In ionic bonding valence electrons are actually TRANSFERRED

Chemical Bonding (Covalent, Ionic, and Metallic) Notes

I. Chemical bond: a mutual electrical attraction between the nuclei and valence electrons of different
atoms that binds the atoms together. Another way to describe a chemical bond is to say the attractive forces
between atoms or ions in compounds. In ionic compounds it is an attractive force between positive and negative ions.

In ionic bonding valence electrons are actually TRANSFERRED between a nonmetal and metal. This
happens because a non-metallic atom is much more electronegative and it can pull electrons away from the less
electronegative metallic atom. In an ionic compound the positive and negative ions combine so that the overall
charge is zero.

Sometimes the more electronegative atom is not “powerful” enough to completely take away the electrons
from another atom so the atoms SHARE electrons. This sharing of electrons is called a covalent bond.

Ionic Bonding occurs between metals and nonmetals. Covalent Bonding occurs between nonmetals.

II. Formation of Ionic Bonds and Ionic Compounds

A. Electron Dot Structures: show the placement and transfer of valence electrons. Rules to remember when drawing
electron dot structures:

1. Only valence electrons are shown. Valence electrons are the electrons in the outermost s and p sublevels.
Transition metals could also have d sublevel valence electrons.

2. Valence electrons are shown as dots and are not drawn randomly! They are arranged around the element's
symbol to correspond to the elements electron configuration. (Only 2 dots or electrons per side.)

3. Follow the Octet Rule which sates that atoms form bonds in order to obtain 0 or 8 valence electrons,
because of this electron dot structures will show no more than 8 electrons for each atom or ion. Another way
to think of the Octet rule: Atoms react by changing the number of their electrons so as to acquire the stable
electron configuration of a noble gas.

B. Electron Dot Structures for Atoms: Write the element's symbol and place the appropriate number of dots to
represent the valence electrons around the symbol. (The electron configuration is given to help you understand the idea
of valence electrons.)

a.) Ca [Ar]4s2 b.) Li [He]2s1

c.) Be [He]2s2 d.) O [He]2s22p4

e.) Br [Ar]4s23d104p5

C. Electron Dot Structures for Ions: Ions form when atoms lose or gain valence electrons.
(1.) cations - these form when atoms have LOST valence electrons. To draw the dot structures write the

symbol, put [ ] around the symbol, and the charge of the ion outside the [ ]. There are NO dots because there
are NO valence electrons! (You may want to write the electron configuration for the atom to help you see what
happens when it ionizes.)

a.) Mg ion b.) Li ion

c.) Al ion d.) Ba ion

(2.) anions-these form when atoms have GAINED valence electrons. To draw the dot structures write
the symbol, draw 8 dots around the symbol, put [ ] around the symbol, and the charge of the ion outside the [ ].
(You may want to write the electron configuration for the atom to help you see what happens when it ionizes.)

a.) S ion b.) Br ion

c.) N ion d.) P ion

(3.) Transition and Inner Transition Elements-the number of valence electrons for these are harder
to predict based on their position on the periodic table because some of these elements have valence electrons
in the d sublevel. Example:

a.) How many valence electrons does an atom of iron have?
To answer this question write the electron configuration for iron:

Are there any unstable electrons in the d level?

When iron ionizes what are the possible ions?

b.) How many valence electrons does an atom of titanium have?
Electron configuration for titanium:

Are there any unstable electrons in the d level?

When iron ionizes what are the possible ions?

D. Pseudo-noble gas electron configuration-elements that cannot acquire a noble gas electron configuration, but can
become somewhat stable with 18 electrons in their outer shell. Examples are: Hg+2, Cd+2, Au+1, Cu+1

E. Electron Dot Structures for Ionic Compounds:
1. Write the electron dot structure for each of the elements involved.
2. Draw arrows from the electrons of the metallic atom to the non-metallic atom. This shows the transfer of electrons.
3. After the Write the dot diagram for the new ionic compound, including charges.
Examples:

a.) Sodium and Chlorine

b.) Magnesium and Oxygen
c.) Aluminum and Oxygen
d.) Calcium and Fluorine
e.) Sodium and Nitrogen

F. Characteristics of ionic compounds (compared to molecular compounds)

-higher melting points -higher boiling points

-generally hard, brittle solids -when melted or dissolved in water they can conduct electricity

-shapes are crystalline in nature (page 177) – square/cube

Lattice Energy: the energy released when one mole of an ionic crystalline compound is formed from gaseous ions.
Negative values for lattice energy mean that energy was released when the ionic crystal is formed.

Formula unit: the simplest collection of atoms from which an ionic compound’s formula can be established. A formula
unit is to an ionic compound as a molecule is to a covalent compound.

III. Formation of Covalent Bonds and Molecular Compounds
A. Covalent Bonds – a bond in which electrons are shared. Which compounds have covalent bonding?

1. Molecular (or covalent) compounds - these are two NON-METALS. These compounds always
have covalent bonding

2. Polyatomic ions (PO-3, NO-1, CN-1). These ions are held together with covalent bonds.

B. Type of Covalent Bonds
l. Nonpolar covalent bond-a covalent bond in which the bonding electrons are shared equally by the
bonded atoms, resulting in an evenly balanced charge. If the difference in electronegativity between
two bonded atoms is less than 0.3 a nonpolar bond will exist.

2. Polar covalent bond-a bond is which the bonded atoms do NOT share the bonding electrons
equally. A polar covalent bond is a bond in which the atoms have an unevenly balanced charge. If the
difference in electronegativity between two bonded atoms is from 0.3 to 1.7, a polar bond will exist If
the difference in EN is less than 0.3 then the bond is nonpolar covalent. The atom with the greater
electronegativity will pull the electrons toward it, giving that atom a slightly negative charge. A partial
negative charge is shown by - and the less electronegative atom will have a partial positive charge,
designated +.

Practice: Find the differences in electronegativity (EN) in the following pairs of atoms. Designate
which, if any, atom is partially negative and partially positive.

a. H and Cl b. F and Br c. S and I d. O and H

D. The Octet Rule and Dot Structures -chemical compounds tend to form so that each atom, by gaining,
losing, or sharing electrons, has an octet of valence electrons. Electron dot structures (also known as Lewis
dot diagrams) show valence electrons as dots around the element’s symbol. Dot structures for molecules
show atoms sharing dots (covalent bonds).
Covalent bonds are single, double, or triple

single bond-two atoms share one pair of electrons (1 sigma bond)
double bond-two atoms share two pair of electrons (1 sigma and 1 pi bond)
triple bond-two atoms share three pair of electrons (1 sigma and 2 pi bonds)

Rules for correctly illustrating the dot structure of a molecule:

1. Add up the TOTAL number of valence electrons in the substance
a.) be sure to subtract 1 electron if it is a positively-charged ion (NH4+1)

b.) be sure to add electrons for each negative charge on an ion (SO4-2)
2. Decide what is the central atom. The central atom is the one that is least represented. (or the least
electronegative)
3. Hook the particles together using a short straight line (or 2 dots) to indicate a covalent bond between
atoms. Each of these "bonds" represents 2 shared electrons.
4. Subtract the number of electrons used in "hooking" the atoms together from the total valence electrons.
5. Use the "leftover" electrons, if any, to fill the octets of the peripheral atoms.
6. Place anymore "leftover" electrons on the central atom (in pairs).
Practice: Draw the electron dot structures for the following:
1. carbon tetrachloride (CCl4)

2. OF2

3. NF4+1

4. PCl3

5. CO2

6. N2

E. Exceptions to the Octet Rule
Hydrogen will have less then an octet. Hydrogen only needs 2 valence electrons
H2

NH3

Covalent Bonding and Nomenclature Notes Part 2

I. Metallic Bonds-a third type of bond. This is what holds pure metal atoms together. Metallic bonding accounts for
many physical properties of metals, such as strength, malleability, ductility, thermal and electrical conductivity, opacity,
and luster.

What happens to form a metallic bond?

1. each metal donates its valence electron(s) to form an electron cloud

2. this leaves positive particles which are "cemented" together with the negative electron cloud, often called a “sea
of electrons.”

II. Polarity - Polar and nonpolar molecules - if a molecule contains a polar bond, is the molecule itself polar also? It depends!!
1. Polar molecules - a polar molecule is positive at one point and negative at another point. For
example, HBr contains a polar bond. As a result the hydrogen side of the molecule is partially
positive and the bromine side of the molecule is partially negative. It "acts like a magnet".

Water is a molecule with two polar bonds. A molecule of water is also polar because there is an
area of positive charge on the hydrogen atoms and an area of partially negative charge on the oxygen. Its
bent shape allows it to act somewhat like a magnet. The presence of 2 unshared pair of electrons is
a main factor for it being a polar molecule.

2. Nonpolar molecules - Carbon tetrachloride has four C-Cl bonds. Each bond is a polar covalent bond.
The molecule itself is nonpolar because of its

1.) shape. It is perfectly symmetrical, and
2.) the partially-positive carbon in the center which is covered by the 4 partially negative chlorine atoms. It cannot

"act like a magnet".

3. Helpful hints and practice:

A. Hints to help you decide if a molecule is POLAR:
1. Does it have at least one polar bond? If so, it's probably polar.
2. Does it have any unshared pairs of electrons around the central atom? If so, it is probably polar.
3. Can the molecule act like a magnet? If so, it is probably polar.

B. Practice: Which of the following molecules are polar and which ones are nonpolar molecules?
If the molecule is polar, tell why it is polar.
1.) SO2

4.) BF3

2.) H2S

5.) CH4

3.) CO2 6.) ClO2-1