Acids, Bases and Salts

• Acids give up hydrogen ions (H+) in a water solution.

• Bases give up hydroxide ions (OH-) in a water solution.

Properties of Bases

• Bitter taste• Feel slippery if you touch them• pH > 7.0• Strong bases:

– Sodium hydroxide (NaOH)– Potassium hydroxide (KOH)

• Weak bases:– Sodium bicarbonate (NaHCO3)

Properties of Acids• Sour taste • pH < 7.0• React with metals to form H2 gas• Strong acids:

– Sulfuric acid (H2SO4)– Hydrochloric acid (HCl)

• Weak acids: – Acetic acid (vinegar – CH3COOH))– Citric acid (in fruits)– Ascorbic acid (Vitamin C)

Naming Acids

• Binary Acid = 2 elements (HCl)– Begin with hydro-.– Use name of 2nd element and end with –ic.– HCl is hydrochloric acid.

• Oxyacid = H + O + 3rd element (H2SO4)

– Usually incorporates the polyatomic ion name into the acid name.

– Written with H first, then the polyatomic ion.

– H2SO4 is sulfuric acid.

Acid Strength• Strong acids ionize completely in water.

– Many H3O+ ions– Strong electrolyte

• Weak acids do not ionize completely in water.– Some H3O+ ions– Weak electrolyte– The ionized H3O+ ions are being bonded back to the

negative species (reverse reaction) at the same time some acid molecules are being ionized to form H3O+ ions .

• Strong and weak bases are similar, except ion involved is OH- instead of H3O+ .

Classifying strength of electrolytes from solubility

(general rules for compounds)• Strong electrolyte

– Water soluble AND ionic– Strong acid, water soluble and not ionic– Strong base, water soluble and not ionic

• Weak electrolyte– Weak acid or base, water soluble and not

ionic

• Nonelectrolyte– None of the above

Cl

The Hydronium Ion• Ionization is the creation of ions from a

molecular compound.– If the force of attraction between solvent molecules

and parts of the solute are stronger than the covalent bonds of the solute, the solute breaks into ions.

– In water, HYDRATION produces heat, which provides energy to break more covalent bonds.

• H + attracts other molecules or ions • In water, H+ becomes a proton bonded to the

oxygen of a water molecule.• The hydronium ion is H3O+.

ClOO+ +-+

H H H

H

H

H

Acid-Base Systems

Type Acid Base Examples

Arrhenius

H+ or H3O+ producer

OH- producer

HNO3 + H2O H3O+ + NO3

-

NH3 + H2O NH4+ + OH-

Brønsted-Lowry

proton (H+) donor

proton (H+) acceptor

HCl + NH3 NH4+ + Cl-

Lewis electron pair acceptor

electron pair donor

BF3 + F- BF4-

BF3 is Lewis acid,

F- is Lewis base

Brønsted-Lowry Acids

• Brønsted-Lowry acids donate a proton.

• A monoprotic acid donates one proton per molecule, such as HCl and HClO4.

• A polyprotic acid donates more than one proton per molecule, such as H2SO4 and H3PO4.

– H2SO4 is a diprotic acid.

– H3PO4 is a triprotic acid.

Conjugate acids and bases• The species remaining after a Brønsted-Lowry

acid gives up its proton is the conjugate base of that acid: Take off one H from the acid.

• The species remaining after a Brønsted-Lowry base accepts its proton is the conjugate acid of that base: Take off one H. Add an H to the base.

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO- (aq)

acid base conjugate conjugate

(proton donor) acid base

Examples of conjugate acids and basesHCO3

- (aq) + H2O(l) H2CO3 (aq) + OH- (aq)

base acid conjugate conjugate

(proton acid base

acceptor)

HF(aq) + H2O(l) H3O+(aq) + F- (aq)

acid base conjugate conjugate

(proton acid base

donor)

Salts

• An acid and base combined together react to neutralize each other.

• Neutralization is a process which produces produces that are not acids or bases.

• H+ and OH- make water.• The remaining reactants combine to make a salt.• A salt is an ionic compound made of a metal

and a nonmetal.• Soap and detergent are organic salts.

Neutralization

HCl + NaOH NaCl + H2O

HNO3 + KOH KNO3 + H2O

H++ OH- H2O

ACID + BASE YIELDS SALT + WATER

saltacid base water

Neutralization• The reaction between an acid and a base which

produces a salt and water. • Recall: A salt is an ionic compound: metal +

nonmetal which is likely to dissociate in water.• Ex: You take a dose of the antacid magnesium

hydroxide, Mg(OH)2, to relieve excess stomach acid, HCl.

• Equations for this neutralization:

2HCl(aq) + Mg(OH)2 (s)MgCl2(aq) + 2H2O(l)

2H+(aq) + Mg(OH)2 (s) Mg2+(aq) + 2H2O(l)

Ionic Equations for Acid-Base Reactions

• H+(aq) + OH-(aq) H2O (l)

• Strong acids ionize completely.

• HCl H+ + Cl-

• Weak acids do not ionize completely.

• HF H+ + F-

Water and ammonia are a weak acid and a weak base, respectively.

• NH3(aq) + H2O (l) NH4+(aq) + OH- (aq)

Measuring Strength of Acids and Bases

• pH is a measurement which indicates acidic or basic strength, measuring the concentration of H+ ions.

• More H+ ions = lower pH = acid.

• An indicator changes color depending on pH of a solution.

• pH paper is treated with indicators to change color when dipped in a solution.

1: acid 7: neutral 14: base

pH and acidity• pH is related to the concept of concentration of

hydronium ions found in water.• Water and all its solutions contain hydronium ions and

hydroxide ions.

• Acidic solutions: More H3O+ ions.

• Basic solutions: Fewer H3O+ ions / more OH- ions.

• If pH = 2:

H3O+ ion concentration = 0.01 M

OH- ion concentration = 0.000000000001 M• If pH = 12:

H3O+ ion concentration = 0.000000000001 M

OH- ion concentration = 0.01 M

The pH scale

• pH stands for “power of hydronium ion.”

pH = -log [H3O+]

• pH value is the exponent on the power of 10 with its sign changed.

pH examples

• Example 1: – 0.001 mol H3O+ = 1 x 10-3 mol H3O+

– Concentration is 0.001M

– pH = -log [H3O+] = -log (1x 10-3 ) = -(-3.0) = 3.0

– pH is 3.0 - This is an acid.

• Example 2: – 0.000 000 01 mol H3O+ = 1 x 10-8 mol H3O+

– Concentration is 0.000 000 01M

– pH = -log [H3O+] = -log (1x 10-8 ) = -(-8.0) = 8.0

– pH is 8.0 - This is a base.

pOH

• pOH is the negative of the logarithm of the hydroxide ion (OH-) concentration.

pOH = -log [OH-]

pH + pOH = 14.0 at 25 deg. C

Buffers• Why is the pH of some lakes unaffected by

acid rain even when they are downwind of big polluters?

• The lakes are surrounded by soils which neutralize the acidic precipitation.

1. One way to neutralize acid is to add a base. Limestone (CaCO3) is a weak base.

2. Another way to neutralize either an acid or a base is to add a buffer.

• A buffer is a substance or combination of substances capable of neutralizing limited quantities of either acids or bases.

Buffers• Calcium carbonate around a lake contaminated

with acid rain would react this way:

CaCO3 + H3O+ Ca 2+ + HCO3- + H2O

• The polyatomic ion HCO3- acts as a buffer:

HCO3- + H3O+ 2H2O + CO2

HCO3- + OH- H2O + CO3

2-

• Another buffer is the hydrogen phosphate ion. Possible sources of this ion are Na2HPO4 and NaH2PO4.

H3O+ + H2PO4 2-

H2PO4 - + H2O