Chapter 2.3 - 2.4
Molecular Compounds
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recall, from Chapter 1, that atomic theory describes electrons moving about the
nucleus of the atom in energy levels, and that the electrons in the outermost energy
level are called the valence electrons.
It is the valence electrons of an atom that form chemical bonds. chemical bond: the
forces of attraction holding atoms or ions together
recall that all matter is either an element or a compound
element: a pure substance that cannot be broken down into simpler substances by
chemical means (empirical definition); a substance composed entirely of one kind
of atom (theoretical definition)
compound: a pure substance that can be broken down by chemical means to produce
two or more pure substances (empirical definition); a substance containing atoms of
more than one element combined in fixed proportions (theoretical definition)
ionic compound: a pure substance formed from a metal and a nonmetal
molecular compound: a pure substance formed from two or more different nonmetals
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Molecular Compounds
Molecular Compounds
Ionic Bonding forming Ionic Compounds
• According to atomic theory, ionic compounds are formed when one or more
valence electrons are transferred from a metal atom to a nonmetal atom.
• This leaves the metal atom as a positive ion, or cation, and the nonmetal atom
as a negative ion, or anion.
• The two oppositely charged ions are attracted to each other by a force called
an ionic bond.
Covalent Bonding forming Molecular Compounds
• According to atomic theory, molecular compounds are formed when one or
more valence electrons are shared between two or more non metals.
• The sharing of electrons produces a covalent bond where the two nuclei of
each atom are attracted the shared pair of electrons
Representing Covalent Bonds
• Recall - Lewis symbol, or electron dot diagram,
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Molecular Compounds
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Two types of “pairs” of electrons in a Lewis structure
Bonding Pair – the electrons that are shared (part of the bond)
Lone Pair – electrons that are not involved in the bonding.
Note that like the ionic molecule the individual atoms achieve a stable octect
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Lewis structures can be simplified
Omit the lones pairs and you get a structural formula
The structural formula for chlorine would be Cl-Cl, and for bromine Br-Br.
Reduced further to produce chemical formulas, Cl2 and Br2
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Lewis Structure <-> Structural Formula <-> Chemical Formula
When an electron dot diagram is used to represent covalent bonding, we
adapt it slightly and call it a Lewis structure (because it illustrates the structure
of the molecule).
A Lewis structure shows the valence electrons surrounding each of the
component atoms as dots, with the exception of the electrons that are shared:
These shared electrons are represented by a dash. In effect, this dash
represents a covalent bond.
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Molecular Compounds
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Numbers of possible covalent bonds
Each pair of shared electrons results in a single bond.
2 pairs of shared electrons results in a double bond
Eg O2 CO2
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? How can we determine the number of bonds that will form?
ANS – Bonding capacity - the number of covalent bonds (shared electron
pairs) that an atom can form.
Note - the bonding capacity is 8-valance
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Molecular Compounds
Molecular Compounds
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Drawing Lewis Structures for molecules with different atoms. (i.e. compounds)
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The central position in the arrangement is often occupied by the element with
the highest bonding capacity.
Carbon and nitrogen, for instance, are commonly at the centre of a structural
formula.
Electronegativity is another means of deciding upon the central atom.
When there is a choice of atoms for the central position in the molecule,
choose the least electronegative element.
Hydrogen is never the central atom since it can only form a single covalent
bond.
Halides and oxygen are also usually not the central atom.
There are exceptions to these generalizations, but they meet our needs in
most cases.
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Coordinate Covalent Bonds
• Many substances contain a combination of covalent and ionic bonding.
ammonium chloride, NH4Cl. white, crystalline solid dissolves rapidly in water
and is an electrolyte—it dissociates to form a cation and an anion.
• It has many of the properties of an ionic compound, but it is composed only of
nonmetals.
• it as an ionic compound composed of a chloride ion, Cl–, and a polyatomic ion,
ammonium, NH4+.
• The bond holding the chloride and ammonium ions together is ionic, but the
bonds within the polyatomic ammonium ion are covalent.
• all of which are covalently bonded groups of atoms carrying an overall charge.
• Consider the formation of the ammonium ion from the regular covalent
molecule ammonia, NH3, and a hydrogen ion, H+. The hydrogen ion does not
bring any electrons with it. To achieve a complete outer shell it can borrow two
electrons from the atom with which it bonds.
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Molecular Compounds
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Molecular Compounds
coordinate covalent bond: a covalent bond in which both of the shared
electrons come from the same atom
H+ has no electrons and borrows 2 from the N
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coordinate covalent bond: a covalent bond in which both of the shared
electrons come from the same atom
H+ has no electrons and borrows 2 from the N
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Molecular Compounds
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Molecular Compounds
The Strength of Covalent Bonds
• Covalent bonds are strong. A large amount of energy is needed to separate
the atoms that make up molecules.
• Molecules tend to be stable at relatively high temperatures: They do not easily
decompose upon heating.
• The stronger the bonds within the molecule, the greater the energy required to
separate them.
• The strength of a bond between two atoms increases as the number of
electron pairs in the bond increases.
• Therefore, triple bonds are stronger than double bonds, which are stronger
than single bonds between the same two atoms.
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Molecular Compounds
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Molecular Compounds
Look at the four pictures below what is similar about a and b, c and d and
different about top 2 and bottom 2.
Use the pictures on the side to help you.
Polar Covalent Bonds
• When electrons are shared between two atoms, a covalent bond results.
• When the atoms are identical, such as in a chlorine molecule, the electrons
are shared equally
• However, this is not the case for a compound like hydrogen chloride, where
electrons are shared between two different elements.
• In this situation, the sharing is unequal, as the bonding electrons spend more
time near one atom than near the other..
• The electrons in the H–Cl bond in a hydrogen chloride molecule spend more
time near the chlorine atom than near the hydrogen atom.
• This is because of chlorine’s greater attraction for electrons.
• Due to this unequal sharing of electrons, the hydrogen atom is, on average,
slightly positively charged while the chlorine atom is slightly negatively
charged
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Molecular Compounds
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Molecular Compounds
We can predict which parts of the molecule will have d+ and d- charges by
comparing the electronegativities of the atoms
The most electronegative atoms will attract the electron pair strongly, and so
will tend to have a d- charge in the covalent compounds they form.
Using Electronegativity to predict
if a bond is ionic, polar
covalent or covalent.
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electronegativities of the
elements involved.
• The absolute value of the
difference in electronegativities of
two bonded atoms provides a
measure of the polarity in the
bond: the greater the difference,
the more polar the bond until it
reaches an ionic bond
• 0 – 0.3 = covalent
• 0.4 – 1.7 = polar
• 1.8 – 4.0 = ionic
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Molecular Compounds
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Molecular Compounds
Polar Molecules
• Molecules of water, ammonia (a reactant in the production of nitrogen
fertilizers), and sulfur dioxide (an industrial pollutant contributing to acid rain
formation) all have polar covalent bonds holding their atoms together.
• If a molecule contains polar covalent bonds, the entire molecule may have a
positive end and a negative end, in which case it would be classified as a polar
molecule.
• Not all molecules containing polar covalent bonds are polar molecules (eg,
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Molecular Compounds
Molecular Compounds
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Intermolecular Forces
Unlike ionic compounds the properties of molecular compounds cannot be
explained simply by covalent bonds.
if covalent bonds were the only forces at work, most molecular compounds would
be gases at STP, as there would be no attraction between the molecules strong
enough to order the molecules into solids or liquids.
intramolecular forces - (“intra” means within) - forces that bond atoms and ions
together within a compound (ionic or molecular)
these are sufficient to explain the existence of ionic and molecular compounds,
and to explain many of the properties of ionic compounds, but they aren’t sufficient
to explain the physical state of molecular compounds.
? Why is water a liquid at SATP and a solid at STP? Why isn’t it a gas at all
temperatures?
ANS - intermolecular forces - (“inter” means between) - forces between
molecules.
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Molecular Compounds
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Molecular Compounds
For molecular solids - these forces are strong enough to cause molecules to
arrange themselves in an orderly fashion to form a lattice structure (similar to that
of ionic solids) (eg. frozen water, dry ice)
For molecular liquids –these forces are strong enough to allow the molecules to
stay close to one another (not be a gas) but weak enough to allow them to flow
around each other and display the properties of a liquid . (liquid water, gasoline)
? are intermolecular forces weaker or stronger when compared to covalent bonds.
ANS – through observation - it is much easier to melt a molecular solid than it is to
cause the same substance to decompose. (break up into it’s elements)
When water is heated from –4°C to 104°C it changes state from a solid, to a
liquid, and then to a gas, but it does not decompose to oxygen and hydrogen.
The energy added in the form of heat is sufficient to overcome the intermolecular
forces between the molecules, but not the covalent bonds between the atoms.
Adding a relatively small amount of heat will cause a solid molecular compound to
change state from a solid to a liquid, and then to a gas, but it takes much more
energy to break the covalent bonds between the atoms in the compound.
Types of Intermolecular Forces
• there are three kinds of intermolecular forces, each with different strengths.
• two of these are classified as van der Waals forces (in honour of Johannes
van der Waals): dipole–dipole forces and London dispersion forces.
• The third intermolecular force—hydrogen bonding—is generally not grouped
with the other two.
• Dipole–dipole forces are the forces of attraction between oppositely charged
ends of polar molecules (e.g., HCl).
• The positive end of each molecule attracts the negative ends of neighbouring
molecules—rather like a weak version of ionic bonds.
• London dispersion forces, exist between all molecules—both polar and
nonpolar
• the only intermolecular forces acting between nonpolar molecules.
• they are the result of temporary displacements of the electron “cloud” around
the atoms in a molecule, resulting in extremely short-lived dipoles.
• the dipoles last for only tiny fractions of a second, the attraction is continually
being lost, so the forces are very weak.
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Hydrogen Bonds
• 1 of 3 of the intermolecular forces
• occurs among highly polar molecules containing F–H, O–H, or N–H bonds.
• the O–H, F-H or N-H bonds are all highly polar covalent.
• The F, O or N carries a slight negative charge while the hydrogen atoms carry
a small positive charge
• As a result, the hydrogen atoms exert a strong force of attraction on the F or O
or N of neighbouring molecules.
• The actual H bond occur between the H’s of one molecule and the F, O or N
of a neighbouring molecule
Hydrogen Bonds
• Although a hydrogen bond is similar to a dipole–dipole force, it is stronger than
any of the van der Waals forces
• the H - bond is a result of the large difference in electronegativity between
hydrogen and F, O or N
• Hydrogen Bond in Water
• Properties of water
Molecular Compounds
Molecular Compounds
– higher than expected melting and boiling points
– high vapour pressure (is the pressure of a vapor in equilibrium with its non-vapor
phases)
– high surface tension (is an effect within the surface layer of a liquid that causes
that layer to behave as an elastic sheet. )
– and the ability to dissolve a large number of substances.
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Water is a polar molecule, consisting of one atom of oxygen bound by single
covalent bonds to two hydrogen atoms.
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Molecular Compounds
Molecular Compounds
Hydrogen Bonds
• high melting point and boiling point – Explanation - Large amounts of energy are
required to break the hydrogen bonds in the solid and liquid states.
• high vapour pressure – Explanation - H bonds “hold” water in the liquid state so
equilibrium is only reached when pressures are high
• High surface tension – H bonds creates strong elastic sheet at the surface.
• Ability to dissolve a large number of solids – polar bonds are very strong vs most ionic
solid attractive forces.
Why does Ice Float
• Most substances are more dense in the solid state than in the liquid state.
• However, solid ice cubes floating in a glass of water.
• If they float, they must be less dense than the liquid.
• ice is less dense than liquid water because it forms an open lattice structure when it
freezes, with a great deal of empty space between the molecules
• There is the same amount of water molecules in a large volume, therefore less dense
• This is one of the many reasons why scientists consider water to be indispensable for
the existence of life.
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Molecular Compounds
Water molecules in liquid state H bonded
vs Water molecules in solid state H bonded
Note the spaces in the frozen water and how the volume is larger.
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