Acids, Bases,

and Salts

Objectives

• Know the fundamental properties of acids and bases.

• Be able to identify an Arrhenius acid.• Be able to write a dissociation equation for an

Arrhenius acid.

Properties of AcidsACIDS• pH < 7• sour taste• electrolytes• react with metals to make

hydrogen gasZn + 2HCl → H2 + ZnCl2

• often formed from non-metal oxide and waterSO3 + H2O →H2SO4

Properties of BasesBASES• pH > 7• bitter taste• electrolytes• feel slippery• often formed from metal

oxide and waterZnO + H2O →Zn(OH)2

Acids and bases “neutralize” each other!HCl + NaOH → NaCl + H2O

Arrhenius AcidsArrhenius discovered that acids are:•molecules that contain H•ionize in water to make H+

•HCl → H+(aq) + Cl–(aq)•H2SO4→ H+(aq) + HSO4

–(aq)

•HBr → ?•H3PO4 → ? Svante Arrhenius

1859-1927

Objectives

• Be able to name acids.• Be able to identify an Arrhenius base and

write a dissociation equation for an Arrhenius base.

• Be able to identify Brønsted-Lowry acids, bases, conjugate acids, and conjugate bases.

• Understand and correctly apply the meaning of the term amphoteric.

Acid Nomenclature• USE YOUR YELLOW SHEET!• use the stem and ending of the anion name

-ide hydro-stem-ic acid-ate stem-ic acid-ite stem-ous acid

• HCl = H+ + Cl– (chlor-ide) = hydrochloric acid• HNO3 = H+ + NO3

– (nitr-ate) = nitric acid• HNO2 = H+ + NO2

– (nitr-ite) = nitrous acid• Common exceptions:

sulfuric (H2SO4) and phosphoric (H3PO4)

Arrhenius Bases• Bases dissociate to form OH- (hydroxide) ions when

aqueous. NaOH(s) → Na+(aq) + OH-(aq)Mg(OH)2(s) → Mg2+(aq) + 2OH-(aq)

Ca(OH)2(s) → ?KOH(s) → ?

• phenolphthalein indicates OH- (pink)• problem: why is ammonia (NH3) basic?

Brønsted-Lowry Acids and Bases

• acid: proton (H+) donor• base: proton (H+) acceptor

HCl(g) + H2O(l) ↔ Cl−(aq) + H3O+(aq) hydronium ionacid base conjugate

baseconjugate

acid

NH3(aq) + H2O(l) ↔ NH4+(aq) + OH−(aq)

acidbase conjugateacid

conjugatebase

HNO3(aq) + NH3(aq) ↔ NO3−(aq) + NH4

+(aq)acid base conj base conj acid

Water: Acid and Base!

• amphoteric: a substance that can act as either an acid or a base (such as water)

hydronium ion H3O+

hydroxide ionOH-

+

+

+ =

= −

H+ is really H3O+ because water bonds with H+base conj. acid

acid conj. base

Objectives

• Understand the process of self-ionization.• Understand how the concentrations of

hydronium and hydroxide ion can vary in water.

• Understand the concept of pH.• Be able to make pH calculations using the log

and 10x functions on a calculator.

Self-Ionization of Water

• H2O + H2O ↔ OH− + H3O+ (reactant strongly favored)• [OH− ] = 10-7 M and [H3O+] = 10-7 M• Kw = [OH− ] x [H3O+] = 10-14 • [OH−] and [H3O+] are inversely proportional

neutral water: Kw = [10-7] x [10-7] = 10-14

acidic: Kw = [10-9] x [10-5]= 10-14

basic: Kw = [10-3] x [10-11]= 10-14

[H3O+]

• ACIDS [H3O+] > 10-7 M• HNO3 (g) + H2O (l) →H3O+ (aq) + NO3

− (aq)• [H3O+] = 10-6 M, 10-5 M, 10-4 M, … 10-1 M or more

• BASES [H3O+] < 10-7 M• NaOH(s) → Na+(aq) + OH−(aq) OH− reduces H3O+

• [H3O+] = 10-8 M, 10-9 M, 10-10 M, … 10-14 M or less

pH Scale• pH = −log[H3O+] • [H3O+] = 10-3 M, pH = 3 (acidic)• [H3O+] = 10-7 M, pH = 7 (neutral)• [H3O+] = 10-11 M, pH = 11 (basic)

• Calculating pH?• [H3O+] = 5.7 x 10-2 M pH = −log(5.7E-2) = 1.2

• Calculating [H3O+]? Use [H3O+] = 10-pH • If pH = 3.8 [H3O+] = 10-3.8 = 1.6 x 10-4 M

Objectives

• Understand how acid precipitation forms.• Understand the effects of acid precipitation

and how they can be reduced.• Understand how acid-base indicators work.

Acid Rain, Acid Fog• acid rain/fog: precipitation with a low pH (< 5)• burning “high-sulfur” coal produce SO2 and SO3 that

react w/ H2O to make H2SO3 and H2SO4

• cars make NOX: reacts w/ H2O to make HNO2 and HNO3

dangerous toorganisms

corrodes metaldecomposes limestone

Acid Rain in the USA

Neutralizing Acid Rain

• Limestone bedrock neutralizes acid, reducing environmental damage.

• Granite does not. Bases such as CaO or CaCO3 must be used to neutralize acids.

H2SO4 + CaCO3 → CaSO4 + H2O + CO2

Acid-Base Indicators• compounds that respond to pH change by changing color• contain a “weak acid” in a chemical equilibrium

indicator anion

H+ indicator anion

+↔

add acid (add H3O+) = clear in low pH add base (removes H3O+) = pink in high pH

universal indicator: mixture, wide pH range

CONJ BASE = pinkACID = clear

H3O++ H2O

Plant Dyes and pH• serviceberry, willow bark, Oregon grape root, are indicators• have been used as natural dyes for skins, feathers, etc.

Objectives

• Understand the concept of KA and how it relates to strong and weak acids.

• Be able to calculate the KA of an acid solution if given the initial molarity and the pH of the solution.

Strengths of Acids• strong acid: completely ionizes in water, products favored

HNO3 (g) + H2O (l) → H3O+(aq) + NO3−(aq)

• weak acid: partially ionizes in water, reactants favoredHC2H3O2(l) + H2O (l) ↔ H3O+(aq) + C2H3O2

−(aq)

Acid Dissociation Constant (KA)

HA ↔ H+ + A−

Note that [H+] = [A− ] * use [H+] = 10-pH

[H+] = [H3O+]

[HA] = initial molarity – [H+]

Strong acids—high KA ( > 1, products favored)Weak acids—low KA ( < 1, reactants favored)

[HA]][A][H

KA

Calculating KA

• Determine [H+] (same value as [A-] )[H+] = 10-pH = 10-1.93 = 0.012 M

• Determine [HA][HA] = initial – [H+] = 0.315 M – 0.012 M = 0.303 M

• Calculate KA

4A 104.7

[0.303]012][0.012][0.

[HA]][A][H

K

KA < 1, weak acid

The initial concentration of an HNO2 solution is 0.315 M. What is the KA of HNO2 if the pH of the solution is 1.93?

Objectives

• Be able to explain the distinction between strong and weak acids versus concentrated and dilute solutions.

• Understand the concept of acid neutralization and be able to determine the products of an acid-base neutralization reaction.

• Be able to calculate either acid or base concentration using data from an acid-base titration.

Strength vs. Concentration

• strength relates to degree of ionization (KA) • concentration relates to amount of solute (M)

strong = product favored weak = reactant favored

concentrated = lots of solute dilute = not much solute

Neutralization

• acid + base → salt + water• H+ + OH− → H2O • salt: ionic compound consisting of a base cation and an

acid anion• HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)• H2SO4 (aq) + 2KOH (aq) → K2SO4 (aq) + 2H2O (l)Try this one…• HNO3 (aq) + Ca(OH)2 (aq) → ? + ?• 2HNO3 (aq) + Ca(OH)2 (aq) → Ca(NO3)2 (aq) + 2H2O (l)

Acid-Base Titration

• standard solution (known concentration) is added to an unknown solution until pH = 7• the concentration of the unknown can be calculated

Titration Calculation

B

BB

A

AA

n

CV

n

CV

What is the concentration of H2SO4 if 10.0 mL is completely neutralized by 14.2 mL of 1.0 M NaOH?

Buffers• buffer: a solution in which the pH remains relatively

constant when a small amount of acid or base is added

• consists of weak acid (or base) and one of its salts• Example: Your blood pH (= 7.2) is maintained by

H2CO3/HCO3− buffer

Add acid: H+ + HCO3− → H2CO3

Add base: H2CO3 + OH− → HCO3− + H2O