Chem. 152
Dr. Saidane
Acids and Bases
Skills you should have mastered
Conceptual
1. Name binary and ternary acids as well as some common bases.
2. Identify the common physical and chemical properties of acids and bases.
3. Distinguish Arrhenius, and Brønsted-Lowry acids and bases definitions.
4. Differentiate among strong and weak electrolytes and strong and weak acids and bases.
Problem-solving
1. Perform stoichiometry calculations that involve acid-base titrations.
2. Write a net ionic equation for an acid-base reaction.
Descriptive
1. Describe the pH scale.
Acid nomenclature
1. A binary acid is an acid that contains only two different elements: hydrogen and one
more electronegative element. Ex: HF, HCl, HBr, and HI.
2. An oxyacid (also called oxoacid) is an acid that is a compound of hydrogen, oxygen and
a third element, usually a nonmetal. Ex: HIO3, H2SO4, HNO2, HCH3CO2. Oxyacids are
one class of ternary acids, which are acids that contain three different elements.
3. A monoprotic acid is an acid that has only one proton (the formula start with one
hydrogen atom.) Ex: HCl and HIO3.
4. A polyprotic acid is an acid that has more than one proton per molecule (the formula of
the acid starts with more than one atom of hydrogen.) Ex: H2SO4, H3PO4, H2CO3.
Properties of Acids and Bases
Acids
 Acids have a sour taste and cause a stinging sensation on sensitive skin.
 Water solutions of acids conduct electricity: electrolytes.
 Acids react vigorously with certain metals such as zinc, iron, and magnesium producing
hydrogen gas.
 Acids react with metal oxides and carbonates.
 Acids affect indicators. Acids change litmus paper from blue to red and phenolphthalein
from red to colorless.
 Common acids are: vinegar (5% acetic acid), citric acid (found in citrus fruits), muriatic acid
(hydrochloric acid), carbonic acid (found in carbonated drinks), lactic acid (found in
buttermilk).
Bases
 Bases taste bitter and feel slippery.
 Water solutions of bases conduct electricity: electrolytes.
 Bases do not typically react with metals.
 Bases react with non metal oxides.
 Bases change litmus paper from red to blue and phenolphthalein from colorless to pink.
 Common bases are: Aqueous ammonia (household cleaner), detergents, and soaps.
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Acids and bases can be ionic as well as molecular compounds.
Some compounds can react as either acids or bases, depending on what the other reactant is.
Acids and Bases Definitions
1. Arrhenius Acids and bases
An Arrhenius acid is a chemical compound that increases the concentration of hydrogen
ions, H+, in aqueous solution. Ex: HCl is an Arrhenius acid:
HCl (aq)  H+ (aq) + Cl-(aq)
An Arrhenius base is a substance that increases the concentration of hydroxide ion, OH-,
in aqueous solution. Ex: NaOH is an Arrhenius base:
NaOH (aq)  Na+ (aq) + OH- (aq)
2. Brønsted-Lowry Acids and Bases
a. A Brønsted-Lowry acid is a molecule or ion that is a proton donor.
b. A Brønsted-Lowry base is a proton acceptor.
Ex: HCl (aq) + H2O  H3O+(aq) + Cl-(aq).
HCl is an acid (proton donor) and water is base (proton acceptor.)
c. In a Brønsted-Lowry acid-base reaction, protons are transferred from one reactant
(the acid) to another (the base.)
d. The species that remains after a Brønsted-Lowry acid has given up a proton is the
conjugate base of that acid. Ex: HF (aq) + H2O (l)  F- (aq) + H3O+ (aq). HF is the
acid and F- is the conjugate base.
e. The species that is formed when a Brønsted-Lowry base gains a proton is the
conjugate acid of that base. Ex: HF (aq) + H2O (l)  F- (aq)+ H3O+ (aq). H2O is the
base and H3O+ is the conjugate acid.
Proton Exchange Between Water Molecules
Water is amphiprotic, meaning that it can act as both an acid (proton donor) and as a base
(proton acceptor).
2 H2O (l)  H3O+ (aq) + OH- (aq)
In aqueous solutions, the molarities of H3O+ and OH- ions are related as follows: [H3O+]
[OH-] = 10-14. If one concentration is increased, then the other must decrease to maintain the
value of 10-14.
Strength of Acids and Bases
Strong acids and bases will completely dissociate in water, resulting in high concentrations
of H3O+ and OH-, respectively. This means that strong acids and bases are strong electrolytes.
Weak acids and bases only slightly ionize in water to form very few H3O+ or OH- ions. This means
that they are weak electrolytes. The strength of an acid or base is a measure of its degree of
dissociation in water.
For binary acids AH:
1. The more polar the H – A bond  the stronger the acid. Ex H - F is stronger than H
– N , because F is more electronegative than N.
2. The weaker the bond is  the easier it is to lose H+  the stronger the acid. Ex: H –
F is stronger than H – Cl, therefore HF is a weaker acid than HCl.
For oxyacids:
1. The greater the number of oxygen atoms attached to the central atom, the stronger
the acid. Ex: HClO is weaker than HClO2, which is weaker than HClO3.
2. For the same number of oxygen atoms attached to the central atom, then the greater
electronegativity of the central atom, the stronger the acid.
Ex: H – O – Cl, is stronger than H – O – Br, is stronger than H – O – I, because Cl is
more electronegative than Br, which is more electronegative than I.
The pH Scale
Chemists use the pH scale to describe the amount of H3O+ in the solution of acids and bases. The
pH of an acidic solution is less than 7. The pH of a basic solution is higher than 7. The pH of a
neutral solution is 7.
pH = -log [H3O+]
Acid Base Titrations
Titration allows us to know how much acid or base is present in a sample. We use a
neutralization reaction in the titration process in which an acid reacts with a base to produce
water and a salt.
In an acid-base titration, the sample to analyze, also called analyte, is placed in a flask
and the titrant is placed in a burette. If the acid is the analyte, a base is used as a titrant. We often
use an acid-base indicator that changes color when all the titrant is consumed in the
neutralization reaction.
Titration allows the detection of the stoichiometric point. At that point, the number of
moles of H3O+ is equal to the number of moles of OH- in the solution. At the stoichiometric
point the solution consist of a salt in water. The solution is acidic if acidic ions are presents, and
basic if basic ions are present.
Example: In the titration of a strong acid with a strong base, we use phenolphtaleine as an
indicator. When all of the acid has been neutralized, the pH is 7. One extra drop of base solution
after the endpoint increases the pH dramatically. When the pH increases above 7, phenolphthalein
indicator changes from colorless to pink indicating the endpoint of the titration.
Acid-base Stoichiometry Calculations
In titrations we can perform stoichiometry calculations as follows:
1. Convert the volume of the analyte solution into mol of analyte using molarity.
2. Use molar ratio from the balanced neutralization reaction to convert moles of analyte to
moles of titrant.
3. Convert moles of titrant into molarity using total volume of solution.