Properties of acids

• Electrolytes: conduct electricity

• React to form salts

• Change the color of an indicator

• Have a sour taste

Properties of Bases

• Bitter taste

• “slippery feel”

• Electrolyte

• React with an acid to form a salt

Acid names:

• ide= hydro______ic acid

• ite == ____________ous acid

• ate == _____________ic acid

Arrhenius Acids and Bases

• An acid contains a Hydrogen ion that easily disassociates

• A base has a hydroxide ion that easily disassociates

Bronsted-Lowry

• An acid is a hydrogen ion donor

• A base is a hydrogen ion acceptor.

• NH3(aq)+ H2O(aq) NH4+

(aq) + OH-(aq)

• Base Acid Conj. Acid Conj. Base

conjugate acids and bases

• These are the reverse reaction “reactants”

• HA(aq) + H2O(l) H3O+(aq) + A-

(aq)

• Acid Base conjugate conjugate

acid base

• HCl + NaOH HOH + Cl-

Lewis acid and bases

• An acid is an electron pair acceptor

• A base is an electron pair donator

base acid• NH3 + BF3 H3N BF3

Acid Strength

A strong acid is one for which the equilibrium lies far to the products side Ka > 1.

A weak acid is one for which the equilibrium lies far to the reactants side Ka < 1.

Table 14.1 on page 659.

•Monoprotic: HCl

•Diprotic: H2SO4

•Triprotic H3PO4

Oxyacids

• Acids where the acidic hydrogen is attached to an oxygen.

• Pictures on pages 658-659.

• Structural formulas given in table 14.8 on page 694.

Organic acids

• Acids with a carbon atom backbone.

• Commonly contain a carboxyl group.

• Acetic acid

• Benzoic acid

Amphoteric Substance

• A substance that can act as both an acid and a base.

• Water

• Ammonia

We can tell an acid from a base by using an indicator

Autoionization of water

• H2O + H2O H3O+ + OH-

• Kw = [H3O+] [OH-] = 1.0 x 10-14

• [H3O+] [OH-] = [H+] [OH-]

• [H+] = [OH-] =1.0 x 10-7 M

• pH = -log [H+]• [H+] = antilog (-pH)

• pOH = -log[OH-]• pK = -log K

• Significant figures in logarithms- the # of decimal

places in the log = # of sig figs in the original number.

• [H+] [OH-] = 1.0 x 10-14

• pH + pOH = 14

• For strong acids the [H+] = the molarity of the acid.

System for solving weak acid equilibrium problems

• List the major species in solution.

• Find any species that can produce H+ and write a balanced equation for the reaction producing H+.

• Use the values for K for the reactions you have written to decide which reaction will dominate.

• Write the equilibrium expression for the dominant reaction.

• List the initial concentrations of the species in the dominant reaction.

• Define the change needed to obtain equilibrium. Define x.

• Write equilibrium concentrations in terms of x.

• Put equilibrium concentrations into equilibrium expression.

• Solve for x the “easy way”• Use the 5% rule to see if

approximation is valid.• Calculate [H+] and pH.

• Calculate the pH of a 0.100M solution of hypochlorous acid.

• Calculate the pH of a solution that contains 1.00M HCN and 1.00M HNO2. Also calculate the concentration of cyanide ions.

% dissociation

• = amount dissociated x100

• initial concentration

• For a given weak acid, the % dissociated increases as the acid becomes more dilute.

• In a 0.100M solution, lactic acid (HC3H5O3)is 3.7% dissociated. Calculate the value of Ka for this acid.

Bases

• B(aq) + H2O(l) HB+(aq) + OH-

(aq)

• Base acid conj conj

• acid base

• Kb = [BH+][OH-]

• [B]

Bases

• Strong bases – hydroxides of group 1A metals and calcium, barium and strontium.

• Weak bases are commonly ammonia and substituted ammonia compounds.

• Table 14.3 on page 678.

• Calculate the pH for a 15.0M solution of NH3. Kb = 1.8 x10-5

Salts that produce neutral solutions

• Salts that consist of the cations of strong bases and the anions of strong acids have no effect on pH.

• KCl, NaNO3 , Ba(HSO4)2

Salts that produce basic solutions

• Salts that consist of cations of strong bases and the anion is the conjugate base of a weak acid.

• NaC2H3O2, KNO2, Sr(CN)2

Salts that produce acidic solutions• Salts that consist of a cation that

is the conjugate acid of a weak base and the anion of a strong acid.

• NH4Cl• A salt that contains a highly

charged metal ion. AlCl3 see sample exercise 14.20.

For any weak acid-conjugate base or weak base-conjugate

acid

• Ka x Kb = Kw

Calculate the pH of:

• 0.10M NaCl

• 0.10M NaF

• 0.10M NH4Cl

• If the anion is the conjugate base of a weak acid and the cation is the conjugate acid of a weak base the Kb must be compared to the Ka. Which ever is greater will dominate.

Predict if the following will be acidic or basic

• NH4C2H3O2

• NH4CN

• NH4NO2

Effect of structure on acid-base properties

Acidic properties depend on two factors: H-X

• The strength of the bond . As the strength increases the acidity decreases

• The polarity of the bond. As polarity increases acidity increases

Strength of oxyacids

• Within a series of oxyacids as the number of oxygens increases the strength of the acid also increases.

• Table 14.8 page 694

Acid-base properties of oxides• Depends on the electronegativity of

the element bonded to oxygen. O-X

• Non-metal oxides in water will form acids. O-X is stronger than H-O in polar water.

• Metal oxides in water form bases. O-H stronger than O-X in polar water.

Common Ion Effect

• The shift in equilibrium position because of the addition of an ion already involved in the equilibrium.

• Equilibrium shifts away from the added component.

Buffered Solutions

• Resists change in pH when either hydrogen or hydroxide ions are added.

• Consist of a solution that contains both a weak acid and its salt or a weak base and its salt.

Important Characteristics are on page 726.

Buffer Capacity

• Represents the amount of hydrogen or hydroxide ions the buffer can absorb without a significant change in pH.

• Best buffers contain a weak acid with a pKa as close as possible to the desired pH.

Indicators

• A substance that changes color depending on the pH of the solution it is in.

• When choosing an indicator we want the indicator end point and the titration equivalence point to be as close as possible.