chapter 17 principles of reactivity: chemistry of acids and bases

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Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases

Acids & Bases Acids are some of peoples' favorite chemicals.

Everyone's favorite soft drink is a dilute acid solution. Your own stomach contains the strong acid : HCl. Citrus fruits contain citric acid. If wine is too aged - exposed to oxygen, it turns sour - it forms acetic acid. Sulfuric Acid is the top commercially produced chemical in the United State. Although much of it is used in the steel and petroleum refining industries, several million tons of sulfuric acid are used to make Jello.

PROPERTIES OF ACIDS AND BASES

ACIDS BASES

Arrhenius definitionacid:

produces hydronium ion (H3O+) in aqueous solution

base:produces hydroxide ion (OH–) in aqueous solution

Brønsted definition

acid:base:

donates a proton (hydrogen ion, H+)accepts a proton

HF + H2O H3O+ + F–

acid conjugateacid

base conjugatebase

A conjugate acid is formed by adding a proton to something.

A conjugate base is formed by removing a proton from something.

NH3 + H2O NH4+ + OH–

base conjugateacid

acid conjugatebase

HA H3O+ + A–

If HA is a stronger acid then A– is a weaker base.If HA is a weaker acid then A– is a stronger base.

B BH+ + OH–

If B is a stronger base then BH+ is a weaker acid.If B is a weaker base then BH+ is a stronger acid.

Relative strengths of conjugate acid-base pairs

You should memorize the names and formulas of the 6 STRONG ACIDs , i.e., HCl, HBr, HI, HClO4, HNO3, and H2SO4.

The organic acid present in vinegar, acetic acid, is a common WEAK ACID.

The common STRONG BASES contain the hydroxide ion (OH-).

Ammonia (NH3), a common WEAK BASE, is that smelly stuff your Grandma used in a dilute solution to clean windows

Ion Product Constant of Water

Water is an important solvent.

Universal solvent Biological solvent Small size Density of water is greater than ice Very polar Hydrogen Bonding

Self-ionization of Water Water is an amphiproticamphiprotic substance that can act either as an

acid or a base.

HC2H3O2(aq) + H2O(l) H3O+ + C2H3O2-(aq)

acid base acid base

H2O(l) + NH3(aq) NH4+

(aq) + OH-(aq)

acid base acid base

Self-ionization of Water When water molecules react with one another to form ions.

H2O(l) + H2O(l) H3O+(aq) + OH-

(aq)

(10-7M) (10-7M)

Kw = [ H3O+ ] [ OH- ]

= 1.0 x 10-14 at 25oC

Note:Note: [H2O] is constant and is already included in Kw.

ion productof water

ion productof water

pH and pOH We need to measure and use acids and bases over a

very large concentration range.

pH and pOH are systems to keep track of these very large ranges.

pH = - log [H3O+]

pOH = - log [OH-]

pH + pOH = 14

pH Calculations

Determine the following. pH = - log [H+] pH of 6.7x10-3 M H+

= 2.2 pH of 5.2x10-12 M H+

= 11.3 [H+], if the pH is 4.5 = 3.2 x 10-5 M H+

pOH Examples

Determine the following. pOH = - log [OH-] = 14 - pH

pOH of 1.7 x 10-4 M NaOH pOH = 3.8 pH = 10.2 pOH of 5.2 x 10-12 M H+

pOH = 2.7 pH = 11.3 [OH-] , if the pH is 4.5 pOH = 9.5 [OH-] = 3.2 x 10-10 M

pH Scale

A log based scale used to keep track of the large change important to acids and bases.14 7 0

10-14 M 10-7 M 1 MVery Neutral VeryBasic Acidic

When you add an acid, the pH gets smaller.

When you add a base, the pH gets larger.

pH of SomeCommon Materials Substance pH

1 M HCl 0.0

Lemon juice 2.3

Coffee 5.0

Pure Water 7.0

Blood 7.35 - 7.45 Milk of Magnesia 10.5

1M NaOH 14.0

HA + H2O H3O+ + A–

Kc =[H3O+][A–][HA][H2O]

~ constant (55 M)

Kc·[H2O] = Ka = [H3O+][A–]

[HA]

pKa = -log Ka if pKa = 5 then Ka = 10–5

if pKa = 8 then Ka = 10–8

stronger acidweaker acid

acid dissociation constant

Definitions Ka, pKa

Acid Ionization Constant, Ka

Acid ionization constants let us define weak, moderate and strong acids.

Ka < 10-3; it is a weak acid.

Ka = 10-3 to 1; it is a moderate acid.

Ka > 1; it is a strong acid.

B + H2O BH+ +OH–

Kc =[BH+][OH–]

[B][H2O] ~ constant

Kc·[H2O] = Kb = [BH+][OH–]

[B]

pKb = -log Kb if pKb = 4 then Kb = 10–4

if pKb = 9 then Kb = 10–9

stronger baseweaker base

base dissociation constant

Definitions Kb, pKb

Ka and Kb Values For weak acids and bases:

Ka and Kb always have values that are smaller than one.

• Acids with a larger Ka are stronger than ones with a smaller Ka.

• Bases with a larger Kb are stronger than ones with a smaller Kb.

• Ka x Kb = Kw

• Most acids and bases are considered weak.

pKa and pKb Concepts The negative logarithms of Ka and Kb are useful in the

same way as pH.

pKa = - log Ka

pKb = - log Kb

pKa + pKb = 14.00

The larger that the value of pKa is, the weaker the acid. The larger that the value of pKb is, the weaker the base.

H2O + H2O H3O+ + OH–

Kw = [H3O+][OH–] = 10–14 (constant at 25ºC)

pKw = pH + pOH = 14 [H3O+] =10–14

[OH–]pH = 14 - pOH

[H3O+] pH [OH–] pOH

neutral 10–7 M 7 10–7 M 7

acidic

basic

Kw: autodissociation of water

HA H3O+ + A– Ka = [H3O+][A–]

[HA]

A– HA + OH– Kb = [HA][OH–]

[A–]

usually only one or the other given in a table

Ka · Kb =[H3O+][A–]

[HA]·

[HA][OH–][A–]

= [H3O+][OH–] = Kw

Ka · Kb = Kw

for a conjugateacid-base pair

or Ka =or Kb =Kw

Ka

Kw

Kb

Ka and Kb for conjugate acid-base pairs

III. pH CalculationsA. Strong acids and bases

100% dissociated

for strong acid: [H3O+]eq = [HA]I

base: [OH–]eq = [B]i

e.g., 1.0 x 10–3 M HCl

[H3O+] = pH =

[OH–] = pOH =

e.g., 2.5 x 10–2 M NaOH

[OH–] = pOH =

[H3O+] = pH =

[OH–] -1.60

+ 14

log

pH = 12.40

III. pH CalculationsB. Weak acids and bases

HA H3O+ + A– Ka = [H3O+][A–]

[HA]

B BH+ + OH– Kb = [BH+][OH–]

[B]

Solve equilibriumexpressions

e.g., What is the pH of 0.10 M HC2H3O2? (Ka = 1.8 x 10–5)

HC2H3O2 H3O+ + C2H3O2–

1.8 x 10–5 =x2

(0.10 - x)

assume x << 0.10

1.8 x 10–5 =x2

(0.10)

x = [H3O+] = 1.3 x 10–3 M

(assumption valid)

pH = 2.87

If [HA]i 400·Ka then x << [HA]i

(If not, then have to solve quadratic.)

Buffers Solutions that resist change in pH when small amounts of

acid or base are added.

Two types:Two types:• weak acid and its salt.• weak base and its salt.

HA(aq) + H2O(l) H3O+(aq) + A-

(aq)

Add OH- Add H+

shift to right shift to left

Based on Le Châtelier’s Principle.

III. pH CalculationsC. Polyprotic acids

H2SO4, H2SO3, H2CO3, etc.

H2A H3O+ + HA– Ka1 = [H3O+][HA–]

[H2A]

HA– H3O+ + A2– Ka2 = [H3O+][A2–]

[HA–]

(Usually, Ka1 >> Ka2)

Assume:1) [H2A], [H3O+], and [HA–] can be determined from the 1st step.(i.e., HA– dissociates only very little.)

2) [A2–] can be determined from the 2nd step.

Lose their protons in separate steps:

Buffers and Blood

Control of blood pH.Control of blood pH.

• Oxygen is transported primarily by hemoglobin in the red blood cells.

• CO2 transported both in plasma and the red blood cells.

CO2 (aq) + 2 H2O H2CO3 (aq)

H3O+(aq) + HCO3

-(aq) Carbonate BufferCarbonate Buffer

Buffers and Blood The amount of CO2 helps control blood pH.

Too much COToo much CO22 - Respiratory arrest, pH goes down, acid level goes up. acidosisacidosis SolutionSolution - ventilate and give bicarbonate

via IV.

Too little COToo little CO22 - Hyperventilation, anxiety, pH goes up, acid level goes down. alkalosisalkalosis SolutionSolution - re-breathe CO2 in paper bag

to raise level.

Quantitative Aspects of Buffers

Ka for a weak acid:

HA H+ + A-

Ka = [H+] [A-] [HA]

Henderson-Hasselbalch Equation:Henderson-Hasselbalch Equation:

pH = pKa + log [anion] [acid]

Neutralization

The reaction of an acid with a base to produce a salt and water.

HCl + NaOH NaCl + H2O

We do this when we use antacids.

Neutralization can be used to determine the amount of acid or base in a sample. - titrations- titrations

Titrations Analytical methods based on measurement of volume.

If the concentration of an acid is known, the concentration of the base can be found.

If we know the concentration of the base, then we can determine the amount of acid.

All that is needed is some calibrated glassware and either an indicator or pH meter.

TItrationsBuretBuret - volumetric glassware used for titrations.

It allows you to add a known amount of your titrant to the solution you are testing.

If a pH meter is used, the equivalence point can be measured.

An indicator will give you the endpoint.

Indicator Examples Acid-base indicators are weak acids that undergo a

color change at a known pH.

bromothymol blue

phenolphthalein methyl red