IB Chemistry
Topic 4: Bonding
4.1 Ionic Bonding
Chemical bond – chemical bonds are made by the interaction of valence electrons. A chemical
bond is like a connection between two atoms.
- Why do bonds form? Because the bond is lower energy, or more stable, than the two atoms
alone.
Electronegativity – the ability to attract bonding electrons.
 Metals - low electronegativity
 Nonmetals - high electronegativity
Three types of chemical bonds
 Ionic
o a large difference of electronegativity
o metal cation with a positive charge and a nonmetal anion
 Covalent
o Both atoms have a high electronegativity
o Bond is between two nonmetals
 Metallic
o Both atoms have low electronegativity
o Made between metals
Define an ionic bond
o An ionic bond is the electrostatic attraction between the positive and negative charge
of the ions.
o Example: Na+ + Cl-  NaCl
o Note: when you write the formula, there should be no charges! Ionic compounds are
neutral.
Electron transfer
 A metal loses electrons and makes a cation (+ charge), and a nonmetal gains electrons to
make an anion (- charge), and an ionic bond forms.
 Example: Sodium loses an electron, chlorine gains an electron, and an ionic bond forms.
Strength of Ionic bonding
 Ionic bonds get stronger with increased charge of the cation or anion and with smaller
ionic radii.
 Example: MgS has a stronger ionic bond than NaCl.
4.2 Covalent bonding
Vocabulary: Covalent bonding, octet rule, Lewis structure, dative bond (coordinate bond)
A covalent bond is the electrostatic attraction of the positively charged nuclei for the shared pairs of
electrons. Example: in H2, each hydrogen atom gives one electron and share the two between them
to make a covalent bond.
Octet rule: Nonmetals form covalent bonds to fill their valence level. Noble gases already have a
filled valence level, so they don’t usually form compounds.
Example: Carbon has a valence level of 2s2, 2p4. To get the valence configuration of neon, it needs
four more electrons. Therefore, carbon usually makes four bonds. Example: CH4
See Lewis structure worksheet for how to draw Lewis structures.
Dative bonds form when one atom donates both electron paris.
Example NH3  NH4+
Bond length and bond strength
Double bonds are shorter and stronger than single bonds. Triple bonds are shorter and stronger than
double bonds.
Single bonds…double bonds…triple bonds
Increasing bond strength 
Decreasing bond length 
Increasing bond energy 
*Bond energy is a measure of bond strength
VSEPR Shape
VSEPR & Valence Bond Theory Examples
E = nonbonding electron pair, A = central atom, X = bonded atoms
TYPE
Arrangement of
Electron Pairs
Angle(s)
Molecular
Structure
Hybrid
Orbital
Example
AX2
linear
180
linear
sp
BeF2
AX3
2
trigonal planar
120
trigonal planar
sp
BF3
AX2E trigonal planar
117
bent
sp2
SO2
3
AX4
tetrahedral
109.5
tetrahedral
sp
CH4
AX3E tetrahedral
107
tripodal
sp3
NH3
AX2E2 tetrahedral
104
bent
sp3
H2O
Polarity
Vocabulary:
Electronegativity values
Electronegativity
Polar bond
Polar molecule
Dipole moment
Nonpolar bonds
Atoms with no difference in electronegativity difference have no overall dipole. Example: hydrogen
molecule H-H
Polar bonds
Electronegativity increases up and to the right of the periodic table. Atoms with a larger difference
in electronegativity have a more polar bond.
Example: H-Cl
Cl is much more electronegative than H
Nonpolar Molecules
Consider carbon tetrachloride CCl4
The bond C-Cl is polar because chlorine is more electronegative than carbon.
The molecule, however, is non-polar because there is no overall dipole moment. CCl4 is
symmetrical and the polar bonds pull equally in each direction. “The dipole cancels”
Polar Molecules
To be a polar molecule, a molecule must
1. Have one or more polar bonds and
2. Be asymmetrical (see below)
If a molecule is symmetrical, the bond polarities cancel and it is therefore nonpolar. Symetrical
shapes: tetrahedral, trigonal bipyrimidal, octahedral, square planar. All other shapes are
asymmetrical.
If a molecule is asymmetrical, the bond polarities pull to one side of a molecule, it has an overall
dipole, and it is therefore polar. When they pull in opposite directions, it is usually nonpolar.
Drawing a dipole: Draw arrows over the H-F bond to show the direction of the negative charge.
Which is more electronegative H or Cl? F is. So F gets the partial negative charge and hydrogen
gets the partial positive charge.
When drawing a molecule for the IB exam, include nonbonding pairs, partial charges with the lower
case delta (δ), and the net dipole (for the whole molecule, not the individual bond polarities.)
Which would you expect to be more polar?
The left molecule () has C-Cl polar bonds that are in opposite directions.
The right molecule has C-Cl bonds that both pull down, which gives the molecule a strong overall
dipole moment.
Network solids
Carbon has three allotropes. Different structural arrangements of the same element are called
allotropes.
 Diamond – Giant molecular structure. Tetrahedral arrangement of carbon atoms—
each carbon is bonded to four others. This very strong arrangement explains why
diamond has such a high melting and boiling point.


Graphite – giant covalent. Each carbon is bonded to three other carbons, making
layers of hexagon shaped rings. Delocalized pi bonds of the sp2 hybridized carbon
atoms give a bond order of 1 1/3, so bond lengths are shorter than that of diamond.
These delocalized electrons allow graphite to conduct electricity. Why is graphite
less dense and softer than diamond? The layers are far apart and can slide easily
because there are only weak Van Der Waals forces between layers.
Fullerene – a large sphere made of five and six-membered carbon rings like a soccer
ball. Sixty carbons (C60). Conducts better than diamond, not as well as graphite. It
can be reduced and form an anion. Addition reactions can happen. It is not a
network solid like graphite and diamond, it is a molecule, so it can dissolve in
nonpolar solvents.
Table1: Summary of the characteristics of the allotropes of carbon.
Allotrope
Structure
Bonding
Electrical
Conductivity
Diamond
Giant
molecular,
tetrahedral
arrangement
Only
covalent
bonds
Poor
conductor
Graphite
Covalent
network,
layers of
hexogonal
rings
Covlaent
bonds,
Van Der
Waals
Forces
between
layers
Fullerene
(C60)
Molecule,
hexogon
and
pentagon
rings like a
soccer ball
Covalent
bonds
Delocalized
electrons?
hardness
No
Hard
because
it is a rigid
structure
Yes
Soft
because
layers
can slide
over each
other
Yes, less
delocalized
electrons
than gaphite
Soft
because
C60
molecules
can slide
over each
other
Good
conductor
Does not
conduct as
well as
graphite
14.1 Shapes of Molecules and Ions *See VSEPR handout on website
VSEPR & Valence Bond Theory Examples
E = nonbonding electron pair, A = central atom, X = bonded atoms
Arrangement of
Molecular
Hybrid
TYPE
Angle(s)
Example
Electron Pairs
Structure
Orbital
AX2
AX3
linear
trigonal planar
180
120
linear
trigonal planar
sp
BeF2
2
BF3
2
sp
AX2E trigonal planar
117
bent
sp
SO2
AX4
109.5
tetrahedral
sp3
CH4
3
tetrahedral
AX3E tetrahedral
107
tripodal
sp
NH3
AX2E2 tetrahedral
104
bent
sp3
H2O
AX5
3
trigonal bipyramidal
90 , 120 trig. bipyramidal sp d
PF5
AX4E trigonal bipyramidal
90 , 117 sawhorse, seesaw sp3d
SF4
AX3E2 trigonal bipyramidal
90
T-shaped
3
ICl3
3
sp d
AX2E3 trigonal bipyramidal
180
linear
sp d
I3-
AX6
90
octahedral
sp3d2
CrCl6
octahedral
3 2
AX5E octahedral
88
sq. pyramidal
sp d
ICl5
AX4E2 octahedral
90
sq. planar
sp3d2
XeF4
Explaining shape –
SL: Give number of bonding pairs and number of nonbonding pairs.
HL: when explaining shape on the exam, give the number of “charge centers” (charge center =
bonded atom or nonbonding pair) and the number of nonbonding pairs.
Also, double and triple bonds push strongly, so angles with a double bond will be bigger than
expected (except for linear molecules, which will be 180).
Resonance structures: When you can draw more than one Lewis structure for a molecule by
moving a double bond, it is called a resonance structure. The real bonding in a molecule with
resonance structures is some combination of all three structures.
 Main point: This makes bond angles and lengths the same for bonds involved in resonance.
Example: In the nitrate ion (NO3-), there is not one double bond. Really, these electrons from the
double bond are ‘delocalized’ (spread out) over all three N-O bonds. This means…
 Bond angles are all 120 degrees
 All N-O bond lengths are equal
Image taken from
http://www.mikeblaber.org/oldwine/chm1045/notes/Bonding/Resonan/Bond07.htm
So the real nitrate molecule might be like this:
*Note: this is not the correct way to draw a Lewis structure for nitrate. Draw resonance structures as
shown above.
Bond order = number of bonding electrons /2.
Single bonds = 1
Double bond = 2
Triple bond = 3
For resonance structures, the extra pi bonding electrons are equally spread around the molecule.
Example: In the nitrate ion (NO3-), there is not one double bond. Really, these pi electrons are
delocalized over all three N-O bonds. Each bond is not single or double. They have a bond order of
1 1/3.
Partial charges: δ+ or δ- (“δ” is the lowercase Greek letter Delta, in case you wanted to know. It is
not a funny ‘s’ which is what my students think. Ask your history teacher who the Greeks were)
14.2 Hybridization
The valence level of carbon is 2s22p2. These 2s and 2p orbitals are atomic orbitals. When a carbon
atom makes a bond with another atom, the atomic orbitals “mix” and make new/different hybrid
orbitals. Only the valence orbitals hybridize.
Draw the electron diagram of the ground state valence electrons in carbon:
How does carbon make four bonds with only two unpaired electrons? One electron moves from the
2s to the empty 2p orbital so that there are four unpaired electrons. These orbitals now have the
same energy and are called sp3 hybrid orbitals. Aren’t they cute.
Hydrogen is the only atom that does not hybridize, since it only has one electron. When bonds form
in CH4, each hydrogen electron pairs with an electron from an sp3 orbital.
In CH4, we say the carbon atom is sp3 hybridized. There are four sp3 hybrid orbitals and they repel
each other to make 109.5° angles, as predicted by VSEPR.
Single bonds are made of sigma bonds (σ bond). Sigma bonds form along the inter-nuclear axis.
Double bonds are made of one sigma bond and one pi bond (π bond). Triple bonds are made from
one sigma bond and two pi bonds. Pi bonds form the sideways interaction of two p orbitals. The
electrons in pi bonds are above and below the inter-nuclear axis.
CH4 has only single bonds, so they are all sigma bonds.
To figure out the hybridization of an atom,
1. Draw the Lewis dot structure
2. Count the number of sigma bonds and nonbonding pairs on the atom. This is the number of
hybrid orbitals that must be made.
2 hybrid orbitals mean sp hybridization, 3 mean sp2 and 4 mean sp3.
Example: NH3
Draw the Lewis dot structure:
3 sigma bonds + 1 nonbonding pair = 4 orbitals = sp3 hybridized
Description: Three sp3 orbitals make sigma bonds with hydrogen, a pair of nonbonding
electrons is in the fourth sp3 orbital
Example: a carbon atom in ethene, C2H4
Draw the Lewis dot structure:
3 sigma bonds + 0 nonbonding pairs = 3 hybrid orbitals = sp2 hybridized
Your turn: Answer the following for the molecule sulfur dihydride: SH2
a. Draw the Lewis structure:
b. What is the hybridization on the sulfur atom? _____
c. Explain how the orbitals “hybridize.”
Sulfur dihydride has _____ sigma bonds and ____ nonbonding pairs, so it needs four hybrid
orbitals.
___ s orbital and ____ p orbitals combine to make ___ hybrid orbitals. Two of
the orbitals are bonding orbitals. Two of the orbitals are _____________ orbitals.
4.3-4.5 Intermolecular Forces and how they influence properties
Memorize Table p133-134
Electrical conductivity - To conduct electricity, a substance must have delocalized electrons that
can move freely.
Volitility – how easily a substance evaporates.
Stronger intermolecular forces = less volatile.
Types of forces, from weakest to strongest:
Strength of
attractive
forces
Strongest
type of
attractive
force
Type of
crystal
Boiling
and
melting
point
Hardness
Example
Weakest
Van Der
Waals
Nonpolar
molecules,
group 8 atomic
gases
Low
Soft
Ar, H2
Dipole-dipole
Polar
Hydrogen
bonds
Polar with H-N,
H-F or H-O
bonds
Covalent
bonds
Network solid
Metallic Bond
Metals
Ionic Bond
Ionic
compounds
↓
Strongest
↓↓
High
Hard
H-Cl
H2O
C(graphite)
Cu
NaCl
Van Der Waals Forces (London, dispersion forces)
London dispersion forces hold group 8 atoms together in a solid. They are made by one atom
inducing a charge on another, so one side of the atom or molecule becomes partially positive and one
side becomes partially negative.
The strength of London dispersion forces increases with
1. the number of electrons (or molar mass) in the atom or molecule.
2. if molecules can be packed more closely – Unbranched hydrocarbons usually have
stronger Van Der Waals forces.
 Example: Branched vs. unbranched
These molecules have the same number of carbon and hydrogen atoms,
but the branched molecule (2-methyl propane) has a lower boiling point.
Image: http://commons.wikimedia.org/wiki/File:Isobutane-n-butane.png

Example: CH3CH3 has a higher melting point than CH4 because there are more electrons in
CH3CH3.
Image from: http://bdml.stanford.edu/twiki/bin/view/Rise/InvestigationOfVanDerWaalsForce
Dipole-dipole forces
 Polar molecules
Dipole-dipole interactions become stronger with a larger difference in electronegativity. Dipoledipole interactions are stronger than London forces, but weaker than chemical bonds (covalent, ionic
or metallic).
Hydrogen Bonds: The strongest dipole dipole interactions are hydrogen bonds, which can be made
between a nonbonding electron pair and a hydrogen that is bonded to N, O or F. Note, this diagram
does not show the nonbonding pairs on oxygen, and that is WRONG. You should write nonbonding
pairs on all Lewis structures for the exam!
Image from: http://chemtools.chem.soton.ac.uk/projects/emalaria/index.php?page=13
Example: NF3 has a higher melting point than NH3.
 Nonpolar molecules
4.4 Metalic bonding Metallic bonding is thought of as cations in a sea of delocalized electrons.
Metallic bonding gets stronger with


Smaller radius of the ion
number of delocalized valence electrons
Example: Sodium, magnesium and aluminum are all metals. They have metallic bonding, in
which positive metal ions are attracted to delocalized electrons. Going from sodium to aluminum:




the ionic radius increases
the number of delocalized electrons increases ...
so the strength of the metallic bonding increases and ...
the melting points and boiling points increase.
Electrical conductivity: metals can conduct electricity because they have many delocalized
electrons.
Metals are malleable and ductile. Malleable means metals can bend without breaking. Ductile
means they can be pulled into wires. Metals are malleable and ductile because there is attraction
between the cations and delocalized electrons, so layers of ions can slide past each other. The
cations are not attracted to each other.
4.2 Network solids
Silica is made of SiO2, where each Si atom is bonded to four O
atoms.
Network solids have higher melting points than other covalent
molecules because to melt a network solid, many covalent
bonds must be broken. The melting point is over a temperature
range Example, SiO2 melting point: 1650(±75) °C.
Example SiO2 has a higher melting point than CO2 because SiO2 is a network solid and many
covalent bonds must be broken when it melts. CO2 is just a nonpolar molecule with weak Van Der
Waals forces. (Note—each Si is bonded to a fourth O, which is either behind or infront of the Si.)
Black = Si, white = O
Image from: http://commons.wikimedia.org/wiki/File:SiO2_-_Glas_-_2D.png
You should know these network solids: SiC, C, Si, SiO2
*See 4.1 above for ionic compound bonding strength.
4.5 Physical Properties
Properties: the main point
Know the following:
1. Know how the strength of each type of force changes in different molecules. Example: are
Van Der Walls forces stronger in He or Ne? Answer: Ne because Van Der Waals forces get
stronger with increasing molar mass. Molar mass of He = 4.00, molar mass of Ne= 20.18
2. Compare molecules of different types with different types of forces. Example: Which has
stronger intermolecular forces, He or H2O? H2O because hydrogen bonding is a stronger
intermolecular force than Van Der Waals forces.
3. Based on the strength of intermolecular forces, predict properties. Example: Which would
have a higher boiling point, He or Ne. Answer: Ne because it has stronger Van Der Waals
forces than He.
Melting and boiling points Stronger IMFs = higher melting and boiling points. Why?
When a substance melts, some of the attractive forces holding the particles together are broken or
loosened so that the particles can move freely around each other but are still close together. The
stronger these forces are, the more energy is needed to overcome them and the higher the melting
temperature.
Lowest boiling points:
The same is true for boiling points. A substance with strong intermolecular forces will have a high
boiling point.
Phases at room temperature (23 ºC)
Melting and boiling points explain phases of each type of compound.
Ionic compounds – solid
Network solids: solid
Metals: Solid except mercury = liquid
Polar molecules: solid, liquid or gas depending on size.
Nonpolar molecules: solid, liquid or gas, depending on size. Know the phase of all elements.
Gases: H2, O2, F2, Cl2, group 8 (as gaseous atoms, not molecules), liquids: Br2, Hg. All others are
solid.
Conductivity –
Ionic compounds can conduct when dissolved in a polar solvent or in the molten phase.
Polar liquids can conduct when strong electrolytes are dissolved.
Network solids: Graphite conducts. All others do not.
Metals: conduct
Nonpolar compounds cannot conduct because their electrons are kept in covalent bonds.
Solubility
Metals
Metals are soluble in other metals. A mixture of metals is called an alloy, which is made my mixing
the metals in the liquid phase and letting them freeze. Alloys are less malleable and ductile than
pure metals because the difference in size of the atoms prevents layers from sliding easily.
If metals are converted to ions, they are soluble in polar solvents.
Ionic compounds
Ionic compounds are insoluble in most polar solvents because the ionic bonding is very strong in
ionic compounds. However, they are usually soluble in water because water is very polar (and
therefore the partial charges are very high in water.) The charges of ions are attracted to the partial
charges of atoms in a polar molecule. Refer to the solubility rules (that you memorized!) for which
compounds are soluble.
Generally, solubility of ionic compounds is greater for:
 Smaller size of ions
 Smaller charge
Image from: grandinetti.org
“Like dissolves like”
Polar substances are soluble in polar substances. Common polar solvent: water.
Nonpolar substances are soluble in nonpolar substances. Common nonpolar solvent: hexane.
Why can’t nonpolar substances dissolve in polar substances? Polar substances form strong dipoledipole interactions that exclude nonpolar substances. Imagine oil and water mixing. The oil gets
pushed out of the web of hydrogen bonding in water.