Chapter 4

Question 1

List out the name
and bond angles of
the molecular
geometry of
molecules with 6
electron domains.

Answer 1

1. Octahedral, no lone pairs, 90 degrees
2. Square pyrimidal, one lone pair, less than 90
   degrees.
3. Square planar, two lone pairs, 90 degrees

Question 2

List out the name
and bond angles of
the molecular
geometry of
molecules with 5
electron domains.

Answer 2

1. trigonal bipyrimidal, no lone pairs, 90 degrees,
   120 degrees.
2. see saw, one lone pair, less than 90 and 120
   degrees.
3. T shaped, two lone pairs, 90 degrees
4. linear, 3 lone pairs, 180 degrees

Question 3

Explain in detail
how covalent bond
is formed

Answer 3

The electrostatic attraction between a shared pair
of electrons and a positive nuclei.
Covalent bonds result from the overlap of atomic
orbitals. A sigma bond (o) is formed by the direct
head-on/end-to-end overlap of atomic orbitals,
resulting in electron density concentrated
between the nuclei of the bonding atoms.
A pi bond (1) is formed by the sideways overlap
of atomic orbitals, resulting in electron density
above and below the plane of the nuclei of the
bonding atoms.

Question 4

State the formula for formal
charge, and how it is used

Answer 4

Formal charge (FC) can be used to decide which Lewis (electron dot) structure
is preferred from several. The FC is the charge an atom would have if all atoms
in the molecule had the same electronegativity. FC = (Number of valence
electrons)-1/2(Number of bonding electrons)-(Number of non-bonding
electrons). The Lewis (electron dot) structure with the atoms having FC values
closest to zero is preferred

Question 5

What are some
exceptions to the
octet rule?

Answer 5

Exceptions to the octet rule include some species
having incomplete octets and expanded octets
Hydrogen: 2 electrons
Boron: 6 electrons
Beryllium: 4 electrons
Sulphur: 6 e domains, 12 electrons
some atoms that can hold 5 electron domains

Question 6

Explain the
wavelength of light
required to
dissociate oxygen
and ozone.

Answer 6

wavelength required to dissociate oxygen: 24lnm
-because of double bond, stronger
wavelength required to dissociate ozone: 330nm
-15 bond order. weaker

Question 7

Describe the
mechanism of the
catalysis of ozone
depletion when
catalysed by CFCs
and NOx.

Answer 7

Catalysis of ozone depletion occurs when chemicals such as CFCs and NOx are released into the atmosphere. CFCs break down in the stratosphere due to UV radiation and release chlorine atoms that can react with ozone molecules, leading to the destruction of the ozone layer. NOx can also react with ozone molecules, leading to the formation of nitrogen oxide and oxygen, which can remove ozone from the atmosphere. Reduction of CFCs and NOx emissions is necessary to protect the ozone layer

CFCs:

CFCl3 + UV radiation → CFCl2 + Cl
Cl + O3 → ClO + O2
ClO + O → Cl + O2
ClO + O3 → Cl + 2O2
NOx:

NO + O3 → NO2 + O2
NO2 + O → NO + O2
NO + O3 → NO2 + O2
NO2 + O3 → NO3 + O2
NO3 + O → NO2 + O2
These reactions result in the depletion of ozone molecules in the stratosphere and are a major environmental concern.

Question 8

Explain how a
hybrid orbital is
formed.

Answer 8

A hybrid orbital results from the mixing of
different types of atomic orbitals on the same
atom.

Question 9

Identify and explain
the relationships
between Lewis
(electron dot)
structures, electron
domains, molecular
geometries and
types of
hybridization.

Answer 9

Sp = linear = 2 electron domains = 180 degrees
sp2 = trigonal planar = 3 electron domains = 120
degrees
sp3 = tetrahedral, trigonal pyrimidal, bent = 4
electron domains = 109.5 degrees, 107 degrees,
104.5 degrees

Question 10

List atoms with 6
and 5 bonding
domains

Answer 10

6 bonding domains:
-S
-Xe
-Br
5 bonding domains:
-S
-P
-Cl
-I

Question 11

Explain the meaning
of electron
delocalization.

Answer 11

pi/ pi-pi.-electrons shared by more than two
atoms/nuclei / a pi/pi–pi-bond/overlapping p-orbitals that extends over more than two atoms/nuclei;

Question 12

Describe the
structure and
bonding of an ionic
compound, and
how it is formed.

Answer 12

The ionic bond is due to electrostatic attraction
between oppositely charged ions. Under normal
conditions, ionic compounds are usually solids
with crystal lattice structure.
lonic compounds are ormed when electrons are
transerred rom one atom to another to orm ions
with complete outer shells o electrons.
Positive ions (cations) form by metals losing
valence electrons.
Negative ions (anions) form by non-metals gaining
electrons.
The number of electrons lost or qained is
determined by the electron configuration of the
atom.

Question 13

List the properties of an ionic compound, and use
the bonding models to explain its properties.

Answer 13

1. Solid at room temperature
2. Have high melting and boiling point, low
   votality (do not form gas easily): to separate
   particles into liquid or gas would require high
   energy to break the strong ionic bonds.
3. Do not conduct electricity as a solid, does
   conduct electricity when liquid or in an aqueous
   solution: no freely moving charged particles.
4. brittle, hard: strong ionic compound do not
   allow the structure to bend or deflect.
5. all dissolves somewhat in water bur solubility
   varies discriminately

Question 14

What is the octet
rule, and what are
the exceptions to
it?

Answer 14

The rule states that all atoms should have in total
of 8 electrons after covalently bonding with other
atoms.
The exceptions to the octet rule include:
1. Hydrogen (2 electrons)
2. Beryllium (4 electrons)
3. Boron (6 electrons)

Question 15

Describe how the
formulas of ionic
compounds are
written.

Answer 15

The ionic compounds are written based on the
charges (2+ and 2-, 3+ and 3-).
Usually the cations come first and then the anion
is written.
There are some exceptions to this:
CH3COONa
CH3COOH
The Na and H are written at the end to indicate
that they are actually bonded to the O.

Question 16

Define covalent
bonding.

Answer 16

The electrostatic attraction between a shared pair
of electrons and positive nuclei.

Question 17

Define bond energy
(bond enthalpy),
and provide some
examples (using
equations).

Answer 17

The energy required to break one mol of a
covalent bond in gaseous molecules.
C12 (g) -> 2Cl (g)
12 (g) -> 21 (g)

Question 18

Outline the energy
required in forming
and breaking
covalent bonds.

Answer 18

Breaking bonds requires energy, so the reaction is
endothermic.
Forming bonds always releases energy so the
reaction is exothermic.

Question 19

Describe the
relationship
between bond
length and bond
strength.

Answer 19

Average bond enthalpy decreases as bond length
increases.
Although there is considerable variation in the
bond lengths and strengths of single bonds in
different compounds, double bonds are generally
much stronger and shorter than single bonds,
because the bond length is shorter. The strongest
covalent bonds are shown by triple bonds, with
the shortest bond length.

Question 20

Define resonance
bonding.

Answer 20

A double bond could appear in 2 (or more)
locations.
When writing the Lewis structures for some
molecules it is possible to write more than one
correct structure.

Question 21

State the valence
shell electron pair
repulsion (VSEPR)
theory.

Answer 21

The theory states that pairs of electrons arrange
themselves around the central atom so that they
are as far apart from
each other as possible. There will be greater
repulsion between non-bonded pairs of electrons
than between bonded pairs.

Question 22

List our what counts
as an electron
domain.

Answer 22

-Lone pair
-Single bond
-Double bond
-Triple bond

Question 23

List out the different
shapes of
molecules and its
corresponding
bond angles.

Answer 23

1. linear, 180 degrees
2. trigonal planar, 120 degrees
3. tetrahedral, 109.5 degrees
4. trigonal pyramidal, 107 degrees
5. bent/ VEE shape, 104.5 degrees

Question 24

Explain how
polarity in covalent
bond

Answer 24

Bond polarity results from the difference in
electronegativities of the bonded atoms.
The difference in electronegativity results in the
formation of an ionic bond, while a SMALLER
electronegativity (0.5-1.7) results in the formation
of a polar covalent bond.
One end of the molecule will thus be more
electron rich than the other end, resulting in a
polar bond. This relatively small difference in
charge is represented by + and -. The bigger the
difference in electronegativities the more polar
the bond. This is called a bond dipole
Covalent bonds can also be non-polar, where the
bonds occur between atoms that have no
different in electronegativity (the “gens”), they are
also called pure covalent bonds. (0.1-0.4)

Question 25

List out giant
covalent/network
covalent structures
of carbon and
silicon and explain
their properties in
terms of their
structure.

Answer 25

Giant covalent/network covalent structures are
covalent substances that don’t exist as discrete
molecules. Some examples of it include: diamond,
silicon, and silicon dioxide.
Diamond:
-made up of purely carbon atoms bonded by
strong covalent bond; each carbon atom is
covalently bonded to four other carbon atoms to
form a giant covalent structure.
as a result, it has a high melting and boiling point.
-because all the electrons are localized it does
not conduct electricity.
Silicon:
-each silicon atom is bonded to 4 other silicon
atoms in a tetrahedral arrangement (like carbon)
-again, because all electrons are held in fixed
position, it is a poor conductor at low
temperature.
Silicon Dioxide (an empirical formula!):
-Silicon dioxide has a bent structure caused by
the lone pairs of electrons on the oxygen atom
-it is a strong covalent bond, a hard substance
with high melting and boiling point and is a poor
conductor of electricity.

Question 26

List out the
allotropes of
carbon and explain
their properties in
terms of their
structure.

Answer 26

Allotropes are different forms of the same
element in the same physical state.
Carbon exists as 3 allotropes: diamond, graphite,
graphene, and fullereness.
Graphite:
-each carbon atom has very strong bonds to
three other carbon atoms
-this gives layers of hexagonal rings consisting of
carbon rings in fused hexagonal rings.
-composed of planar sheets of hexagonally
arranged C atoms stacked on top of each other
-layers are held together by relatively low LDF
-because only bonded to 3 other C atoms, each C
atom has an electron which becomes delocalized
across the plane; so it conducts electricity
-trigonal planar structure; sp2
-soft and slippery because layers are able to slide
over one another due to weak IMF
Graphene:
-a single layer of hexagonally arranged carbon
atoms, i.e. it is essentially a form of graphite which
is just one atom thick; have high tensile strength
(thin)
-extremely light, functions as a semiconductor
and is 200 times stronger than steel
-trigonal planar
-it is the most chemically reactive carbon
allotrope
C60 fullerene:
-each C atom is covalently bonded to three other
C atoms
-carbons bonded together in 20 hexagons (6C
rings) and 12 pentagon (5C ring); gives a geodesic
spherical structure
-presence of delocalized electrons = conducts
electricity

Question 27

Explain coordinate
bonding.

Answer 27

A coordinate covalent bond is where the pair of
shared electrons only come from one atom.
The most important thing to note that is when
drawn into a Lewis dot diagram, coordinate bond
looks exactly like a regular bond

Question 28

List and explain the
3 types of
intermolecular
forces.

Answer 28

London Dispersion Force:
-refers to instantaneous induced dipole forces
that exist between any atoms or groups of atoms
and should be used for non-polar entities
-made up of both instantaneous dipoles and
induced dipoles (caused by changes in electron
density and movement of electrons within an
atom or molecule)
-instantaneous dipoles induce nearby atoms to
make it into an induced dipole
-ALL atoms and molecules have LDF (both polar
and non-polar)
-larger molar mass = larger surface area =
increase in LDF
Dipole dipole forces:
-occur only between polar molecules (molecules
that have a net dipole movement)
-dipole dipole force is the electrostatic attraction
between a partial positive charge on one
molecule, and a partial negative charge on the
other.
-they ALSO have LDF!
-The strength of the dipole-dipole forces
depends on the overall polarity of a molecule.
the more polar the molecule, the stronger the dipole-dipole force
-lower electronegativity = partial positive charge
-higher electronegativity = partial negative charge
Hydrogen Bonding:
-occurs between molecules that have an
electronegative nitrogen, oxygen, or fluorine
atom directly bonded to a hydrogen atom
-The strongly electronegative atom pulls the
(only) electron away from the hydrogen nucleus,
“exposing” the nucleus.
-this exposed hydrogen nucleus is in turn strongly
attracted to the lone pairs on nearby molecules.

Question 29

Define a metallic
bond and explain
their properties and
strength.

Answer 29

A metallic bond is the electrostatic attraction
between a lattice of positive ions and delocalized
electrons.
The positive ions are arranged into a lattice
structure, and there are valence electrons that
create a sea of delocalized electrons which are
not bound to one atom but can float around and
are attracted to any nearby nucleus.
This creates non-directional bonds between metal
atoms that explains why metals are malleable
(Since there are no strong attractions “planes” of
atoms can slide over each other easily.); good
conductors of electricity
Their properties include:
-good conductors
-ductile (can be made into wires)
-malleable (can be bent into shape)
-shiny when polished
The strength of metallic bond depends on the
charge of the ions and the radius of metal ion.
greater ionic charge (more delocalized electrons)
= smaller ionic radius = greater strength = hight
melting and boiling points
So towards the left of the periodic table, the melting and boiling point increases.
Moving down the periodic table, the melting and
boiling point decreases as it increases in atomic
size.

Question 30

Describe what
alloys are and
explain their uses.

Answer 30

Alloys usually contain more than one metal
(homogenous mixtures composed of 2 or more
metals or a metal and a non-metal) and have
enhanced properties.
Alloys can be created with the addition of carbon
or phosphorus (non-metals) in some cases.
**They are often stronger and have more
resistance to corrosion.
The addition of different sized atoms in the alloy
means that they layers cannot slide over each
other as easily (originally due to non-directional
bonding. This results in the alloy being harder
than its component metals.
Examples!:
Brass is an alloy of copper and zinc.
Steel is an alloy of iron with carbon and some
other elements such as nickel, tungsten, or
molybdenum.
Different grades and types of steel are created by
changing the amount and type of alloying agents.