• acids have sour taste and turn litmus paper red
• bases have a bitter taste and feel slippery
• Svante Arrhenius (1859-1927)
• Acids associated with H⁺ ions
• Bases associated with OH^- ions
• Solution is acidic if there is more H⁺ than OH^-
• Solution is basic if there is more OH^- than H⁺

• autoionization of water - dissociation of H₂O molecules to H⁺ and OH^- ions
• at room temperature only 1 out of 10⁹ molecules are ionized
• exclude water from equilibrium expressions involving aqueous solutions
• ion-product constant
• k_w = k[H₂O] = [H⁺][OH^-] = 1.0 x 10^-14 (at 25° C)
• solution is neutral when [H⁺] = [OH^-]
• solution is acidic when [H⁺] > [OH^-]
• solution is basic when [H⁺] < [OH^-]

• H⁺ ion is a proton with no valence electrons
• H⁺ ion react with H₂O molecule to form H₃O⁺, hydronium ion
• H₃O⁺ ion can bond with other H₂O molecules to form hydrated hydrogen ions
• H⁺ and H₃O⁺ used interchangeably

• concentration of [H⁺] expressed in terms of pH
• pH = -log [H⁺]
• acidic solutions
• [H⁺] > 1.0 x 10^-7
• [OH^-] < 1.0 x 10^-7
• pH < 7.00
• neutral solutions
• [H⁺] = [OH^-] = 1.0 x 10^-7
• pH = 7
• basic solutions
• [H⁺] < 1.0 x 10^-7
• [OH^-] > 1.0 x 10^-7
• pH > 7

• pOH = -log [OH^-]
• pH + pOH = -log K_w = 14.00

• pH meter
• has a pair of electrodes connected to a meter that measures in millivolts
• voltage generated when electrodes placed in solution, and is measured by meter
• red litmus paper for pH of 5 or lower
• blue litmus paper for pH of 8 or higher

• Arrhenius definition of acids and bases
• Acids when dissolved in water increase H⁺ concentration
• Bases when dissolved in water increase OH^- concentration

• Brq nsted-Lowry definition of acids an bases
• Acid is a proton donor
• Base is a proton acceptor
• Can be applied to non-aqueous solutions
• Brq nsted-Lowry acid must be able to lose a H⁺ ion
• Brq nsted-Lowry base must have a non bonding pair of electrons to bind to H⁺ ion
• Amphoteric - substance that can act as an acid or base

• conjugate acid - product formed by adding a proton to base
• conjugate base - product formed by removal of a proton from acid

• the stronger the acid, the weaker the conjugate base
• the stronger the base, the weaker the conjugate acid
• equilibrium favors transfer of proton from stronger acid to stronger base

• strong acids and bases are strong electrolytes

• strongest monoprotic acids
• HCl, HBr, HI, HNO₃, HclO₃, HclO₄, and diprotic H₂SO₄
• For strong monoprotic acid concentration of [H⁺] equals the original concentration of the acid

• most common strong bases are ionic hydroxides of alkali metals and the heavier alkaline-earth metals
• complete dissociation

• HA_(aq) + H₂O_(l) ß à H₃O⁺ + A^-_(aq)
• HA_(aq) ß à H⁺_(aq) + A^-_(aq)
• 
• K_a = acid - dissociation constant
• The lager the K_a the stronger the acid
• K_a usually less than 10^-3

• 1) write ionization equilibrium
• 2) write equilibrium expression
• 3) I.C.E. Table
• 4) substitute equilibrium concentrations into equilibrium expression
• 
• in weak acids [H⁺] is small fraction of concentration of acid
• percent ionization depends on temperature, identity of acid and concentration
• as percent ionization decreases, concentration increases

• more than one ionizable H atom
• easier to remove first proton than second
• acid dissociation constants are K_a1, K_a2, etc…
• K_a values usually differ by 10³

• base-dissociation constant, K_b
• equilibrium at which base reacts with H₂O to form a conjugate acid and OH^-
• contain 1 or more lone pair of electrons

• weak bases have NH₃ and anions of weak acids

• when two reactions are added together than equilibrium constant of third reaction is equal to the product of the equilibrium constants of the added reactions
• reaction 1 + reaction 2 = reaction 3
• K₁ x K₂ = K₃
• K_a x K_b = [H⁺][OH^-] = K_w
• Acid-dissociation constant times base-dissociation constant equals the ion-product constant for water
• K_a x K_b = K_w = 1.0 x 10^-14
• pK_a x pK_b = pK_w = 14; (pK_a= -log K_a and pK_b = -log K_b)

• hydrolysis - ions reacting with water to produce H⁺ and OH^- ions
• anions from weak acids react with water to produce OH^- ions which is basic
• anions of strong acids are not basic and do not influence pH
• anions that have ionizable protons are amphoteric
• behavior depends on K_a and K_b
• all cations except those of alkali metals and heavier alkaline earth (Ca²⁺, Sr²⁺ and Ba²⁺) are weak acids in water
• alkali metal and alkaline earth cations do not hydrolyze
• do not affect pH
• strengths of acids and bases from salts
• 1) salts derived from strong acid and base
• no hydrolysis and solution has pH of 7
• 2) salts derived from strong base and weak acid
• strong conjugate base
• anion hydrolyzes and produces OH^- ions
• cation does not hydrolyze
• pH greater than 7
• 3) salts derived from weak base and strong acids
• cation is strong conjugate acid
• cation hydrolyzes to produce H⁺
• anion does not hydrolyze
• solution has pH below 7
• 4) salts derived from weak acid and base
• both cation and anion hydrolyze
• pH depends on extent on hydrolysis of each ion]

16.10.1 Factors that Affect Acid Strength

• strength of acid depends on:
• 1) polarity of H-X bond
• 2) strength of H-X bond
• 3) stability of conjugate base, X^-
• molecule will transfer proton if H-X bond is polarized
• in ionic hydrides H^- acts as proton acceptor because of negative charge
• nonpolar bonds produce neither acidic or basic solutions
• strong bonds less easily dissociated that weak bonds
• the greater the stability of conjugate base, the stronger the acid]

• metal hydrides are basic or have no acid-base properties in water
• nonmetal hydrides can be between having no acid-base properties to being acidic
• in each group of nonmetallic elements, acidity increases with increasing atomic number
• bond strengths decrease as central atom gets larger and overlap of orbitals get smaller

• Y-O-H bond
• Oxyacids - have OH bonded to central atom
• Base if bonded to a metal because pair of electrons shared between Y-O is completely transferred to O
• Ionic compound with OH^- is formed
• When bonded to nonmetal the bond is covalent and compounds are acidic or neutral
• As electronegativity of Y increases , acidity also increases
• O-H bond becomes more polar
• Conjugate base usually an anion and stability increases as electronegativity of Y increases
• Relating acid strengths of oxyacids to electronegativity of Y and to number of groups attached to Y
• 1) same number of oxygen atoms, acid strength increases as electronegativity of central atom increases
• 2) same central atom Y, acid strength increases with increasing number of bonded oxygen atoms to central atom
• acidity increases as oxidation number of central atom increases

• carboxyl group - COOH
• acidic behavior of carboxylic acids
• addition oxygen atom in carboxyl group draws density from O-H bond which increases the polarity
• conjugate base ion have resonance forms
• acidity increases as number of electronegative atoms in acid increases

• Lewis acid - electron pair acceptor
• Lewis base - electron pair donor
• Any Brq nsted-Lowry is a Lewis base
• Lewis acids have molecules that have incomplete octets

• hydration - attraction of metal ions to water molecules
• metal ion acts as Lewis acid
• water molecule acts as Lewis base
• electron density drawn from oxygen atom to water molecule
• O-H bond becomes more polarized
• For hydrolysis reactions K_a increases with increasing charge and decreasing radius of ion