Chapter 4 Notes Dr. Floyd Beckford

In the periodic table, elements are arranged in order of increasing atomic number, which is essentially a statement of the periodic law: The properties of the elements are periodic functions of their atomic numbers.

In the periodic table, the vertical columns are called groups and the horizontal rows are called periods. All members of a group have similar chemical and physical properties; in a period the members change properties progressively across the table.

Some groups have common names:
Alkali metals: Group IA; Li, Na, K, Rb, Cs
Alkaline earth metals: Group IIA; Be, Mg, Ca, Sr, Ba
Halogens: Group VIIIA; F, Cl, Br, I
Noble (also called rare or inert) gases: He, Ne, Ar, Kr, Xe

Classification of Elements
There are three primary classifications of the elements: metals, nonmetals and metalloids (semi-metals)

The metallic elements, with the exception of mercury, are solids at room temperature. Most metals have a high melting point and a high density.

Nonmetals can be solid, liquid or gas. Solid or liquid nonmetals tend to have low melting points and densities.

A number of the elements (B, Si, Ge, As, Sb and Po) have properties intermediate between metals and nonmetals. The elements are the metalloids. Aluminum is sometimes classified as a metalloid.

Metallic character increases down a group and decreases across a period. Nonmetallic character increases up a group and across a period.

AQUEOUS SOLUTIONS
Many reactions take place in solution. A solution is a homogeneous mixture composed of a solute dissolved in a solvent.

Water is the most common solvent. Solutions that have water as the solvent are called aqueous solutions. Solutes in aqueous solutions are of two major kinds:
1. Electrolytes: These are substances whose dilute aqueous solutions are conductors of electricity.
(a) Strong electrolytes are substances whose aqueous solutions are very good conductors of electricity.
(b) Aqueous solutions of weak electrolytes conduct electricity poorly.
1. Nonelectrolytes: Solutions of these substances do not conduct electricity.

The major difference between electrolytes and nonelectrolytes is that electrolytes contains ions or are capable of producing ions in aqueous solution, whereas nonelectrolytes do not have any of these properties.

The formation of ions in solution occur in two ways:
(a) Dissociation: In the dissociation of a salt, the salt already contains ions that separate or dissociate when the substance is dissolved in water.

We have mentioned before that electrolytes can be strong or weak. The strength of an electrolyte depends on the degree of dissociation or ionization. Strong electrolytes are essentially 100% “ionized”. Weak electrolytes are only partially “ionized”.

Examples of strong electrolytes include strong acids, strong bases and most salts.
An acid is a substance that produces hydrogen ions, H⁺, in aqueous solutions.
A base is a substance that produces hydroxide ions, OH^-, in aqueous solutions.
A salt is a compound that contains a cation other than H⁺ and an anion other than OH^- or O^2-.

The H⁺ is nothing more than a bare proton; it does not exist by itself in aqueous solution. Rather, in water, a proton combines with a water molecule to form a hydrated hydrogen ion, H3O+. This hydrated ion is called a hydronium ion.

SOLUBILITY
The term solubility is used to describe the amount of one substance that will dissolve in another. While there is no sharp dividing line between “soluble” and “insoluble” compounds, they can be used relatively describe the solubility of a substance.

Solubility Rules: 1. All common compounds of the alkali metals as well as the ammonium ion, NH₄⁺, are soluble in water.
2. All nitrates, NO₃^-; perchlorates, ClO₄^-; acetates, CH₃OO^- are soluble in water.
3. The common halides (Cl^-, Br^-, I^-) are soluble in water except the silver, mercury and lead halides.
4. The common sulfates, SO₄2-, are soluble in water except PbSO₄, BaSO₄ and HgSO₄.
5. The common carbonates, CO₃^2- and phosphates, PO₄^2-, are insoluble in water except those of the alkali metals and NH₄⁺.
6. The common sulfides, S^2-, are insoluble in water except for the alkali metals, the alkaline earth metals and NH₄⁺.
7. The common metal hydroxides, OH^-, are insoluble except those of the alkali metals, Ca, Ba and Sr.

REACTIONS IN AQUEOUS SOLUTIONS
In chemical reactions involving ionic compounds, some of the ions do not react. The non-reacting ions are called spectator ions.

The state of the potassium ion, K⁺, remains unchanged during the reaction. K⁺ is a spectator ion. Similarly, NO₃^- is also a spectator ion.

Since equations are written to represent chemical reactions, we can write equations to various kind of information about the process taking place. Three major kinds of equations are written:

1. Formula unit equation (or molecular equation): In these equations, the complete formulas for all compounds are shown.

2. Total ionic equation: Compounds are written in the form in which they are predominantly present. Strong electrolytes in solution are written in their ionic form. Nonelectrolytes, weak electrolytes, precipitates and gases are written in their molecular forms.

3. Net ionic equation: Only those molecules or ion that have changed (reacted) are included. The spectator ions are omitted.
H⁺(aq) + OH^-(aq) ---> H₂O(l)

Example: The reaction of Na₂CO₃ with H₂SO₄
1. Na₂CO₃(aq) + H₂SO₄(aq) ---> Na₂SO₄(aq) + H₂O(l) + CO₂(g)
2. 2Na⁺(aq) + CO₃^2-(aq) + 2H⁺(aq) + SO₄^2-(aq) ---> 2Na⁺(aq) + SO₄^2-(aq) + H₂O(l) + CO₂(g)
3. CO₃^2-(aq) + 2H⁺(aq) ---> H₂O(l) + CO₂(g)

Example: Reaction of NaOH with acetic acid, CH₃COOH
1. CH₃COOH(aq) + NaOH(aq) ---> CH₃COONa(aq) + H₂O(l)
2. CH₃COOH(aq) + Na⁺(aq) + OH^-(aq) ---> CH₃COO-(aq) + Na⁺(aq) + H₂O(l)
3. CH₃COOH(aq) + OH⁺(aq) ---> CH₃COO^-(aq) + H₂O(l)
In this example acetic acid, a weak acid, is written in the molecular form.

A summary of rules to observe when writing ionic equations:
(a) Strong electrolytes are written in their ionic form.
(b) Nonelectrolyte, weak electrolytes, gases and precipitates are written in their molecular form.
(c) The net ionic equation should only include those substances that have undergone a chemical change.
(d) Equations must be balanced, both in atoms and in charge.

I. PRECIPITATION REACTIONS.
In precipitation reactions, ions in solution combine to form an insoluble product, the precipitate, that falls out of solution.

II. ACID-BASE REACTIONS
In simple cases:

In general, for a weak acid-strong base reaction,

OXIDATION NUMBERS
The oxidation number (o.n.) of an atom (sometimes called its oxidation state) represents the number of electrons lost, gained or shared unequally by the atom.

Rules for assigning oxidation numbers:
1. All elements in their free state (uncombined with other elements) have an oxidation number of 0. E.g. N in N₂ = 0.
2. The o.n. of a monatomic ion is the same as its charge. e.g. Cl in Cl^- = -1, Al in Al³⁺ = +3
3. The algebraic sum of the o.n. of the elements of neutral compound is 0.
4. The algebraic sum of the o.n. of the elements in a polyatomic ion is equal to the charge on the ion.
5. (i) In all its compounds H has o.n. = +1 except when combine to a metal in which case it is -1. (ii) In all its compounds O has o.n. = -2 except in peroxides in which case it is -1 and when combined to F (o.n. = +2). (iii) In all its compounds F has o.n. = -1.

I. OXIDATION-REDUCTION REACTIONS
These are usually referred to as redox reactions because oxidation and reduction are simultaneous processes.

In oxidation, the o.n. of an element increases as a result of losing electron(s). In reduction, the o.n. of an element decreases as a result of gaining electrons.

OXIDATION	REDUCTION
Increase in o.n.	Decrease in o.n.
Loss of electron(s)	Gain of electron(s)

Obviously in redox reactions, the oxidized substance is the reducing agent and the reduced substance is the oxidizing agent.

Zn gets oxidized to Zn²⁺ and is therefore the reducing agent.

Example: Identify the oxidizing and reducing agent in the following reaction.

II. DISPLACEMENT REACTIONS
As the name implies, displacement reactions are those in which one element displaces another from a compound in solution. These reactions are always redox equations.

The activity of a metal relates to its tendency to form positive ions. Active metals displace less active metals or hydrogen from their compounds. This statement is the basis of the Activity Series.

Principles for using the Activity Series:
1. Reactivity decreases from top to bottom.
2. A free metal will displace those below it from aqueous solutions
3. Free metals above hydrogen react with nonoxidizing acids in solution to liberate hydrogen gas.

NAMING INORGANIC COMPOUNDS
Binary compounds
Binary compounds consist of two elements. The rule is to name the metallic element first and the less metallic element second.
· The less metallic element is named by adding “ide” to the stem of the element.

Some common stems:

C – carb		O – ox			S – sulf
N – nitr		H – hydr		F – fluor
Cl – chlor		Br – brom		I – iod

Examples:	NaCl – sodium chloride
		KF – potassium fluoride		MgO – magnesium oxide
		Ca3N2 – calcium nitride

The method above can be used to name binary ionic compounds where the metals exhibit only one oxidation other than zero. Some metals can form two or more binary compounds with the same nonmetal.

To distinguish between the different oxidation numbers, a Roman numeral in parentheses is written immediately after the name of the metal. The Roman numeral is used only in the name, not in the formula, of the compound.

Examples: 	FeCl2	Iron(II) chloride	o.n. = +2
		FeCl3	Iron(III) chloride	o.n. = +3
		CuCl	Copper(I) chloride	o.n. = +1
		CuCl2	Copper(II) chloride	o.n. = +2	

In binary molecular compounds, elemental proportions are indicated by using a common prefix system for both elements. The prefixes usually used are mono, di, tri, tetra, penta, hexa, hepta, octa, nona and deca to represent compositions of 1 –10 respectively. The prefix mono- is usually omitted.

Examples:	N2O – dinitrogen oxide		PCl3 – phosphorus trichloride
		N2O3 – dinitrogen trioxide	SF6 – sulfur hexafluoride

Ternary compounds
Ternary compounds contains three elements:
(a) a cation, normally a metal or hydrogen
(b) a polyatomic ion containing oxygen and another nonmetal

The rules for naming these compounds are similar to those for naming binary compounds. I.e. the cation is named first followed by the anion.

When the cation is H⁺, the ternary compounds are called oxoacids. To name these acids we place the endings “-ic” and “-ous” after the stem of the nonmetal. The endings are used to indicate the different oxidation states of the nonmetal. The “ic” ending is arbitrarily designated to the higher oxidation state.

	H2SO4 – sulfuric acid 		o.n. of S = +6
	H2SO3 – sulfurous acid 		o.n. of S = +4
	HNO3 – nitric acid 		o.n. of N = +5
	HNO2 – nitrous acid 		o.n. of N = +3

If more than two oxidation states are possible, when the o.n. is higher than in the “ic” acid, the prefix “per” is placed before the stem of the nonmetal. When the o.n. is lower than the “ous” acid, the prefix “hypo” is placed before the stem.

Examples:	HClO		hypochlorous acid	o.n. Cl = +1
		HClO2		chlorous acid		o.n. Cl = +3
		HClO3		chloric acid		o.n. Cl = +5
		HClO4		perchloric acid		o.n. Cl = +7

Note the regular changes in going from one acid to another. Can you explain why?

To name the anion of ternary salts (when the cation is a metal), the ending “ite” and “ate” is added to the stem of the nonmetal for the “ous” and “ic” acids respectively. So sulfuric acid gives sulfate salts and sulfurous acid gives sulfite salts.

Example:	NaNO3 – sodium nitrate (from nitric acid)
		KClO4 – potassium perchlorate (from perchloric acid)
		Mg(OCl)2 – magnesium hypochlorite (from hypochlorous acid)