, and back during recharge.
- Light weight, high power batteries
Corrosion
the natural redox process that returns refined metals to their more stable oxides
ex) rust where Fe2+ is oxidized to Fe2O3
How can you prevent or slow the process of corrosion?
“Galvanizing” an iron object with zinc metal prevents corrosion of the iron because the Zn is preferentiallty oxidized.
Endothermic
delta E = negative
w = negative
q = negative
delta H = positive
Exothermic
delta E = positive
w = positive
q = positive
delta H = negative
A reaction is spontaneous, nonspontaneous or at equilibrium if delta G is:
spontaneous = - nonspontaneous = + equilibrium = 0
Which value of specific heat capacity will have the greatest temperature change?
lowest specific heat capacity value
- Which of the following statements is true of a dynamic equilibrium process:
a) The rate of the forward reaction is equal to the rate of the reverse reaction
b) The concentrations (or pressures) of the products equal to that of the reactants
c) Neither the products or the reactants are completely used up
d) Two of the above statements about dynamic equilibrium are true
e) All three of the above statements about dynamic equilibrium are
a and c are both true
For the exothermic reaction: S (s) + 3/2 O2 (g) SO3 (g) Which of the following will result in the formation of more products for this reaction? (a) increase the pressure of the container (b) add more SO3 to the reaction mixture (c) add a catalyst (d) heat the reaction mixture (e) add N2 gas
a
- Suppose we have a solution of AgBr at equilibrium. What will be the effect on this equilibrium when solid of Ag2CO3 is added:
a) Precipitate AgBr will form
b) More of the AgBr will dissolve
c) The pH of the solution will decrease
d) There will be no effect on equilibrium of the AgBr, since Ag2CO3 is a solid
e) The reaction mixture may explode…
a
When the equilibrium constant is greater than 1, which side is favoured? What if it’s smaller than 1?
greater than 1= products are favoured
smaller than 1= reactants are favoured
