Acids and Bases Summary
Properties of Aqueous Acids:
1. Acids are electrolytes: aqueous solutions of acids ionize when added to water and conduct electric
current.
2. Traditional Acids contain hydrogen and react with active metals to produce hydrogen gas (H2) by
single replacement reactions.
3. Acids change the color of dyes known as acid base indicators; i.e. acids turn blue litmus red, turns
bromothymol blue to yellow, and leaves phenolphthalein clear.
4. Neutralizes bases: acids react with bases to produce salts and water. Equivalent amounts of acid and
base are neutralized (neutralization reactions) to produce water and a salt (ionic compound).
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(aq)
5. Sour Taste - Weak acids taste sour, i.e.: citrus fruits (lemons, oranges, grapefruit) contain citric
acid; apples - malic acid; vinegar - acetic acid; sour milk - lactic acid; wine - tartaric acid;
carbonated beverages - phosphoric acid.
6. Acids have a pH less than 7.0
7. Strong Acids: Acids that are completely or nearly completely ionized (dissolved) including HCl,
H2SO4, HBr, HNO3, etc. Strong acids in concentrated solutions are very corrosive.
Weak Acids: those acids that are only slightly ionized (dissolved) including CH3COOH,
H3PO4, HF, etc.
Properties of Aqueous Bases:
1. Bases are electrolytes: aqueous solutions of bases conduct electricity.
2. Bases, like acids, change the color of acid base indicators: i.e. bases turn red litmus blue, leaves
bromothymol blue, blue, and turns phenolphthalein pink.
3. Bases neutralize acids to produce salts and water. NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(aq)
4. Bases have a bitter taste.
5. Dilute aqueous solutions of bases feel slippery to the skin (soap)
6. Bases have a pH greater than 7.0
7. Strong Bases: those bases that are completely or nearly completely ionized including KOH, NaOH,
Ba(OH)2, Sr(OH)2, and Ca(OH)2. All active metal hydroxides (group 1 and 2 metals)
form strong bases. Strong bases like strong acids are very corrosive.
Weak Bases: those bases that are only slightly ionized including NH4OH, and all others.
Aqueous Acids
1.
“Acid”: traditionally refers to either a substance that produces hydrogen ions in water or to a
solution containing hydrogen ions (H+; H3O+). The terms “hydrogen ion (H+) and
hydronium ion (H3O+)” are used interchangeably and are considered to be the same thing.
All pure acids are molecular electrolytes, i.e.: sufficiently polar so that water molecules
attract away hydrogen ions leaving behind anions, negatively charged ions.
2.
Strong acid: an acid that is a strong electrolyte and ionizes (dissociates) 100% in aqueous
solutions.
HClO3
HI
HIO4
HBr*
HNO3*
HMnO4
HCl*
HClO4
H2SO4*
3. Weak acids: an acid that is a weak electrolyte: i.e. a solute that yields a relatively low concentration
of ions in aqueous solution.
CH3COOH*
HF*
HNO2
H2CO3*
H2S
H3PO4*
HCN
HClO
H2SO3
4.
Monoprotic acids - an acid that can donate 1 proton per molecule; one mole H+/mole acid (HCl)
5. Polyprotic acid - an acid that can donate more than 1 proton per molecule.
a. Diprotic acid - an acid that can donate 2 protons per molecule. (H2SO4)
b. Triprotic acid - an acid that can donate 3 protons per molecule. (H3PO4)
H2SO4 is the only strong polyprotic acid.
6. The strength of an acid depends upon the degree of ionization (dissolving) in aqueous solution and
not on the amount of hydrogen in the molecule. Any similar acid (monoprotic, diprotic, etc.) with
100% ionization has the same strength because the acid is the H+ ion (H3O+).
7. Organic acid - an organic compound that contains an acid group, the carboxyl group (-COOH).
CH3COOH (acetic acid); HCOOH methanoic acid); CH3CH2CH2COOH (butanoic acid)
Types of Acids and Bases:
1. Arrhenius Acids and Base: (Traditional definition)
a. Arrhenius acid: any substance that produces hydrogen ions or increases the hydrogen ion
concentration of a solution. (H+, H3O+) (HCl, HC2H3O2)
HCl  H+ + ClH2SO4  2H+ + SO4-2
b. Arrhenius base: any substance that produces hydroxide ions (OH- ) or increases the hydroxide
ion concentration of a solution. (NaOH, NH4OH)
NaOH  Na+ + OHProducts of an Arrhenius acid-base reactions are typically a salt (ionic compound) and water.
HCl + NaOH  NaCl + H2O
2. Bronsted – Lowry Acids and Bases:
a. Bronsted – Lowry acid: any substance capable of donating a proton (proton donor). The
proton is exactly the same as the H+ (H3O+) ion in the Arrhenius acid.
b. Bronsted – Lowry base: any substance capable of accepting a proton (proton acceptor).
c. The products that result from a Bronsted–Lowry acid-base reaction are called the conjugate
acid and the conjugate base.
HF(aq) + HCO3-(aq)  H2CO3(aq) + F-(aq)
acid
base
conj. acid
conj. base
The conjugate base is the particle that remains after one proton is donated by an acid. The
conjugate acid is formed when a base accepts one proton from an acid. The only difference
between an acid and its conjugate base is one proton (H+/H3O+).
3. Lewis Acids and Bases
a. Lewis acid: any substance that can accept a pair of electrons to form a coordinate
covalent bond. Ex: H+, BF3
b. Lewis base: any substance that can donate a pair of electrons to form a coordinate covalent
bond (OH-, HSO4-, O2-); i.e. substances that have one or more pairs of unshared
electrons. (H2O, NH3). May include substances that do not contain either
hydrogen or hydroxide ions.
c. The product of a Lewis acid-base reaction is normally a single substance formed by the
combining of the acid and base.
Na2O + SO3  Na2SO4
BF3 + NH3  F3BNH3
H+ + NH3  NH4+
4. Aqueous (Water) Solutions
a. Electrolytes - substances whose water solutions conduct electricity (i.e. contains ions).
1. Strong Electrolytes: very good conductors of electricity; nearly 100% of the solute is
dissolved (dissociated). Ex: salt/water
2. Weak Electrolytes: very weak conductors of electricity; solute only partially dissolved.
Ex: acetic acid (CH3COOH)
3. Non-electrolyte: a substance whose water solution does not conduct electricity (does not
contain ions). Ex: sugar/water
b. Types of Electrolytes
1. Ionic Substances - bonded by ionic bonds, ionic substances dissociate in water.
NaCl(s) + H2O  Na+ + Cl-
c. Molecular (Covalent) Substances - bonded by covalent bonds, covalent substances ionize in
water.
HCl + H2O  H+(H3O+) + ClHC2H3O2(l) + H2O(l)  H3O+(aq) + C2H3O2-(aq)
d. Amphoteric (amphiprotic): substances such as HCO3-, HSO4-, H2O, and NH3 that can act as
either proton donors (acids) or proton acceptors (bases) depending upon what else is present.
1. In the presence of a strong acid, the HSO4- ion acts as a base (proton acceptor).
H3O+ + HSO4-  H2SO4 + H2O
2. In the presence of a strong base, the HSO4- ion acts as a proton donor (acid).
HSO4- + OH-  H2O + SO425. pH, pOH, H+, and OH-.
a. Self Ionization of Water: even in very pure water, water is very slightly ionized yielding
one mole of H+ ions and one mole of OH- ions per 107 liters.
H2O(l) + H2O(l)  H3O+ + OH- + + H2O(l); since H+ = OH- = 10-7M
KW = H+OH- = 10-7 10-7 = 10-14
The equilibrium constant for water is KW = H+OH- = 10-7 x 10-7 = 10-14. KW remains
constant in water and dilute aqueous solutions at constant temperature. An increase in the
concentration of one ion will cause a decrease in concentration of the other.
b. pH: A convenient method of expressing the acidity or basicity of a solution in terms of values
related to the hydrogen ion concentration. A neutral solution has a pH of 7.0. Acidic solutions
have values below 7.0, and basic solutions have values above 7.0.
Solutions are acidic if: H+  OH- or pH  pOH
Solutions are basic if: H+  OH- or pH  pOH
c. pH: the negative of the common log (base 10) of the hydrogen ion concentration (pH = logH+).
d. calculations involving pH, pOH, hydrogen ion concentrations H+, and hydroxide ion
concentrations OH- are related to the following equalities:
H+OH- = 10-14;
pH + pOH = 14;
H+ = 10-14/OH-;
pH = -logH+;
a) To find H+ given pH: pH  H+
b) To find pH given H+: H+  pH
OH- = 10-14/H+
pOH = -logOH-
pH = -logH+
H+ = -antilog pH
Note: Logarithms are another method of writing exponential forms.
If b and N are positive numbers, then logb N = k if and only if bk = N.
Exponential Form
Logarithmic Form
23 = 8
log2 8 = 3
2
6 = 36
log6 36 = 2
10-3 = 0.001
log10 (0.001) = -3
Example:
a. Calculate pH for H+ = 0.0000001M
because pH = -logH+ = -log1.0 x 10-7 = --7 = 7
b. Calculate H+, OH-, pH, and pOH of a 0.0001M solution of HNO3.
because HNO3 is an acid, a 0.0001M solution of HNO3 has
a. H+ = 0.0001M = 1.0 x 10-4M; since H+OH- = 10-14
b. OH- = 1.0 x 10-14 = 1 x 10-10M
1.0 x 10-4
c. pH = -logH+ = -log(10-4) = -(-4)= 4
d. pOH = -logOH- = -log(10-10) = -(-10) = 10; or since pH = 4 then pOH = 14 - 4 = 10
c. When the concentration is in the form H+ = 1.0 x 10-x M, the concentration = molarity
and pH = x. (10-pH)
Calculate the pH of a 0.00001M NaOH solution.
OH- = 10-5M
pOH = 5 and pH = 9
d. When the concentration is in the form n x 10-x M, the concentration = molarity and
pH = x – log n. This can be calculated in one step on your calculator.
a. Calculate the pH of an acid solution whose concentration is 3.4 x 10-5M.
pH = 3.4 EE 5  Log  enter = 4.5
If pH = 4.5 calculate H+; since H+ = -antilog pH, then 4.5  2nd 10x = 3.4 x 10-5
b. Calculate the pH and pOH of a 6.4 x 10-4M solution of HCl.
OH- = 1.0 x 10-14M2 = 1.6 x 10-11M
6.4 x 10-4M
pOH =1.6 EE 11  Log  enter = 10.8; or pOH = 14 – pH = 14 – 3.2 = 10.8
e. Given the pH, H+ can be calculated:
pH = -logH+ and H+ = antilog(-pH) = -antilog pH
If pH = 8.58, find H+
H+ = -antilog pH = 8.58  2nd 10x = 2.63 x 10-9
f. If pH = 4.36, calculate H+ and OH-.
H+ = -antilog pH = 4.36  2nd 10x = 4.37 x 10-5
OH- = 1 x 10-14 = 2.29 x 10-10 or pOH = 14 - 4.36 = 9.64;
4.37 x 10-5
and, OH- = -antilog pOH = -9.64  2nd 10x = 2.29 x 10-10.
6. Chemical Equivalents – the quantity of solutes that have equivalent combining capacity.
HCl + KOH  KCl + H2O
H2SO4 + KOH  1/2K2SO4 + H2O
mole + 1 mole  1 mole + 1 mole
1/2 mole + 1 mole 1/2 mole + 1 mole
1/2 mole H2SO4 and 1 mole HCl are chemical equivalents in neutralizing 1 mole of KOH
a. Equivalent mass of an acid – the quantity, in grams, that supplies one mole of protons (H+).
b. Equivalent mass of a base – the quantity, in grams, that supplies one mole of hydroxide ions
(OH-).
7. Normality – the number of equivalents of solute per liter of solution.
a. molality = moles of solute/kg solvent
b. Molarity = moles of solute/liter solution
c. Normality = the number of equivalents/liter solution
N = nM (n = # of equivalents mole )
MaVa = MbVb neutralization/titration
NaVa = NbVb
nMaVa = nMbVb
1.0M H2SO4 = 2N (1M x 2 equivalents/mole) and 1 equivalent H2SO4 = 49 g H2SO4.
0.5M H3PO4 = 1.5N (0.5M x 3 equivalents/mole) and 1 equivalent H3PO4 = 32.7 g.
3.2M Mg(OH)2 = 6.4N (3.2M x 2 equivalents/mole) and 1 equivalent Mg(OH)2 = 29.2 g.
Naming Acids:
1. Binary Acids: an acid made up of hydrogen and a second element. The acid is named by adding the
prefix “hydro-” to the root name of the second element, adding the suffix “-ic” and ending in
acid.
Prefix
Root Name
Suffix
HBr
hydro
brom
ic
acid.
HCl
hydro
chlor
ic
acid.
H2S
hydro
sulfur
ic
acid.
2. Oxyacids: an acid containing hydrogen, oxygen, and a third element. Oxyacids are named as
follows:
a) if the anion name ends in “-ate”, use the anion root name plus the suffix “-ic” and end with
“acid.”
b) if the anion name ends in “-ite”, use the anion root name plus the suffix “-ous” and end with
“acid.”
c) when there are more than two ions for a particular oxyanion, hypo- is used to designate the ion
with the fewest oxygen atoms, and per- designates the ion with the most oxygen atoms.
Ion
Root Name
Suffix
-2
H2SO4
SO4
sulfur
ic
acid
HNO3
NO3nitr
ic
acid
H2SO3
SO3-2
sulfur
ous
acid
HClO
ClOhypo chlor
ous
acid
HClO2
ClO2
chlor
ous
acid
HClO3
ClO3chlor
ic
acid
HClO4
ClO4per
chlor
ic
acid