A.P. Chemistry

Chapter 14Acid- Base Chemistry

14.1Arrhenius

Acid- an acid is any substance that dissolves in water to produce H+ (H3O+) ionsBase- a base is a substance that dissolves in water to produce OH- ions

Bronsted- LowryAcid- an acid is a proton (H+) donorBase- a base is a proton acceptor

Hydronium Ion- the ion that foms when water accepts a proton (H+) H3O+

General reaction that occurs when an acid is dissolved in waterHA (aq) + H2O(l) H3O+(aq) + A-(aq)

Conjugate base- everything that remains of an acid molecule after a proton is lost

Conjugate acid- the molecule which is formed when the proton is transferred to the base

Conjugate acid- base pair consists of two substances related to each other by the donating and accepting of a single proton.

Equilibrium expression for a general reaction when an acid is dissolved in water would be Ka = [H3O+][A-] = [H+][A-]

[HA] [HA] (water not included; pure liquid)Ka is called the Acid Dissociation Constant

Example: Write the formula of the conjugate base of H2SO4. (a conjugate base differs from its acid by the lack of a proton: HSO4

-)

Example: For the reaction HSO4

-(aq) + HCO3-(aq) SO4

2-(aq) + H2CO3(aq)

• identify the acids and bases for the forward and reverse reactions

(forward: HSO4- is proton donor, acid; HCO3- is proton acceptor,

a base.)• identify the conjugate acid-base pairs

14.2 Acid StrengthStrong Acid- one in which the equilibrium lies far to the

right; it has a large value of Ka. Strong acids also yield weak conjugate bases.

Weak Acids- one in which the equilibrium lies far to the left; it has a small value of Ka. Weak acids yield relatively strong conjugate bases.

See Table 14.1, p. 662Example: The strengths of the following acids increase in

the order HCN < HF < HNO3. Arrange the conjugate bases of these acids in order of increasing base strength.(CN- > F- > NO3

-)

Diprotic Acid- an acid having two acidic protons; Examples: H2SO4

Oxyacids- the acidic proton is attached to an oxygen atom, Examples:

Organic Acids- those acids with a carbon atom backbone; commonly contain the

carboxylic group (-COOH, or ) Examples: Monoprotic acid- acids having one acidic proton.

Examples:

Amphoteric- can behave either as an acid or a base; ex., water.Autoionization- involves the transfer of a proton from one water

molecule to another to produce a hydroxide ion and a hydronium ion.2H2O(l) H3O+(aq) + OH-(aq)

Ion-product Constant Kw (or dissociation constant for water)-

always refers to the autoionization of water. Using the above reaction

Kw = [H3O+][OH-] = [H+][OH-]

At 25oC, [H+] = [OH-] = 1.0 x 10-7 M therefore, at 25oC, Kw = (1.0 x 10-7M)2 = 1.0 x 10-14 mol2L-2

It is important to recognize the meaning of Kw: The product of [H+] and [OH-] must always equal 1.0 x 10-14. This lends itself to three possibilities:• A neutral solution, where [H+] = [OH-]• An acidic solution, where [H+] > [OH-]• A basic solution, where [H+] < [OH-] Ka x Kb = Kw = 1.0 x 10-14

Example:The H+ concentration in a solution is 5.0 x 10-5M. What is the OH- ion concentration?

(2.0 x 10-10M; don’t forget unit when asked for a concentration!)

14.3 pH The pH scale is used to determine the strength of an acid or a base.

Traditionally it ranges from 0-14. Acids have pH< 7; bases have pH > 7 and 7 on the scale is considered to be neutral. pH is defined as

pH = -log[H3O+(aq)] or pH = -log[H+(aq)]

In the case of strong acids and bases dissociation is complete and therefore

the concentration of the H+ ion or OH-ion can be determined directly from the stoichiometric ratio in the balanced equation and the concentration of the acid or the base. Therefore we can simply plug in the numbers to the equation above, or, use the fact that, @ 25oC

logKw = log[H+] + log[OH-] = 7 + 7 = 14 = pH + pOH

And then use

pOH = -log[OH-(aq)]

***Significant Figures for logarithms: the number of decimal places in the log is

equal to the number of significant figures in the original

number.***

Example: The OH- ion concentration in a certain ammonia solution is 7.2 x 10-4M. What is the pOH and pH? (Answer: 3.14; 10.86)

Problem: What is the pH and pOH of a 0.001 M solution of HNO3?

Problem: What is the pH of a 0.045 M solution of HCl? Problem: What is the pH and pOH of a 0.02 M solution of

H2SO4? Example: The pH in many cola-type soft drinks is about 3.0.

How many times greater is the H+ concentration in these drinks than in neutral water? (104 or 10,000X)

(Working backwards: If you know pH and want to find concentration, use the antilog.)

[H+] = 10-pH [OH-] = 10-pOH

Another useful relationship in calculations is

pKa = -logKa and pKb = -logKb

and pKw = pH + pOH Problem: What is the concentration of a HCl solution for

which the pH is 4.32?

14.4 Calculating the pH of Strong Acid SolutionsSee Table for List of strong and weak acids and

bases. Know the List!Example: What is the pH of a 0.1 M HCl solution?

Of a 0.0001 M HCl solution?Example: What is the pOH of a 0.1 M NaOH

solution? The pH?Example: What is the pOH of a 0.00001 M NaOH

solution? The pH?

14.5 Calculating the pH of a Weak Acid Solution VERY IMPORTANT!!!Step 1- List the major species in the solutionStep 2- Choose the species that can produce H+, and write a balanced

equation for the reaction producing H+Step 3- Using the values of the equilibrium constant for the reactions you

have written, decide which equilibrium will dominate in producing H+.Step 4- write the equilibrium expression for the dominant equilibrium.Step 5- List the initial concentrations of the species participating in the

dominant equilibriumStep 6- Define the change needed to achieve equilibrium; that is, define x.Step 7- Write the equilibrium concentrations in terms of x.Step 8- Substitute the equilibrium concentrations into the equilibrium

expressionStep 9- Solve for x the “easy” way; that is, assume [HA]o - x = [HA]o

(assume x is negligible)Step 10- Using the 5% rule, verify is that approximation is validStep 11- calculate [H+] and pH.

5% Rule: x/[HA] x 100% < 5% then the approximation is acceptable Percent Dissociation: % dissociation = amount dissociated x 100%

Initial concentration For a given weak acid, the percent dissociation

increases as the acid becomes more dilute. For solutions of any weak acid HA, [H+] decreases as [HA]o decreases, but the percent dissociation increases as [HA]o decreases.

Example: Determine the pH of a of a 0.10 M solution of CH3COOH. Ka for acetic acid is 1.8 x 10-5

Determine the percent ionization of acetic acid in the example.

Problem: Calculate the concentrations of H+, F-, and HF in a 0.31 M HF solution. (Ka for HF is 7.1 x 10-4) Calculate the % ionization.

(Answer: [H+] = [F-] = 1.5 x 10-2 M[HF] = 0.31 M - 0.015 M = 0.30 M; 4.8%)

14.6 BasesStrong Base – one which dissociates completely in water and has a

large Kb value.General reaction for a base B and water is

B(aq) + H2O(l) BH+(aq) + OH-(aq) Base acid conjugate conjugate

acid baseThe equilibrium constant, Kb, for this equation is

Kb = [BH+][OH-] [B]

Weak bases have small values of Kb. (see Table 14.3, p. 685)

Example: What is the H+ concentration in 0.0025 M NaOH? (4.0 x 10-12 M)

Example: What is the pH of a 0.010 M C5H5N (pyridine) solution?(8.62)

14.7 Polyprotic AcidsPolyprotic acids- acids which can furnish more than 1 proton.

(See Table 14.4, p. 689)Characteristics of Weak Polyprotic Acids (p. 694)Typically, successive Ka values are so much smaller than the first value, that

only the first dissociation step makes a significant contribution to the equilibrium concentration of H+. This means that the calculation of the pH for a solution of a typical weak polyprotic acid is identical to that for a solution of a weak monoprotic acid.

Sulfuric acid is unique in being a strong acid in its first dissociation step and a weak acid in its second dissociation step. For relatively concentrated solutions of sulfuric acid (1.0 M or higher) the large concentration of H+ ions from the first dissociation step represses the second dissociation step (which can then be neglected as a contributor of H+ ions). For dilute solutions of sulfuric acid, the second step does make a significant contribution, and the quadratic equation must be used to obtain the total H+ concentration.

Example: Calculate the concentrations of H2A, HA-, A2-, and H+ in a 1.0 M H2A solution.

14.8 Acid-Base Properties of SaltsSalt- another name for an ionic compoundSalts that contain highly charged metal ions

produce an acidic solution. The metal ion becomes hydrated which then causes the solution to become acidic. See Table 14.6, p. 700.

Example: Write net ionic equations to show which of the following ions hydrolyze in aqueous solution? NO3

-, NO2-, NH4

+

Example: calculate Kb for NO3

- (2.2 x 10-11) Example: calculate Ka for NH4

+(5.6 x 10-10)

Example: Determine the pH of a 0.10 M solution of NH4Cl. (Kb for NH3 is 1.8 x 10-5)

Example: Predict whether the following aqueous solutions will be

acidic, basic, or neutral. KI (neutral) NH4I (acidic) CH3COOK (basic)

14.9 The Effect of Structure on Acid-Base PropertiesThere are two main factors that determine whether a molecule

containing an X-H bond will behave as a Bronsted-Lowry acid: the strength of the bond and the polarity of the bond. Increased polarity and high electron density typically lends to large Ka values (strong acids). (homework problem; has been on AP exam!!)

HF?Example: Which is the stronger acid? HCl or HBr? (HBr) HCl or H2S

(HCl) Example: Which is the stronger acid? HClO3 or HBrO3? (HClO3) H3PO3

or H3PO4 (H3PO4)

14.10 Acid-Base Properties of Oxides (Important in Equation Writing!!)

A compound containing the H-O-X group will produce an acidic solution in water if the O-X bond is strong and covalent. If the O-X bond is ionic, the compound will produce a basic solution in water. (homework problem; has been on AP exam!!)

Acidic oxides- a covalent oxide dissolves in water, and an acidic solution forms.

Basic oxides- an ionic oxide dissolves in water, and oxide has a great affinity for H+, causing basic solutions.

14.11 Lewis Acid-Base Model (encompasses the B-L model, but the reverse is not true!)

(Be able to draw the Lewis structure of this reaction)Acid- an acid acts as an electron pair acceptorBase- a base acts as an electron pair donor Example: identify the Lewis acids and bases in each of the following

reactions:Ag+(aq) + Cl-(aq) AgCl(s) (acid: Ag+)BF3(g) + NF3(g) F3N-BF3(s) (acid: BF3)SO2(g) + H2O(l) H2SO3(aq) (acid: SO2)