A Chemist’s View of
Explosives:
Chemical Bonding 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.
http://www.visionlearning.com/library/module_vie
wer.php?mid=55
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.
http://web.visionlearning.com/custom/che
mistry/animations/CHE1.7-anH2Obond.shtml
Ionic Bonding occurs between metals and
nonmetals.
Covalent Bonding occurs between
nonmetals.
Bonds (and
compounds) form
in order to obtain
an electron
configuration like
that of noble gases!
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 that right the dot diagram for the new
ionic compound, including charges.
http://www.beyondbooks.com/psc92/3b.asp
A. Lewis Dot Structures for Ionic Compounds
(compounds held together by ionic bonds –
usually a M
Example: a.) Sodium and Chlorine
•
Na
+
•• •
• Cl •
••

[ Na
]+
+ [
•• •
•• Cl
• ]••
Examples:
b.) Magnesium and Oxygen
Examples:
c.) Aluminum and Oxygen
Examples:
d.) Calcium and Fluorine
Examples:
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
Copper sulfate has a
triclinic crystal structure.
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.
http://www.visionlearning.com/library/module_
viewer.php?mid=55
Remember what happens when an
ionic bond forms?
• One or more electrons
from 1 atom are
removed and
attached to another
atom, resulting in a
cation and an ion
which attract each
other
Write this down!!
Ionic Compounds
- never exist as individual
formula units, are solids
Molecular Compounds
-can exist as an individual
molecule, are usually liquids or gases
http://web.jjay.cuny.edu/~acarpi/NSC/6-react.htm
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.
Nonpolar Covalent Bond (equal sharing)
Oxygen Atom
Oxygen Atom
Oxygen Molecule (O2)
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 +.
Bond Polarity
• Electronegativity
–Attraction an atom has for a shared
pair of electrons.
–higher e-neg atom  –lower e-neg atom +
Polar Covalent Bonds: Unevenly
matched, but willing to share.
The H-O bonds in water are polar covalent because
oxygen is more electronegative than hydrogen, and
therefore electrons are pulled closer to oxygen.
Practice: Find the differences in
electronegativity (EN page 151 or on
calculator) 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
Another Example: cesium-fluorine
bonding
Cs EN = 0.7
F EN = 4
4 - 0.7 = 3.3
Ionic Bond
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).
Ar
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).
http://chemsite.lsrhs.net/d_bonding/flashLewis.html
Practice: Draw the electron dot
structures for the following:
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
a. Hydrogen atoms have less then an
octet.
Hydrogen only needs 2 valence electrons.
H2
HN3
F. Resonance – a concept in which two or
more Lewis structures for the same
arrangement of atoms (resonance
structures) are used to describe the bonding
in a molecule or ion. To show resonance, a
double-headed arrow is placed between a
molecule’s resonance structures.
Example: Ozone
G. Coordinate Covalent Bond is formed
when one atom contributes BOTH bonding
electrons in a covalent bond.
Examples:
carbon monoxide
SO42-
Chemical Bonding 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.”
Metallic Bonds: Mellow dogs with plenty
of bones to go around.
A Sea of Electrons
Metals Form Alloys
Metals do not combine with metals. They form
Alloys which is a solution of a metal in a metal.
Examples are steel, brass, bronze and pewter.
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”.
Cl
Cl
Cl
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
2.) H2S
3.) CO2
4.) BF3
5.) CH4
6.) ClO2-1
7.) CH3Cl
8.) PO4-
9.) MgCl2
III. Hydrates: Some compounds trap water inside their
crystal structure and are known as hydrates. You will
not be able to predict which compounds will form
hydrates. CuSO4 5H20 is an example of a hydrate.
This says that one formula unit of cupric sulfate will trap
5 molecules of water inside its crystal.
Hydrates are named by naming the ionic compound by the
regular rules and then adding (as a second word) a
prefix indicating the number of water molecules. You
will use the word “hydrate” to indicate water. The above
compound would be called cupric sulfate pentahydrate.
To find the formula mass of a hydrate, simply find the mass
of the ionic compound by itself and then ADD the mass
of water molecule(s) to that mass.
Practice: What is the formula mass of barium chloride
dihydrate?
1. What is the formula mass of aluminum sulfate
octahydrate?