image Bonding and Naming Chemical Compounds

Georgia Perimeter College


  1. Distinguish between atoms and molecules.
  2. Distinguish between elements and compounds.
  3. Distinguish between pure substances (elements and compounds) and mixtures.
  4. Name simple inorganic ionic compounds.
This section addresses, in whole or in part, the following Georgia GPS standard(s):
  • S8P1. Students will examine the scientific view of the nature of matter.
    a. Distinguish between atoms and molecules.
    b. Describe the difference between pure substances (elements and compounds) and mixtures.
    f. Recognize that there are more than 100 elements and some have similar properties as shown on the Periodic Table of Elements.

This section addresses, in whole or in part, the following Benchmarks for Scientific Literacy:
  • All matter is made up of atoms, which are far too small to see directly through a microscope. The atoms of any element are alike but are different from atoms of other elements. Atoms may stick together in well-defined molecules or may be packed together in large arrays. Different arrangements of atoms into groups compose all substances.
  • Scientific ideas about elements were borrowed from some Greek philosophers of 2,000 years earlier, who believed that everything was made from four basic substances: air, earth, fire, and water. It was the combinations of these "elements" in different proportions that gave other substances their observable properties. The Greeks were wrong about those four, but now over 100 different elements have been identified, some rare and some plentiful, out of which everything is made. Because most elements tend to combine with others, few elements are found in their pure form.

This section addresses, in whole or in part, the following National Science Education Standards:
  • A substance has characteristic properties, such as density, a boiling point, and solubility, all of which are independent of the amount of the sample. A mixture of substances often can be separated into the original substances using one or more of the characteristic properties.
  • Chemical elements do not break down during normal laboratory reactions involving such treatments as heating, exposure to electric current, or reaction with acids. There are more than 100 known elements that combine in a multitude of ways to produce compounds, which account for the living and nonliving substances that we encounter.


A compound is a group of atoms with a specific number and type of atoms arranged in a specific way. Exactly the same elements in exactly the same proportions are in every bit of the compound.

Example: Water is a compound composed of one oxygen atom and two hydrogen atoms. Each hydrogen atom is attached to an oxygen atom by a chemical bond. H2O is the formula for the compound, water.

If any other elements are attached, it is not water. For example, H2S is hydrogen sulfide. Hydrogen sulfide does not have the same types of atoms as water, so it is a different compound.

If a different number of atoms of hydrogen or oxygen are attached, it is not water. H2O2 is the formula for hydrogen peroxide. It might have the right elements in it to be water, but it does not have them in the right proportion.

A molecule is a single formula of a compound joined by covalent bonds. The Law of Constant Proportions states that a given compound always contains the same proportion by weight of the same elements.


Electron Configuration and Valence Electrons

In a stable atom, the number of electrons is equal to the number of protons.

Electrons in atoms are present in discrete orbits or "shells" around the nucleus of the atom.

There is a ranking or heirarchy of the shells, with the shells further from the nucleus having a higher energy.

The innermost electron shell holds only two electrons.

The outermost shell contains the valence electrons. The maximum number of electrons that can occupy the outer shell is eight. When there are eight electrons in the outer shell, it is said to have an octet of electrons.

The valence of an atom is the likely charge it will take on as an ion.

A valence is the amount of positive or negative charge on an ion of an element.

Example: Hydrogen only has one electron. It can lose an electron to become H+, a hydrogen ion, or it can gain an electron to become H-, a hydride ion.

The Octet Rule

The octet rule states that atoms are most stable when they have a full shell of 8 electrons in the outside electron shell.

Octet = 8

An atom with eight electrons in the outer shell is more stable than an atom which as fewer electrons in the outer shell.

The exception to this is Helium (atomic number 2) which only has two electrons in its outer shell. It has a full shell, so it is a stable inert element.

Valence electrons are the only electrons involved in chemical bonds.

Atoms will form chemical bonds with other atoms by either sharing electrons, or by transferring electrons in order to complete their octet and get 8 electrons in the outer shell.


In a stable atom, the number of electrons is equal to the number of protons.

An atom which has a different number of electrons than it does protons is called an ion.

Ions are charged particles. Types of ions:

  1. Cation - A positively charged ion.
    A cation is an atom or group of atoms with a net positive charge, caused by the loss of one or more electrons.
    Examples: Na+, NH4+, Mg+2


  3. Anion - a negatively charged ion.
    An anion is an atom or group of atoms with a net negative charge, caused by the gain of one or more electrons.
    Examples: F-, S2-, NO3-


  5. Polyatomic ion - a group of atoms which function as a group and which has a net positive or negative charge (cation or anion).
    Examples: NH4+ or NO3-

The Periodic Chart can show how the octet rule works. All of the Group I elements have one electron in the outside shell and they all have a valence of plus one. Group I elements will lose that one electron in the outside shell, to become a single positive ion with a full electron shell of eight electrons (an octet) in the s and p subshells under it.


A bond is an attachment among atoms. Atoms may be held together for any of several reasons, but all bonds have to do with the electrons (particularly the outside electrons) of atoms.

There are several types of bonds:

  1. Ionic bonds occur due to a full electrical charge difference attraction.
  2. Covalent bonds occur due to sharing electrons.
  3. There are bonds that come about from partial charges or the position or shape of electrons about an atom.

Image courtesy of NASA.

Ionic Bonds

The attraction between a positive ion and a negative ion is an ionic bond.

Some atoms (such as metals) tend to lose electrons to make the outside ring of electrons more stable. When an atom loses electrons it becomes a positive ion (or cation) because the number of protons exceeds the number of electrons.

Other atoms tend to gain electrons to complete the outside electron ring. The non-metal ions tend to gain electrons to fill out the outer shell. When the number of electrons exceeds the number of protons, the ion is negative. (Non-metal ions and most of the polyatomic ions have a negative charge.)

Ionic Bond Animation

Ionic bond animation created by Michelle, a student at Buford High School,
 and provided courtesy of Debbie Valdez,
Technology Coordinator, Buford High School, Buford, GA.

Ionic compounds - composed of cations and anions which are ionically bonded to each other due to attractions of opposite charges

1. Cations and anions combine in a ratio that produces a neutral compound; smallest whole number ratio is used for formula of an ionic compound.

e.g., Na+ + Cl- --> NaCl (one of each is needed to balance the charges: +1 and -1)

Mg+2 + Cl- ---> MgCl2
(two Cl's are needed to balance the charges since Cl is -1 and Mg is +2 charge)

2. Cation is listed first, then anion in the formula

Nomenclature (Naming rules) of Simple Inorganic Ionic Compounds

(Click here for handout)

1. The names of the ions of metal elements with only one valence, such as the Group 1 or Group 2 elements, is the same as the name of the element.

Name of cation (except omit "ion") followed by name of anion (omit "ion") -- but now we need to look at how to name cations and anions.

2. Naming Type I monatomic cations: (Type I includes Groups IA, IIA, Al+3, Ga+3, Zn+2 and Ag+) Because of their predictable charges, these cations are easy to name:

a) Name of cation element followed by "ion"
e.g., Na+ is the sodium ion
Mg+2 is the magnesium ion

3. Naming Type II monatomic cations: (All metals that are not Type I are Type II cations)

a) There are a number of elements, usually transition elements that having more than one valence, that have a name for each ion. Type II cations can have several different possible charges. It is necessary to specify the charge of this type of cation in the particular situation when naming the cation or the compound containing the cation.

Example: Iron can have various positive charges, so if we simply said "iron cation", how could anyone know which one we were talking about? Fe+2 vs. Fe+3
Ferric iron is an iron ion with a positive three charge.
Ferrous iron is an iron ion with a charge of plus two.

Example: Copper can be Cu+ or Cu+2 etc.

b) To name Type II cations, we must specify the charge by putting the charge in roman numerals in parentheses after the name of the element followed by the word "ion":

Fe+2 is iron(II) ion; Cu+ is copper(I) ion; Mn+3 is manganese(III) ion, etc.

4. Naming monatomic anions: luckily, monatomic anions have predictable charges, so naming them is easy:

a) first part of element name
b) -ide ending
c) word "ion"

The names of the ions of nonmetal elements (anions) develop an -ide on the end of the name of the element. For instance, fluorine ion is fluoride, oxygen ion is oxide, and iodine ion is iodide.

e.g., Cl- is chloride ion, S2- is sulfide ion

Common monatomic anion charges: (You need to remember these)

  • Group VIIA: -1 (F, Cl, Br, I, At)
  • Group VIA: -2 for Oxygen and Sulfur and usually Se and Te
  • Group VA: -3 for N, P and often As
  • Also, H is -1 in combination with a metal cation (but is +1 in combination with a nonmetal anion)

Give systematic names: (HINT: Remember that the object here is to first identify whether the cation is type I or II so that you can decide whether you need to specify the charge. If it is type II, you will need to figure out the charge of the cation deductively -- by looking at the charge of the anion and remembering that the compound formula is written so that the charges balance to equal zero.)

a) CuCl
b) MgBr2
c) Fe2O3
d) MnI4
e) SnO2
f) NaCl
g) AgBr


Try writing formulas given the systematic names: (HINT: remember that the formulas must be written so that the charges are balanced to equal zero)

a) Iron (II) chloride
b) Barium nitride
c) Cobalt (II) oxide


5. There are a number of common groups of atoms that have a charge for the whole group. Such a group is called a polyatomic ion or radical.

a) Polyatomic ions do not have systematic names; therefore, names of common polyatomic ions must simply be learned. Although there are numerous polyatomic ions, we will restrict ourselves to the following most common ones: (LEARN THESE)
  • Acetate CH3COO- (or C2H3O2- )
  • Ammonium NH4+
  • Nitrate NO3-
  • Nitrite NO2-
  • Carbonate CO32-
  • Chromate CrO42-
  • Dichromate Cr2O72-
  • Permanganate MnO4-
  • Cyanide CN-
  • Phosphate PO43-
  • Sulfate SO42-
  • Hydroxide OH-
  • Perchlorate ClO4-

b) Once you know the names of the polyatomic ions, ionic compounds containing polyatomic ions are named the same way as those with only monatomic ions.

Examples: Give systematic names for the following:

  • NH4Cl
  • Ca(OH)2
  • Ba(NO3)2
  • CuSO4
  • Cr(CO3)3

Write formulas for the following compounds:

  • Potassium dichromate
  • Iron(II) phosphate
  • Ammonium sulfate


Covalent Bonds



Covalent bond animation created by Michelle, a student at Buford High School,
 and provided courtesy of Debbie Valdez,
Technology Coordinator, Buford High School, Buford, GA.


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Content provided by Dr. Margaret Venable, Dr. Michael Denniston, and Dr. David Wilner.

Some of this information is used with permission of Chemtutor.

Page created by Pamela J.W. Gore
Georgia Perimeter College,
Clarkston, GA

Page created November 16 - December 19, 2006