Chemistry
The origin of chemistry can be traced to alchemy, or the
art of converting metals like copper to gold.
Ionic Bonding
Bonds
 What is a chemical bond?
 Electrostatic forces of attractions (between 2 atoms) between
the nuclei of one atom and the electrons of the other atom.
 Why do atoms bond?
 Three kinds of bonding:
 Ionic
 Covalent
 Metallic
Noble Gas
 Also called the inert gases or rare gases: He, Ne, Ar,




Kr, Xe and Rn.
Noble gases are unreactive.
Exist as individual atoms, monatomic.
Is there a need to bond?
Are there noble gas compounds?
Noble gas compounds
 Examples:
 Xenon tetrafluoride (XeF4)
 Xenon tetroxide (XeO4)
 Krypton difluoride (KrF2)
 Radon difluoride (RnF2)
 Xenon trioxide (XeO2)
 Notice what these
compounds contain?
Noble Gas Structure
 Duplet or octet configurations are most stable.
 Also known as a noble gas configuration.
 Common feature – fully filled valence electron shell
He
Ne
Why atoms combine
 Atoms WANT to achieve the noble gas configuration.
 How do atoms achieve the noble gas structure?
 Transferring or sharing electrons.
Recap – ions
 Normally, an atom is electrically neutral
 number of protons = number of electrons
 An ion is formed when an atom loses or gains
electrons.
 An ion is a charged particle formed from an atom
or a group of atoms by the loss or gain of electrons.
Cations
 Positive ions (cations) are formed by removing/
losing electrons from atoms.
 Loss of electrons tends to occur in atoms with few
valence electrons (e.g. 1, 2 & 3)
 Notice that these are METALS.
Cations
Na atom
Na+ ion
Number of
protons
11
11
Number of
electrons
11
10
 Electronic configuration: 1s22s22p63s1  1s22s22p6
 Na  Na++ e-
Na
 Na + + e-
Anions
 Negative ions (cations) are formed by gaining
electrons from atoms.
 Gain of electrons tends to occur in atoms with larger
number of valence electrons (e.g. 5, 6 & 7)
 Notice that these are NON-METALS.
Anions
Cl atom
Cl- ion
Number of
protons
17
17
Number of
electrons
17
18
 Electronic configuration: 1s22s22p63s23p5 
1s22s22p63s23p6
 Cl + e- Cl-
Cl + e- 
Cl
Discuss
 Consider sodium atom and sodium ion. Which do
you think is bigger and why?
 Consider chlorine atom and chloride ion. Which do
you think is bigger and why?
Na
Na+
Cl
Cl-
186 pm
106 pm
100 pm
181 pm
Trends in radii
 Atoms are always larger than any of their cations.
 Atoms are always smaller than any of their anions.
http://chewtychem.wiki.hci.edu.sg/Ionic+Bonding
Quickcheck
 Is Mg+ or Mg2+ bigger?
 Is O- or O2- bigger?
Ionic Bond
 An ionic bond is a chemical bond formed by the
electrostatic attraction between the positive and
negative ions.
 Formation of an ionic bond can be viewed as a
transfer of electrons from a metallic atom to a
non-metallic atom.
 Both will gain a duplet or octet configuration.
Video
 Link to video
Ionic bond
 Two processes occuring:
Na
 Na + + e-
Cl + e- 
Cl
Dot and Cross Diagram
-
+
Na
Cl
Na
Cl
Na+
Cl-
1s22s22p63s1
1s22s22p63s23p5
1s22s22p6
1s22s22p63s23p6
Isoelectronic
with
Isoelectronic
with
Neon:
Argon:
1s22s22p6
1s22s22p63s23p6
Dot and Cross Diagram
Magnesium Chloride
-
Cl
2+
Mg
-
Cl
Aluminium Oxide
3+
2 Al
2-
3 O
Is this possible?
Na
+ 7e- 
1s22s22p63s1
Cl
7-
Na
1s22s22p63s23p6 (Ar)
 Cl
1s22s22p63s23p5
7+
1s22s22p6 (Ne)
+ 7e-
Chemical formulae
 Take for example the following:
 An ionic compound made of Magnesium and
Fluorine.
Mg2+
Mg2+
FF-
F-
Chemical formulae
 For Magnesium fluoride, the ions present are Mg2+




and F-.
Mg2+ has 2 positive charges
F- has 1 negative charge.
To make the overall compound electrically
neutral there must be two F- to balance one Mg2+.
The formula is MgF2.
Chemical formulae
Mg2+ FMg1 F2
Na+
MgF2
OH-
Na1 OH1
NaOH
Charge on ion
+1
+2
+3
-1
-2
Name of ion
Formula
Sodium
Na+
Potassium
K+
Silver
Ag+
Ammonium
NH4+
Hydrogen
H+
Magnesium
Mg2+
Calcium
Ca2+
Iron(II)
Fe2+
Zinc
Zn2+
Aluminium
Al3+
Iron (III)
Fe3+
Chloride
Cl-
Fluoride
F-
Hydroxide
OH-
Nitrate
NO3-
Carbonate
CO32-
Oxide
O2-
Sulfate
SO42-
Find the chemical formula
Magnesium oxide
2. Zinc chloride
3. Calcium hydroxide
4. Iron(II) fluoride
5. Iron(III) sulfate
6. Ammonium nitrate
7. Silver chloride
8. Potassium iodide
9. Manganese(IV) oxide
1.
General knowledge
 General knowledge
 正离子 = cation
 负离子 = anion
 离子键 = ionic bonding
 共价键 = covalent bonding
 金属键 = metallic bonding
 Now you can hao4 lian4 to your friends!
Ionic lattices
 In ionic compounds, ions are held in fixed positions
in an orderly arrangement by strong electrostatic
forces (or ionic bonds) between the cations and
anions.
Ionic lattices
 Attractions are maximised in this structure while
repulsions are minimised.
 How many Cl- surround one Na+ and vice versa?
 The coordination number is the number of nearest
neighbours (atoms, ions or molecules). What are the
coordination numbers of Na+ and Cl-?
Ionic Lattices
 Different ionic lattice structures exist.
 http://www.avogadro.co.uk/structure/chemstruc/io
nic/g-ionic.htm
Properties
 High Melting and Boiling Points
 Hard and brittle
 Conducts electricity when dissolved in water or when
molten
 Many ionic compounds are soluble in water or polar
solvents like alcohol, but insoluble in most non-polar
solvents like hexane
Properties - NaCl
High Boiling and
Melting Point
due to breaking
of strong ionic
bonds
Is brittle due to repulsion
between similarly charged ions
Conductor of
electricity in liquid
state and when in
solution due to
presence of free
moving charges.
Linus Pauling
 Only person to have won
two unshared Nobel
Prizes (Chemistry and
Peace)
 The Nature of the
Chemical Bond and the
Structure of Molecules
and Crystals
 Pauling
Electronegativity Scale
Electronegativity
Click this link for the electronegativity table.
http://chemwiki.ucdavis.edu/@api/deki/files/4756/=electronegativity_chart.png
 Every atom has an attraction for the electrons shared
in a bond. Why?
Electronegativity
 Degree of attraction can be related to the
electronegativity of the atom.
 The higher the electronegativity, the more the
electrons in a chemical bond are attracted to the
atom.
The high affinity for electrons of fluorine leads it to direct reactions with all
other elements in which the reaction has been attempted, except for helium and
neon.
What patterns/ trends do you see?
Metals generally have
Electronegativity
values
generally
low
generally
electronegativity
increases
decrease
acrossdown
values,
a period.
the
while
group.
non-metals
have higher electronegativity values.
Predicting Ionic compounds
 What do you notice about ionic compounds?
 What are they made up of?
Metallic Bonding
METALLIC BONDS
Metallic names
 Ever wondered why some metals have weird symbols
in the Periodic table?
 E.g. Au (aurum) and Hg (hydrargyrum)
 Aurum actually means ‘shining dawn’.
 Hydragyrum means watery silver.
Comparing Ionic and Metallic Bonding
+
Li
-
Cl
Li
Li
Metallic bonding (I)
 Metals exist as giant structures too.
 In metals, atoms are packed closely together in
regular three-dimensional patterns to form a giant
lattice.
Metallic bonding (II)
Platinum atoms
Image originally created by IBM Corporation.
Metallic bonding (III)
 Mobile/ delocalised
e
e
e
e
Li+
Li+
Li+
Li+
e
Li+
e
Li+
e
Li+
e
Li+
e
Li+
e
Li+
e
Li+
e
Li+
electrons.
 ‘Sea of electrons’
surrounding the
positively charged
metal cations.
 Opposites attract i.e.
every positive ion is
attracted to the ‘sea
of electrons’.
Metallic bonding (III)
 Forces of attraction between positively charged ions
and negatively charged electrons – metallic bonding.
 Only found in metals! Not ionic or covalent
bonding!
Metallic Bonding
PROPERTIES OF METALS
Malleability and Ductility (I)
 A malleable substance is one which can be bent or
hammered out of shape without breaking.
 A ductile substance is one which can be stretched or
drawn into thin wires without breaking.
Malleability and Ductility (II)
 The following video shows the malleability of metals.
 Metals can be forged into different shapes and sizes
by beating and hammering
 Especially useful in making tools and machines.
Malleability and Ductility (III)
 The following video shows the ductility of metals.
 Notice that an iron strip can be twisted many times
before it finally breaks.
 Useful in making wires and cables.
 Why are metals malleable and ductile?
Malleability and Ductility (IV)
 Metal atoms are arranged in orderly layers.
 Application of a force causes metal atoms to slide
over each other easily.
 Why does the metal not break easily?
 Look at the following animation to understand
better.
Malleability and Ductility (V)
 What is this property useful for?
Sculpting
Wire
Malleability
Tubing
Ductility
Aircraft
Architecture
Machinery
Melting and Boiling Points (I)
 Metals generally have high melting and boiling
points.
 Most metals are packed closely together and the
strong forces of attraction between the positively
charged metal ions and the ‘sea of electrons’ result in
strong metallic bonding.
 A lot of energy is required to separate the metal
atoms.
Melting and Boiling Points (II)
 Can you name some exceptions?
 Just for fun: The following metals might melt at
Singapore’s room temperature! Francium (27oC),
Caesium (28oC) and Gallium (30oC).
Mercury
Melting and Boiling Points (III)
 What is this property useful for?
Lights
Electrical
appliances
Aircraft
High melting
point
Electrical Conductivity (I)
 ‘Sea of electrons’ surround metal cations.
 Mobile or delocalised electrons in the metal structure
allows conduction of electricity.
Electrical Conductivity (II)
 Metals conduct electricity due to mobile delocalised
electrons.
 Can they conduct electricity in the molten state?
 Can ionic compounds conduct electricity in the solid
state?
 Using electrical conductivity, how can we determine
whether a substance is an ionic compound or metal?
Electrical Conductivity (IV)
Sodium
chloride
Iron
Solid
Does not
conduct
electricity
Conducts
electricity
Molten
Conducts
electricity
Conducts
electricity
Aqueous
Conducts
electricity
(Insoluble)
Electrical Conductivity (V)
 What is this property useful for?
Wires
Electrical
appliances
Lightning
rod
Good
conductor of
electricity
Heat Conductivity (I)
 Metals conduct heat well due to mobile delocalised
electrons as well.
 When you heat one end of a metal, delocalised
electrons gain energy, move faster, and collide with
neighbouring electrons, thus transferring heat from
one end to the other.
 Look at the following animation to observe what
happens when metal is heated.
Heat Conductivity (II)
 Compare the movement of free electrons when a
metal conducts heat and electricity. What are the
differences?
Heat Conductivity (III)
 What is this property useful for?
Heating coil
Good
conductor of
heat
Refrigerator
Aircon
Metals as a whole
In general, metals…
Reason(s)
Have high melting and boiling points.
Strong metallic bonding.
Are good conductors of electricity.
Mobile/ delocalised electron that can
carry electric charges.
Are good conductors of heat.
Mobile/ delocalised electrons that
collide and transfer heat.
Are malleable and ductile.
Atoms arranged in orderly layers that
can slide past one another.
Have high densities.
For you to find out.
Are shiny in appearance.
For those who are interested.
Covalent Bonding
Covalent Bonding
 Consider the following (chlorine):
 We know that chlorine is often written as Cl2.
 Is this possible?
-
Cl
-
Cl
Covalent Bonding
 Covalent Bonding is the mutual electrostatic
attraction between the nuclei of atoms and their shared
electrons.
 Normally occurs between 2 non-metals only.
 For example, when two Hydrogen atoms meet, they will
each share one electron to get a duplet configuration.
Covalent Bonding
 Let’s take a look at Hydrogen.
 As two H atoms approach, the electron on each atom
is attracted to the nucleus of the other, i.e. there are
forces of attraction (between what?).
 What happens if they are too close?
 Because of repulsion between nuclei and
attractions between electrons and nuclei, there is a
distance between the two atoms where the molecule
is most stable. This distance is called the bond
length.
Lennard-Jones Potential
H
H
H
H
Bond
length
Distance between two atoms
http://upload.wikimedia.org/wikipedia/commons/5/5a/12-6-Lennard-Jones-
Covalent Bonding
 Bond length or bond distance is the average
distance between nuclei of two bonded atoms in a
molecule.
Covalent Bonding
 Now take a look at two chlorine atoms. How do they
share electrons?
Cl
Cl
 How many electrons are there around each chlorine
atom now?
Covalent bonding
 Covalent bonds exist for compounds too.
 What about the water molecule?
H
O
H
Covalent bonding
 How many electrons does an oxygen atom require?
O
O
 In this case, oxygen requires 2 electrons each. They
can achieve the noble gas configuration by sharing 2
electrons each. The resulting bond is a double
bond.
Nitrogen
 Try drawing the covalent molecule, nitrogen.
N
N
 In this case, nitrogen requires 3 electrons each. They
can achieve the noble gas configuration by sharing 3
electrons each. The resulting bond is a triple bond.
Try these!
 NH3
 CO2
 SiCl4
 CH3I
Structural Formula
Molecule
Chemical Formula
Structural Formula
Hydrogen
H2
H–H
Oxygen
O2
O=O
Nitrogen
N2
NN
Water
H2O
H–O–H
Methane
CH4
H
H–C–H
H
Polyatomic ions
O
O- N+
O- C
O-
Nitrate (NO3-)
O
O
S
O-
O
OCarbonate (CO32-)
O- H
O-
Sulfate (SO42-)
Hydroxide (OH-)
Sigma Bonds
 Head-on overlap of orbitals (s, p or d-orbitals)
 The resultant electron cloud is called a sigma bond
 Strongest kind of covalent bond
1s
1s
H atomH2 moleculeH atom
p
p
Pi bond
 Side-on overlap of orbitals (p or d orbitals)
 Weaker than sigma bond because of less overlap.
 Can a molecule with only single bonds have pi
bonds? E.g. Cl2?
Sigma bond vs Pi bond
 Pi bonds can only form after a sigma bond is formed.
 Sigma bonds can be found in all covalent
compounds.
 Pi bonds are only found in double bonds or triple
bonds.
Identify the bonds in the given molecules
F
F
O
O
N
N
H
H
C
C
H
H
H
C
C
H
Hydrogen
 Let’s take a look at the hydrogen molecule.
 = 2.1
 = 2.1
H
H
 Each H atom has the same attraction for the shared
electrons .
 Thus the electron density is evenly distributed
over the whole molecule.
 The molecule is said to be non-polar.
Hydrogen chloride
 Now consider the hydrogen chloride molecule.
 = 2.1
 = 3.0
H
Cl
 Chlorine is much more electronegative than
hydrogen, hence, it attracts the bonding pair of
electrons more strongly.
 Thus the electron density of the bond in HCl is
pulled towards the Cl end of the molecule.
Hydrogen chloride
 Now consider the hydrogen chloride molecule.
 = 2.1
 = 3.0
H
Cl
+
-
 This results in a separation of charge.
 The molecule is said to be polar and the bond is a
polar covalent bond.
 The molecule is said to have a permanent dipole
moment.
Polar bonds
 What happens if one atom is very much more
electronegative than the other?
+
-
+
-
 If one atom is very much more electronegative than
the other, it pulls away the bonding electrons such
that it becomes an ionic compound.
Fluorine
 Fluorine is so electronegative that it can bond to
almost all atoms, even some noble gas atoms.
 It does so because it has very strong attractions for
electrons, such that even the electrons of noble gas
compounds can be attracted by fluorine, thus
forming a bond!
Electronegativity and bond polarity
 As a rule of thumb:
 If difference in electronegativity ranges from 0 to 0.5, it is
considered a non-polar covalent bond.
 If difference in electronegativity ranges from 0.6 to 1.6, it is
considered a polar covalent bond.
 If difference in electronegativity is above 2.0, it is considered
an ionic bond.
 For 1.7 to 1.9, if a metal and non-metal is present, it is
considered ionic; if two non-metals are involved, it is
considered polar covalent.
Try these!
 Predict whether the following are covalent or ionic.
 sodium bromide
 hydrogen fluoride
 aluminium oxide
 aluminium chloride
 beryllium chloride
 caesium fluoride
Overall dipole
 Vector sum of all dipole moments on the molecule
Individual dipole
Overall dipole
O
-
H
+
H
+
Polar or non-polar?
 Not all molecules with a polar bond has a dipole!
 For example, carbon dioxide.
O=C=O
-
+
Cancel each other
-
Try these!
 Predict whether the following are polar molecules.
 NH3
 H2S
 CO
 CH4
 CH3Cl
Bond strength
 The strength of the covalent bond between atoms is
known as the bond strength.
 The higher the bond strength, the more energy is
required to break the bond.
 Which molecule has the highest bond strength,
nitrogen, oxygen or hydrogen?
Bond strength
Bond
Bond Length
Bond Energy
C–C
154 pm
348 kJ/mol
C=C
134 pm
614 kJ/mol
CC
120 pm
839 kJ/mol
266 kJ/mol
225 kJ/mol
 一根竹竿容易弯，一把筷子难折断
 Strengthtriple bond>Strengthdouble bond>Strength single bond
 What is the relationship between bond length and
bond strength?
Bond strength
Across a period:
Down a group:
Bond
Bond
Bond Energy
Length
Bond
Bond
Bond Energy
Length
C–F
135 pm
488 kJ/mol
C–F
135 pm
488 kJ/mol
C–O
143 pm
360 kJ/mol
C – Cl
177 pm
330 kJ/mol
C–N
147 pm
308 kJ/mol
C – Br
194 pm
288 kJ/mol
C–I
214 pm
216 kJ/mol
 In general, the shorter the bond length, the greater the
bond strength.
 A number of factors affect bond strength. These include,
number of bonds, size of atoms, electronegativity of atoms
etc.
Intermolecular Forces
 So far the attractive forces we have covered hold
atoms or ions together (intramolecular). What holds
MOLECULES together (intermolecular)?


van der Waals’ Forces
hydrogen-bonding
van der Waals’ forces
 3 kinds
 Permanent dipole-permanent dipole
 Permanent dipole-induced dipole
 Instantaneous dipole-induced dipole
Permanent dipole-permanent dipole
 Consider the hydrochloride molecule
 If two HCl molecules approach each other, they will
tend to arrange themselves such that the positive end
of one molecule is close to the negative end of the
other.
dipole-dipole attractions
H
+
Cl
-
H
+
Cl
-
Permanent dipole-permanent dipole
 Found in molecules with a permanent dipole
moment i.e. a polar molecule.
 Stronger the dipole moment, stronger the
intermolecular force. Why?
Strong covalent bond
 Much weaker than a covalent bond.
(intramolecular)
Weak dipole-dipole attractions
(intermolecular)
H
+
Cl
-
H
+
Cl
-
Polar molecules
 When exposed to an electric field, polar molecules
align themselves to the positive and negative
terminals
http://witcombe.sbc.edu/water/chemistrystructure.html
Instantaneous dipole-induced dipole
 What then holds non-polar molecules together?
 Instantaneous dipole-induced dipole, also known as
London dispersion forces.
 Exists between ALL molecules and atoms.
 Only kind of intermolecular attraction possible
between non-polar molecules.
Instantaneous dipole-induced dipole
-
Electron cloud
distribution is
symmetrical
+
Electron cloud
distribution becomes
unsymmetrical for an
instant
-
+
-
+
Neighbouring electron
cloud experiences an
induced dipole
NOTE: This kind of attraction is very short-lived because electrons are
always moving and the dipoles will disappear very quickly!
Instantaneous dipole-induced dipole
 What does this London dispersion forces depend on?
 Number of electrons/ Electron cloud size
 Shape/ Surface area
B.P. vs Number of electrons
Boiling point/ oC vs Number of electrons
0
0
10
20
30
40
50
60
70
80
90
-50
Rn
Boiling point/ oC
-100
Xe
-150
Kr
-200
Ar
-250
Ne
He
-300
Number of electrons
100
Helium
 Has a boiling point of -
268.93oC. Lowest
among all elements!
 Remains at a liquid
even at -273.15oC
(absolute zero)
 Can only become a
solid at 25 bar
pressure at -272.2oC
 Because of the weak
van der Waals’ forces!
Recap
 3 kinds
 Permanent dipole-permanent dipole
 Permanent dipole-induced dipole
 Instantaneous dipole-induced dipole
Recap
 How does pd-pd and id-id work?
 Under what conditions does pd-pd happen?
 Does id-id work on only non-polar molecules?
Instantaneous dipole-induced dipole
 Instantaneous dipole-induced dipole becomes more
important than permanent dipole-permanent dipole
in determining the boiling point as the molecule or
atom’s electron cloud increases.
Topics
 van der Waals’ in action
 Hydrogen-bonding
 Properties of Hydrogen-bonded molecules
 Hydrogen-bonding in real life
 Properties of simple covalent molecules
 (Properties of giant covalent structures)
Geckos and Spiders
 How do they stay on walls and even ceilings?
 Let’s take a look at this video.
 Then look at this website.
Geckos
 Geckos - Geckos have millions of setae--microscopic
hairs on the bottom of their feet. These tiny setae are
only as long as two diameters of a human hair. That's
100 millionth of a meter long. Each seta ends with 1,000
even tinier pads at the tip. These tips, called spatulae,
are only 200 billionths of a meter wide-below the
wavelength of visible light.
 A single seta can lift the weight of an ant. A million
setae, which could easily fit onto the area of a dime,
could lift a 45-pound child. If a gecko used all of its setae
at the same time, it could support 280 pounds.
 Geckos cannot stick to teflon (non-stick coating on
cooking pans). Go and find out why!
http://www.sciencedaily.com/releases/2002/08/020828063412.htm
Spiders
 On each of the spider's feet there are hair-like tufts,
called scopulae, … it was discovered that a single
scopula is itself composed of many, many, much
smaller, single hairs… The number of setules per
foot is estimated to be 78,000 each, and since
spiders have eight feet, they have upwards of
600,000 individual points of contact with any given
surface.
 The total adhesive force is extremely powerful, up to
170 times the weight of the spider, if all eight legs
are in contact.
http://www.istl.org/05-summer/article3.html
Hmmm…
 If scientists manage to come up with materials that
are like a gecko’s legs…,
 What uses can they be used for?
 What issues will crop up?
Intermolecular
 Are the only intermolecular forces van der Waals’?
Group IV hydrogen compounds
Group IV
110
CH4 has the lowest boiling
point as it is the smallest
molecule.
Boiling Point/ oC
60
10
-40
0
50
150
Group IV
SnH4
-90
GeH4
SiH4
-140
-190
100
CH4
Molecular mass
Larger molecules have
greater number of
electrons and thus, greater
intermolecular forces. As a
result, they have higher
boiling points.
Group V hydrogen compounds
Group V
110
NH3 is the smallest
molecule and is expected to
have the lowest boiling
point. Some additional
force must be present!
Boiling point/ oC
60
10
0
20
40
NH3
60
80
SbH3
AsH3
-140
120
140
Group V
-40
-90
100
PH3
Molecular Mass
Group VI hydrogen compounds
Group VI
110
H2O
H2O is the smallest
molecule and is
expected to have the
lowest boiling point.
Some additional force
must be present!
90
70
Boiling point/ oC
50
30
Group VI
10
-10 0
20
40
60
80
100
120
140
H2Te
-30
-50
-70
H2Se
H2S
Molecular Mass
Group VII hydrogen compounds
Group VII
Boiling point/ oC
110
HF is the smallest
molecule and is
expected to have the
lowest boiling point.
Some additional force
must be present!
60
HF
10
0
50
100
-40
-90
-140
150
HI
HBr
HCl
Molecular Mass
Group
VII
Hydrogen compounds
H2O
110
Hydrogen compounds
60
The additional forces
existing between NH3, H2O
and HF molecules are
called hydrogen bonds.
HF
H2Te
Boiling Point/ oC
10
0
-40
20
40
60
80
H2Se
H2S
NH3
AsH3
HBr
HCl
-90
PH3
GeH4
SiH4
-140
CH4
-190
100
Molecular mass
120
SbH3
HI
SnH4
140
Group IV
Group V
Group VI
Group VII
Hydrogen bonding
 Hydrogen bonding occurs only when molecules
contain an H atom covalently bonded to a very
small, highly electronegative atom with lone
pairs of electrons, i.e. F, O and H.
How H-bonding works
 When H is covalently bonded to an extremely
electronegative atom, F, O or N, the electronegative
atom will attract the electron cloud strongly, leaving
the H nucleus almost bare.
 Thus when another molecule containing an F, O
or N with lone pair of electrons approaches, it can
get very close to the H atom, thus the intermolecular
force is much stronger.
 This accounts for the high boiling points of water,
ammonia and hydrogen fluoride.
Conditions for H-bonding
 The H-atom must be covalently bonded to
either N, O or F, the 3 most electronegative
elements.
 There must be a lone pair on N, O or F of the
neighbouring molecule which can attract the
partial positive charge on the H-atom.
Hydrogen bonding
Weaker
(intermolecular)
Hydrogen bonds
O
H
H
Why must it be hydrogen???
Strong
(intramolecular)
covalent bonds
Strength of H-bonds
 Strength of H bonds: H – F > H – O > H – N
 Why does water have a much higher boiling point
than HF when HF forms stronger H-bonds?
Hydrogen Bonding in water
 Why does ice float on water?
 When most substances freeze, the particles are closer
to one another as they are in the solid state as
compared to the liquid state.
 Water, however, has maximum density at 4oC. Recall
that density = mass/volume. When water cools down
from 4oC to 0oC, the formation of solid ice actually
forces the water molecules to be fixed in a
tetrahedral shape.
 Click here for animation.
Hydrogen Bonding in water
 Similarly when ice melts, the solid structure actually
collapses, so that the water molecules are closer
together. They are closest at the temperature of 4oC.
Why is this important for freshwater fish in winter?
Ice (00C)
Water (10C)
Water (20C)
Water (30C)
Water (40C)
http://v.ku6.com/show/SXoq6mSJK2Ysgaga.html
Solubility in water
 Polar molecules are able to dissolve in water, which
is a highly polar molecule, due to H-bonding.
H H
H
C
C
H H
Ethanol
O H
Solubility in water
 Sugar (sucrose) is highly polar and can dissolve very well
in water.
 Carboxylic acids like ethanoic acid (vinegar) are able to
form H-bonds as well in water.
H
H
C
H
Sucrose
O
C
H
O
Ethanoic acid
Solubility in water
 Some ionic compounds dissolve in water as well.
 See animation.
H-bonds in life
 DNA (deoxyribonucleic acid) consists of two strands
of polymers (very long molecules).
 The 2 strands are held together by H-bonds.
 H-bonds can be broken by heating to high
temperatures.
http://en.wikipedia.org/wiki/File:DNA_chemical_structure.svg
http://en.wikipedia.org/wiki/File:DNA_orbit_animated_static_thumb.png
H-bonds in Life
 Proteins are made up of amino acids.
 Primary, secondary, tertiary and quaternary
structure.
 Secondary – alpha helix and beta-pleated sheet.
Alpha Helix
Beta-pleated Sheet
Properties of covalent molecules
 We have learnt that van der waals’ forces and
hydrogen bonds hold molecules together. These
forces are relatively weak as compared to covalent
bonds.
 What are the properties of such compounds?
Melting point/boiling point
 Simple covalent molecules have high volatility, i.e.


-
they have low boiling point.
Note: No breaking of covalent bonds required!
Reasons for low melting/boiling point:
Strong covalent bonds within the molecules but weak
van der Waals’ forces between the molecules
Little energy is required to overcome the weak
intermolecular forces.
Melting point/boiling point
 In an iodine molecule, the two atoms are held by a
strong covalent bond.
 Weak van der Waals’ forces hold the iodine
Strong covalent bond
molecules together.
 When heat is supplied, the weak intermolecular
forces break and iodine sublimes.
Weak van der Waals’ forces
between molecules of iodine
Iodine
molecule, I2
Solubility
 Simple covalent molecules are generally insoluble in
water or polar solvents UNLESS they are able to
form hydrogen bonds or can dissociate (later).
 Rule: “Like dissolves like.”
 For instance if you try to dissolve a non-polar
molecule like oil in water, water molecules will prefer
to bond to water molecules (they have H-bonds)
whereas the oil molecules will prefer clump together
because energy is needed to break the H-bonds.
Hence oil does not dissolve in water.
Electrical conductivity
 Simple covalent molecules do not conduct electricity.
 Absence of free moving electrons/ ions.
N
 What about water?
N
Electrical conductivity
 Some simple covalent molecules (often acids or
bases) when dissolved in water produce free moving
ions in the solution which can conduct electricity.
H+
H
Cl
H
Cl
Cl-
Checklist – Covalent or not?
 Commonly formed between non-metals
 Form bonds by sharing electrons
 After bonding, each atom achieves noble gas
configuration
Iodine? Sand?
Are they covalent?
Do they have similar physical
properties?
Note: Sand is made up mostly of silicon dioxide, SiO2 and is a main
component of glass.
Iodine vs Glass
 What do you think happens to glass (75% SiO2)?
 What about iodine?
Watch the Youtube Video on heating iodine in
glass: http://www.youtube.com/watch?v=Efs9OwE9Y0&NR=1
 After gentle heating,
- Physical form of glass remains intact.
- Iodine sublimed (Solid → Gas)
Classification of covalent substances
Covalent substances
Simple molecular
structure
Giant covalent
structure
Structure of SiO2
 SiO2 has a giant covalent structure.
 All the atoms are held
together by strong covalent
bonds that extend
throughout the structure.
 No separate molecules
(covalent bonds hold the
atoms together, not van der
Waals’ forces)
Note: This structure is not really accurate. The O
atoms are supposed to form a ‘V’ shape with the
silicon instead of a straight line.
O
Si
Structure of SiO2
 Each silicon atom is
bonded to 4 oxygen
atoms in a tetrahedral
shape.
 Each oxygen atom is in
turn bonded to 2 silicon
atoms.
 Each tetrahedral is
arranged in a repeating
pattern extending in
three dimensions.
O
O
Si
O
O
Structure of SiO2
 To melt this solid, a great
deal of heat is required
to break the covalent
bonds
 High melting point
Discuss
 What other physical properties does sand have,
besides high melting point, and why?
- Electrical conductivity?
- Brittleness?
- Solubility? In water? In organic solvent?
Carbon
Eight forms of elemental carbon: a) Diamond, b) Graphite, c) Lonsdaleite, d) C60
(Buckminsterfullerene or buckyball), e) C540, f) C70, g) Amorphous carbon, and h)
single-walled carbon nanotube or buckytube.
Diamond (from the ancient Greek αδάμας – adámas "unbreakable")
 Allotrope of carbon
C
 Giant covalent structure
 Tetrahedral
arrangement: Each
carbon atom is bonded
tetrahedrally to four
other carbon atoms
 Strong covalent bonding
in all directions
C
C
C
C
C
C
C
C
C
C
C
C
Strong covalent
bonds
C
Diamond
C
C
C
C
C
C
C
C
C
C
C
C
C
C
C
C
C
C
C
See Simulation of Diamond: http://www.worldofmolecules.com/3D/graphite.htm
Predict the properties of diamond
 Melting point

Extremely high melting point due to strong covalent bonds, Carbon
in all its allotropes has the highest melting point of all elements.
 Hardness

Hardest natural material due to strong covalent bonds, 10 on the
Mohs Scale of hardness
 Electrical conductivity

Does not conduct electricity because all valence electrons are used up
for bonding i.e. no delocalised electrons
 Solubility

Not soluble in polar or non-polar solvents due to very strong covalent
bonds.
Some uses of diamond
 Rock drill
 Jewelry
 Memory??? Watch this video.
 http://www.lifegem.com/
The carat
 You have often heard about the carat in
advertisements or movies. How much exactly is a
carat?
 1 carat = 200 mg
carat vs carat
 Is this the same as the carat used for gold?
 此carat非彼carat
 For gold, 24 carat is 100% carat.
 So for instance, 18 carat gold means it is 18/24 ×
100% = 75% gold.
Graphite
 Another allotrope of
carbon
 Trigonal planar
arrangement with
respect to each carbon
atom
 Layers of carbon atoms
 Each carbon covalently
bonded to 3 other
carbon atoms
Graphite
 Draw the dot-and-cross diagram of each carbon
atom. What do you notice?
One extra
valence
electron!
C
C
C
C
C C
C
C
One extra
valence
electron!
Graphite
Pictures taken from: http://en.wikipedia.org/wiki/Graphite
Weak van
der Waals’
forces
Layers of
carbon
atoms
Side view
 The layers are held together by weak van der Waals’ forces.
 Within the layers, atoms are covalently bonded in repeating
pattern of hexagons (6-membered rings).
See Simulation of Graphite: http://www.worldofmolecules.com/3D/graphite.htm
Properties
 Hard or soft?
 Within the layers, bonded strongly by covalent
bonds.
 But between layers, weak van der Waals’ forces can
be easily overcome upon stress.
Weak van
der Waals’
forces
Properties
 Does it conduct electricity?
 YES! Remember the extra electron from the dot-and-
cross diagram?
 Presence of delocalised electrons which can move
along the layers in the presence of an electric current
Delocalised
valence
electrons
Properties
 Solubility?
 Melting/boiling point?
 Same as diamond.
Uses of Graphite
 As a lubricant for hot machines
 Pencil lead
 Some batteries
Find out the uses of the following (Enrichment)
 Carbon nanotubes
 Graphene
 Carbon fibre
Covalent Substances
Simple molecular structure
Giant covalent structure
Arrangement
- Strong covalent bonds between the
atoms within each molecule
- Weak van der Waals’ forces/ H-bonds
acting between the molecules
- Strong covalent bonds between the
atoms
- Consist of three-dimensional repeating
patterns but no separate units.
Examples: sand, diamond, graphite
Physical properties
-Volatile, low melting & boiling point
-Usually liquids or gases at room
temperature
-Insoluble in water (There are
exceptions); Most are soluble in organic
solvents.
- Non-conductors of electricity
-Non-volatile, high melting & boiling
point
-Solids at room temperature
-Insoluble in all solvents
-Non-conductors of electricity (except
graphite)
Examples: iodine, carbon dioxide,
methane
Examples: sand, diamond, graphite