Chapter 18: Chemical Bonds

• Section 1: Stability in Bonding

• Section 2: Types of Bonds

• Section 3: Writing Formulas and Naming Compounds

Section 1: Stability in Bonding

Why do atoms combine to form compounds?• Because the atoms are “looking” for a stable structure• Stability is achieved when the outermost energy level of the

atom is full With the exception of hydrogen and helium atoms, an

atom is structurally and chemically stable when there are 8 electrons in the outer energy level

• The group 8A elements have a full outer energy level and are stable and do not need to combine with atoms

Atoms achieve stability through chemical bonding • Chemical bond – the force that holds the atoms in a

compound together• There are two types of chemical bonds

Ionic – the giving or accepting of electrons to achieve a full energy level

Covalent – the sharing of one or more pair of electrons to achieve a full energy level

ONLY THE ELECTRONS IN AN ATOM’S OUTERMOST ENERGY LEVEL ARE AVAILABLE TO FORM CHEMICAL BONDS

Section 2: Types of Bonds

Ionic bonding• When atoms gain or lose electrons they acquire a charge

Ion – an atom that carries a charge Typically, metal will give up electron(s) and become ions

with a positive Nonmetals will accept electron(s) and become ions with a

negative charge On your Periodic Table, the main group elements (the A’s),

the number before the letter “A” indicates the number of electrons in the outermost energy level. Exs.: Calcium – Group = 2A – 2 electrons Oxygen – Group = 6A – 6 electrons

The B group, or transition, elements can have a varying number of valence electrons

• The formation of ionic bonds Example: the formation of salt (NaCl)

According to the Periodic Table, sodium (Na) is a 1A metal and chlorine (Cl) is a 7Anonmetal

Na will give up 1 electron to gain a noble gas configuration while chlorine will accept 1 electron to achieve the full outer energy level

By giving up a electron, sodium becomes an ion with a + charge, Na+, while chorine, by accepting an electron, becomes an ion with a negative charge, Cl-

Because the two ions have opposing charges, they attract and form an ionic compound

The compound NaCl is an ionic compound

Section 2: Types of Bonds

Ionic bonding• Conclusions in the formation of ionic compounds

metals bond with nonmetals metals give up electrons, nonmetals accept electrons the compound formed is electrically neutral – the (+) charges and

the (–) charges add up to 0

Covalent bonding• A covalent bond forms when two atoms share one or more pair of

valance electrons Only nonmetals form compounds through covalent bonds

The reason for this is that nonmetals have nearly complete outermost energy levels so it is easier for them to share electrons than it is to give up or add electrons when bonding with other nonmetals

Example: the formation of the gas fluorine (F2) According to the Periodic Table, fluorine is a 7A nonmetal, and

needs one electron to achieve stability If the two fluorine atoms bond ionically, one fluorine atom

would give up an electron and the other accept it. As a result, one atom would be stable but the other would be two electrons short of having a full outer level. This is not good. Remember, the whole goal of chemical bonding is that all the atoms involved in bonding achieve stability

Solution: each fluorine atom shares one electron with the other

This sharing of electrons to form stable octets is called covalent bonding, and the smallest particle of a covalent compound is called a molecule

Section 3: Writing Formulas andNaming Compounds

Naming compounds1. Ionic compounds

Use the name of the metal Use the root of the nonmetal with the “ide” suffix Example: MgBr2 - magnesium `bromide Use the name of the polyatomic (table 4, page 619) if

present Example: CaCO3 – calcium carbonate

2. Binary covalent compounds Use the name of the first nonmetal followed by the root of

the second with the “ide” suffix Use a prefix to indicate how many of each element is in a

molecule Always use prefix with 2nd nonmetal Only use prefix with 1st nonmetal when there is more

than one atom in the formula Example: CO2 = carbon dioxide (refer to table 5, page 621

for common prefixes)

Section 3: Writing Formulas andNaming Compounds

The chemical formula contains the symbol for each atom in the compound, and tells something about number of atoms of each element in the compound• The formula for an ionic compound gives the simplest ratio of

the different ions in the pound. Example: Calcium chloride

Formula: CaCl2 The subscript 2 tells you that for every Ca2+ ion there are two Cl- ions

• Covalent compounds Formulas for covalent compound tell the exact number of

each kind of atom in the molecule Example: glucose and sucrose

1. both are sugars – glucose is plant sugar used by plants in photosynthesis and sucrose is ordinary table sugar

2. both are composed of atoms of carbon, oxygen, and hydrogen covalently bonded together

3. the formulas: glucose: sucrose:

Section 3: Writing Formulas andNaming Compounds

Formulas of Ionic CompoundsIn any ionic compound the net charge must equal zero• Use the criss-cross method to determine the formula for

ionic compounds Ex 1: Magnesium fluoride

Mg → 2A metal – will give up 2 electrons F → 7A nonmetal – will accept 1 electron

1. write symbols of the ions showing the charge of each

2. criss-cross the superscript numerals3. the numerals now become the subscript

designating how many of which ion is in the compound4. disregard the + and - signs5. when writing the formula never show the subscript “1”

Section 3: Writing Formulas andNaming Compounds

Formulas of Ionic Compounds

Ex. 2: calcium hydroxide Ca → 2A metal, will lose 2 electrons OH → a polyatomic ion with a net charge of –1

1. criss-cross numeral superscripts2. ignore the signs, numbers become subscripts3. when writing the formula enclose the

OH ion in parentheses to indicate that it is a single unit

Oxidation number – A positive or negative number that indicates how many electrons an atom has gained, lost, or shared to become stable• The oxidation number for the “A” group elements is easy to

calculate Metals – the oxidation # is number with the letter

Ex: the 1A metals have an oxidation number of +1 Nonmetals – the oxidation # is the number that must be

added to reach 8 Ex: Oxygen, 6A, so: 6 + x = 8, x = 8-6, x = 2, the 6A

nonmetals have an oxidation # of -2

Section 3: Writing Formulas andNaming Compounds

• For the “B” group elements, known as the transition metals, determining the oxidation number can be more difficult The atoms of the transition elements can have

their valence electrons distributed between two sublevels

During the formation of an ionic compound, the transfer of electrons may involve one or both of the sublevels

A varying number of electrons can be lost, any number between 2 and 7

It is impossible to predict the oxidation number of the element

• Analysis of the ionic compound can give the simplest whole number ratio of the metal and nonmetal ions Ex: what is the oxidation # of iron in the compound

FeCl3

1. we know that chlorine always forms an ion with a -1 charge

2. we know the net charge of the compound must = 0so: total (-) = -1x3 = -3, and: total (+) = total (-)so: the oxidation # for Fe = +3

Section 3: Writing Formulas andNaming Compounds

• The oxidation number of the transition metal is given as a roman numeral in the compound name Ex: Iron (II) oxide and Iron (III) oxide

• We can now determine the formula for those compounds using the criss-cross method:

Iron (II) OxideIron (III) Oxide