CLASS CONCEPTS

1. Introduction to Chemistry

2. The Periodic Table

3. Quantum Numbers

4. Electron Configuration

5. Chemical Families

6. Oxidation Numbers

7. Chemical Formulas

8. Chemical Names

9. Formula Mass

10. Percentage Composition

11. Reaction Types

12. Balancing Equations

13. The Mole Concept

14. Solution Concentration

15. Stoichiometry

16. Kinetic Theory

17. The Gas Laws

18. Enthalpy & Heat

19. Reaction Rates

20. Acids & Bases

21. pH Scale

22. Salts

23. Net Ionic Equations

24. Redox Reactions

25. Organic Chemistry

26. Nuclear Chemistry

20. Acids and Bases

 
Today three different theoriesWWW are used to define acids and bases. Each of these theories focuses on a slightly different property. Each theory broadens the definitions to include a wider range of substances.

1. The Arrhenius Theory:

  • Acid - produces hydrogen ions, H+ in water solution.

    In an equation, the positive part of a Arrhenius acid will be hydrogen.

  • Base - produces hydroxide ions, OH in water solution.

    In an equation, the negative part of a Arrhenius base will be hydroxide.

Examples of Arrhenius acids: HCl, HNO3, H2SO4

Examples of Arrhenius bases: NaOH, Ba(OH)2, Al(OH)3

Arrhenius acids, bases, and soluble salts are called electrolytes. When each dissolves, electricity-conducting ions are released.

 
2. The Bronsted - Lowry Theory:

  • Acid - proton donor.

    In an equation, a Bronsted - Lowry acid must have hydrogen in its formula.

  • Base - proton acceptor.

    A Bronsted - Lowry base is hard to generalize for all equations. It may be a negative ion.

    You may have to look at the products. Find one that contains hydrogen. If the negative part of this product was in a reactant that did not contain hydrogen, that reactant is most likely the base.

The Bronsted - Lowry Theory focuses on the action of protons in reactions. Since protons are in the nucleus of an atom, the hydrogen ion is the only source of protons in a normal chemical reaction.

B-L broadens the definition of acids and bases. Although hydrogen is required to produce the proton, no specific ions must be formed.

Three important terms are used in association with this theory:

  • Conjugate base - the particle that remains after an acid gives up a proton.
     
  • Conjugate acid - the particle formed when a base accepts a proton.

Remove a proton from a B-L acid to get its conjugate base.
Add a proton to a B-L base to get its conjugate acid.

Acid     Base
HCl     Cl
H2SO4     HSO4
HI     I
H2O     OH
H3O+     H2O

Examples of B-L acids and bases:

1. HCl + H2O → Cl + H3O+
  • HCl is the B-L acid   ....   Cl is its conjugate base
  • H2O is the B-L base   ....   H3O+ is its conjugate acid

2. CO3−2 + HS → HCO3 + S−2

  • HS is the B-L acid   ....   S−2 is its conjugate base
  • CO3−2 is the B-L base   ....   HCO3 is its conjugate acid
Hydronium ion - formed by a hydrogen ion and a water molecule - H3O+

Since a hydrogen ion is nothing more than a proton (a bare positive charge), when formed, this proton is immediately attracted to a polar water molecule forming a hydronium ion. For this reason, hydrogen ions never actually exist in water solution.

 
B-L Acids and Bases

 
3. Lewis Theory:

  • Acid - electron-pair acceptor.

    In an equation, a Lewis acid gets more negative from the left side to the right.

    Positive ions usually act as Lewis acids.

  • Base - electron-pair donor.

    In an equation, a Lewis base gets more positive from the left side to the right.

    Negative ions usually act as Lewis bases.

The Lewis Theory is the broadest of all. According to this theory, ANY reaction involving the exchange of a pair of electrons will have an acid and base.

Examples of Lewis acids and bases:

1. BrO3&minus + MnO2 → Br + MnO4&minus
  • BrO3&minus is the Lewis acid
  • MnO2 is the Lewis base
2. BF3 + F → BF4
  • BF3 is the Lewis acid
  • F is the Lewis base
Other important terms associated with the Lewis Theory:

  • Complex ion - a central positive ion surrounded by bonded ligands.

    The central ion has empty orbitals and can act as an electron pair acceptor, (Lewis Acid).

  • Ligand - a negative ion or polar molecule bonded to the central ion in a complex.

    Ligands have unshared electron pairs to donate, (Lewis Base).

Identifying Acids and Bases

 
Naming Acids

Binary acids are made up of only two elements - hydrogen and a nonmetal.

Naming binary acids:

  • Begin with the prefix hydro-.
     
  • Write the stem, a part of the name of the element that combines with hydrogen.
     
  • Add the suffix -ic.
Examples of binary acids:
    HCl - hydro chlor ic - hydrochloric acid
    HBr - hydro brom ic - hydrobromic acid
Ternary acids are made up of three elements - hydrogen, oxygen, and another element. Naming ternary acids:
  • Write the stem, a part of the name of the third element.
     
  • The most common acid is given the suffix -ic.
     
  • Add the prefix per- for the acid with one more oxygen.
     
  • The suffix -ous is given to the acid with one less oxygen.
     
  • Add the prefix hypo- for the acid with two less oxygen atoms.
Examples of ternary acids:
    HClO4 - per chlor ic - perchloric acid - one more oxygen.
    HClO3 - chlor ic - chloric acid - most common form.
    HClO2 - chlor ous - chlorous acid - one less oxygen.
    HClO - hypo chlor ous - hypochlorous acid - two less oxygens.

Aqua regia (Latin for "king's water") is a highly corrosive, fuming yellow solution formed by mixing one part concentrated nitric acid with three parts concentrated hydrochloric acid. It was named by alchemists because of its ability to dissolve gold, the "king's metal", that no acid could do by itself.

 
Read a historical story about the use of aqua regia.

 
Other important terms related to acids:

  • Amphoteric - a substance that acts as either acid or base, depending on what it reacts with.

    Water is the most common amphoteric substance. In the presence of a proton donor, it acts like a base. In the presence of a proton acceptor, it acts like an acid.

  • Anhydrous - without water.

    Anhydrides are substances that have had water removed.

    Example: Taking the water out of Ba(OH)2 leaves BaO.

  • Acid anhydride - an oxide that produces an acid when dissolved in water.

    Oxides of nonmetals are acid anhydrides.

    Example: SO2 + H2O → H2SO3 (an acid)

  • Basic anhydride - an oxide that produces a base in when dissolved in water.

    Oxides of metals are basic anhydrides.

    Example: Na2O + H2O → 2NaOH (a base)

 
Strong acids and bases ionize completely in water solution.

This rule-of-thumb can be used in our class:

  • HCl, HBr, and HI are the only strong binary acids.
     
  • In strong ternary acids, the number of oxygen atoms exceeds the number of hydrogen atoms by two or more. Examples are H2SO4 and HNO3
     
  • Hydroxides of groups 1 and 2, except Be, are strong bases.
Chemists define the "strength" of acids and bases by their ionization, NOT by how corrosive they are. Hydrofluoric acid is one of the most corrosive of all substances, but is considered a "weak" acid because it does not ionize completely.

Weak acids and bases ionize only slightly in water solution.

This rule-of-thumb can be used in our class:

  • Any binary acid not listed above is weak.
     
  • A ternary acid is weak if the ratio of oxygen to hydrogen is less than two to one. An example is H3PO4
     
  • Any hydroxide not listed above is a weak base.
 
Naming Acids and Bases

 
Acids and Bases