GEN CHEM 2 | Acids & Bases

H⁺

• bare proton

• cannot exist on its own water

• will bond with H₂O to form a hydronium ion (H₃O⁺)

Taste of Acids

They taste sour

Taste of Bases

They taste bitter

Three Definitions of Acids and Bases

Most Inclusive

1. Lewis

2. Bronsted-Lowry

3. Arrhenius

Least Inclusive

Arrhenius

• first systematic theory on acids and bases

• limited to aqueous solutions

Arrhenius Acids

increases [H₃O⁺] or [H⁺] in water

Arrhenius Bases

increases [OH^-] in water

Arrhenius Limitations

• only accounts for reactions in water

• only accounts for bases that already have hydroxide

• it assumes H⁺ ions float alone

Examples of Arrhenius Acids

• HClO3

• HNO3

• HClO4

• CHCOOH

• HBr

• H3PO4

• H2SO4

• H2SO3

• HCl

Examples of Arrhenius Bases

• NaOH

• KOH

• LiOH

• Mg(OH)2

• Ca(OH)2

• Al(OH)3

Bronsted-Lowry Theory

focuses on proton transfer

Bronsted-Lowry Acid

proton donor

Bronsted-Lowry Base

proton acceptor

Conjugate Acid-Base Pairs

a pair of substances that differ by exactly one H⁺ ion

Bronsted Conjugate Base

what remains after an acid donates its H⁺ ion

Bronsted Conjugate Acid

what forms after a base accepts an H⁺ ion

Bronsted-Lowry Theory

focuses on proton transfer

Bronsted-Lowry Limitations

• limited to proton transfer, does not explain substances without hydrogen

• does not apply to aprotic solvents (no protons) and gas phase

• does not explain acidic and basic oxides

• does not recognize acidic nature of proton-less compounds

Amphoteric Substances

• can act as an acid or a base, or is able to donate and accept proton transfer

• usually an acid if its in a reversible reaction

Lewis Theory

focuses on electron pair behavior

Lewis Acid

electron pair acceptor

Lewis Base

electron pair donor, also known as a ligand

Lewis Theory Importance

• explains metal-ligand complexes

• covers reactions without proton transfer

• connects to directly to organic mechanisms

• provides the most general acid-base definition

Lewis Theory Limitations

• does not explain the behavior of protonic acids that do not form coordination compounds with bases

• could not explain the relative strengths of acids and bases

• violates the octet rule for molecules that do not have exactly eight electrons

Litmus Indicator of Acid

blue → red

Litmus Indicator of Base

red → blue

Phenolphthalein Indicator of Acid

colorless (also for neutral solvents)

Phenolphthalein Indicator of Base

pink

Universal Indicator of Acid

red/orange/yellow (pH < 7)

Universal Indicator of Base

blue/green/purple (pH > 7)

Acids’ reaction with Metals

• they produce H2 gas, shown through fizzing, bubbles, and the metal slowly decomposing

Acids’ reaction with Carbonates/Bicarbonates

• they produce carbon dioxide gas through vigorous fizzing or effervescence

Conductivity of Acids/Bases

• acids and bases in water usually form ions to conduct electricity,

• the stronger the acid/base, the stronger the electric current is

Strength of Acids and Bases

how completely an acid/base ionizes in water

Concentration of Acids and Bases

how much of solute (in moles) per liters of solution

Strong Acids and Bases

• strong electrolytes that ionize completely in water

• acids have a pH of 0 to 1

• bases have a pH of 12 to 14

• its conjugate will be weak

Weak Acids and Bases

• do not completely dissociate in an aqueous solution

• they produce H⁺ and OH^-

• acids have a pH of 5 to 7

• bases have a pH of 7 to 10

• its conjugate will be strong

Water in Acid-Base Reaction

• undergoes tiny auto-ionization

• H₃O⁺ and OH^- are always present

• acidity and basicity are always about the balance and imbalance of these two components

• neutral water has equal concentration of hydronium ion and hydroxide ion

pH

• also referred as acidity/basicity

• historically denotes “potential of hydrogen”

• ranges from 0 to 14

• logarithmic in nature, where in log_x10ⁿ, n = pH level

• determines the hydrogen ion or hydronium ion concentration in a solution

• quantifies the acidity/alkalinity of a solution

pH Equations

• pH = -log[H⁺]

• [H⁺] = 10^-pH

• [H⁺] = [H₃O⁺]

pOH

• potential of hydroxide ion

• determines the hydroxide ion concentration in a solution

• quantifies the basicity of a solution

• p)H = -log[OH^-]

• [OH^-] = 10^-pOH

Relating pH and pOH

pH + pOH = 14

Hydronium vs Hydroxide Concentration of Acids

• [H₃O⁺] > [OH^-]

• [H₃O⁺] > 1.0 × 10^-7

• [OH^-] < 1.0 × 10^-7

Hydronium vs Hydroxide Concentration of Bases

• [H₃O⁺] < [OH^-]

• [H₃O⁺] < 1.0 × 10^-7

• [OH^-] > 1.0 × 10^-7