Chemical Bonds
Chemical bonds are interactions of electrons leading to strong forces of attraction which
holds atoms together in molecules and compounds. Atoms may transfer or share electrons, and
either process may provide for a stable arrangement of electrons between the atoms that results in
the formation of molecules.
Rules for Electron Dot Structures and Bonding Structures
The central atom follows the Octet Rule (usually) and in most cases the least electronegative
nonmetal is surrounded by the other atoms. Other atoms follow the Octet Rule whenever electrons
are available (see exceptions in class). Check to see that every atom has the influence of 8 electrons
and the total number of electrons is correct for that molecule.
Drawing bonding structures (called Lewis structures)
1. select a reasonable "skeleton" for the molecule or polyatomic ion
a. the LEAST electronegative element is usually the central element, except that hydrogen never
is
example: S C S in compound CS2
b. oxygen atoms do not bond to each other except in a few cases such as O2 and O3
2. calculate the total number of outer shell electrons available in all the atoms of the molecule or ion
3. draw a single bond to represent each pair of shared electrons in the skeleton
4. allowing 2 electrons for each shared pair, subtract the total number of electrons already used
5. distribute the remaining electrons in such a fashion as to give each element an octet, if possible
6. for ions, be sure to add (for negative ions) or to subtract (for positive ions) the number of
electrons indicated by the charge on the ion
7. remember that you can use double or triple bonds in order to give elements an octet, but only
when necessary
8. if there are any electrons "left over", place these additional unshared (lone) pairs of electrons into
the skeleton to fill the octet of every group 1,2, 13, 14, 15, 16, 17 element (except hydrogen, which
can only share 2 electrons).
Ionic Bonds
- metals and nonmetals react chemically by the TRANSFER of electrons (from metals to
nonmetals)
- metals form positive ions by losing valence electrons to the nonmetals which then form
negative ions
- positive ions are strongly attracted to the negative ions by the electrostatic attraction
that exist
between unlike charges
- the new substance formed does not resemble either of the original atoms
- this attraction binding unlike ions together is called ionic bonding
example: CaF2
see classroom drawing
Covalent Bonds
- two or more atoms both of which tend to gain electrons during reactions (nonmetals)
may combine by sharing 1 or 2 or 3 pairs of electrons
- the force holding the atoms together is due to the attraction of each atom for the
electrons that are held jointly (a stable condition)
- HYDROGEN, CARBON, NITROGEN, AND OXYGEN are noted for forming covalent
bonds
single covalent bond:
see examples in classroom
double covalent bond:
triple covalent bond:
Homework/Test Problems
First determine if the molecule is ionic or covalently bonded. Then draw the electron dot
structures showing an acceptable bonding structure.
1. H2S
2. F2
3. HF
6. MgO
7. NH3
11. CO2
MgCl2
12. K2S
13. CH4
14. C2H2
15.
16. SiO2
17. NF3
18. HCl
19. CHCl3
20. C2F2
21. C2H6
22. C2H4
23. CHN
24. Si2F4
*25. BF3
8. PBr3
Bonding Information
4. H2O
9. CCl4
5. AlF3
10. CS2
Extra Bonding Notes Bonding Review
The difference in the electronegativities of two elements can be used to predict the
nature of the bond. When this difference is small, the bond is primarily covalent. As the
difference increases, the covalent bonds become increasingly polar. When the difference
becomes even greater, the bond becomes ionic.
Generally the line is drawn at 1.7. When the differences in electronegativities is greater
than 1.7 the bond is ionic (and less than 1.7 is covalent). Another boundary often is
drawn at a difference of 1.0 (sometimes 0.8) to separate polar bonds from nonpolar bonds.
When a molecule behaves as if one end were negative and the opposite end positive, the
molecule is said to be polar. Polar molecules are known as dipoles. A molecule is polar
when there is an uneven distribution of electrons in the molecule.
When two atoms of the same element form a molecule, the shared electrons are
equidistant from the nuclei of the two atoms. This makes the bond nonpolar.
HCl is an example of a two-atom polar molecule. The shared electron pair is attracted
toward the highly electronegative chlorine atom and away from the hydrogen atom. The
resulting concentration of negative charge is closer to the chlorine atom and that end of
the molecule will be slightly negative. The other end will be slightly positive but the
molecule as a whole will be neutral.
Summary:
1) compounds or bonded atoms in molecules are polar is the center of positive charge
does not coincide with the center of negative charge.
2) when a covalent bond is formed between atoms of different electronegativities, the pair
of electrons will be more closely associated with the more electronegative atom, and the
resulting covalent bond will be somewhat polar.
3) the greater the difference between the electronegativities of the atoms involved in the
bond, the greater the polarity of the bond.
4) if the difference in electronegativity is too large, the electrons will be transferred and
ionic bonding will result instead.
5. if both atoms in covalent bond have identical ionization potentials and
electronegativities, no ions are formed and there is no polarity.
Hydrogen bonds: In compounds such as water, ammonia (NH3), and hydrogen fluoride
(HF), the hydrogen atoms are bonded to small atoms of high electronegativity (oxygen,
nitrogen, and fluorine, respectively). The hydrogen atom has only a very small share of
the electron pair that forms the bond. Such molecules are highly polar. In fact, each
hydrogen atom acts largely as exposed proton. It can be attracted to, and form a weak
bond with, the highly electronegative atom of a neighboring molecule. This is called a
hydrogen bond. It is more than just an electrostatic attraction between opposite charges.
It actually has some covalent character.
Hydrogen bonding is responsible for a number of unusual properties. Hydrogen bonding
occurs between water molecules. Water must therefore be raised to a much higher
temperature before the kinetic energy of its molecules becomes great enough to break the
hydrogen bonds between the molecules. Breaking these hydrogen bonds is necessary in
order to boil water. X ray studies show that the three-dimensional structure caused by
hydrogen bonding gives ice crystals a crystalline arrangement with many hexagonal
openings. This open structure accounts for the low density of ice.
Metallic Bonds
Most metals have only one or two valence electrons and low ionization energies. The valence
electrons do not seem to belong to any individual atom but move easily from one atom to another.
Metals can be thought of as positive ions immersed in a “sea” of mobile electrons. The attractive
forces that bind metals atoms together are called metallic bonds. The ease with which the valence
electrons move within the crystal distinguishes the metallic bond from ionic or covalent bonds.
a) metals are good conductors of heat and electricity because of the mobility of their valence
electrons.
b) High luster of metals is the result of the way in which valence electrons absorb and re-emit
light energy that strikes them
c) Metals can be flattened out or stretched out into a wire because the electrons and ions can move
into other positions without breaking up the essential structure.
Summary:
The forces between ions are very strong; so that ionically bonded substances have high
melting and boiling points, and are usually solids at room temperature. Water is usually
capable of dissolving them.
Atoms in covalently bonded substances are electrically neutral, do not conduct electricity,
have low melting and boiling points, and are gases or volatile liquids at room temperatures.
Organic solvents will often dissolve them.
Extension Information on Bonds
Ionic Bonds
(have large differences in electronegativity
Ionic Crystals – electrostatic attractions between ions, NO MOLECULES. Nondirectional
bonds; localized electrons on ions. Examples: NaCl, K2SO4, NH4Cl, (NH4)2SO4
Crystal properties:
1. medium high melting point (600 - 2000 C)
2. medium high boiling points
3. hard and brittle
4. nonconductor of electricity
5. poor conductor of heat
Molecular Crystals – small individual molecules held internally together by covalent
directional bonds. The electrons are localized on molecules. The molecules are attracted to
each other by (1) dipole attraction (2) Van der Waal forces (3) hydrogen bonds. Examples:
HCl, SO2, CO2, CH4, H2SO4, H2O
Crystal properties:
1. very low melting point (-370 to 300 C)
2. very low boiling point
3. soft
4. nonconductor of electricity
5. poor conductor of heat
Covalent Bonds (only very small differences in electronegativity)
Covalent Crystals – all atoms in the crystal are inter bonded by covalent bonds to make one large
crystal. The electrons are usually localized in the bonds. Examples: diamond, SiC, SiO2, graphite
Crystal properties:
1. very high melting point (1200 - 3500 C)
2. very high boiling point
3. very hard and brittle
4. usually a nonconductor of electricity
5. usually a poor conductor of heat
Metallic Bonds
Metallic Crystals – positive nuclei lattice in a cloud of delocalized electrons. Examples: Hg,
Cu, Au, Fe, alloys.
Crystal properties:
1. very low to very high melting point
2. very low to very high boiling point
3. very soft to very hard
4. ductile and malleable
5. good conductor of heat and electricity
Extra Notes on Bonding
Bonding Geometry
Bonding
Pairs
(central
Total Pairs atom
excluding
double
bonds
1
1
2
2
Nonbonding
pairs
Overall
(central
Geometry
atom)
0
0
3
0
2
1
3
linear
linear
trigonal
planar
Molecular
Geometry
linear
linear
trigonal
planar
bent
Example
Hybridization
H2
BeF2
none
sp
BF3
sp2
(SO2)
4
0
tetrahedral
CH4
3
1
pyramidal
NH3
4
sp3
tetrahedral
2
2
bent
H2O
1
3
HCl
5
0
linear
trigonal
bipyramid
4
1
distorted
tetrahedron
SF4
5
3
2
trigonal
bypramid
PF5
dsp3
ClF
3
T-shaped
2
3
XeF2
linear
octahedral
6
6
0
5
1
4
2
SF6
octahedral
square
pyramid
square
linear
ClF5
XeF4
d2sp3
Structure of Atoms
Atoms are the smallest unit of an element. Any smaller, and you’d see that elements are all the same
inside - they’re made of neutrons, protons and electrons. At this point, you wouldn’t be able to
differentiate one atom from another, any more than you could differentiate your lungs from
someone else’s lungs.
So what does a generic atom look like? It has a nucleus consisting of protons and neutrons. These 2
sub-atomic particles like to stay at ‘home’, which is the nucleus. Electrons, on the other hand, are
wild creatures who roam around in orbitals/shells. They have a tendency to escape, or join other
atoms to form ions.
The Inside of an Atom
Electrons have a negative charge of -1, while protons have a positive charge of +1. Neutrons are
neutral, as their name suggests. An atom always has equal numbers of protons and electrons so that
the charges are balanced.
Protons and neutrons each have a relative mass of 1, while electrons weigh practically nothing
when compared to these 2. What this means is that the mass of an atom is decided by its number of
protons and neutrons, while electrons are negligible.
But how are you supposed to know how many protons, neutrons and electrons an atom has? This is
where the symbol comes in:
Carbon-12
The ‘12′ on the top is known as the “mass number” or “nucleon number”. It’s basically how
much the atom weighs. Sometimes it’s represented by an “A”.
The ‘6′ at the bottom tells you how many protons the element has. It’s the “atomic number” or
“proton number”. This is the magic number that identifies each element, like an IC number. It can
be represented by “Z”.
“Z” will tell you the number of protons in the atom. Since the number of protons = number of
electrons, you can find out the electron number from Z.
A = p+n, so A-Z = n. So in order to find out how many neutrons an atom has, it’s just top number bottom number. From one little symbol, you can already find out so many things! This knowledge
is necessary, as a lot of questions will be set to test your understanding of such diagrams and
symbols. Among other things, you will be required to deduce the numbers of protons, neutrons and
electrons in both atoms and ions, and you can’t do this without a solid grounding in atomic structure!
Isotopes
Atoms are very special in that most, if not all of them, have “twins”. Human identical twins have
the same genes, and the same is true for atoms. Except that they don’t have genes, they have proton
numbers! Isotopes are atoms with the same proton number but different neutron numbers. It’s
a good idea to memorise that sentence, as it is tested pretty frequently. You will also be required to
“spot the twins”, but it’s easy once you know what to look at - their proton numbers!
Isotopes have similar chemical properties because they have the same number of protons and
electrons. However, they will have slightly different physical properties (e.g. density) because of
the difference in neutron number. It’s like having one fat twin and one skinny twin. They have the
same genes, but one is heavier.
Electronic configuration
The last thing you need to know about atoms is their electronic structure. Electrons are arranged in a
2.8.8.x system. For example, our good friend Carbon up there has 6 electrons, so its configuration is
2.4, because 2+4 = 6. Calcium has 20 electrons, so it has a 2.8.8.2 structure. The last number in the
configuration tells you the number of valence electrons. These are the electrons on the outermost
shell. They are of great importance, because they’re the ones involved in forming compounds.
The elements of the periodic table are all arranged according to a system based on their proton
numbers, so once you learn this system, you won’t have to worry about needing to memorise
electronic configurations. You can just figure it out!
Ions
Ions are charged particles. They are formed from atoms by the loss of electrons to form positive
ions, or by the gain of electrons to form negative ions. Why would atoms want to lose or gain
electrons? By doing so, they can achieve the maximum number of valence electrons (i.e. 8), and
have a stable noble gas configuration. Having a full outermost shell of electrons is the key to
stability!
Metals have a tendency to lose electrons and form cations (positive ions) because they have
fewer than 4 valence electrons. Non-metals prefer to gain electrons to form anions (negative ions)
because they have 4 or more valence electrons.
Structure and Properties
What does the structure of something have to do with its properties? Answer: everything.
Before we go on to structure, let’s learn to tell the difference between elements, mixtures, and
compounds.
Element
An element is made up of only one type of atoms. It doesn’t matter what the atoms are, or how they
like to group themselves, as long as they’re all the same type. All those things in your Periodic
Table? Elements.
Mixture
“Mixture” is just a catch-all category. The particles can be atoms or molecules, as long as they’re
not all chemically combined together. So it doesn’t always have to be single atoms as in the picture
above. An example of a mixture would be crude oil, or petroleum.
Compound
A compound has 2 or more types of atoms chemically combined together. You’ll know it’s a
compound if it’s stuck together, like the above. If there are different kinds of atoms but they’re not
stuck together, it’s a mixture. Example of a compound: water. The 2 hydrogen atoms are “stuck” to
the oxygen atom.
How else can you differentiate between a mixture and a compound?
1. Compounds have a fixed ratio, while mixtures do not. For example, water always has 2
hydrogen atoms to one oxygen.
2. Compounds cannot be separated by physical methods such as distillation and filtration,
while mixtures can.
3. Pure compounds have a fixed boiling point. Mixtures boil over a range of temperatures.
Same goes for melting points.
Ionic Bonding
A lot of the structural properties of matter is due to the kind of bonding in the substance. As we
cover the 3 different kinds of bonding, pay attention to the formation of bonds, the properties of
such compounds, and how to deduce the physical and chemical properties of a substance given the
kind of bonding present.
First up is the ionic bond, which is formed between a metal and a non-metal. It involves the
transfer of electrons from a metal to a non-metal so that both can have noble gas configuration.
Ionic Bonding
In the above example, Na is the metal and Cl is the non-metal. You can find this out just by looking
at your periodic table. Look at the last ring of electrons around Cl-. Did you notice that one of the
electrons is represented by a cross instead of a dot? This is because it originally belonged to Na. Na
has 11 electrons to have a configuration of 2.8.1. In order for it to achieve noble gas configuration,
it loses that 1 extra electron to get a configuration of 2.8 and form an ion, Na+.
This 1 extra electron wanders over to the chlorine (Cl) atom, which gladly accepts it so that its 2.8.7
configuration becomes 2.8.8, which is a noble gas configuration. Thus Cl- is formed.
So now both Na and Cl are happy. Since they are now ions with opposite charges, they naturally
come together to form the ionic compound, NaCl.
Ionic compounds have the following properties:
1. High melting and boiling points due to the strong electrostatic forces between oppositely
charged ions, which require a lot of energy to overcome. This is the model answer you need
to memorise!
2. Conducts electricity in molten (l) and aqueous (aq) states, but not in solid state. This is
because in solid state, the ions are held in fixed positions in a lattice and cannot move
around to conduct electricity. In molten and aqueous states, however, the ions are free to
move around and carry charges.
3. Soluble in water but not in organic solvents such as chloromethane, chloroform, ethanol
and methylbenzene.
So the summary for ionic bonding is: Ionic bonding is between a metal and a non-metal, it involves
transfer of electrons, and ionic compounds have high melting points and boiling points and can
conduct electricity in certain states. You can definitely count on ionic bonding to come out in the
exams, so you need to know it and know it WELL.
Covalent Bonding
Covalent bonding is about sharing of valence electrons. It’s in the name: co-valent. In covalent
compounds, non-metals share their electrons so that each atom can have a noble gas configuration.
This means they either have 2 or 8 electrons in their outermost shell.
Covalent Bonding
As in the case of the hydrogen H2 molecule above, each atom is sharing its one valence electron so
that both of them can claim the other’s electron to complete their own shell. The shared electrons
are doing double duty because they “belong” to both the atoms. The attraction of the nuclei of both
atoms for the electrons is holding the entire molecule together.
Covalent bonds themselves are very strong. However, covalent compounds have low melting and
boiling points. This is due to the weak intermolecular forces between the molecules, which
require little energy to overcome. Please differentiate between the covalent bond, which is INSIDE
the molecule, and intermolecular forces, which are OUTSIDE the molecule. When you heat a
covalent compound, you are separating one molecule from another, but you are NOT breaking up
the molecule.
An example to illustrate this is the boiling of water. If you had water in a pot, it would consist of
water molecules (H2O). Upon heating, the water molecules move away from each other to become
water vapour. If you had a pair of nanoscopic chopsticks and plucked a water vapour particle from
the air, you would see that it’s still H2O. You didn’t break up the molecule, you just separated it
from other molecules by breaking the weak intermolecular forces. The strong covalent forces are
still holding the water molecule together.
Covalent compounds cannot conduct electricity. They have no ions and no free electrons. Certain
covalent compounds, however, can conduct electricity when dissolved in water. An example of
such compounds would be acids. Hydrogen chloride (HCl) is a gas at room temperature, but
dissolve it in water and you get our good friend, hydrochloric acid. However, generally speaking,
covalent compounds cannot conduct electricity in whatever state.
Finally, covalent compounds don’t dissolve in water (with the exception of acids). They do
dissolve in organic solvents.
In summary, covalent bonds are formed by the sharing of electrons. Covalent compounds have low
melting and boiling points, and do not conduct electricity.