The Arrhenius Theory Acid - a substance that produces H+ ion

when dissolved in water; Base - a substance that produces OH- ion when dissolved in water; Strong acid - ionizes essentially completely into H+ and an anion;

Strong base - dissociates nearly completely

into OH- and a cation; Weak acid and base - ionize reversibly; Limitation of Arrhenius Theory.

The Brnsted-Lowry Theory Acid proton donor Base proton acceptor Conjugate acid and base, HA/A-, differ by

one proton. The conjugate acid of a base is the base plus the attached proton (base + H+). The conjugate base of an acid is the acid minus the proton (acid H+).

A substance that can act either as an acid or

a base is amphiprotic. For weak acids and bases, equations can be written to describe equilibrium conditions.

Ionization of Ammonia

Strengths of Conjugate Acid-Base Pairs The stronger an acid, the weaker is its

conjugate base. The stronger a base, the weaker is its conjugate acid. An acid-base reaction is favored in the direction from the stronger member to the weaker member of each conjugate acid-base pair.

Ka and Kb values are used to compare the

strengths of weak acids and bases. Water has a leveling effect; when the strong acids are dissolved in water, they all completely ionize to the hydronium ion.

Relative AcidBase Pair Strength

Acid and Base Ionization Constantsweak acid: CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)[H3O+][CH3COO-] Acid ionization constant: Ka =

[CH3COOH]weak base: NH3 (aq) + H2O (l) NH4 + (aq) + OH- (aq) [NH4+][OH-] Base ionization constant: Kb = [NH3]

Acid and base ionization constants are the

measure of the strengths of acids and bases.

Relative Strengths of Binary AcidsHX The greater the tendency for the transfer of a proton from HX to H2O, the more the forward reaction is favored and the stronger the acid.

In a periodic group: Bond-dissociation energy is inversely proportional to acid strength. The weaker the bond, the stronger the acid. Anion radius is directly proportional to acid strength. The larger the resultant anions radius, the stronger is the acid.

The strengths of binary acids increase from

top to bottom in a group of the periodic table.

Relative Strengths of Binary Acids HXIn a periodic group: Bond dissociation energy: the weaker the bond, the stronger the acid.Bond dissociation energy 569 > 431 > 368 > 297 (kJ/mol) HF HCl HBr HI Acid strength Ka 6.6x10-4 < ~106 < ~108 < ~109

Anion radius: the larger the anions radius, the stronger the acid.Anion radius (ppm) (kJ/mol) Acid strength Ka 136 < 181 < 195 < 216 HF HCl HBr HI 6.6 x 10-4 < ~106 < ~108 < ~109

The strength of binary acids increase from top to bottom in a group of the periodic table.

Relative Strengths of Binary Acids H XIn a period:

Electronegativity difference between H and

X () EN) The larger the electronegativity difference between H and X, the more easily the proton is removed and the stronger is the acid.EN Acid strength 0.4 < CH4 < 0.9 NH3 < 1.4 < < H2O < 1.9 HF

The strengths of binary acids increase from

left to right across a period of the periodic table.

Representative Trends in Strengths of Binary Acids

Strengths of OxoacidsHOE Two factors: electronegativity of the central atom (E); number of terminal oxygen atoms. As the electronegativity of the central atom (E) increases and as the number of terminal oxygen atoms increases, the acid strength also increases.

Strengths of OxoacidsAs the electronegativity of the central atom (E) increases the acid strength increases.ElectronegativityAcid strength Ka2.5 HOI < 2.8 < 3.0 HOBr HOCl

2.3 x 10-11 < 2.5 x 10-9 < 2.9 x 10-8

As the number of terminal oxygen atoms increases, the acid strength also increases.# of terminal O atoms 0H-O-Cl

1H-O-Cl-O

2 O

3 O

H-O-Cl-O H-O-Cl-O

Acid strength Ka

O 2.9 x 10< 1.1x10-2 < ~1000 < ~10

Strengths of Carboxylic AcidsR

O C O

H

Carboxylic acids all have the -COOH group in common. Therefore, differences in acid strength must come from differences in the R group attached to the carboxyl group.

In general, the more that electronegative

atoms are attached in the R group, the stronger the acid.

Strengths of Carboxylic AcidsIn general, the more that electronegative atoms are attached in the R group, the stronger the acid.I-CHCHCOOH Cl-CHCHCOOH CH-CHClCOOH CHCClCOOH

Ka

8.3x10-5

7.0

The pH Scale

Equilibrium in Solutions of Weak Acids and Weak Bases Weak acid: HA (aq) + H2O (l) H3O+ (aq) + A- (aq)[H3O+][A-] Ka = [HA]

Weak base: B (aq) + H2O (l) HB+ (aq) + OH- (aq)[HB+][OH-] Kb = [B]

You need to be able to write acid and base ionization equations!!!

pKa and pKbpKa = -logKa pKb = -logKblarger Ka => smaller pKa => stronger acid larger Kb => smaller pKb => stronger base

Some Acid-Base Equilibrium Calculations An equation is written for the reversible

reaction, data are organized under this equation, the changes that occur in establishing equilibrium are assessed, and finally calculations of equilibrium concentrations are done. When Macid/Ka > 100 or Mbase/Kb > 100, the calculations can be simplified.

Acid-Base Properties of Salt SolutionsSalts are ionic compounds formed in the reaction between an acid and a base. Are aqueous solutions of salts acidic, basic, or neutral?Let us consider three salts: NaCl, KCN, NH4Cl

Acid-Base Properties of Salt Solutions1. NaCl

Na+ is from NaOH , a strong base Cl- is from HCl, a strong acid;H2O

NaCl (s) Na+ (aq) + Cl- (aq)

Na+ and Cl- ions do not react with water. The solution is neutral.

2. KCN

Acid-Base Properties of Salt SolutionsH2O

K+ is from KOH , a strong base CN- is from HCN, a weak acid

KCN (s) K+ (aq) + CN- (aq) K+ ions do not react with water, but CN- ions do. CN- (aq) + H2O (l) HCN (aq) + OH- (aq) hydrolysisThe OH- ions are produced, so the solution is basic.

Acid-Base Properties of Salt Solutions3. NH4ClNH4 + is from NH3 , a weak base Cl- is from HCl, a strong acid H2O NH4Cl (s) NH4 + (aq) + Cl- (aq) Cl ions do not react with water, but NH4 + ions do. NH4 + (aq) + H2O (l) H3O+ (aq) + NH3 (aq) hydrolysis The H3O+ ions are produced, so the solution is acidic.

HydrolysisThe hydrolysis of an ion is the reaction of an ion with water to produce the conjugate acid and hydroxide ion or the conjugate base and hydrogen ion.You need to be able to write equation for hydrolysis reaction!

Ions As Acids and BasesCertain ions can cause an aqueous solution to become acidic or basic due to hydrolysis. Salts of strong acids and strong bases form

neutral solutions. Salts of weak acids and strong bases form basic solutions.

Salts of strong acids and weak bases form

acidic solutions. Salts of weak acids and weak bases form solutions that are acidic in some cases, neutral or basic in others.

Strong Acids and Strong BasesStrong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4Strong bases: Group IA and IIA hydroxides - Memorize!!

Prediction of Whether a Salt Solution is Acidic, Basic or Neutral - An ExampleIndicate whether the solutions (a) Na2S and (b) KClO4 are acidic, basic or neutral.

The pH of a Salt SolutionWhat is the pH of 1.0 M NaCN solution?Hydrolysis of CN- ions: CN- + H2O HCN + OHCN-is a conjugate base of HCN. Ka of HCN can be found. What is Kb for CN-? Ka x Kb = Kw so, Kb = Kw/Ka

Common Ion Effect Illustrated

The Common Ion EffectCH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO- (aq)adding CH3COOequilibrium shifts to left Le Chteliers Principle

The Common Ion Effect If one solution contains a weak acid and

another contains the same acid and its conjugate base as a second solute, the two solutions have different pH values. The solution containing both the weak acid and its conjugate base has a pH much higher than the solution containing only the weak acid.

The conjugate base is referred to as a

common ion because it is found in both the weak acid and the anion. The common ion effect is the suppression of the ionization of a weak acid or a weak base by the presence of a common ion from a strong electrolyte.

The Common Ion Effect - An ExampleCalculate the pH of 1.00 M CH3COOH - 1.00 M CH3COONa solution. Comparison: 1.00 M CH3COOH pH = 2.38 1.00 M CH3COOH -1.00 M CH3COONa pH = 4.74 The common ion effect!!!

Depicting Buffer Action

Buffer Solutions A buffer solution is a solution that changes

Ph only slightly when small amounts of a strong acid or a strong base are added. A buffer contains (i) a weak acid with its salt (conjugate base) or (ii) a weak base with its salt (conjugate acid) CH3COOH/CH3COONa NH3/NH4Cl

How a Buffer Solution Works The acid component of the buffer can

neutralize small added amounts of OH-, and the basic component can neutralize small added amounts of H3O+. Pure water does not buffer at all.

Henderson-Hasselbalch Equation for Buff Solutions[conjugate base] pH = pKa + log [weak acid] If [weak acid] = [conjugate base], pH = pKaRequirements: -The ratio of [conjugate base] to [weak acid] is between 0.10 and 10; -[conjugate base]/Ka > 100, [weak acid]/Ka > 100.

Buffer Capacity and Buffer Range There is a limit to the capacity of a buffer

solution to neutralize added acid or base, and this limit is reached before all of one of the buffer components has been consumed. In general, the more concentrated the buffer components in a solution, the more added acid or base the solution can neutralize.

As a rule, a buffer is most effective if the

concentrations of the buffer acid and its conjugate base are equal.

Acid-Base Indicators An acid-base indicator is a weak acid having

one color and the conjugate base of the acid having a different color. One of the colors may be colorless. HIn + H2O H3O+ + In Color 1 Color 2

Acid-base indicators are often used for

applications in which a precise pH reading isnt necessary. A common indicator used in introductory chemistry laboratories is litmus.

Several Common Indicators

Types of Calculations in Acid-Base Equilibria pH, pOH Kw= [H3O+][OH-]; pKw = pH + pOH Weak acid and weak base acid or base ionization,

Ka or Kb A salt aqueous solution hydrolysis, KaKb = Kw Common ion effect; Buffer

solution

HendersonHasselbalch

equation.

Neutralization Reactions Neutralization is the reaction of an acid and a

base. Titrationis a common technique for conducting a neutralization. At the equivalence point in a titration, the acid and base have been brought together in exact stoichiometric proportions. The point in the titration at which the indicator changes color is called the end point.

The indicator endpoint and the equivalence

point for a neutralization reaction can be best matched by plotting a titration curve, a graph of pH versus volume of titrant. In a typical titration, 50 mL or less of titrant that is 1 M or less is used.

Titration Curve for Strong Acid - Strong Base

Features of Titration Curve for Strong Acid - Strong Base pH is low at the beginning; pH changes slowly until just before equivalence

point; pH changes sharply around equivalence point; pH = 7.0 at equivalence point; Further beyond equivalence point, pH changes slowly; Any indicator whose color changes in pH range of 410 can be used in titration.

Titration Curve for Weak Acid - Strong Base

Features of Titration Curve for Weak Acid - Strong Base The initial pH is higher because weak acid is partially ionized. At the half-neutralization point, pH = pKa.

pH is greater than 7 at equivalence point because the anion of the weak acid hydrolyzes. The steep portion of titration curve around

equivalence point has a smaller pH range. The choice of indicators for the titration is more limited.

Types of Calculations in Acid-Base Equilibria pH, pOH Kw= [H3O+][OH-]; pKw = pH + pOH Weak acid and weak base; A salt aqueous solution hydrolysis, KaKb = Kw;

Common ion effect; Buffer

solution

HendersonHasselbalch

equation; Neutralization, titration curve.

Lewis Acids And Bases There are reactions in non-aqueous solvents,

in the gaseous state, and even in the solid state that can be considered acid-base reactions in which Brnsted- Lowry theory is not adequate to explain. A Lewis acid is a species that is an electron-pair acceptor and a Lewis base is a species that is an electron-pair donor.

In organic chemistry, Lewis acids are often

called electrophiles and Lewis bases are often called nucleophiles.

Summary In the Brnsted-Lowry theory an acid is a

proton donor and a base is a proton acceptor. If an acid is strong, its conjugate base is weak; and if a base is strong, its conjugate acid is weak. Water is amphiprotic: it can be either an acid or a base. It undergoes limited selfionization producing H3O+ and OH-.

pH = -log[H3O+] pOH = -log[OH-] pKw = -logKw The pH in both pure water and in neutral

solutions is 7. Acidic solutions have a pH less than 7 and basic solutions have a pH greater than 7. In aqueous solutions at 25 oC, pH + pOH = 14.00. Hydrolysis reactions cause certain salt solutions to be either acidic or basic.

A strong electrolyte that produces an ion

common to the ionization equilibrium of a weak acid or a weak base suppresses the ionization of the weak electrolyte. Acid-base indicators are weak acids for which the acid and its conjugate base have different colors.

In Lewis acid-base theory, a Lewis acid

accepts an electron pair and a Lewis base donates an electron pair.