1. Chapter 17: Acids and Bases
Acid-base reactions involve proton (hydrogen ion, H+) transfer
The generalization of the Arrhenius definition of acids and bases is called the Brønsted-Lowry definitions:
An acid is a proton donor
A base is a proton acceptor
This allows for gas phase acid-base reactions
Arrehenius definiton: Acids contain H+ and Bases contain OH-

2. Species that differ by a proton, like H2O and H3O+, are called conjugate acid-base pairs
The reaction of HCl and H2O. HCl is the acid because it donates a proton. Water is the base because it accepts a proton.

3. (a) Formic acid transfers a proton to a water molecule
(a) Formic acid transfers a proton to a water molecule. HCHO2 is the acid and H2O is the base. (b) When a hydronium ion transfers a proton to the CHO2- ion, H3O+ is the acid and formate ion is the base.

4. An amphoteric substances can act as either an acid or base
These are also called amphiprotic, and can be either molecules or ions
For example, the hydrogen carbonate ion:

5. Acids that are less than completely ionized are called weak acids
The strength of an acid is a measure of its ability to transfer a proton
Acids that react completely with water (like HCl and HNO3) are classified as strong
Acids that are less than completely ionized are called weak acids
Bases can be classified in a similar fashion:
Strong bases, like the oxide ion, react completely
Weak bases, like NH3, undergo incomplete reactions

6. The strongest acid in water is the hydronium ion
If a more powerful proton donor is added to water, it quantitatively reacts with water to produce H3O+
Similarly, the strongest base that can be found in water is the hydroxide ion, because more powerful proton acceptors react quantitatively with water to produce OH-

7. Acetic acid (HC2H3O2) is a weak acid It ionizes only slightly in water
The hydronium ion is a better proton donor than acetic acid (it is a stronger acid)
The acetate ion is a better proton acceptor than water (it is a stronger base)
The position of an acid-base equilibrium favors the weaker acid and base

8. This can be generalized:
Stronger acids and bases tend to react with each other to produce their weaker conjugates
The stronger a Brønsted acid is, the weaker is its conjugate base
The weaker a Brønsted acid is, the stronger is its conjugate base
These ideas can be applied to the binary acids (acids made from hydrogen and one other element)

9. The strengths of the binary acids increases from left to right within the same period
For example, HCl is stronger acid than H2S which is a stronger acid than PH3
The strengths of the binary acids increase from top to bottom within a group
For example, HI is a stronger acid than HBr which is a stronger acid than HCl
Trends are also present in the oxoacids (acids of hydrogen, oxygen, and one other element)

10. When the central atom holds the same number of oxygen atoms, the acid strength increases from the bottom to top within a group and from left to right within a period
Acid strength: HClO4 > HBrO4 > HIO4
Acid strength: HClO4 > H2SO4 > H3PO4
For a given central atom, the acid strength of an oxoacid increases with the number of oxygens held by the central atom
Acid strength: H2SO4 > H2SO3

11. The strength of an acid can be analyzed in terms the the basicity of the anion formed during the ionization
The basicity is the willingness of the anion to accept a proton from the hydronium ion
Consider H2SO4 and H3PO4:

12. The anions are:
In oxoanions, the lone oxygens carry most of the negative charge, making the hydrogen sulfate ion a weaker base than the hydrogen phosphate ion

13. In terms of the percentage of molecules that are ionized, sulfuric acid is a stronger acid than phosphoric acid
There is a third definition for acid and bases
It is a further generalization, or broadening, of the species that can be classified as either an acid or base
The definitions are based on electron pairs and are called Lewis acids and bases

14. A Lewis acid is any ionic or molecular species that can accept a pair of electrons in the formation of a coordinate covalent bond
A Lewis base is any ionic or molecular species that can donate a pair of electrons in the formation of a coordinate covalent bond
Neutralization is the formation of a coordinate covalent bond between the donor (base) and acceptor (acid)

15. NH3 (a Lewis base) forms a coordinate covalent bond with BF3 (a Lewis acid) during neutralization. NH3BF3 is called an addition compound because it was made by joining two smaller molecules.

16. Carbon dioxide (a Lewis acid) reacts with hydroxide ion (a Lewis base) in solution to form the bicarbonate ion. The electrons in the coordinate covalent bond come from the oxygen atom in the hydroxide ion.

17. Lewis acids:
Molecules or ions with incomplete valence shells (for example BF3 or H+)
Molecules or ions with complete valence shells, but with multiple bonds that can be shifted to make room for more electrons (for example CO2)
Molecules or ions that have central atoms capable of holding additional electrons (usually, atoms of elements in Period 3 and below, for example SO2)

18. Consider a proton transfer from the Lewis perspective
Lewis bases:
Molecules or ions that have unshared pairs of electrons and that have complete shells (for example O2- or NH3)
All Brønsted acids and bases are Lewis acids and bases, just like all Arrhenius acids and bases are Brønsted acids and bases
Consider a proton transfer from the Lewis perspective

19. For example, the proton transfer between the hydronium ion and ammonia:

20. An amphoteric species is capable of acting as both an acid and base
Aluminum oxide is an example of an amphoteric compound
SKIP

21. This equilibrium is described by the ion product of water
Water undergoes self-ionization or autoionization making it a weak electrolyte
This equilibrium is described by the ion product of water
What is Kc?
In pure water, and even dilute aqueous solutions, the molar concentratation of water is essentially a constant, wither a value of 55.6 M.

22. At 25oC in pure water it has been found that
In any aqueous solution, the product of [H+] and [OH-] equals Kw
This provides an alternate a way to define the acidity or basicity of a solution
Solutes can change concentrations of hydrogen ion and hydroxide, but they can’t alter Kw.

23. Neutral solutions: [H3O+] = [OH-] or [H+] = [OH-]
Acidic solutions: [H3O+] > [OH-] or [H+] > [OH-]
Basic solutions: [H3O+] < [OH-] or [H+] < [OH-]
To make the comparison of small values of [H+] easier, the pH was defined:
In terms of the pH:
Neutral solutions: pH = 7.00
Acidic solutions: pH < 7.00
Basic solutions: pH > 7.00
Do sum of pH + pOH = 14.

24. The pH of some common solutions
The pH of some common solutions. [H+] decreases, while [OH-] increases, from top to bottom.

25. An acid-base indicator is a species that changes color based on the pH
The pH of a solution can be measured with a pH meter or estimated using a visual acid-base indicator (see Table 17.3, page 762)
An acid-base indicator is a species that changes color based on the pH
Calculating the pH of a strong acid or base is “easy” because they are 100% dissociated in aqueous
For example, the pH of 0.10 M HCl is 1.00 and the pH of 0.10 M NaOH is 13.00
Example 17.6

26. In the last example it was assumed that the total concentration of [H+] was due to the strong acid (HCl) and [OH-] was due to the strong base (NaOH)
This assumption is valid because the autoionization of water is suppressed in strongly acidic or strongly basic solutions
This assumption fails for very dilute solutions of acids or bases (less than 10-6 M)