Ionic Bonds,
Covalent Bonds &
Molecular Structure
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Chemical Bonding
•What is chemical bonding?
•It is a strong attraction or force which holds atoms
or ions together in a chemical compound.
•Why do atoms form bonds?
•Octet Rule says that atoms want a full valence shell of 8 e•It is the valence e- which are responsible for chemical bonds
•So by reacting, they may fulfill the octet rule
•But most importantly, they form stable compounds!
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Ionic Bonds
• You know that metals tend to have low IE values
and so lose electrons fairly readily to form cations
• You also know that nonmetals tend to have more
negative EA values and so attract electrons fairly
readily to form anions
• So what happens when a metal atom collides with a
nonmetal atom?
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Ionic Bonds
• The metal atom with its small IE gives an electron
(or more) to the nonmetal with its negative EA
• The cation and anion have achieved a Noble Gas
electron configuration
• And the cation and anion are held together by
electrostatic forces (opposite charges attract). This
is the ionic bond.
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Ionic Bonds
•Ionic Bonds: a chemical bond between ions of opposite
charge (classically, a metal cation bonded to a nonmetal
anion).
•Electrons are transferred from the metal to the
nonmetal.
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Ionic Bonds
• In an ionic solid like NaCl, you can’t separate out
individual Na-Cl ionic bonds, instead it is a 3-D
network of Na+ and Cl- ions which are
interconnected.
• This network is the crystal lattice.
• Let’s delve in deeper!
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Born-Haber Cycle
• How exactly does sodium metal combine with
chlorine gas to produce sodium chloride?
• F irst, write the overall equation.
• Although the rxn occurs simultaneously, we can
break the overall rxn into 5 distinct steps.
• These steps are called the Born-Haber Cycle
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Step 1: Convert Metal to Gas
• As sodium metal is a solid, we must first convert it
to the gaseous state:
Na(s) -> Na(g)
• This is the heat of sublimation for sodium, Hsub. It
is also called the heat of formation of Na(g), or Hf.
Energy is always required in this step as the gas
state is higher energy.
• We will call this H1
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Step 2: Convert Cl2 to Cl Atoms
• As chlorine gas is a diatomic element, we must
break the Cl-Cl bond to form Cl(g) atoms:
1/2 Cl2(g) → Cl(g)
• This is the heat of formation of Cl(g), or Hf, OR
we may use 1/2 the Cl-Cl bond energy (the energy
required to break a Cl-Cl bond), D(Cl-Cl). Breaking
bonds ALWAYS takes energy.
• We will call this H2
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Step 3: IE1 for Na(g)
• Now we ionize the sodium gaseous atom:
Na(g) → Na+(g) + e• This is simply the IE1 for sodium. This requires
energy.
• We will call this H3
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Step 4: EA1 for Cl(g)
• Now we ionize the chlorine gaseous atom:
Cl(g) + e- → Cl-(g)
• This is simply the EA1 for chlorine. This releases
energy. (the first step so far to release energy)
• We will call this DH4
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Step 5: -Lattice Energy
• Now we form the ionic solid sodium chloride from
the gaseous ions:
Cl-(g) + Na+(g) → NaCl(s)
• This step releases energy, as bonds are formed.
• Energy is always released when bonds are formed.
• We will call this DH5
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Step 5: -Lattice Energy
• The energy released in Step 5 is the negative or
reverse of what we call the Lattice Energy of an
ionic compound.
• Lattice Energy is abbreviated LE, U, or HLatt.
• Lattice Energy is DEFINED as the energy
REQUIRED to separate a mole of a solid ionic
compound into its gaseous ions.
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Ionic Compound Formation
• What do you notice about the 5 steps?
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Lattice Energy Factors
• What factors affect the Lattice Energy?
• This is Physics!
•Charge on ions
•Distance between ions (size of ions)
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Ionic Compound Formation
• How would you draw the Born-Haber Cycle for
MgCl2?
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Ionic Compounds
•The attractive force between full opposite charges is
very strong.
•So ionic bonds are very strong.
•Therefore, ionic compounds have very high melting
points, boiling points, and high lattice energies.
•NaCl melts at 804°C
•LE are in thousands of kJ/mol, so it takes a lot of
energy to break apart a solid ionic crystal. (But it does
happen, does salt dissolve in water?)
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Covalent Bonding
• Covalent Bonding and Covalent Compounds: A bond where
atoms share electrons.
• Remember that it is difficult for atoms to gain or lose 3 or more
electrons.
• So many atoms share electrons in order to have 8 valence
electrons.
• It’s like sharing a room with someone, it’s both your room, but
you’re sharing.
• When atoms share one or more electrons, a covalent bond is
formed because both nuclei are attracted to the shared
electrons.
• Compounds which contain covalent bonds are called covalent
compounds or molecular compounds or molecules.
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Covalent Bonding
• The two atoms share one or more electrons; with the shared
electrons having a high probability of being found between the
two nuclei.
• The above figure represents the hydrogen molecule, where 2
electrons are shared equally between the two atoms.
• The 7 diatomic elemental molecules share electrons equally just
as the above figure shows.
• These covalent bonds where the electrons are shared equally are
also called nonpolar bonds or nonpolar covalent bonds.
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Covalent Bonding
• But as many of us know from personal experience, not everyone
shares equally!
• Just as there are greedy people, some atoms are more electrongreedy than others and take more than their fair share of
electrons.
• When two atoms share electrons unequally, a polar covalent
bond results.
• The atom which has a stronger attraction for electrons will pull
the shared electrons towards its nuclei.
• Thus, the unequally shared electrons will tend to be closer to
the electron-greedy atom.
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Covalent Bonding
• In the figure below, the HF molecule would look like the left
picture if the two atoms shared two electrons equally.
• But F is extremely electron-greedy so it pulls the shared
electrons towards its nuclei as in the picture on the right.
• Thus, the HF molecule is polar covalent.
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Covalent Bonding
• As electrons have a negative charge, the electron-greedy atom
will have what we call a partial charge.
• It doesn’t have a full negative charge as it is still sharing
electrons, but as it has more than its fair share, it has a slight or
partial negative charge.
• If the one atom has a partial negative charge, what type of
charge do you think the other atom has?
• How do we show these partial charges?
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Covalent Bonding
• We show these partial charges like this: + for a partial positive
charge, and - for a partial negative charge.
• The HF molecule can be drawn to show these partial charges
(Note that we show the bond between the H and F atoms with a
horizontal line between them.):
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Water and Polar Covalent Bonding
• Water has two polar covalent bonds.
• Oxygen has a partial negative charge, while both hydrogens
have a partial positive charge.
• As there is a positive end and a negative end of the bond, there
is a charge separation or a dipole.
• So water has 2 polar covalent bonds and 2 bond dipoles.
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Water and Polar Covalent Bonding
• In water, these 2 bond dipoles make water a very polar
molecule.
• This is why water has some very special properties including its
high freezing and boiling points.
• It is also why ionic compounds tend to be water soluble.
• Life on our planet would be very different (and might be
nonexistent) if water were not a polar molecule.
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Water and Polar Covalent Bonding
• However, there are many nonpolar molecules with polar
covalent bonds.
• Sometimes, due to the shape or geometry of the molecule, bond
dipoles cancel out.
• Carbon dioxide is an example of a nonpolar molecule with polar
covalent bonds.
• If the bond dipoles cancel, the molecule is nonpolar with polar
covalent bonds.
• If the bond dipoles do not cancel, the molecule is polar with
polar covalent bonds.
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The Bonding Continuum
Here’s a picture of covalent bonding, polar covalent bonding, and
ionic bonding.
In (a), electrons are shared equally as in H2.
In (b), electrons are shared unequally as in HF.
In (c), electrons have been transferred from one atom to another,
as in NaCl.
Notice that a polar covalent bond is between a covalent bond and
an ionic bond, so sometimes we say that a polar covalent bond
has partial ionic character.
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Covalent Bond Lengths & Strengths
• Let’s go back to the hydrogen molecule!
• You can see why the H2 molecule forms as the attractive forces
between the nuclei and shared electrons overcome the repulsive
forces.
• But there is an optimum distance between the 2 nuclei where
the bond between the 2 nuclei and electrons is greatest.
• This optimum distance is called the bond length.
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Covalent Bond Lengths & Strengths
• In the following figure, you can see what happens if the H atoms
get too close or too far apart: the bond is unstable and the
molecule falls apart!
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Covalent Bond Lengths & Strengths
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Covalent Bond Lengths & Strengths
• So the bond length is the optimal distance between 2 atoms.
• As this is the distance where the bond strength is greatest, it is
also the distance at which the most energy is required to break
the bond.
• The energy required to break a covalent bond is called the bond
energy or the bond dissociation energy.
• Every type of bond has its own characteristic bond energy, D,
but it always takes energy to break a bond.
• But this means that energy is released when a bond forms!
• There are Tables of bond energies for many different types of
bonds and you’ll use them later.
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Covalent Bond Lengths & Strengths
• As a rough idea of how strong covalent bonds are, bond
energies range from about 100 kJ/mol to over 600 kJ/mol.
• By comparison, Lattice Energies for ionic compounds were
thousands of kJ/mol!
• Here are a few things to point out:
•Double and triple bonds are shorter and have higher bond
energies than single bonds.
•There may be many bonds in a molecule, so it may take a lot
of energy to break ALL of the bonds.
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Covalent Bond Lengths & Strengths
Bond
C–C
C=C
CºC
C–N
C=N
CºN
C–O
C=O
CºO
Bond L ength (10 -10 m)
1.54
1.34
1.20
1.43
1.38
1.16
1.43
1.23
1.13
Bond
N –N
Bond L ength (10 -10 m)
1.47
N =N
N ºN
N –O
N =O
1.24
1.10
1.36
1.22
O –O
O =O
1.48
1.21
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Covalent Bond Lengths & Strengths
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Properties of Ionic & Molecular Cmpds
•You’ve seen how ionic and molecular
compounds form.
•Do they have different properties?
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Properties of Compounds
Molecular Compounds
Ionic Compounds
Composed of molecules
Composed of ions
nonmetals bonded to nonmetals
metals bonded to nonmetals
gases, liquids, or solids
solids
nonconductors
conductors when melted or
dissolved
generally low melting points
high melting points
generally low boiling points
high boiling points
tend to be insoluble in water
tend to be water soluble
tend to be soluble in organic
solvents
tend to be insoluble in organic
compounds
Energy to break bonds: 100’s
kJ/mol
Energy to break Lattice: 1000’s
kJ/mol
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Predicting Bond Types: Electronegativity
• Ionic and molecular compounds have different properties.
• How can you predict whether a compound is ionic or covalent?
• Several ways:
•Nonmetal bonded to nonmetal equals covalent (But is it
polar or nonpolar?)
•Metal bonded to nonmetal equals ionic
•Use electronegativity values
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Predicting Bond Types: Electronegativity
• As was mentioned earlier, some atoms are electron-greedy or
are very strongly attracted to electrons in a bond.
• Electronegativity is a measure of the ability of an atom in a
molecule to attract electrons towards itself within a chemical
bond.
• Chemist Linus Pauling developed the electronegativity scale for
the elements.
• Fluorine is the most electronegative element, with its
electronegativity set at 4.0.
• As F is the most electronegative element, the electronegativity of
the elements increases going across from left to right across a
period, and it would decrease going down a group.
• If F is the most electronegative element, what is the LEAST
electronegative element (or most electropositive)?
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Predicting Bond Types: Electronegativity
• Why aren’t the Noble Gases on the Table?
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Predicting Bond Types: Electronegativity
• To determine the type of bond, the difference in
electronegativity values must be calculated.
• Ex: H-F bond:
• Your Turn: Determine the electronegativity difference in the
Na-Cl bond.
• Your Turn: Determine the electronegativity difference in the CH bond.
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Predicting Bond Types: Electronegativity
• After the electronegativity difference has been calculated, the
bond type can easily be determined.
• If the difference is ≥0.0 but ≤ 0.4, the bond is covalent.
• If the difference is > 0.4 but < 2.1, the bond is polar covalent.
• If the difference is ≥ 2.1, the bond is ionic.
Ex: Determine the bond type in H-Cl.
1) Cl = 3.0; H = 2.1, so the difference is 3.0 - 2.1 = 0.9
2) Thus, the H-Cl bond is polar covalent.
Your Turn: Determine the following bond types:
1) O-O
2) Na-F
3) Si-N
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Bond Energies, D, & Enthalpy Changes
• We can use Bond energies to approximate the
enthalpy change for a reaction.
• If we know how much energy it takes to break
or make a chemical bond, we can calculate the
energy change for a rxn.
• Now this is an energy, not an enthalpy, but the
difference may be less than 1%, so we usually
ignore this and just say that it is a bond
enthalpy.
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Bond Energies & Enthalpy Changes
• There are Tables and Tables of bond energies
for many different types of bonds.
• We can use these to find H°rxn if we don’t have
the right H°f values or we can’t use Hess’s
Law.
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Bond Energies, D, & Enthalpy Changes
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Bond Energies & Enthalpy Changes
• To use D values to find the enthalpy change for a
rxn, we use the following equation:
DH rxn = å D(bonds broken) - å D(bonds formed)
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Bond Energies & Enthalpy Changes
• Why does this seem backwards from before?
Because D is defined as the energy required to
break a bond, all bond energies are positive.
• By subtracting bonds formed from the bonds
broken, we are actually giving the bond
enthalpies the correct sign.
• Example: calculate the enthalpy change for the
following rxn:
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