Chemical Bonding

Report
Chemical Bonds
Ionic
(Metal & NonMetal)
Covalent
(2 or > nonmetals)
Metallic
Intermolecular
forces
Hydrogen bond
Van Der Waals
forces
Substances
Giant Structure
- High melting & boiling
points
Ionic compounds
– forms giant ionic
structures
Graphite
Giant covalent structure
- Atoms are held by
strong covalent bonds
-No van der waals forces
Diamond
Silicon dioxide
Metals
Simple molecular
structure
- Low melting & boiling
points
Covalent
- Molecules are held by
weak van der waals
forces
 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 electro-negativity
Electro-negativity
 ability of an atom in a covalent bond to
shared paired of electrons
attract ___________________________
to
itself.
 Metals
generally have low electronegativity
values, while non-metals have relatively
high electronegativity values.
 Fluorine, which has the greatest attraction
for electrons in bond-forming situations
(highest E value).
Type of Bond
Electronegativity Difference
Ionic
2.0 – 4.0
Polar Covalent
0.5 – 1.9
Nonpolar covalent
0 – 0.4
Polar covalent bonds are covalent bonds with
ionic character.
Ionic bond
Non polar
(pure covalent)
bond
Polar
covalent bond
Electrons are not shared.
E.g. Na+ Cl- , electron is transferred.
Electrons are equally shared.
E.g.Cl-Cl
Electrons are not equally shared.
E.g.  +  -
H  Cl
Atoms have different
electronegativity values
Type of Bond
Electronegativity Difference
Ionic
2.0 – 4.0
Polar Covalent
0.5 – 1.9
Nonpolar covalent
0 – 0.4
What type of bond is the following?
(a) N (3.0) and H (2.1)
(b) H (2.1) and H(2.1)
(c) Ca(1.0) and Cl(3.0)
(d) Al (1.5) and Cl(3.0)
(e) H (2.1) and F(4.0)
No bond is purely ionic or covalent..they have a
little bit of both characters.
 When there is unequal sharing of electrons a
dipole exists.

Dipole
- is a molecule that has 2 poles or regions with
opposite charges.
- is represented by a dipole arrow towards the
more negative end.
Besides ionic, metallic, and covalent bonds,
there are also attractions between
molecules
Intermolecular attractions are weaker than
ionic, covalent, and metallic bonds
There are 2 main types of attractions between
molecules: Van der Waals and Hydrogen

Van der Waals forces consists of the two weak
attractions between molecules
1. dipole
interactions
– polar
molecules
attracted to
one another
2. dispersion
forces – caused
by the motion of
electrons
(weakest of all
forces)

Hydrogen Bonds are forces where a hydrogen
atom is weakly attracted to an unshared electron
pair of another atom

This other atom may be in the same molecule or in a
nearby molecule, but always has to include hydrogen.
INTERMOLECULAR HYDROGEN BONDING
INTRAMOLECULAR HYDROGEN BONDING


Hydrogen Bonds have about 5% of the strength of an
average covalent bond
Hydrogen Bond is the strongest of all intermolecular
forces

A few solids that consist of molecules do not melt until the
temperature reaches 1000ºC or higher called network
solids (Example: diamond, silicon carbide)

A Network Solid contains atoms that are all covalently
bonded to each other
•
Melting a network solid would require breaking bonds
throughout the solid (which is difficult to do)
http://library.thinkquest.org/C006669/data/Chem/bonding/inter.html






The bonding pair of electrons spends most of its time
between the two atomic nuclei.
screening the positive charges from one another and
enabling the nuclei to come closer together.
Negative charge on the electron pair attracts both nuclei
and holds them together in a covalent bond.
When two atoms are chemically bonded, the two atoms
close together have less energy and therefore are more
stable than when separated.
Energy is given off by the atoms to form a bond, and
energy must be supplied (absorbed) to break the bond.
A covalent bond is the result of electrostatic attraction
between the nuclei of the 2 atoms and the pair of shared
electrons.
Number of bonds elements prefers depending on the number of valence
electrons. In general -

Fam i l y
H alog en s
F, B r, C l, I
C alcog en s
O,S
N itrog en
N,P
C arb on
C,Si
# C ov al en t B on d s*
1 bond
O


N

3 bond
C

4 bond
X
2 bond
often
often
often
always
When compounds are formed they tend to follow the Octet
Rule.
Octet Rule: Atoms will share electrons (e-) until it is surrounded by eight
valence electrons.
Rules of the (VSEPR) gamei) O.R. works mostly for second period elements.
Many exceptions especially with 3rd period elements (d-orbitals)
ii) H prefers 2 e.- (electron deficient)
.
..
.
iii) :C:
N:
:O:
:F:
.
.
.
4 unpaired 3unpaired 2unpaired
1unpaired up = unpaired e4 bonds
3 bonds
2 bonds
1 bond
O=C=O
NN
O=O
F-F
iv) H & F are terminal in the structural formula (Never central)
The atomic arrangement for a molecule is usually given.
HNO3
CH2ClF
Cl
H
C
O
F
N
CH3COOH
O
O
H
H
H C
H
H
H2Se
O
H
Se
C
H2SO4
O3
O
H
H O
S
O
O
O H
O H
O
In general when there is a single central atom in the molecule, CH2ClF, SeCl2, O3
(CO2, NH3, PO43-), the central atom is the first atom in the chemical formula.
Except when the first atom in the chemical formula is Hydrogen (H) or fluorine
(F). In which case the central atom is the second atom in the chemical
formula.
Find the central atom for the following:
1) H2O
a) H
b) O
2) PCl3
a) P
b) Cl
3) SO3
a) S
b) O
4) CO32-
a) C
b) O
5) BeH2
a) Be
b) H
6) IO3-
a) I
b) O
O
Count
(i)
(ii)
the no. of valence electrons.
If the species has a –n charge, add n to the electrons
If the species has a +n charge, subtract n from the electrons
Draw a skeletal structure.
(i)
If C is present, place C at the centre.
(ii)
If C is not present, place the LEAST electronegative atom at the centre.
Note: H is never the in the center.


Complete the octets of the outer atoms (except for H) by adding lone pairs
of electrons (including the 2 electrons shared with the central atom)

If there are any electrons left over, place them on the central atom as lone
pairs.

If the central atom does not have a complete octet, rearrange lone pairs on
the outer atoms to form double bonds between the central and outer atoms.
Continue doing until the central atom’s octet is satisfied.

If the species is charged, place it inside brackets and write the charge
outside the brackets.
Bond pair _____
Lone pair
..
Draw Lewis structures of the following molecules:
(a) H2O , NH3 , CO2 , OCl2 , PCl4+
(b) SO2 , NO+ , OCN- , COF2 , CO32- , NO2- , O3
(c) BeCl2 , BH3 , PCl5
In which of the above obey the octet rule?
1. Molecules with an odd number of electrons
2. Molecules in which an atom has less than an octet
3. Molecules in which an atom has more than an octet
1. Odd Number of Electrons
NO
Number of valence electrons = 11
N O
N O
R eso nace Arro w s
Resonance occurs when more than one valid Lewis structure can be
written for a particular molecule (i.e. rearrange electrons)
NO2
Number of valence electrons = 17
O N O
O N O
O N O
Molecules and atoms which are neutral (contain no formal charge) and with an
unpaired electron are called Radicals
O2
O O
O O
O x yg e n is a g ro u n d s ta te
"d ira d ic a l"
2. Less than an Octet
Includes Lewis acids such as halides of B, Al and compounds of Be
BCl3
Cl
Group 3A atom only has six electrons around it
B
Cl
Cl
However, Lewis acids “accept” a pair of electrons readily from
Lewis bases to establish a stable octet
Cl
Cl
Al
Cl
L e w is a c id
H
+
N
H
H
L e w is b a s e
Cl
Cl
_
Al
H
Cl
H
N
s a lt
+
H
AlX3
Aluminium chloride is an ionic solid in which Al3+ is surrounded by
six Cl-.
However, it sublimes at 192°C to vapour Al2Cl6 molecules
B2H6
Cl
Cl
Cl
Al
Al
Cl
Cl
Cl
A Lewis structure cannot be written for diborane.
This is explained by a three-centre bond – single electron is
delocalized over a B-H-B
H
B
B
H
H
H
H
H
Octet Rule Always Applies to the Second Period = n2 ;
number of orbitals
2s, 2px, 2py, 2pz ---orbitals cannot hold more than two electrons
Ne [He]; 2s2, 2px2, 2py2, 2pz2
n=2
n=3
Ar [Ne]; 3s2, 3px2, 3py2, 3pz2 3d0 3d0 3d0 3d0 3d0
n=3
3. More than an Octet
Elements from the third Period and beyond, have ns, np and unfilled nd
orbitals which can be used in bonding
PCl5
P : (Ne) 3s2 3p3 3d0
Number of valence electrons = 5 + (5 x 7) = 40
10 electrons around the phosphorus
Cl
Cl
Cl
P
Cl
SF4
S : (Ne) 3s2 3p4 3d0
Number of valence electrons = 6 + (4 x 7) = 34
F
F
S
The Larger the central atom, the more atoms you
can bond to it – usually small atoms such as F, Cl
and O allow central atoms such as P and S to
expand their valency.
F
F
Cl
 Bond
Strength
Triple bonds > Double bonds > Single bonds
Page 93
(1)The attraction between the 2 nuclei for 3 electron pairs in a
triple bond is > that for 2 electron pairs in a double bond which is
> than that for 1 electron pair in a single bond.
(2) Triple bonds are shorter due to greater attraction between the
bonding electrons and the nuclei with more electrons in the bond.
 Strength
Triple bonds > Double bonds > Single bonds
 Length
Single bonds > Double bonds > Triple bonds
Bond Type
Length
(nm)
Strength
(kJmol-1 )
C-C
0.154
348
C=C
0.134
612
CΞC
0.120
837
 In
some molecules and polyatomic ions, both
electrons to be shared come from the same
atom forming 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.
How do you draw the Lewis structure?
Page 94
 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
In BF3, there are only 6 electrons in the outer shell of boron.
There is space for the B to accept a pair of electrons.

Due to difference in electronegativity value between the 2
atoms in the bond.
Element
F
O
N
Cl
C
H
Electronegativity
4.0
3.5
3.5
3.0
2.5
2.1
Unequal distribution of electron density results in small
charges on the atoms
( δ+ and δ- )
Example

A dipole is established when two electrical charge of opposite sign
are separated by a small distance.
Dipole moment
A
molecule can possess polar bonds and still
non-polar.
 Check the Geometry of the molecule:
 The polar bonds are arranged symmetrically
so as to give zero net direction of charge.
i.e. Overall dipoles cancel so that there is no
overall dipole.
Non-polar Covalent bond
 No difference in
electronegativity value –
bond consists of
2 ____________ atoms.
 _______ net charge.
Examples :
Polar Covalent bond
 Due to the difference in
electronegativity value –
bond consists of
2 ____________ atoms.
 _______ net charge.
Examples:
 In
(i)
(ii)
the water molecule,
O-H bonds are significantly polar
The bent structure makes the distribution
of those polar bonds asymmetrical.
 Some
molecules have very low polarity - so
low as to be regarded as non-polar,
Name of molecule
Formula
Polarity of molecule
Hydrogen chloride
HCl
Polar
Water
H2O
Polar
Ammonia
NH3
Polar
Benzene
C6H6
Non-polar
Boron trichloride
BCl3
Non-polar
Methane
CH4
Non-polar
Bromobenzene
C6H5Br
Polar
Carbon dioxide
CO2
Non-polar
Sulfur dioxide
SO2
Polar
Tetrachloromethane
CCl4
Non-polar
For CO2 each C-O bond is polar since O is more
electronegative than C.
Why is the molecule non-polar?
 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
 Find
the number of electron pairs / charge
centres in the valence shell of the central
atom.
 Electron pairs / charge centres repel each
other to the positions of minimum energy in
order to gain maximum stability.
 Pairs forming a double or triple bond act as a
single bond
 Non-bonding pairs repel more than bonding
pairs.
 Methane
(CH4) – tetrahedral
 Ammonia (NH3) – pyramidal
 Water (H2O) – bent
 Carbon Dioxide (CO2) - linear
Valence Shell Electron-Pair Repulsion Theory
(VSEPR)
Procedure
1.
Sum the total Number of Valence Electrons
Drawing the Lewis Structure
2.
The atom usually written first in the chemical formula is the Central atom in the Lewis
structure
3.
Complete the octet bonded to the Central atom. However, elements in the third row have
empty d-orbitals which can be used for bonding.
4.
If there are not enough electrons to give the central atom an octet try multiple bonds.
Predicting the Shape of the Molecule
5.
Sum the Number of Electron Domains around the Central Atom in the Lewis Structure;
Single = Double = Triple Bonds = Non-Bonding Lone Pair of Electrons = One Electron Domain
6.
From the Total Number of Electron Domains, Predict the Geometry and Bond Angle(s); 2
(Linear = 180º); 3 (Trigonal Planar = 120º); 4 (Tetrahedral = 109.5º); 5 (Trigonal
Bipyramidal = 120º and 90º); 6 (Octahedral = 90º)
7.
Lone Pair Electron Domains exert a greater repulsive force than Bonding Domains. Electron
Domains of Multiple Bonds exert a greater repulsive force than Single Bonds. Thus they
tend to compress the bond angle.
Lone pairs are held closer to the nucleus than the bonding pairs.
The distance between the lone pair electrons and the bonding pairs of
electrons is shorter than the distance between the bonding pairs to
each other.
Repulsion due to lone pairs causes the bond angles to become smaller
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
(1050)
Consider
NH4+ , H3O+
,
NO2-
NH4+
As the 4 negative charge centres repel each other
and take up positions in space to be as far apart as
possible, the electron pairs are distributed in
tetrahedral arrangement.
H3O+
As the 4 negative charge centres repel each other and take
up positions in space to be as far apart as possible, the
electron pairs are distributed in tetrahedral arrangement.
With one lone pair of electrons, the actual structure is
trigonal pyramid with a bond angle of 1070 for H-O-H bond.
Read page 104
NO2As the 3 negative charge centres repel each other and take
up positions in space to be as far apart as possible, the
electron pairs are distributed in trigonal planar arrangement.
With one lone pair of electrons, the actual structure of the
ion is bent with a bond angle of about 1170 for O-N-O bond.
Consider
N2H4 , C2H2
 Forces
between molecules.
 Does not exist in giant structure (ionic cpds,
metals & giant covalent structure)
Intermolecular
forces
Van der Waals
forces
Hydrogen bond
http://chemtools.chem.soton.ac.uk/projects/emalaria/index.php?page=13
Intermolecular Forces:
are generally much weaker than covalent
or ionic bonds. Less energy is thus required to vaporize a liquid or melt a
solid. Boiling points can be used to reflect the strengths of intermolecular
forces (the higher the Bpt, the stronger the forces)
Hydrogen Bonding : the attractive force between hydrogen in a
polar bond (particularly H-F, H-O, H-N bond) and an unshared
electron pair on a nearby small electronegative atom or ion
Very polar bond in H-F.
The other hydrogen halides don’t form
hydrogen bonds, since H-X bond is less
polar. As well as that, their lone pairs are
at higher energy levels. That makes the
lone pairs bigger, and so they don't carry
such an intensely concentrated negative
charge for the hydrogens to be attracted
to.
Hydrogen Bonding & Water




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’ forces
(London Dispersion Forces).
Van der Waals’ forces increases
with increasing mass.
London Dispersion Forces –
significant only when molecules are close to each other
Due to electron repulsion, a temporary dipole on one atom
can induce a similar dipole on a neighboring atom
Prof. Fritz London
The ease with which an external electric field can induce a dipole
(alter the electron distribution) with a molecule is referred to as the
"polarizability" of that molecule
The greater the polarizability of a molecule the easier it is to induce
a momentary dipole and the stronger the dispersion forces
Larger molecules tend to have greater polarizability
Their electrons are further away from the nucleus (any
asymmetric distribution produces a larger dipole due to larger
charge separation)
The number of electrons is greater (higher probability of
asymmetric distribution)
thus, dispersion forces tend to increase with increasing molecular
mass
Dispersion forces are also present between polar/non-polar and
polar/polar molecules (i.e. between all molecules)
 The
strength of the intermolecular forces
determines how easily the molecules will
separate and hence the melting and boiling
points.
Why is there an increasing boiling points of the noble gases as
you go down the group?
The boiling points of the noble gases are
Helium
Neon
Argon
Krypton
Xenon
Radon
-269°C
-246°C
-186°C
-152°C
-108°C
-62°C
Because electrons are always moving around very
quickly, the charges switch around all the time.
 the more electrons in a molecule / atom, the
stronger these Van der Waals or London forces
are.
 This is seen in the increasing boiling points of the
noble gases as you go down the group.

The boiling points of the noble gases are
Helium
Neon
Argon
Krypton
Xenon
Radon
-269°C
-246°C
-186°C
-152°C
-108°C
-62°C
Cl2 : gas
I2 : solid
Iodine molecule is made up of larger atoms
with more electrons compared to chlorine.
With more electrons moving around, the
temporary dipole will be larger.
The larger atoms in the molecule means that
the valence electrons are less strongly held,
Hence the induced dipoles will be larger.
van der Waals forces are present between
covalent molecules with no H atom attached to
N, O or F.
 E.g. van der Waals forces are present between
HCl molecules.
 There’s also other intermolecular forces
beween the molecules :
permanent dipole - permanent dipole

 These
intermolecular forces between polar
molecules are stronger than between nonpolar molecules. (all things being equal)
 For polar substances with similar RMM, the
higher the dipole moment, the stronger the
dipole-dipole attractions and the higher the
boiling points.
Propane (C3H8) and ethanal (CH3CHO) both with RMM
= 44
 Ethanal has a higher bp.
Ethanal
- is a polar molecule
- has stronger intermolecular forces (van der waals &
dipole-dipole interactions) between the molecules of
ethanal than between the propane molecules.

It is not true that polar molecules have stronger
intermolecular forces and hence higher bp than
non-polar molecules.
Non-polar molecules with higher RMM might have higher bp.
Read further for a few exceptions
Boiling points increase for polar molecules of similar mass, but increasing
dipole:
Substance
Molecular Mass Dipole moment,
(amu)
u (D)
Boiling Point
(°K)
Propane C3H8
44
0.1
231
Dimethyl ether
CH3OCH3
46
1.3
248
Methyl chloride
CH3Cl
50
2.0
249
Acetaldehyde
CH3CHO
44
2.7
294
Acetonitrile
CH3CN
41
3.9
355

A hydrogen bond is a weak type of force that
forms a special type of dipole-dipole attraction
which occurs when a hydrogen atom bonded to a
strongly electronegative atom exists in the
vicinity of another electronegative atom with a
lone pair of electrons. These bonds are generally
stronger than ordinary dipole-dipole and
dispersion forces, but weaker than true covalent
and ionic bonds.
 Hydrogen
bonding is present between
covalent molecules with H atoms attached to
O, N and F
 Strength
of H bond
The larger the electronegativity of H and the
other atom (N, O or F), the stronger the H
bond.
Strength F > O > N
 Number of them that can be formed between
neighbouring molecules
Although the strength is such F > O > N, HF can only form 1 H bond to 1
neighbour.
H2O can form 2, thus promoting more intermolecular interactions .
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.
Ammonia molecule has 3 N-H bonds. N is
larger and < electronegative than F, has
far weaker H bonds due to lower electron
density on the N atom (only 1 lone pair)
compared to O and F.
When the RMM is large, we expect the boiling
point to be high because larger molecules have
more space for electron distribution and more
possibilities for instantaneous dipole moment.
 However,

Compound
RMM
Boiling pt (K)
H2O
18
373
HF
20
292.5
NH3
17
239.8
H2S
34
212
HCl
36.4
187
PH3
34
185.2
Greater intermolecular
force because H2O, HF,
NH3 all exhibit hydrogen
bonding .
tend to have higher viscosity
than those that do not have
H bond.
Substances which have multiple
H bonds exhibit even higher
Viscosity.
 Electronegativity
Cannot occur without significant
electronegativity difference between H and
the atom it is boded to.
E.g. Both PH3 and NH3
have trigonal pyramidal shape but only NH3 has
H bonding.
 Atomic size
When the radii of the 2 atoms differ greatly,
their nuclei cannot achieve close proximity
when they interact resulting in weak
interaction.
Is there any hydrogen bonding between the
molecules if CH3F?
H
H
C
H
H
F
H C F
H
H is not joined directly to F in each molecule,
hence no hydrogen bonding between the molecules.
Is there any hydrogen bonding between the
molecules of ethanol?
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.
Van der Waals forces are made of dipole-dipole and London dispersion forces
Group 4A hydrides
Groups 4, 5, 6A hydrides
Types of Covalent Substances
Covalent substances can be divided into 2
categories as shown in the table below:
Simple Covalent
Molecules
Macromolecules
Examples: Hydrogen,
Examples: Diamond,
Oxygen, Water, Carbon Graphite and sand
Dioxide and Methane
(Silicon Dioxide)
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.

Macromolecules
Molecules with giant molecular structures are called macromolecules.
Diamond
Graphite
Giant Molecular 1. Each carbon atom has 4 1. Three valence electrons
Structure
valence electrons.
in each carbon atom are
used for covalent bonding.
1. Each carbon atom is
2. The fourth electron is not
joined to 4 other carbon
used in chemical
atoms by strong
bonding.
covalent bonds in a
3. This gives hexagonal rings
of six atoms that join
tetrahedral
together to from flat layers.
arrangement.
4. These layers of carbon
atoms, lie on top of each
other and are held together
by van der Waal’s forces.
Diamond
Diagrammatical
Representation
Graphite
Structure of Graphite
Covalent Bonding
Properties
Uses
Diamond
1. Hard
2. Very high melting
point and boiling
point
3. Non-conductor of
electricity
Graphite
1. Soft and slippery
2. High melting point
and boiling point
3. Good conductor of
electricity
1. Gemstones
2. As tips of grinding,
cutting and
polishing tools.
1. In pencils
2. As a dry lubricant
3. Brushes for electric
motors
Fullerene
 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
isolated.
 This has led to a new branch of science called
nanotechnology.
Recall:
In ionic compounds, the ions are held together by
strong ionic bonds in a giant ionic lattice.
In simple covalent molecules, the attractive
forces between the molecules are known as
intermolecular forces or van der Waal’s forces,
which is weaker than the ionic bonds.
In giant covalent molecules, the atoms are held
together by strong covalent bonds in a giant
covalent lattice.
4
Metallic Bonds
Metals consist of positive ions surrounded by a 'sea of moving
electrons'.
The negative 'sea of electrons' attracts all the positive ions and
cements everything together.
Metallic bonds are the results of the strong forces of attraction
between the negative electrons and the positive ions.
Hence, metals have high melting points and high boiling points.
Physical Properties of Metals
Physical
Properties
High Density
Explanation
The close packing of
atoms explains why
most metals have a
high density
Good Conductor Metals are good
of Electricity
conductors of
electricity in the solid
and molten state.
This is due to the sea
of delocalized
electrons.
Diagram
Physical
Properties
Malleable and
Ductile
Explanation
Metals are malleable
(can be bent or
flattened) and ductile
(can be drawn into
wires) because the
layers of metal ions can
slide over each other
when a force is applied.
High melting and Metallic bonds are
strong bonds.
boiling point
Except for: Mercury has
a low melting point and
is a liquid at room
temperature.
Similarly, sodium and
potassium have low
melting and boiling
point.
Diagram
Explanation

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.
‘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
tetrachloromethane.
 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
ions.
Example,
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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