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