Hydrogen bond
Van Der Waals
 Occurs
when one or more electrons are
transferred from the outer shell of one atom
to the outer shell of another atom (both to
gain stable duplet or octet structure).
 Results from the electrostatic attraction
between cations and anions.
Sodium atom Na 2,8,1
Chlorine atom Cl
Sodium ion
Chloride ion
 [Na]+ + e-
(Lewis structure)
Cl + e-  [ Cl ]Na
+ Cl  [Na]+[ Cl x]-
Sodium metal reacts with chlorine gas in a violent exothermic
reaction to produce NaCl.
What ions will be formed by these elements?
Na , Mg , Al , P , S and Cl
Draw the dot and cross diagram for magnesium
 The
transition elements form more than one
stable positive ion.
Name of Transition Element
Positive Charge
Fe2+ , Fe3+
Cu+ , Cu2+
Mn2+ , Mn3+ , Mn4+
Cr2+ , Cr3+ (not stable in air)
 Ionic
bonding typically occurs between metal
and non-metal. E.g. Barium fluoride, BaF2
 The reactivity of metals and non-metals can
be assessed using electronegativity (ability of
an atom in a covalent bond to attract shared
pairs of electrons to itself).
Fluorine, which has the greatest attraction for electrons in
bond-forming situations is assigned the highest value on
this scale. All other atoms are assigned values less than
that of fluorine as shown above.
Note the following trends:
Metals generally have low electronegativity values, while
non-metals have relatively high electronegativity values.
Electronegativity values generally increase from left to
right within the Periodic Table of the elements.
Electronegativity values generally decrease from top to
within each family of elements within the Periodic Table.
If the difference in E values is > 1.8 => ionic bond
 If the difference in E values is 0, non-polar
covalent bond
 If the difference is 0 – 1.8, polar covalent bond
Polar covalent bonds are covalent bonds with
ionic character.
Ionic bond
Non polar
covalent bond
covalent bond
Electrons are not shared.
E.g. Na+ Cl- , electron is transferred.
Electrons are equally shared.
Electrons are not equally shared.
E.g.  +  -
H  Cl
Atoms have different
electronegativity values
 Use
the table above to predict the type of
bonding beween
Fluorine , F2
Hydrogen iodide, HI and
Lithium fluoride, LiF
Name of ion
Formula Example of compound
NH4Cl, ammonium chloride
H3O+Cl-, hydrochloric acid
MgSO4 , magnesium sulfate
Hydrogencarbonate HCO3-
KHCO3 , potassium hydrogencarbonate
AgNO3 , silver nitrate
K3PO4 , potassium phosphate
NaOH , sodium hydroxide
Na2CO3 , sodium hydroxide
 In
an ionic compound, constituent ions are
held in fixed positions in an orderly
arrangementt by strong ionic bonds.
Lattice structure
consisting of a regular
array of positively and
negatively harged ions.
 Formed
by equal sharing of electrons between
non-metallic elements to achieve the stable
electronic configuration of noble gases
 the shared electrons are localised between the
two nuclei the attraction between the
localised shared electrons and the nuclei is
known as a covalent bond.
 In the Hydrogen molecule, a bond between
two atoms is formed by the sharing of
electrons between the atoms.
The bonding pair of electrons spends most of its time
between the two atomic nuclei, thereby screening the
positive charges from one another and enabling the nuclei
to come closer together than if the bonding electrons were
absent. Negative charge on the electron pair attracts both
nuclei and holds them together in a covalent bond.
From an energy standpoint, when we say two atoms are
chemically bonded we mean the two atoms close together
have less energy and therefore are more stable than when
Energy given off by the atoms form a bond, and energy
must be supplied to pull them apart.
A covalent bond is the result of electrostatic attraction
between the nuclei of the 2 atoms and the pair of shared
Draw Lewis structures of the following molecules:
Chlorine, Cl2
Hydrogen chloride, HCl
Methane, CH4
Oxygen, O2
Nitrogen, N2
Carbon dioxide, CO2
Water, H2O
 In
some molecules and polyatomic ions, both
electrons to be shared come from the same
atom. The covalent formed is called the
coordinate or dative bond.
 Carbon monoxide (CO) can be viewed as
containing one coordinate bond and two
"normal" covalent bonds between the C atom
and the O atom.
thick white smoke of solid ammonium
chloride is formed in the reaction below:
 Ammonium
ions, NH4+, are formed by the
transfer of a hydrogen ion from the hydrogen
chloride to the lone pair of electrons on the
ammonia molecule.
When the ammonium ion, NH4+, is formed, the
fourth hydrogen is attached by a dative covalent
bond, because only the hydrogen's nucleus is
transferred from the chlorine to the nitrogen.
The hydrogen's electron is left behind on the
chlorine to form a negative chloride ion.
 Once the ammonium ion has been formed it is
impossible to tell any difference between the
dative covalent and the ordinary covalent bonds.
 Something
similar happens. A hydrogen ion
(H+) is transferred from the chlorine to one
of the lone pairs on the oxygen atom.
 The
H3O+ ion is variously called the
hydroxonium ion.
Other examples:
 The reaction between ammonia and boron
trifluoride, BF3
 Calculate the total no. of valence electrons for all
atoms in the molecule or ion.
 Arrange all the atoms surrounding the central atom
by using a pair of electrons per bond.The central
atom is usually the atom that is least
electronegative. [not H]
 Assign the remaining electrons to the terminal atoms
so that each terminal atom has 8 electrons. [ H will
only have 2 ]
 Place any electrons left over on the central atom. [P
and S elets from period 3 may have > 8 electrons]
 Form multiple bonds if there are not enough
electrons to give the central atom an octet of
Write the Lewis structure (electron dot
diagram) for hydrogen cyanide, HCN
 Strength
Triple bonds > Double bonds > Single bonds
 Length
Single bonds > Double bonds > Triple bonds
Bond Type
(kJmol-1 )
 In
diatomic molecules (e.g. H2 ,Cl2) both
atoms exert an identical attraction.
 When the atoms are different (e.g. HCl) with
one more electronegative than the other, a
polar bond is formed.
 Relative polarity is predicted from
electronegativity values.
 C-O
is more polar than C-Cl since the
difference in E value for C-O is greater than
that for C-Cl.
 The
shapes of simple molecules and ions can
be determined by using the Valence Shell
Electron Repulsion (VSEPR) theory.
- Electron pairs around the central atom repel
each other
- Bonding pairs and lone pairs arrange
themselves to be as far apart as possible
- All electrons in a multiple bond must lie in
the same direction, hence double and triple
bonds count as 1 pair of electrons.
The theory refer to negative charge centres ( =
pairs of electrons)
The 5 basic
molecular shapes
show the
arrangement of
the electron pairs
(charge centres)
that result in
minimum repulsion
between the
bonding and lone
pairs of electrons.
Order of repulsion :
lone pair-lone pair > lone pair- bonding pair > bonding pair – bonding pair
Methane, CH4
Bond angle is 109.50
Ammonia, NH3
Greater repulsion by lone
pair of electrons.
Bond angle is smaller
than 109.50(1050)
Water, H2O
Even greater repulsion by
two lone pair of electrons.
Bond angle is even smaller
dipole is established when two electrical
charge of opposite sign are separated by a
small distance.
 Polar molecules are formed between 2
atoms of different electronegativities.
 Polarity of a molecule depends on the
- the relative electronegativities of the atoms
in the molecule and
- the shape of the molecule
molecule with atoms of different
electronegativities may be non-polar even
though there are polar bonds in the molecule
as the dipoles may cancel each other and the
overall dipole moment is 0.
Name of molecule
Polarity of molecule
Hydrogen chloride
Boron trichloride
Carbon dioxide
Sulfur dioxide
 When
there is an uneven spread of electrons
making a covalent bond, the bond is called a
polar covalent bond. When there is an
uneven spread of electrons over a molecule
giving an unequal spread of charge over the
molecule, the molecule is said to be polar.
Water is an example of a polar molecule.
Like bond polarity, molecular polarity is also
shown using δ- and δ+.
 Consequently,
polar molecules arise when
there is a net direction of charge over the
molecule: polar bonds arranged in such a
way as to give a net direction of charge.
However, there are some instances when the
polar bonds are arranged symmetrically so as
to give zero net direction of charge; this is a
non-polar molecule. For example,
carbon dioxide and carbon tetrachloride,
 Note
that non-polar bonds can never give rise
to polar molecules. Some molecules have
very low polarity - so low as to be regarded
as non-polar,
 Usually
consists of a 3-D lattice of covalently
bonded atoms .
 The atoms can be either same like silicon
and carbon (graphite and diamond) or of 2
different elements such as silicon dioxide.
 Allotropes are two (or more) crystalline
forms of the same element, in which the
atoms ( or molecules) are bonded differently.
 Each C atom is tetrahedrally bonded
to 4 other C atoms by single
covalent bonds.
 Very strong C-C covalent bonds
have to be broken before melting occurs.
It is very hard and has high very high melting point
 All the electrons are used up in bonding (held
tightly between) the atoms, and are not mobile.
Hence, the electrons are localized, it does not
conduct electricity.
 Each C atom is covalently bonded to only 3
other C atoms to give layers of hexagonal rings.
Weak van der Waals’ force operates between
the layers, due to the large surface area.
 The layers can slide over each other so it is an
excellent lubricant .
 Each C has a spare electron which become
delocalized along the plane. Hence, graphite is
a good conductor of electricity.
 60 C atoms are arranged in
hexagons and pentagons to give
a geodesic spherical structure
similar to a football.
 Following
the discover of Buckminsterfullerene
, many other similar carbon molecules have been
 This has led to a new branch of science called
Van der Waals’ forces
 Chance charge separation - Electrons can at any
moment be unevenly spread producing a
temporary instantaneous (fluctuating) dipole.
 An instantaneous dipole can induce another
dipole in a neighbouring particle resulting in a
weak attraction between the two particles.
 The forces of attraction between temporary or
induced dipoles are known as Van der Waals’
 Van der Waals’ forces increases with increasing
Dipole-dipole forces
 Polar molecules are attracted to each other
by electrostatic forces.
 Although still relatively weak the attraction
is stronger than van der Waals’ forces but
weaker than ionic or covalent bonds.
 For polar substances with similar relative
molecular masses, the higher the dipole
moment, the stronger the dipole-dipole
attractions and the higher the boiling points.
Hydrogen Bonding
If two molecules of hydrogen
fluoride are close to one
another, the H atom of one
molecule will be attracted
to the F of the other molecule,
because of the electrostatic attraction between
the partial charge on the hydrogen atom and the
particla charge on the fluorine atom.
 This charge separation or dipole exists because F
is more electronegative than H.
 The electrostatic attraction that holds the H
atom of one molecule to the fluorine of another
molecule is an example of hydrogen bond.
Hydrogen Bonding
The essential requirement for its formation are a H
atom directly attached to O, N or F and a lone pair of
electrons on the electronegative atom.
In ammonia molecule, the
N atom h 1 lone pair of
electrons. Each NH3
molecule van form 1 H bond.
N is larger and < electronegative
than F, hence the H bonding is
weaker than that formed by HF
Hydrogen Bonding
Each the water molecule has 2 lone pairs of
electrons which can form H bonds with 2 other
water molecules.
 The collective strength of the H bonds in water
is greater than the strength of the H bonds in HF
because each O atom (with 2 lone pairs) in the
water molecule can form 2 H bonds with 2 other
water molecules, whereas each F atom in HF
molecule can only form 1 H bond with another
HF molecule.
Hydrogen bonding affects
 the boiling points of water, ammonia,
hydorgen fluoride and other molecules
 the solubility of simple covalent molecules
such as ammonia, methanol and ethanoic
acid in water
 the density of water and ice.
 the viscosity of liquids, e.g. the alcohols.
 The
valence electrons in
metals become detached
from the individual atoms
so that metals consist of a
close packed lattice of
positive ions in a sea of
delocalized electrons.
metallic bond is the attraction that two
neighbouring positive ions have for the
dolocalized electrons between them.
Metals are
 malleable, that is, they can be bent and
reshaped under pressure.
 ductile, which means they can be drawn out
into a wire.
 The valence electrons do not belong to any
particular atom, hence, if sufficient force is
applied to the metal, 1 layer of metals can
slide over another without disrupting the
metallic bonding.
The valence electrons do not belong to any particular
atom, hence, if sufficient force is applied to the
metal, 1 layer of metals can slide over another
without disrupting the metallic bonding.
The metallic bonding in metal is strong and flexible
and so metals can be hammered into thin sheets
(malleability) or drawn into lonng wires (ductility)
without breaking.
If atoms of other elements are added by alloying, the
layers of ions will not slide over each other so readily.
The alloy is thus less malleable and ductile and
consequently harder and stronger.
Compare properties of metallic, giant
covalent, simple molecular & ionic
‘Like tends to dissolve like’. Polar substances
tend to dissolve in polar solvents, such as water,
whereas non-polar substances tend to dissolve in
non-polar solvents, such as heptane or
 Organic molecules often contain a polar head
and a non-polar carbon chain tail. As the nonpolar carbon chain length increases in an
homologous series the molecules become less
soluble in water.
 Ethanol is a good solvent for other substances as
it contains both polar and non-polar ends.
 Water
will mix with polar liquids such as
ethanol. The oppositely charged ends of the
different molecules attract one another
forming hydrogen bonds.
 Gases
are generally slightly soluble in water.
 A small number of gases are highly soluble
because they react with water to release
SO2(g) + H2O (g)
H+(aq) + HSO3-(aq)
This solution is known as sulfurous acid , a
major component of acid rain

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