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27 Cards in this Set
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- Back
- 3rd side (hint)
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Dynamic equilibrium |
Many reactions are readily reversible and can reach a state of dynamic equilibrium in which: ■ the rate of the forward reaction is equal to the rate of the backward reaction ■ the concentrations of reactants and products remain constant |
Only happens in a closed system Temperature and pressure must be constant |
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Types of Equilibria |
Homogenous equilibria - system in which everything is in the same physical state Heterogeneous equilibria - system in which everything is not in the same physical state |
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Catalysts |
They don't influence the concentrations of reactants or products (yield) but increase the rate of reaction so equilibrium position is reached quicker. They don't affect the position of the equilibrium, so have no affect on the equilibrium constants. |
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Le Chatelier's principle |
When a system is subjected to a change in conditions, the equilibrium position will shift to counteract the change and restore equilibrium |
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Le Chatelier's principle: concentration |
Increasing concentration of reactants: ▪ equilibrium position shifts in forward direction ▪ increases rate of forward reaction ▪ reducing concentration of reactants Decreasing concentration of reactants: ▪ equilibrium position shifts in reverse direction ▪ increases rate of backwards reaction ▪ increasing concentration of reactants |
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Le Chatelier's principle: pressure |
Increasing pressure: ▪ equilibrium position shifts in direction with fewer gas molecules ▪ increases rate of that reaction ▪ fewer molecules reduces pressure Decreasing pressure: ▪ equilibrium position shifts in direction with more gas molecules ▪ increases rate of that reaction ▪ more molecules increases pressure |
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Le Chatelier's principle: temperature |
Increasing temperature: ▪ equilibrium position shifts in endothermic direction ▪ increases rate of that reaction ▪ more energy is absorbed ▪ reducing temperature Decreasing temperature: ▪ equilibrium position shifts in exothermic direction ▪ increases rate of that reaction ▪ more energy is released ▪ increasing temperature |
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Compromises |
Compromises are made between high yield, high rate and cost. • low temperature and pressure may increase yield but reduce rate of reaction • high temperature and pressure may increase yield and rate but are expensive to maintain |
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Kc |
Solids and pure liquids aren't included in the Kc expression as their concentrations remain constant throughout the reaction. A particular value of Kc will remain constant for a given temperature. |
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Kp |
Only gases are included in the expression, using their partial pressures at equilibrium. |
The total pressure of a gas mixture is the sum of the partial pressures of the individual gases. |
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Temperature and equilibrium constants |
● if changing the temperature favours the forward reaction, then the equilibrium position will shift to the right, causing more product to form and so the equilibrium constant will increase ● if changing the temperature favours the reverse reaction, then the equilibrium position will shift to the left, causing more reactant to form and so the equilibrium constant will decrease |
Only temperature affects the equilibrium constants, as it changes both the amounts of products and reactants at equilibrium and the ratio of products to reactants. Pressure and concentration affect the amounts of products and reactants at equilibrium, but the ratio of products to reactants remains the same, so equilibrium constants aren't affected. |
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Acids and bases |
A Brønsted–Lowry acid is a proton donor A Brønsted–Lowry base is a proton acceptor Acid-base reactions involve the transfer of protons, forming acid-base conjugate pairs |
Conjugate pairs are species linked by the transfer of protons |
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Strong and weak acids and bases |
Strong acids and strong bases completely dissociate in water Weak acids and weak bases partially dissociate in water |
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pH |
pH is a measure of the hydrogen ion concentration and can tell you the degree of dissociation, and thus the nature of substances dissolved in solution. pH = -log[H+] |
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pH of strong acids |
Strong, monoprotic acids release 1 proton per molecule of acid when they fully dissociate. So [H+] = [HA] Polyprotic acids can release more than 1 proton into solution per molecule when they dissociate. The lower the pH, the stronger the acid. |
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pH of weak acids |
Weak acids don't fully dissociate in solution, and instead form an equilibrium which lies to the far left. The equilibrium constant for a particular acid at a specific temperature is Ka. As only a tiny amount of HA dissociates, we can assume that: ● [HA] >> [H+], so [HA]start = [HA]eqb ● as the dissociation of acid >> dissociation of water, we can assume that [H+] = [A-]. The larger the Ka, the stronger the acid. |
The assumption that [HA]start = [HA]eqb doesn't apply to strong acids as they fully dissociate [HA] remaining is significantly different to that at the start |
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Kw |
Water can act as both a base (by accepting a proton to form H3O+) and an acid (by donating a proton to form OH-), so will form an equilibrium between hydroxide and hydroxonium ions, which lies to the far left since water only slightly dissociates. When writing an equilibrium expression, we can assume that: ● [H2O] >> [H+] ● [H2O] >> [OH-] ● [H2O] is constant, which, when multiplied by Kc gives Ka, the ionic product of water. Kw always has the same value for both pure water and aqueous solution at a given temperature. Under standard conditions, Kw = 1×10^-14 mol^2/dm^6 |
For pure water, [H+] = [OH-] due to the dissociation |
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pKw |
pKw = -log(Kw) Under standard conditions, it's 14 |
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pKa |
pKa = -log(Ka) The smaller the pKa, the stronger the acid. |
Finding pH at the half equivalence point of the titration curve for a weak acid/ Strong base titration. Half equivalence is the stage of a titration when half the equivalence volume of strong base has been added, and half the acid has been neutralized. At the half equivalence point, [HA] = [A-]. So for a weak acid: ● Ka = [H+] ● pKa = pH |
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Titration curves |
Initial and final pH depends on the strength of the acid/ base in the solution in the conical flask. All graphs have an equivalence point where [H+] = [OH-]. Neutralization occurs. The addition of reactant in burette causes rapid change in pH. Near the beginning and end of the graph, the addition of acid/base causes little change to pH. |
Weak acid, weak base titration curves don't have a vertical equivalence line, so there's no sharp pH change. There aren't any indicators that can be used in this case. |
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Choosing an indicator |
For a titration, an indicator must be chosen which changes colour over a narrow pH range which lies entirely within the equivalence line of the titration curve. |
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Buffer solution |
A solution that resists changes in pH when small amounts of acid or base are added. A buffer solution has to contain things which will remove any H+ ions or OH- ions that you might add to it, otherwise the pH will change. |
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Acidic buffers |
Buffer with a pH less than 7 and is made from setting up an equilibrium between a weak acid and it's conjugate base (from one of it's salts) of equal molar concentrations. The solution will therefore contain: ■ lots of undissociated acid, HA ■ lots of A- ions from the salt ■ enough H+ ions to make the solution acidic |
Adding acid increases H+ concentration: ● additional H+ combines with the A- to make HA ● equilibrium position shifts left ● H+ concentration falls close to normal, so pH doesn't change Adding base increases OH- concentration: ● additional OH- combines with the H+ to make H2O ● H+ concentration falls ● more HA dissociates to replace the H+ ● equilibrium position shifts right ● H+ concentration rises close to normal, so pH doesn't change |
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Alkaline buffers |
Buffer with a pH more than 7 and is made from setting up an equilibrium between a weak base and it's conjugate acid (from one of it's salts) of equal molar concentrations.The solution will therefore contain: ■ lots of unreacted base, B ■ lots of HB+ from the salt ■ enough OH- ions to make the solution alkaline |
Adding acid increases H+ concentration: ● additional H+ combines with the B to make HB+ ● equilibrium position shifts left ● H+ concentration falls close to normal, so pH doesn't change Adding base increases OH- concentration: ● additional OH- combines with the H+ to make H2O ● H+ concentration falls ● more HB dissociates to replace the H+ ● equilibrium position shifts right ● H+ concentration rises close to normal, so pH doesn't change |
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pH of a buffer |
We assume that: ● salt of an acid completely dissociates, so [A-] = [salt] ● HA only slightly dissociates, so [HA]initial = [HA]eqb |
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Blood and buffer solutions |
Blood pH needs to be kept around 7.4, so this is done using carbonic acid (H2CO3) which dissociates into hydrogen carbonate (HCO3) and H+. If H+ concentration increases in the blood: ● additional H+ combines with the HCO3 to make H2CO3 ● equilibrium position shifts left ● H+ concentration falls close to normal, so pH doesn't change If H+ concentration falls: ● more H2CO3 dissociates to replace the H+ ● equilibrium position shifts right ● H+ concentration rises close to normal, so pH doesn't change |
The levels of HCO3 are controlled by the kidneys by excretion, while the levels of H2CO3 are controlled by respiration by exhaling CO2, as H2CO3 dissociates as such: H2CO2 <-> H2O + CO2 |
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Titration curves and buffer activity |
For titrations of strong bases into weak acids or strong acids into weak bases: • at the start of the curve there's a rapid change in pH • the curve then levels off as a buffer solution is made which resists changes in the pH • past the equivalence point, the excess of strong acid or base uses up all the weak base/acid |
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