1. Chapter 18: Chemical Bonds
Section 1: Stability in Bonding
Section 2: Types of Bonds
Section 3: Writing Formulas and Naming Compounds

2. 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

3. 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

4. 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

5. Section 3: Writing Formulas and Naming Compounds
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
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)

6. Section 3: Writing Formulas and Naming 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: CaCl 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
both are sugars – glucose is plant sugar used by plants in photosynthesis and sucrose is ordinary table sugar
both are composed of atoms of carbon, oxygen, and hydrogen covalently bonded together
the formulas: glucose: 𝑪 𝟔 𝑯 𝟏𝟐 𝑶 𝟔 sucrose: 𝑪 𝟏𝟐 𝑯 𝟐𝟒 𝑶 𝟏𝟐

7. Section 3: Writing Formulas and Naming Compounds
Formulas of Ionic Compounds
In 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 electrons F → 7A nonmetal – will accept electron
write symbols of the ions showing the charge of each
criss-cross the superscript numerals
the numerals now become the subscript designating how many of which ion is in the compound
disregard the + and - signs
when writing the formula never show the subscript “1”

8. Section 3: Writing Formulas and Naming 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
criss-cross numeral superscripts
ignore the signs, numbers become subscripts
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

9. Section 3: Writing Formulas and Naming 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
we know that chlorine always forms an ion with a -1 charge
we know the net charge of the compound must = 0 so: total (-) = -1x3 = -3, and: total (+) = total (-) so: the oxidation # for Fe = +3

10. Section 3: Writing Formulas and Naming 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 (III) Oxide
Iron (II) Oxide