Download Chapter 6 Class Notes CHEM

CHEMICAL BONDING
Chapter 6 – Sections 2 - 5 – Pages 178 - 207
What is a molecule?
A neutral group of atoms that are held together by
covalent bonds. A single molecule of a chemical
compound is an individual unit capable of existing on its
own.
Bond length and bond energy between atoms in a
compound are related. The shorter the bond length
between two atoms in a compound, the higher the bond
energy will be. If two atoms have a small bond length,
then they will be expected to have a high bond energy.
This means that the smaller the bond length, then larger
energy is required to break this bond - more energy is
required when bond length small. When bond length is
large, less energy is required to break the bond between
the two atoms.
Digital Insert
“Bond Energy”
How to know whether a
compound is ionic or
covalent
Generally, ionic bonds occur between atoms of the
elements included in Groups 1 and 2, the alkali metals
and alkaline earth metals. As we’ve discussed, these
elements form cations meaning that they give up one
or two electrons giving them a +1 or a +2 charge.
And between atoms of the elements included from
groups 16 and 17, the chalcogens and the halogens.
These elements form anions meaning that they take
on one or two electrons giving them a -1 or -2 charge.
Ionic compounds
Examples include NaCl, CaO, MgO, LiCl, KCl, MgCl2,
CaF2
When ionic bonds are formed, there is a balance of
charge between the two atoms in the compound.
Take NaCl for example: Na forms a +1 ion while Cl
forms a -1 ion. Now look at CaF2: Calcium forms a 2+
ion while Fluorine only forms a -1 ion meaning that
there must be TWO atoms of Fluorine present in the
compound to balance the ionic charge between the
two atoms of the compound.
Covalent Compounds
Are usually compounds of two nonmetals – that is,
elements that are found on the right hand side of the
periodic table.
Digital Insert
“Comparing Ionic and Covalent Compounds”
Resonance
Is defined by bonding molecules or ions that cannot
be correctly represented by a single Lewis structure.
Resonance
Is indicated by a double-headed arrow between a
molecule’s structures.
This is a sulfite ion. From the depiction above, you can see that the structure
Exists in three different forms. This is due to resonance.
Ionic bonds are different from covalent bonds in that they have
transferred valance electrons to satisfy their octet. Their bonds are
stronger than covalent forces. Because of this, they exhibit higher
melting points and boiling points. They form what is called a crystal
lattice. From this crystal lattice, every ionic compound has a lattice
energy. When a chemical bond is formed, each individual atom
participating in bonding lowers its potential chemical energy by
bonding with another atom. All atoms “seek” a noble gas electron
configuration because this is most STABLE meaning this is the lowest
POTENTIAL ENERGY state.
An ionic compound
looks like…
An ionic crystal lattice
How an ionic compound
forms…
Ionic compounds vs
Covalent compounds
IONIC SUBSTANCES
COVALENT SUBSTANCES
MELTING POINT
high
low
BOILING POINT
high
low
SOLUBILITY IN WATER
high
low
ODOR
Non-detectable
detectable
ELECTRICAL CONDUCTIVITY
conductive
Non-conductive
But why do these
differences exist?
One word…INTERMOLECULAR FORCES.
These forces are defined as the forces of attraction ‘between’
molecules.
Ionic forces are stronger than covalent forces.
“Inter-” = between
“Intra-” = within
Intramolecular forces include electronegativity and polarity
Some INTERmolecular
forces
Dipole-Dipole forces
Hydrogen bonding forces
London dispersion forces
Refer to your textbook and be sure you are familiar
with the definitions of these three intermolecular
forces!!!
Digital Insert
“Dipole-Dipole Forces”
“Hydrogen Bonding”
“London Dispersion Force”
A polar-covalent
molecule - water
A non-polar covalent
bond vs a polar covalent
bond
MOLECULAR POLARITY
What determines whether a molecule exhibits
polarity? ________________
When a compound consists of atoms with greater
electronegativity than other atoms in the compound,
polarity exists.
Polyatomic ions are defined as a charged group of
covalently bonded atoms.
Metallic Bonding
The electron-sea model: overlapping of orbitals allows outer
electrons of atoms to roam freely throughout metals – which is
called electron delocalization – this gives the crystal lattice a very
stable structure. This is why metals have high melting points.
Good conductors – electron delocalization.
Malleable and ductile – all bonds are pulling in the same
direction.
Lustrous – absorb and reflect light at the same
wavelength/frequency due to electron configuration.
Digital Insert
“Metallic Bonding”
MOLECULAR
GEOMETRY
VSEPR (valence-shell, electron-pair repulsion) theory
– repulsion between the sets of valence-level
electrons causes the sets to be as far apart as
possible.
This is because negative charges repel neighboring
negative charges.
Lone pair electrons have a greater force of repulsion
to other electrons than do bonding electrons.
Some Common
Molecular Geometries
Hybridization
The second theory explaining molecular geometry. It is the
mixing of two or more atomic orbitals of similar energies on
the same atom to produce a new hybrid atomic orbital of equal
energy.
Linear hybridizes sp
Trigonal planar hybridizes sp2
Tetrahedral hybridizes sp3
Digital Insert
“Hybrid Orbitals”
Chapter 6 Quiz
1. Why do atoms form bonds with other atoms?
2. What is the relationship between bond energy and bond length?
3. What type of bond would you expect to exist between an atom of Calcium, Ca and an
atom of Oxygen, O?
4. What type of bond would you expect to exist between two atoms of Hydrogen and an
atom of Oxygen?
5. Which type of substance, ionic or covalent, is expected to have a higher melting point?
6. Briefly explain the VSEPR theory.
7. What is the hybridization for a molecule of methane, CH4?
8. Draw the Lewis structure for a molecule of Nitrogen gas, N2.
9. What are the two theories that help to explain the geometry of molecules?
10. What effect do intermolecular forces have on physical properties such as melting
point, boiling point, solubility and conductivity?
ANSWERS
1. To lower their individual atomic potential energies.
2. Inverse proportionality – when one goes up, the other goes down.
3. Ionic – Ca is from group 2; O is group 16.
4. Covalent – valence electrons are shared between atoms.
5. Ionic – the intermolecular forces are stronger.
6. Valence electrons will be as far apart as possible about an atom.
7. sp3
8. see next slide
9. VSEPR and hybridization
10. All of these properties tend to be high in an ionic compound due to
strong intermolecular forces.