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
16. Kinetic Theory
17. The Gas Laws
18. Enthalpy & Heat
19. Reaction Rates
20. Acids & Bases
21. pH Scale
23. Net Ionic Equations
24. Redox Reactions
25. Organic Chemistry
26. Nuclear Chemistry
20. Acids and Bases
Today three different theories 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:
Examples of Arrhenius acids: HCl, HNO3, H2SO4
- 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 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:
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.
- 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.
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.
Examples of B-L acids and bases:
1. HCl + H2O → Cl− + H3O+
Hydronium ion - formed by a hydrogen ion and a water molecule - 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
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:
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.
- 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.
Examples of Lewis acids and bases:
1. BrO3&minus + MnO2 → Br− + MnO4&minus
Other important terms associated with the Lewis Theory:
2. BF3 + F− → BF4−
- BrO3&minus is the Lewis acid
- MnO2 is the Lewis base
- BF3 is the Lewis acid
- F− is the Lewis base
Identifying Acids and Bases
- 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).
Binary acids are made up of only two elements - hydrogen and a nonmetal.
Naming binary acids:
Examples of 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.
HCl - hydro chlor ic - hydrochloric acid
Ternary acids are made up of three elements - hydrogen, oxygen, and another element.
Naming ternary acids:
HBr - hydro brom ic - hydrobromic acid
Examples of 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.
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:
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.
- 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.
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