Acids and Bases

Properties of Acids

Sour taste React w/ metals to form H2

Most contain hydrogen Are electrolytes Change color in the presence of

indicators (turns litmus red) Has a pH lower than 7

Two Types of Acids

Strong acids– Any acid that dissociates completely in

aqueous sol’n Weak acids

– Any acid that partially dissociates in aqueous sol’n

Properties of Bases

Bitter taste Slippery feel Are electrolytes Change color in the presence of

indicators (turns litmus blue) Has a pH higher than 7

Types of Bases

Strong Base– Any base that dissociates completely in

aqueous sol’n Weak Base

– Any base that partially dissociates in aqueous sol’n

Neutralization

Neutralization rxn: a rxn of an acid and a base in aqueous sol’n to produce a salt and water

Salt: compound formed from the positive ion of a base and a negative ion of an acid

Properties of the acid and base cancel each other

Arrhenius Model of Acids and Bases

Proposed the model in 1887 Acid: any compound that produces H+ ions in

aqueous (water) sol’n Base: any compound that produces OH-

(hydroxide) ion in aqueous sol’n Offers an explanation of why acids and bases

neutralize each other (H+ + OH- = H2O)

Problems with Model

Restricts acids and bases to water sol’ns (similar reactions occur in the gas phase)

Does not include certain compounds that have characteristics of bases (e.g., ammonia)

Brønsted-Lowry Model of Acids and Bases

Brønsted acid: a hydrogen ion donor (H+, or proton)

Brønsted base: a hydrogen ion acceptor Defines acids and bases independently of

how they behave in water Amphiprotic: having the property of behaving

as an acid and a base– Also called amphoteric, e.g., water

Conjugate Acid-Base Pairs

The rxn between Brønsted-Lowry acids and bases can proceed in the reverse direction (reversible reactions)

HX (aq) + H2O (l) H3O+ (aq) + X- (aq) The water molecule becomes a hydronium ion

(H3O+), and is an acid because it has an extra H+ to donate

The acid HX, after donating the H+, becomes a base X-

Conjugate Acids and Bases

HX (aq) + H2O (l) H3O+ (aq) + X- (aq)Acid

Base Conjugate Acid

Conjugate Base

Forward reaction: Acid and base

Reverse reaction: Conjugate acid and conjugate base

Conjugate Acid: species produced when a base accepts a hydrogen ion from an acid

Conjugate Base: species produced when an acid donates a hydrogen ion to a base

Conjugate Acid-Base Pair: two substances related to each other by the donating and accepting of a single hydrogen ion

Types of Acids

Monoprotic acids: acids that contain only 1 hydrogen; e.g., HCl

Diprotic acids: acids that contain 2 hydrogens; e.g. H2CO3

Triprotic acids: acids that contain 3 hydrogens; e.g. H3PO4

More Types of Acids

Binary acids: acids that contain only 2 elements; e.g. HF

Polyatomic acids: acids that contain more than 2 elements; e.g. H2SO4

– These acids contain polyatomic ions– Also called ternary or oxy- acids

Naming Binary Acids

Start with the prefix hydro- Put it in front of the root word of the anion

(- charged ion) Add –ic to the end Examples

– Hydrobromic (HBr)– Hydrofluoric (HF)– Hydroiodic (HI)– Hydrochloric (HCl)

Naming Polyatomic Acids

Start with the root word of the name of the polyatomic ion

Add –ous if name ends in –ite Add -ic if name ends in –ate Examples:

– Chlorous (from chlorite, ClO2-)

– Nitric (from nitrate, NO3-)

– Sulfurous (from sulfite, SO3-2)

pH and [H3O+]

pH: number that is derived from the concentration of hydronium ions ([H3O+]) in sol’n– pH = -log [H3O+]– As pH increases, [H3O+] decreases

Scale ranges from 0 – 14– pH = 7 is neutral– pH < 7 is acidic– pH > 7 is basic

p[OH]

pOH = - log [OH-] pH + pOH = 14.00 Calculating ion concentrations from pH

[H+] = antilog (-pH) [OH-] = antilog (-pOH)

Dissociation Constants

Acid dissociation constant: (Ka): the equilibrium constant for the rxn of an aqueous weak acid and water

Base dissociation constant: (Kb): the equilibrium constant for the rxn of an aqueous weak base w/ water

Both are derived from the ratio of the concentration of the products and reactants at equilibrium

Acid Dissociation Constant

Ka = [H3O+] [A-]

[HA] Ka is a measure of the strength of an acid

Ka values for weak acids are always less than one

Used mostly w/ weak acids because the Ka values for strong acids approach infinity

Examples

HMnO4 (aq) + H2O (l)

H2S (aq) + H2O (l)

Base Dissociation Constant

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

Kb is a measure of the strength of a base

Kb values for weak bases are always less than 1

Kb values for strong bases approach infinity

Examples

H2NOH (aq) + H2O (l)

NH3 (aq) + H2O (l)

Water

Water can dissociate into its component ions, H+ and OH-

– 2H2O (l) H3O+ (aq) + OH- (aq) One water molecule acts as a weak acid, and

the other acts as a weak base The ions are present in such small amounts

they can’t be detected by a conductivity apparatus

In pure water, [H3O+] =1.0 x 10 –7 M and [OH-] = 1.0 x 10-7 M

Dissociation Constant for Water

It is defined as Kw: the ion product constant for water

Kw = [H3O+] [OH-] Kw = (1.0 x 10-7)(1.0 x 10-7) Kw = 1.0 x 10-14

The value of Kw can always be used to find the concentration of either H3O+ or OH- given the concentration of the other

Examples

What is the pH of a 0.001 M sol’n of HCl, a strong acid?

Examples

What is the pH of a sol’n if [H3O+] = 3.4 x 10-5 M?

Examples

The pH of a sol’n is measured with a pH meter and determined to be 9.00. What is the [H3O+]?

Examples

The pH o f a sol’n is measured with a pH meter and determined to be 7.52. What is [H3O+]?

Calculating Ka

In these problems, remember that the concentration of the [H3O+] ions will equal the concentration of the conjugate base ions.– This is because for every molecule of

weak acid that dissociates, there will be an equal number of H3O+ ions and base ions

Example

Assume that enough lactic acid is dissolved in sour milk to give a solution concentration of 0.100 M lactic acid. A pH meter shows that the pH of the sour milk is 2.43. Calculate Ka for the lactic acid equilibrium system.

Titrations

An analytical procedure used to determine the concentration of a sample by reacting it with a standard sol’n

In a titration, an indicator is used to determine the end point

Standard sol’n: a sol’n of precisely known concentration

Indicator: any substance in sol’n that changes color as it reacts with either an acid or a base

Titrations

Each indicator changes its color over a particular range of pH values (transition interval)

An unknown acid sol’n will be titrated with a standard sol’n that is a strong base

An unknown base sol’n will be titrated with a standard sol’n that is a strong acid

Titrations

Equivalence point: point at which the concentration of H3O+ ions is the same as the concentration of OH- ions; [H3O+ ] = [OH-]

Endpoint: the point at which the indicator changes color

Titration curve: graph that shows how pH changes in a titration

Titrations

The equivalence point is at the center of the steep, vertical region of the titration curve

At the equivalence point, pH increases greatly w/ only a few drops

Example Problem 1

What is the molarity of a CsOH solution if 20.0 mL of the solution is neutralized by 26.4 mL of 0.250M HBr solution?

HBr + CsOH → H2O + CsBr

Example Problem 2

What is the molarity of a nitric acid solution if 43.33 mL 0.200M KOH solution is needed to neutralize 20.00 mL of unknown solution?

Example Problem 3

What is the concentration of a household ammonia cleaning solution (NH4OH) if 49.90 mL of 0.5900M H2SO4 is required to neutralize 25.00 mL solution?