Zusammenfassung der Ressource
2.2 Electrons, Bonding and Structure
- 2.2.1 Shells and Orbitals
- s-orbitals
- Spherical shape
- From n=1, each shell
contains 1 s orbital
- Giving 2s
electrons in
each shell
- p-orbitals
- has a 3D dumb-bell shape
- Present from n=2 upwards
- Each shell has 3 p-orbitals
at right angles to each
other
- Each p-orbital can
hold 2 electrons,
giving a possible
total of 6p
electrons
- d-orbitals and f-orbitals
- From n=3 upwards, each shell contains
5 d-orbitals, giving a possiblesum of 10d
electrons
- From n=4 upwards each shell contains 7
f-orbitals, giving a possible total of 14f
electrons
- 2.2.2 Sub-shells and energy levels
- Sub-shells
- Have the same
principle quantum
number
- There are s,p,d and f subshells as each is
only made of one type of atomic orbital
- Filling shells and sub-shells
- Lowest energy filled first
- An energy level must be full before
a higher starts to fill
- Sub-shells are made of several orbitals, each
with the same energy level
- Each orbital is filled singly before pairing starts
- The 4s is at a slightly lower energy level than 3d, so 4s fills first
- Electron configuration
- Written in the form nx^y
- n=shell number
- x=type of orbital
- y=the number of electrons in the orbitals
making up the sub-shell
- Electron configuration of ions
- Ions are formed by gaining or loosing electrons
- If they form a positive ion then the electrons found in the
highest energy levels are lost first
- 2.2.3 Chemical bonding
- Chemical reactions are accompanied by electron
transfers and often bem=come more stable by
transferring electrons or combining electrons with
other atoms
- The most stable and unreactive elements
are the noble gases as they have a full
outer shell, other elements react to try
and achieve the same configuration
- The 8 outer shell electrons are
made from 2 in the s and 6 in
the p
- 3 main types of chemical bonding
- Ionic bonding
- Metal and a non-metal
- Electrons are transferred from the metal to the non-metal, to
form oppositely charged ions that attract each other
- Metallic bonding
- Occurs in metals
- Electrons are shared between all the atoms
- Covalent bonding
- 2 non-metals
- Electrons shared between the atoms and are
attracted to the nuclei of both bonded atoms
- 2.2.4 Ionic bonding
- Ionic bonds
- Electrons are transferred
- Oppositely charged ions are formed
- Metal ion is positive
- Non-metal ion is positive
- Giant ionic lattices
- Each ion is surrounded
by oppositely charged
ions
- Ions attract from all directions,
forming a 3D lattice
- 2.2.5
Structures of
ionic
compounds
- Due to the huge amounts of electrostatic attraction, they have uniques properties
- High
melting
and boiling
points
- Solids at room temperature as a large amount
of energy is needed to break the electrostatic
bonds that hold the opposingly charged ions
together in the lattice
- The greater the charge, the stronger
the electrostatic forces between the
ions, meaning more energy is required
to break up the ionic lattice during
melting
- Electrical conductivity
- When solid they don't conduct electricity as the
ions are held in fixed positions and so cannot
move and therefore cannot conduct electricity
- When dissolved or melted the solid lattice breaks down and the
ions are free to move, so it can now conduct electricity
- Solubility
- An ionic lattice
dissolves in polar
solvents such as
water, polar
substances
contain
substances that
have polar bonds
- Polar water molecules break down an ionic
lattice by surrounding each ion to form a
solution
- The slight charges
of the polar
substances are able
to attract the
charged ions in the
giant ionic lattice,
meaning the lattice
is disrupted and
ions are pulled out
of it
- 2.2.6 Covalent bonding
- An electron occupies
the space between the
2 atoms' nuclei, this
shared pair of
electrons forms a
covalent bond
- The attraction overcomes the
repulsion between the two
positively charged nuclei
- Single covalent bonding
- If atoms are bonded by
one shared pair of
electrons then it is
known as a single bond
- Ionic bonds act
in all directions,
covalent only in
one
- Lone pairs
- A pair of outer-shell
electrons that aren't
involved in chemical
bonding
- Average bond enthalpy
- Not all covalent bonds are the
same strength
- 2.2.7
Dative
covalent
bonding
- One
atom
supplies
both
the
shared
electrons
in
a
covalent
bond
- Examples
- The
ammonium
ion, NH4+
- 3 covalent bonds and 1 dative
- One N has a lone
pair and bonds to
an H+
- Overall ion has a 1+
charge
- The oxonium
ion, H3O+
- Acid+water
forms
oxonium ions
- One of
the lone
pairs
from
the O in
H2Oforms
a
dative
covalent
bond
- 2.2.8 Structures of covalent
compounnds
- 2 possible structures
- Giant covalent
- High melting
and boiling
point as high
temperatures
are needed to
break the
strong
covalent
bonds within
the lattice
- They cannot
conduct
electricity as
they have
no free
charged
particles
(except for
graphite)
- They are
insoluble in
both polar
and
non-polar
solvents as
the covalent
bonds are
too strong
to be broken
by either
polar or
non-polar
solvents
- Simple molecular
- Low melting and
boiling points are the
intermolecular
forces are weak and
so relatively little
energy is needed to
break them
- Not conductors of
electricity as there are
no free charged
particles
- Generally soluble
in non-polar
solvents such as
hexane. Weak
London forces
are able to form
between
covalent
molecules and
the solvents,
helping the
molecular lattice
to break down
and the
substance
dissolves
- 2.2.9 Shapes of molecules and ions
- Electron repulsion
theory
- The shape of a
molecule or
ion is
determined by
the number of
electron pairs
in the outer
shell
surrounding
the central
atom. These
pairs can be
bonding or
lone pairs
- All electrons
have a
negative
charge, each
electron pair
repels other
electron
pairs
- The shape
adopted will
be the shape
that allows
all the pairs
of electrons
to be as far
apart as
possible
- Bonded pairs: 1
Name of shape: linear
Angle:
- Bonded pairs: 2
Name of shape: linear
Angle:180
- Bonded pairs: 3
Name of shape:
trigonal planar
Angle: 120
- Bonded pairs: 4
Name of shape:
tetrahedral Angle:
109.5
- Bonded pairs: 5
Name of shape:
trigonal bipyramidial
Angle: 90 and 120
- Bonded pairs:6
Name of shape:
octahedral
Angle: 90
- Bonded pairs: 3
Lone pairs: 1
Name of shape:
pyramidal
Angle: 107
- Bonded pairs: 2
Lone pairs: 2
Name of shape:
non-linear
Angle: 1045
- 2.2.10 Electronegativity
and bond polarity
- Electronegativity
measures the
attraction of a
bonded atom for
the pair of
electrons in a
covalent bond
- Electronegativity
increases
towards the top
right corner of
the periodic
table, fluorine
being the most
electronegative
element
- If atoms in a
covalent bond are
different, one of
the atoms is likely
to be more
strongly attracted
to the bonding
electrons. The
bonding atom
with a greater
attraction for the
electron pair is
said to be more
electronegative
than the other
atom
- This creates a
small charge
difference
across the
bond. This
difference is
always present
and is called a
permanent
dipole. this is
now a polar
covalent bond
- Polar molecules have
polar bonds
- If a molecule is
non-symmetrical it can create
a charge difference across the
whole molecule. In
symmetrical molecules the
dipoles of the bonds cancel
out
- Oxygen is one of the most
electronegative elements and
bonds containing oxygen will be
polar, but this doesn't mean the
molecule will be polar
- 2.2.11 Intermolecular forces
- Intermolecular forces occur due to
constant random movements of
the electrons within the shells of
the atoms in molecules
- 2 main types of
intermolecular
force are hydrogen
bonds and van der
Waals forces
- Van der Waals forces:
- London (dispersion) forces
- They are caused by
random
movements of
electrons in atoms '
shells. This
movement
unbalances the
distribution of
charge within the
electron shells
- At any moment
there will be an
instantaneous
dipole across the
molecule
- The
instantaneous
dipole induces
a dipole in
neighbouring
molecules,
which can
further induce
dipoles in
neighbouring
molecules
- Increases with
the number of
electrons
- The effect of London
forces on boiling
points;
- As the
number of
electrons
increases,
so does the
strength of
the London
forces
- Permanent dipole-
permanent dipole
interactions
- Molecules with
permanent dipoles
will be attracted to
other molecules with
permanent dipoles.
E.g. like a magnet
- Permanent
dipole-induced
dipole
interactions
- Permanent dipoles
can come from
polar bonds being
present. when this
is near to a
non-polar bond it
can cause electrons
in the shells of the
nearby molecules
to shift slightly,
causing the
non-polar molecule
to become slightly
polar and the
attraction then
occurs
- 2.2.12 Hydrogen bonding
- A strong
permenant
dipole-permenant
dipole interaction
between an
electron deficient
hydrogen and the
lone pair of a N,F
or O molecule
- Allows water to be
denser than ice as the
hydrogen bonds become
fixed in an open
structure in ice
- Water has a high
boiling point as the
hydrogen bonds are
much stronger than
other intermolecular
forces
- Can explain high surface tension and viscosity of water