Equilibrium Flashcards
What is meant by chemical equilibrium?
State in which the rate of forward reaction is equal to rate of reverse reaction
Why is chemical equilibrium described as a dynamic state?
Both reactions are continuously occurring, both continue
What is the law of mass action?
The rate of a chemical reaction is proportional to the active masses of the reacting substances.
What are ‘‘active masses’’
regard as concentrations of the participating
species or partial pressures in the case of gases.
What is the Equilibrium constant?
concentrations of products to the power of moles/concentrations of reactants to the power of moles
[C]c [D]d / [A]a [B]b
For each reaction, how does the value for Keg vary?
For each reaction, there is only one value for Keq at a specific temperature.
What are the characteristics of equilibrium?
Characteristics of equilibrium:
- Dynamic: balance of reversible reactions
- All reactants and products are present (both reactions can occur)
- Move to equilibrium is spontaneous – if disturbed then it returns to the same equilibrium point.
- Represents a compromise between DH (change in enthalpy taking the enthalpy to a minimum DS (change in entropy) taking the entropy to a maximum.
- constant only at a specific temperature
Is the move to equilibrium spontaneous?
Yes, if disturbed then it returns to the same equilibrium point.
What does the value of Keq tell us?
The extent to which reactants are converted into products.
Keq < 1: Reactants are favored at equilibrium
Keq > 1: Products are favored at equilibrium
For gases the concentration is proportional to what?
the partial pressure at a fixed temperature
If gases nA + nB are mixing, with a ratio of 2:1, what will happen?
Eventually will mix (two gases)
What is the molar concentration of the gas equal to?
its partial pressure divided by RT – and RT is constant at a given temperature. (n/V = P/RT)
Consider the following reaction involving gases:
2SO2(g) + O2(g) ⇄ 2SO3(g)
What is kP?
[Pso3]2 / [Pso2]2 [Po2]
What is the relationship between Kp and Kc
Kp = Kc (RT) ∆n
What is the relationship between the free energy change for a reaction G and the equilibrium constant for that reaction, K?
ΔG = - RTln K
Where ΔGo is the standard free energy change and may be written ΔGo = ΔHo- TΔS
o and ΔHo and ΔSo are standard enthalpy and entropy changes and may be determined experimentally.
Write the equilibrium constant expression of
N2O4 ->2NO2
[NO2]^2
———
[N2O4]
State Le Châtelier’s principle
If “stress” is applied to system at equilibrium, system reacts to relieve stress applied (Le Chatelier’s Principle ignores the presence of catalysts)
What are the factors that influence equilibrium?
Concentration, temperature, and partial pressure (for
gaseous)
A + B - C + D
What occurs when concentration of A increases?
System alters to minimise effect and equilibrium shifts to right and more of C and D produced
A + B - C + D
What occurs when concentration of C increases?
System alters to minimise effect, equilibrium shifts to left therefore more of A and B produced
A + B - C + D
What occurs when concentration of A decreases?
System alters to minimise effect, equilibrium shifts to left therefore more of A and B produced
A + B - C + D
What occurs when concentration of C decreases?
System alters to minimise effect, equilibrium shifts to right therefore more of C and D produced
What effect does an increase in pressure have?
[opposite - side with fewer molecules]
System alters to oppose stress to side with fewer molecules
What effect does a decrease in pressure have on an equilibrium reaction?
[think of opposite - side with more molecules]
The system alters to oppose stress, and moves to side with more molecules
What happens if we increase the volume of the container?
Causes decrease in pressure
What happens if we decrease the volume of the container?
Increases the presssure
PCl3 + Cl2 - PCl5
What would the effect on concentration of PCl5 in equilibrium mixture when pressure is increased?
The equilibrium shifts to the aids with smaller amount of molecules to oppose stress, therefore is a forward reaction and PCl5 concentration increases
Determine whether the following reactions favor high or low pressures?
- 2SO2(g) + O2(g) ⇄ 2 SO3(g);
- PCl5(g) ⇄ PCl3(g) + Cl2(g);
- CO(g) + 2H2(g) ⇄ CH3OH(g);
- N2O4(g) ⇄ 2 NO2(g);
- H2(g) + F2(g) ⇄ 2 HF(g);
- High
- Low
- High
- Low
- Same no. moles on each side – nothing will happen
When you increase the temperature what does it favour?
Endothermic (prefer high temperatures because they need heat to be absorbed)
When you decrease the temperature what does it favour
Exothermic (prefer low temperatures because they are releasing heat)
H2 + I2 -> 2HI
Use le châtelier’s principle to predict and explain the effect of a decrease in temperature on the yield of hydrogen iodide
Decrease in yield
Reaction tries to minimise effect of reduced temperature and equilibrium shifts in reverse
H2 + I2 -> 2HI
Colourless purple colourless
Colourless purple colourless
Use le châtelier’s principle to predict and explain the effect of a decrease in temperature on the intensity of colour of the equilibrium mixture
More intense, darker
Reaction tries to minimise effect of reduced temperature and equilibrium shifts in reverse
Determine whether the following reactions favors high or low temperature?
- 2SO2(g) + O2(g) ⇄ 2 SO3(g); DH = -180 kJ
- CO(g) + H2O(g) ⇄ CO2(g) + H2(g); DH = -46 kJ
- CO(g) + Cl2(g) ⇄ COCl2(g); DH = -108 kJ
- N2O4(g) ⇄ 2 NO2(g); DH = +57 kJ
- CO(g) + 2H2(g) ⇄ CH3OH(g); DH = -270 kJ
- Low
- Low
- Low
- High
- Low
The colour of equilibrium mixture is paler at 0 degrees celsius than at T, where T > 0 degrees celcius
Explaining your reason, deduce whether the decomposition of dinitrogen tetroxide into nidtogen dioxide is an exothermic or an endothermic reaction
N204 -> 2NO2
Colourless Dark Brown
Endothermic
At lower temperature, rxns try to minimise effect of reduced temp so equilibrium shifts in direction that produces heat, which is exothermic which is a forward reverse reaction, [more N2O4 is produced]
At higher temperature, rxns try to minimise effect of higher temperature so equilibrium shifts in direction that absorbs heat. Since forward reaction is exothermic, more NO2 is made
Would there be a change in the value of Kc at T If a different initial concentration is used?
No, KC constant is at given temperature
It will only change if temperature is changed
How is gibbs free energy related to the equilibrium constant?
DG = -RTlnK.
where K is the equilibrium constant for a chemical reaction
What is the difference between
a) DG
b) DGo
a) can be measured or calculated from tables of standard values
b) reactants and products in standard states
What are standard state conditions?
Standard state conditions are defined by Standard Temperature & Pressure (STP) with a temperature of 0 oC or 273.15 Kelvin (K) and a pressure of 1 atmosphere (1 atm = 101 325 Pa), temperature.
If DGo is large and negative (i.e. < -10kJ), then what kind of reaction is it?
the reaction is spontaneous as written –
reactants will go almost completely to products if equilibrium is reached.
If DGo is large and positive (i.e. >10kJ), then what kind of reaction is it?
the reaction is not spontaneous as written –
reactants will not give any significant amount of products at equilibrium even if a
catalyst is added
If DGo is small negative or positive (i.e. – 10 kJ < DGo
rxn < 10kJ then what kind of reaction is it?
the reaction mixture at equilibrium will contain significant amounts of both reactants and products.
Does thermodynamics tell us anything about the rate?
Thermodynamics cannot tell us anything concerning the rate at which equilibrium is attained
What is a catalyst?
A catalyst is a chemical that affects the rate of a chemical reaction. It does not affect the position of equilibrium because it affects the forward and backward reactions equally.
Why is proton transfer equilibria important?
An enormously important biological aspect of chemical
equilibrium is that involving the transfer of protons (hydrogen ions, H+) between species in aqueous environments, such as living cells. Even small drifts in the equilibrium concentration of H+ can result in disease, cell damage, and death!
What is the Brønsted–Lowry theory of acid and bases?
An acid is a proton donor and a base is a proton acceptor
Are acids and bases in water in equilibrium?
The proton, H+, is highly mobile and acids and bases in water are always in equilibrium with their deprotonated and protonated counterparts and hydronium ions (H3O+).
Define a conjugate acid-base pair according to Brønsted Lowry
An acid and base species that differ only by one proton
What is the conjugate acid and conjugate base of H2O?
acid - H30+
base - OH-
Define PH
-log10[H+]
The higher the pH, the ____ the concentration of hydronium ions H3O+ in the solution!
lower
How does a change in pH by 1 unit corresponds to a change in their molar concentration
corresponds to a 10-fold change
What does kA tell us?
The value of the acidity constant indicates the extent to which proton transfer occurs at equilibrium in aqueous solution.
If the value of kA is small, is the value of pkA large or small?
The smaller the value of Ka (for instance 10-8 compared with 10-6) and therefore the larger the value of pKa (for instance, 8 compared with 6), the lower is the concentration of deprotonated molecules.
If an acid has pKa > 0, what does this indicate?
indicating only a small extent of deprotonation in water. These acids are classified as weak acids.
What is the ka of most acids?
Ka < 1 (and usually much less than 1),
If Ka > 1, what does this indicate?
A few acids, most notably, in aqueous solution, HCl, HBr, HI, HNO3, H2SO4, and HClO4, are classified as strong acids and are commonly regarded as being
completely deprotonated in aqueous solution.
How do you find
a) PkA
b) PkB
a) pKa = -logKa
b) pKb = -logKb
What is the the autoprotolysis constant for water, Kw?
At 25 degrees, [H30+] [OH-] = 1.0 x 10^-14
(pkW = -logkW =14.00
ka+kb = kw
As Ka increases, what happens to Kb?
Kb decreases to maintain a product equal to the constant Kw - As the strength of a base decreases, the strength of its conjugate acid increases and vice versa
How do you get Pkw?
PH + POH = Pkw
How do we calculate the pH of a solution of a weak acid?
Organize the necessary work into a table with columns headed by the species present in the mixture (ignoring H2O) and, in successive rows write:
- The initial molar concentrations of the species.
- The changes in these quantities that must take place for the system to reach equilibrium.
- The resulting equilibrium values.
What is pKa, pKb, pKw?
It is used to determine the strength of a base or alkaline solution.
Acetic acid lends a sour taste to vinegar and is produced by aerobic oxidation of ethanol by bacteria in fermented beverages, such as wine and cider:
CH3CH2OH(aq) + O2(g) ⇄ CH3COOH(aq) + H2O(l).
Ka = 1.8 x 10-5
We draw up an equilibrium table based on the proton transfer equilibrium
CH3COOH(aq) + H2O(l) ⇄ H3O+ (aq) + CH3CO2-
1) INITIAL CONCENTRATION
CH3COOH H30+ CH3CO2-
.15 0 0
2) CHANGE TO REACH EQUILIBRIUM (mol/dcm3)
CH3COOH H30+ CH3CO2-
-x x x
3) EQUILIBRIUM CONSTANT
CH3COOH H30+ CH3CO2-
0.15-x x x
The value of x is found by inserting the equilibrium concentrations into the expression for the acidity constant:
xx / .15 - x
We can simplify the equation = 0.15 x Ka = x^2
(.15 x 1.8 x 10^-5)^1/2 = 1.6 x 10^-3
-log10(1.6 x 10^-3)
PH = 2.8
Calculate the pH of an 0.20 M aqueous solution of methylamine, CH3NH2, for which pKb =3.44.
We draw up the equilibrium table based on the proton transfer equilibrium:
CH3NH2(aq) +H2O(l) ⇄ CH3NH3 + (aq) +OH-(aq)
CH3NH3 H30+ CH3NH3+ .20 0 0 -x +x +x .20-x x x x2/.20-x = x2/.30 = 3.6 x 10^-4 (.20 x 3.6 x 10^-4) = x2 .0085 POH = -log10(.0085) = 2.07 14.00 - 2.07 = 11.93
What occurs when a salt is added into water (in terms of acids and bases)
The ions present when a salt is added to water may themselves be either acids or
bases and consequently affect the pH of the solution.
Give an example of a salt added to water and what occurs?
- ammonium chloride is added to water, it provides both an acid (NH4+) and a base (Cl-).
The solution consists of a weak acid (NH4+) and a very weak base (Cl-).
The net effect is that the solution is acidic
-a solution of sodium acetate consists of an ion that is neither acidic nor basic (the Na+ion) and a base (CH3CO2-).
The net effect is that the solution is basic, and its pH is greater than 7.
What is a polyprotic acid?
is a molecular compound that can donate more than one proton, best considered to be a molecular species that can give rise to a series of acids as it donates its succession of protons.
Give an example of a polyprotic acid in biology
Enzymes are polyprotic acids, for they possess many protons that can be donated to a substrate molecule or to the surrounding aqueous medium of the cell.
Give chemical examples of polyprotic acids and state why they’re polyprotic
Sulfuric acid is the parent of two acids, H2SO4 itself and HSO4-
Phosphoric acid is the parent of three acids, namely H3PO4, H2PO4- and HPO4 2-
When dealing with polyprotic acids, what must we consider in terms of it’s ka?
Its first and second deprotonation
example : H2A + H20 -> H30+ + HA-
1) [H30+] [HA-] / [H2A]
2) [H30+] [A2-] / [HA-]
What are Amphiprotic systems?
Can either donate or accept a proton (H+).
Give an example of an amphiprotic system
Examples include amino acids, amino acids, which can
act as proton donors by virtue of their carboxyl groups and as bases by virtue of their amino groups.
For instance, HCO3 − can act as an acid (to form CO3 2−) and as a base (to form H2CO3).
How do you find the PH of an amphiprotic system?
PH = 1/2 (pKa1 + pKa2)
What is a buffer solution?
is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa.
What is the function of buffer solutions?
It is able to neutralize small amounts of added acid or base, thus maintaining the pH of the solution relatively stable. This is important for processes and/or reactions which require specific and stable pH ranges.
What is the henderson-hasselbach equation?
PH = pKa - log [acid]
[base]
What is the equation when the concentrations of the conjugate acid and base are equal (at the half-stoichiometric point)
the second term on the right of the equation is log 1 =0, so under these conditions: pH =pKa
What is an acid buffer and how is it prepared?
An acid buffer stabilizes the solution at a pH below 7 is typically prepared by making a solution of a weak acid (such as acetic acid) and a salt that supplies its conjugate base (such as sodium acetate).
What is a base buffer and how is it prepared?
A base buffer stabilizes a solution at a pH above 7 is prepared by making a solution of a weak base (such as ammonia) and a salt that supplies its conjugate acid (such as ammonium chloride).
Give an example of a biological buffer
Physiological buffers are responsible for maintaining the pH of blood within a narrow range of 7.37 to 7.43, thereby stabilizing the active conformations of biological macromolecules and optimizing the rates of biochemical reactions
How does an acid/base buffer stabilise the solution?
-An acid buffer stabilizes the pH of a solution because the abundant supply of Aions (from the salt) can remove any H3O+ ions brought by additional strong acid;
- At the same time, the abundant supply of HA molecules (from the acid component of the buffer) can provide H3O+ ions to react with any strong base that is added.
- Similarly, in a base buffer the weak base B can accept protons when a strong acid
is added
- Its conjugate acid BH+ can supply protons if a strong base is added.
What is the most significant difference between the solution of an electrolyte and a nonelectrolyte?
there are long-range Coulombic interactions between the ions in electrolytes
Why do electrolyte solutions exhibit non-ideal behaviour even at very low concentrations?
because the solute particles, the ions, do not move independently of one another
The thermodynamic tendency to transport a species A through a biological cell membrane is partially determined by what?
the concentration gradient across the membrane, which results in a difference in molar Gibbs energy between the inside and the outside of the cell
RT In [A]in/[A]out
When is transport into the cell of either neutral or charged species thermodynamically favorable?
if [A]in < [A]out.
List redox reactions in biology
Electron transfer in plant photosynthesis and the oxidative breakdown of glucose are all redox processes.
What is an electrochemical cell?
A device that consists of two electronic conductors (metal or graphite, for instance) dipping into an electrolyte (an ionic conductor), which may be a solution, a liquid, or a solid.
A reaction is reversible because:
(a) reactants are reactive
(b) products are reactive
(c) products are stable
(d) reactants are stable
(b) products are reactive
A large value of Kc means that at equilibrium:
(a) less reactants and more products
(b) more reactants and less product
(c) same amount
(d) none
(a) less reactants and more products
Which statement about the following equilibrium is correct 2SO2 (g) + O2 (g) ⇄ 2SO3 (g) ΔH = – 188.3 kJ mol–1
(a) the value of Kp falls with a rise in temp (move forward = eq constant will increase)
(b) the value of Kp falls with increasing pressure (3 moles of gas and 2 moles)
(c) adding V2O5 catalyst increase the equilibrium yield of sulphur trioxide (don’t care about catalysts)
(d) the value of Kp is equal to Kc (same equilibrium constant, but its just expressed as a pressure or concentration)
(b) the value of Kp falls with increasing pressure (3 moles of gas and 2 moles)
In a chemical reaction, equilibrium is said to have established When:
(a) opposing reactions stops
(b) concentrations of reactants and products are equal
(c) rate constants of opposing reactions are equal
a) (In equilibrium, NOTHING STOPS)
b) NOT THE CASE, you can have more products or more reactants
Answer = C)
Which one of following solution have zero pH:
(a) 1M HCl
(b) 0.5 M H2SO4
(c) 0.1 M HNO3
(d) 1M CH3COOH
- log10 (concentration of H+)
a) -log10(.1) = 0
Solubility of Ca(OH)2 is exothermic. Solubility will increase:
(a) at high temp
(b) at low temp
(c) temp independent
(d) none
(b) at low temp (exothermic!)
Strength of an acid can be determined by:
(a) pKa
(b) pKp
(c) pOH
(d) pKw
(a) pKa
not b) = (p = -log10 of equilibrium constant expressed as a function of pressure)
not d) (equilibrium constant for the dissociation of water)
In an exothermic reversible reaction increase in temp shifts the equilibrium to:
(a) reactant side
(b) product side
(c) remains unchanged
(d) none of the above
(a) reactant side
The extent to 2HI ⇄ H2 + I2 (DH = 51.0 kJ) can be increased by: (amount of products you can get)
(a) increasing pressure
(b) increasing product
(c) increasing temp
(d) adding a catalyst
(c) increasing temp
not a) ( bc moles is same)
not b) (no because reactants would increase)
not d) (catalyst doesn’t change rate)
The solution having zero pH will be:
(a) basic
(b) high basic
(c) neutral
(d) highly acidic
(d) highly acidic
A solution have H+ ions concentration 1 x 10–7 its pH will be: (a) acid (b) basic (c) neutral (d) zero
(c) neutral
Which pH is considered as basic:
(a) 1
(b) 7
(c) 2
(d) 11
(d) 11
pH of pure water is:
(a) 3.2
(b) 4.2
(c) 7.0
(d) 0
(c) 7.0
Which of following change will favour the formation of more SO3 at equilibrium: 2SO2 + O2 ⇄ 2SO3 + heat (a) by adding 2SO3 at equilibrium (b) by increasing temp (c) by decreasing temp (d) by decreasing pressure
(c) by decreasing temp (heat is released)
Which of following change will favour the formation of more HI in the given reaction (DH = -51.0 kJ): H2 + I2 ⇄ 2HI (a) increasing pressure (b) decreasing pressure (c) by adding more HI (d) by adding more H2 and I
(d) by adding more H2 and I
