logo imageChapter 10 Notes
Dr. Floyd Beckford

CHAPTER 10
AQUEOUS REACTIONS: I

General properties of acids
1. They have a sour taste
2. They change the color of indicators. Acids turn blue litmus red.
3. Nonoxidizing acids reacts with metals above hydrogen in the activity series.
4. They neutralize metal oxides and hydroxides.
5. They conduct an electric current in dilute aqueous solutions.

General properties of bases
1. They have a bitter taste.
2. They have a slippery feeling
3. They change the colors of indicators. Bases turn red litmus blue.
4. They conduct an electric current in aqueous solutions.
5. They neutralize acids.

THE ARRHENIUS THEORY
Protonic acids are those containing acidic hydrogen atoms. According to the Arrhenius
· An acid is a substance that contains hydrogen atoms and produces protons (H+ ions). in aqueous solutions.
· A base is substance that contains the hydroxide group and produces hydroxide ions, OH-, in aqueous solutions.

· Neutralization (acid-base reaction) is defined as the combination of H+ ions with OH- ions to form H2O molecules.

H+(aq) + OH-(aq) ---> H2O(l) Neutralization

The proton, H+, does not exist as such in aqueous solution. It attracts the slightly negative oxygen atom of a water molecule. The process is called hydration and we say that the proton is hydrated.

H+(aq) + H2O(l) ---> H3O+(aq)

The H3O++ ion {H(H2O)n} is called the hydronium ion. The degree of hydration varies (usually n < 3).

It is the hydrated ion that gives aqueous solutions of acids their characteristic acidic properties.

THE BRONSTED-LOWRY THEORY
Acidity and basicity arise from the donation and acceptance respectively of protons. This exchange occurs in aqueous solution.

A Bronsted acid is any substance that can donate a proton to any other substance. Acids can be compounds with no ionic charge such as nitric acid or they can be cations or anions.

	HNO3(aq) + H2O(l) <---> H3O+(aq) + NO3-(aq)

	NH4+(aq) + H2O(l) <---> H3O+(aq) + NH3(aq)

	H2PO4-(aq) + H2O(aq) <---> H3O+(aq) + HPO42-(aq)
A Bronsted base is a substance that can accept a proton from any other substance. As with the case of acids, these substances can be compounds with no charge such as ammonia or anions.

	NH3(aq) + H2O(l) <---> NH4+(aq) + OH-(aq)

	CO32-(aq) + H2O(l) <---> HCO3-(aq) + OH-(aq)
An acid-base reaction according to this theory, is the transfer from an acid to base. Neutralization reactions can be described in terms of conjugate acid-base pairs.

When an acid reacts it loses a proton. The product formed from the reaction could itself accept a proton. That is, the product of an acid reacting in turn act as a base. Conversely, the product of the reaction of a base can as an acid.

Most acid-base reactions are reversible in this way.

In the forward reaction, the acid is HCO3- and the base is H2O. In the reverse reaction, the acid is H3O+ and the base is CO32-. Notice that the base in the reverse reaction is the product formed from the acid in the forward reaction. The acid HCO3- and the base CO32- is called a conjugate acid-base pair. All Bronsted acid-base reactions involve two conjugate acid-base pairs. For example, two pairs are HF/F- and NH4-/NH3

In general a conjugate acid-base act pair are species that differ by the presence of one H+.

Two water molecules can interact to produce one hydronium ion and one hydroxide ion by transfer of a proton from one water molecule to another.

2H2O(l) <---> H3O+(aq) + OH-(aq)
This process is called autoionization (self-ionization). In this way water can be said to act as both an acid and as a base.

ACID STRENGTH
The hydronium ion is the strongest acid that can exist in water. Any acid stronger than H3O+ reacts with water to form H3O+ and a weak base.

HNO3(aq) + H2O(l) ---> H3O+(aq) + NO3-(aq)

Likewise, no base is stronger than OH- in water.

H-(aq) + H2O(l) ---> H2(g) + OH-(g)

Recall that the strength of an acid depends on the degree of ionization or dissociation in aqueous solutions. The ease of ionization in turn depends on
1. The bond strength of H-A
2. The stability of the resulting ions in solution

For binary acids, the stronger the H-El bond the weaker the acid. So for the series H-X (X = halogens),

(weakest) HF << HCl < HBr < HI (strongest)

Likewise for the Group VIA elements,

(weakest) H2O << H2S < H2Se < H2Te(strongest)

For ternary acids containing the same central element increase with increasing oxidation state of the central element and with increasing numbers of oxygen atoms.

Down a group, acid strengths increase with increasing electronegativity of the central element (having the same oxidation state).

	e.g.	HBrO4 < HClO4

		H2SeO4 < H2SO4

		HIO3 < HBrO3
ACID-BASE PROPERTIES OF SALTS
· A normal salt contains no unreacted H+ and OH- ions e.g NaCl, KNO3, MgI2

A polyprotic acid contains more than one replaceable protons. Suppose we a polyprotic acid such as H3PO4 reacting with NaOH to form sodium phosphate (a normal salt).

H3PO4 + 3NaOH ---> Na3PO4 + 3H2O

Suppose we use 2 moles of NaOH. Then we form sodium hydrogen phosphate (acidic salt)

H3PO4 + 2NaOH ---> Na2HPO4 + 2H2O

The salt formed can still react with bases because it has a transferable proton and so it is called an acidic salt. Likewise we can form sodium dihydrogen phosphate (acidic salt)

H3PO4 + NaOH ---> NaH2PO4 + H2O

Similarly, we can form basic salts when we use less than stoichiometric amounts of acid and a polyhydroxy base.

	
	Al(OH)3(s) + HCl(aq) ---> Al(OH)2Cl(s) + H2O(l)
				    Basic salt
	Al(OH)3(s) + 2HCl(aq) ---> Al(OH)Cl2(s) + 2H2O(l)
				    Basic salt
	Al(OH)3(s) + 3HCl(aq) ---> AlCl3(aq) + 3H2O(l)	(Stoichiometric amounts)
				    Normal salt
Amphoterism is the term used to describe substances that can act as an acid or as a base. E.g. aluminum hydroxide

	Al(OH)3(s) + 3HCl(aq) ---> AlCl3(aq) + 3H2O(l)		BASE
	Al(OH)3(s) + KOH(aq) ---> KAl(OH)4(s) 			ACID
				    Potassium aluminate

Preparation of acids
The preparation of acids can be generalized as follows:

Nonvolatile acid + salt of volatile acid ---> volatile acid + salt of nonvolatile acid

Volatility is implied by low boiling point.

	H3PO4(l) + NaBr(s) ---> NaH2PO4(s) + HBr(g)
	b.p. 213 °C				b.p. -67.0 °C	
	H2SO4(l) + 2KCl(s) ---> K2SO4(s) + 2HCl(g)
The ternary acids are formed by reacting their acidic oxides with water.
	Cl2O7(l) + H2O(l) ---> 2HClO4(aq)
	PX5 + 4H2O ---> H3PO4 + 5HX	(X = halogens)

THE LEWIS THEORY
This theory is based on the sharing of electron pairs between acids and bases.

· A Lewis acid is a substance that can accept a pair of electrons.
A Lewis acid must have a vacant valence orbital with which to accept the electron pair. This can be done, if needed by expanding the valence shell. Typical Lewis acids are metal cations or electron deficient compounds such as BF3.

· A Lewis base is a substance that can donate a pair of electrons.
Typical Lewis bases are neutral or anionic species based on Groups VA and VIA elements. e.g. NH3

A Lewis acid-base (neutralization) reaction is defined as the formation of a coordinate covalent bond so called to indicate that both electrons in the bond originated on only one of bound atoms.