Part 2 Acids and Bases Topics we still

Part 2 – Acids and Bases • Topics we still need to remember and review and WILL BE ON THIS TOPIC 18 EXAM: • p. H and p. OH • [H+] and [OH-] • Strong acids and bases • Conjugate pairs

18. 1 Lewis acids and bases • A Lewis acid is a lone pair acceptor and a Lewis base is a lone pair donor. • When a Lewis base reacts with a Lewis acid a coordinate bond is formed. • A nucleophile is a Lewis base and an electrophile is a Lewis acid

Lewis Acid/Base Chemistry • Lewis acids are electron pair acceptors. • Lewis bases are electron pair donors. • All Brønsted–Lowry acids and bases are also called Lewis acids and bases. • There are compounds which do not meet the Brønsted–Lowry definition which meet the Lewis definition.

Lewis acids + bases • Coordinate covalent bond- formed between a Lewis acid and Lewis base when both electrons in a covalent bond are from the Lewis base

More examples

Comparing Ammonia’s Reaction with H+ and BF 3

Nucleophiles • Nucleophile (“likes nucleus”) = electron-rich species that DONATES a lone pair to form a new (coordinate) covalent bond (Lewis BASE)

Electrophile • Electrophile (“likes electrons”) is an electron-deficient species that ACCEPTS a lone pair from another reactant to form a new (coordinate) covalent bond (Lewis ACID)

18. 2 Calculations involving acids and bases • The expression for the dissociation constant of a weak acid (Ka) and a weak base (Kb) • For a conjugate acid-base pair, pair Ka x Kb = Kw • The relationship between Ka and p. Ka is p. Ka = -log. Ka and between Kb and p. Kb is p. Kb = -log. Kb.

Kw is temperature dependent Write the equation for the auto-ionization of water, then write the expression for the equilibrium constant

What’s up with water?

p. H and p. OH scales are inter-related

Summary

Dissociation constants • Weak acids and bases DO NOT fully dissociate in solution • You have to write the equilibrium expression and solve for the [H+] or [OH-] • Ka = acid dissociation constant • Kb = base dissociation constant

Generic Examples

Important Note • The values of Ka and Kb are constant for a specific acid and base respectively, at a specified temperature. Their values give a measure of the strength of the acid or base. • The higher the value of Ka or Kb at a particular temperature, the greater the ionization and so the acid or base is stronger.

Calculations with Ka and Kb

RICE box time Write expression and solve

p. H + p. OH = 14 p. OH = 2. 20 [OH-] = 10 -2. 20= 6. 3 x 10 -3 RICE box time Write expression and solve

RICE box time Write expression and solve

RICE box time Write expression and solve

p. Ka and p. Kb • To make the numbers easier to deal with, we convert the Ka and Kb values into their negative logarithms to the base 10, p. Ka and p. Kb.

Notes about p. Ka and p. Kb • p. Ka and p. Kb numbers are usually positive and have no units • The relationship beween Ka and p. Ka and between Kb and p. Kb is inverse • A change of one unit in p. Ka or p. Ka represents a ten-fold change in the value of Ka or Kb • p. Ka and p. Kb must be quoted at a specific temperature. LIKE p. H!!!!!!

Conjugate Acid/Base Pairs • The weaker an acid is, the lower its Ka value is, which means its Kb value is higher and its conjugate base is stronger. • Weak acid = strong conjugate base • Similarly, the weaker a base is, the stronger its conjugate acid. The stronger a base is, the weaker its conjugate acid.

Conjugate Acid/Base Pairs

18. 3 p. H curves • The characteristics of the p. H curves produced by the different combinations of strong and weak acids and bases. • An acid-base indicator is a weak acid or a weak base where the components of the conjugate acid-base pair have different colours. • The relationship between the p. H range of an acid-base indicator, which is a weak acid, and its p. Ka value. • The buffer region on the p. H curve represents the region where small additions of acid or base result in little or no change in p. H. • The composition and action of a buffer solution.

Buffer solutions • A buffer solution is resistant to changes in p. H on the addition of small amounts of acid or alkali • (non buffer solution below)

Buffers • Solutions of a weak conjugate acid–base pair that resist drastic changes in p. H are called buffers. • These solutions contain relatively high concentrations (10– 3 M or more) of both the acid and base. Their concentrations are approximately equal.

How do buffers work? • Two main types • Acidic buffers maintain the p. H below 7 • Basic buffers maintain the p. H above 7 • Mixture of two solutions that contain both species in a conjugate acid -base pair so that they can react and neutralize either acids or bases added to the solutions • For example….

Acidic buffers • Made by mixing an aqueous solution of a weak acid with a solution of its salt of a strong alkali • Therefore, the mixture has a relatively high concentrations of both CH 3 COOH and CH 3 COO-, an acid and its conjugate base, which are ready to react and neutralize added OH- and H+

Basic buffers • Made by mixing an aqueous solution of a weak base with a solution of its salt of a strong acid • Therefore, the mixture has a relatively high concentrations of both CH 3 COOH and CH 3 COO-, an acid and its conjugate base, which are ready to react and neutralize added OH- and H+

• Can you write equations to show these species would neutralize any added H+ or OH-?

Making buffer solutions • The p. H of a buffer is determined by the interactions of its components. Specifically, the p. Ka or p. Ka of the acid/base and the ratio of the initial concentrations of acid and salt, or base and salt, used in its preparation. • Bison

Ways to Make a Buffer 1) Mix a weak acid and a salt of its conjugate base or a weak base and a salt of its conjugate acid. 2) Add strong acid and partially neutralize a weak base or add strong base and partially neutralize a weak acid.

How a Buffer Works • Adding a small amount of acid or base only slightly neutralizes one component of the buffer, so the p. H doesn’t change much.

Calculating the p. H of a Buffer • For a weak acid: Ka = [H+][A–]/[HA] • Take –log of both sides: –log Ka = –log[H+] + –log([A–]/[HA]) • Rearrange: –log[H+] = –log Ka +log([A–]/[HA]) • Which is: p. H = p. Ka + log([A–]/[HA]) • This equation is known as the Henderson– Hasselbalch equation. This applies only to buffers.

Henderson–Hasselbalch Equation What is the p. H of a buffer that is 0. 12 M in lactic acid, CH 3 CH(OH)COOH, and 0. 10 M in sodium lactate? Ka for lactic acid is 1. 4 × 10 4. Øp. H = p. Ka + log([A–]/[HA]) = –log(1. 4 × 10– 4) + log[(0. 10 M)/(0. 12 M)] = 3. 85 + (– 0. 08) = 3. 77

Buffer Capacity • The amount of acid or base the buffer can neutralize before the p. H begins to change to an appreciable degree • Using the Henderson–Hasselbalch equation, p. H will be the same for a conjugate acid–base pair of 1 M each or 0. 1 M each; however, the buffer which is 1 M can neutralize more acid or base before the p. H changes. • 2 Days Older

p. H Range • The range of p. H values over which a buffer system works effectively • Optimal p. H: where p. H = p. Ka([HA] = [A–]) • If one concentration is more than 10 times the other, the buffering action is poor; this means that the p. H range of a buffer is usually ± 1 p. H unit from p. Ka.

Addition of a Strong Acid or a Strong Base to a Buffer 1) Adding of the strong acid or base is a neutralization reaction; calculate like stoichiometry problem to find [HA] and [A–] when all of the added acid or base reacts. 2) Use the Henderson–Hasselbalch equation to find p. H. © 2015 Pearson Education, Inc.

Salt hydrolysis • Salt of a strong acid + strong base = no hydrolysis • Salt of a weak acid + strong base = anion hydrolysis • Salt of a strong acid + weak base = cation hydrolysis • Salt of a weak acid + weak base = depends on Ka and Kb of species involved

Salt hydrolysis summary

Titration • In this technique, an acid (or base) solution of known concentration is slowly added to a base (or acid) solution of unknown concentration. • A p. H meter or indicators are used to determine when the solution has reached the equivalence point: The amount of acid equals that exactly neutralizes the base. © 2015 Pearson Education, Inc.

Titration of a Strong Acid with a Strong Base • From the start of the titration to near the equivalence point, the p. H goes up slowly. • Just before (and after) the equivalence point, the p. H rises rapidly. • At the equivalence point, p. H = 7. • As more base is added, the p. H again levels off. © 2015 Pearson Education, Inc.

Titration of a Strong Base with a Strong Acid • It looks like you “flipped over” the strong acid being titrated by a strong base. • Start with a high p. H (basic solution); the p. H = 7 at the equivalence point; low p. H to end.

Titration of a Weak Acid with a Strong Base • Use Ka to find initial p. H. • Find the p. H in the “buffer region” using stoichiometry followed by the Henderson–Hasselbalch equation. • At the equivalence point the p. H is >7. Use the conjugate base of the weak acid to determine the p. H. • As more base is added, the p. H levels off. This is exactly the same as for strong acids.

Ways That a Weak Acid Titration Differs from a Strong Acid Titration 1) A solution of weak acid has a higher initial p. H than a strong acid. 2) The p. H change near the equivalence point is smaller for a weak acid. (This is at least partly due to the buffer region. ) 3) The p. H at the equivalence point is greater than 7 for a weak acid.

Use of Indicators • Indicators are weak acids that have a different color than their conjugate base form. • Each indicator has its own p. H range over which it changes color. • An indicator can be used to find the equivalence point in a titration as long as it changes color in the small volume change region where the p. H rapidly changes.

Indicator Choice Can Be Critical!

Indicators summary

1. strong acid and strong base

2. Weak acid and strong base

3. Strong acid and weak base

4. Weak acid and weak base

8. 5 Acid Deposition • Rain is naturally acidic because of dissolved CO 2 and has a p. H of 5. 6. Acid deposition has a p. H below 5. 6. • Acid deposition is formed when nitrogen or sulfur oxides dissolve in water to form HNO 3, HNO 2, H 2 SO 4, and H 2 SO 3. • Sources of the oxides of sulfur and nitrogen and the effects of acid deposition should be covered.

Causes of acid deposition • All rain water is naturally acidic due to the presence of dissolved carbon dioxide (carbonic acid). • Acid rain refers to rain water with a p. H below 5. 6 due to the additional presence of acidic pollutants. • Acid deposition = all of the precipitation • Wet acid deposition- rain, snow, sleet, hail, fog, mist, dew fall to ground as aqueous precipitates • Dry acid deposition = acidifying particles or gases fall to the ground as dust or smoke and later dissolve in water to form acids.

Sulfur oxides • Produced from burning fossil fuels, particularly coal and heavy oil in power plants. Also smelting metals.

Nitrogen oxides • Produced mainly from internal combustion engines. Burning the fuel provides the heat energy necessary for nitrogen and oxygen in air to combine.

Impact on materials • (write some equations, list the effects, and the impact)

What happens to plants?

What happens to water?

What happens to human health?

List some things that are being done to help reduce these problems.