Experiment 8
Models of Chemical Bonding
OBJECTIVES

To learn about ionic, covalent, and metallic bonds
To describe the general characteristics of these bonds, specifically
names of compounds, type of bond, melting points, and electrical
conductivity

To learn about periodicity and investigate physical properties of
substances

DISCUSSION
For a review, check http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/bond.html
and be sure to refer to chapter 9 of your text book.
Compounds are pure substances that are composed of two or more elements. The full
range of substances can be divided basically into three broad categories, two of which
are ionic and covalent compounds. The bonding between atoms in metals is different
from these two general types and is called a metallic bond.
The type of chemical bond that occurs in a substance in part defines its properties.
Ionic bonds are simply the attractive forces between oppositely-charged ions. Ionic
compounds contain cations (which can be either monatomic or polyatomic) and anions
(which can be either monatomic or polyatomic.) A simple view of the formation of an
ionic compound is that the electrons lost by the cation are gained by the anion. A
formal transfer of electrons can be envisioned which means that the total positive
charge from the cations must be balanced by the total negative charge from the anions.
There are no free electrons floating around. As an example, this transfer of electrons
gives rise to the formula Na2O for sodium oxide rather than NaO, NaO2, or any other
incorrect formula.
Compounds with ionic bonds will generally show properties of conductivity in the
molten state or in an aqueous solution. They have somewhat high melting points, are
soluble in water (polar solvent), insoluble in nonpolar solvent (like chloroform, mineral
oil), and a low vapor pressure (no odor). These properties are explained by the presence
of strong ionic bonds in the compound, formed by the attractions of oppositely charged
ions. When melting or dissolving, these ions are freed from their crystalline structure
or become solvated, thus accounting for their conductivity.
Covalent bonds are bonds formed by the sharing of electrons, in contrast to the
transfer which can be envisioned for an ionic bond. Covalent compounds consist of
single molecules. A molecule is “the smallest part of any substance which has the
qualities of that substance, and which can exist alone in a free state. As an example, a
molecule of water consists of two atoms of hydrogen and one of oxygen”.
Covalent substances can have a variety of properties, depending on how polar these
substances are. The concept of electronegativity allows one to determine if the
covalent bond is polar, or nonpolar. It is a measure of the tendency of an atom to
attract the bonding electrons in a chemical bond toward itself. It is not a preciselydefined quantity like ionization potential or electron affinity, which can be measured
accurately. In fact, there are more than a dozen scales of electronegativity. Luckily, for
introductory courses only one scale is really needed to put the concept to use: this is
the Pauling scale, named after the two-time Nobel Prize winning chemist Linus Pauling.
According to this scale, fluorine is the most electronegative element, with a Pauling
electronegativity of 4.0, followed by oxygen with an electronegativity of 3.5.
Covalent bonds are formed when the electronegativity difference between the two
elements involved in the bond is smaller than 1.7 indicating that the two atoms have
similar attractions for the electrons. This value can be used as a ballpark dividing line
between ionic and covalent bonds. If the difference in electronegativities between the
two atoms is greater than 1.7, then the bond is predominately ionic.
It is unusual, however, that you'll need to make such quantitative determinations,
especially since they are only approximate. For example, the electronegativity difference
between Li and I is 2.5-1.0, or 1.5. This would imply more covalent than ionic
character, but most chemists would classify lithium iodide as an ionic compound
because it is composed of a metal and a nonmetal. A rule of thumb is that ionic
compounds are formed between metals and nonmetals, whereas covalent compounds
are formed between nonmetals. Big exceptions to this rule are compound based on the
ammonium ion. Compounds of these ions are ionic, since by definition an ionic bond is
the electrostatic attraction between oppositely-charged ions in the crystal.
Molecular compounds are classified as polar covalent or nonpolar covalent, based on
their polarity. These two types of molecular substances differ because of the difference
in their molecular polarity. The nonpolar solids are not soluble in water (a polar
solvent), whereas the polar solids are water-soluble. Since there are no charged units
in either type of solid they do not conduct electricity. Nonpolar molecular solids usually
have high vapor pressures and low melting points. The vapor pressure of most polar
molecular compounds is lower, so most do not have an odor.
A metallic bond is generated in metals. The atoms in a metal are arranged in a regular
pattern or lattice.
Bonding in metals can be described as resulting from the electrical attractions among
positively charged metal ions and the ‘sea of electrons’ belonging to the lattice as a
whole. The electrons become the property of ALL atoms and the electrons can move
freely within the molecular orbitals. In a way each electron becomes detached from its
parent atom and the electrons are said to be delocalized. The metal is held together by
the forces of the attraction between the valence electrons that are free to move through
orbits and the positive nuclei. Imagine this to extend over the entire lattice of a metal,
thus accounting for conductivity in both the solid and liquid states. Substances with
metallic bonds, for example: iron, copper, lead and copper-zinc alloy, will generally
show properties of conductivity in the solid and liquid states, insolubility in both types
of solvents, high melting points and low vapor pressures.
The following image is from your textbook and nicely summarizes the three types of
bonding and shows the difference in bonding on a microscale level.
The Three Models
Chemical Bonding
of
MATERIALS AND EQUIPMENT
Sodium chloride
Microwell plates
Tin
Disposable Pipets
Glycerol
Conductivity apparatus
Vitamin C
Melting point apparatus
Iron
Magnet
Carbon
Potassium bromide
Copper(II) sulfate pentahydrate
Ethanol
Aspirin, C9H8O4
Magnesium
Copper
Benzoic acid, C6H5COOH
Iodine
Zinc nitrate
Sulfur
Sodium iodide
PRELABORATORY ASSIGNMENT: Bonding in Chemistry
Questions
1. Describe the bonding that occurs in ammonium sulfate.
2. Explain the basis for conductivity in metals.
3. Metals exhibit conductivity in the solid and molten state, whereas ionic compounds
show conductivity in the molten state, but not in the solid state. Explain.
4. List two chemical characteristics that lead to an elements being classified as a metal.
5. Solutions of ionic substances conduit electricity. Please explain.
6. Solutions of some covalent substances also conduct electricity. Please categorize
covalent substance as to which ones do conduct electricity and which ones don’t.
7. Draw a picture of a working conductivity apparatus.
Procedure
Record all data and observations directly into your notebook.
1. Devise a table, in which you make a list of all elements/compounds provided to you.
2. Enter the chemical formula, structure (check your lecture text –you might have to go
to chapter 10 for this one - or check http://chemfinder.cambridgesoft.com/) and
systematic name for each substance into your table.
3. Predict the type of bonding and organize the substances according to type of bond.
4. Use the microwell plates for testing of solids and solutions.
Add a small amount (spatula tip) of each substance to each well and observe with a
magnifying glass. Enter the appearance into your table.
5. Using the disposable pipets, add a little water to each well (from the previous step).
Stir to dissolve. Does it dissolve? Is it completely soluble?
6. Test the conductivity of each solution you obtained.
7. Repeat the same with ethanol as your solvent.
8. For substances that did not dissolve, check the conductivity in the solid state. Use a
9-V conductivity apparatus to test the solid for electrical conductivity. Do the test by
simultaneously touching the bare wire electrodes to the sample while observing the
light bulb.
9. Bring a magnet close to each solid substance and record your observations.
10. Predict the melting points (~100°C, ~500 °C, ~1000°C, …..). For each class of
compounds, chose one substance, add a spatula tip to a small test tube and gently heat
it over a Bunsen burner. Does it melt? If a substance does not readily melt, increase
the temperature.
11. Using a melt-temp apparatus, determine the melting point of any one substance
available to you (Our melt-temp apparatus can measure temperatures to ~ 250°C).
12. Malleability: Which classes of substances are malleable, which ones are not?
Devise a plan for testing.
Websites
Cool pictures of molecules at http://www.liv.ac.uk/Chemistry/ArtGallery.html
Well, this is a little on the funny side….
http://www.chm.bris.ac.uk/sillymolecules/sillymols.htm
Molecule of the Month: http://www.chm.bris.ac.uk/motm/motm.htm
WASTE DISPOSAL
Dispose of all solutions in the waste containers located in the hood. Solids should be put
into waste container when cool and disposed of with solid waste.
Analysis and Conclusions
1. Look carefully at your results. Do you see any patterns? Summarize your results.
2. Make generalizations on the solubility, conductivity, and melting point for covalent,
ionic, and metallic bonds.
3. Predict the following, based on the patterns established in this experiment
• solubility of sodium iodide (NaI) in water
• relative melting point of sodium iodide
• electrical conductivity of sodium iodide
4. Is the dissolution of sodium iodide in water a physical or a chemical process? Please
explain.