Co-ordinate (dative) covalent bonds
Work out which molecules have dative covalent bonds
and draw the structure
Specification reference 3.1.3
Pages 52-53
Learning outcomes
 State what is meant by the term electronegativity
 State what makes one atom more electronegative than
another
 State what the symbols δ+ and δ- mean when placed
above atoms in a covalent bond
Polar and non-polar
bonds
 Covalent bond - ?
 Ionic bond - ?
Vote now...
 Which of the three molecules is polar?
 Why?
Non-polar
 Covalent bond – shared pair of electrons
 In hydrogen, the two bonding atoms are identical
 They have an equal share of the pair of electrons in the
bond which results in a perfect 100% covalent bond
 Hydrogen is a non-polar molecule, the electrons are
evenly distributed between the atoms that make up
that bond
 What happens if the atoms are different?
Polar
 If the atoms bonding are different, one of the atoms is
likely to attract the bonding electrons more
 The bonding atom with the greater attraction is more
electronegative than the other atom
 Example, HCl
 The Cl atom is more electronegative than the H atom
 The Cl atom has a greater attraction to the bonding pair
of electrons that the H atom
 The bonding electrons are closer to the Cl atom than the H
atom
Trends in
electronegativity
Increasing electronegativity
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
Br
2.8
Key definition – Electronegativity
Electronegativity is the power of an atom to attract the electron density in a
covalent bond towards itself.
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Electronegativity
 Electron density is often used to describe the way the
negative charge is distributed in a molecule
 The Pauling scale is used a measure of electronegativity
(research opportunity…?)
 Why do the noble gases not have a number assigned?
 They do not (in general) form covalent bonds
Electronegativity
 Electronegativity depends on three things:
 The nuclear charge
 The distance between the nucleus and the outer shell
electrons
 The shielding of the nuclear charge by electrons in inner
shells
 Note – the smaller the atom, the closer to the nucleus
the shared outer shell electrons ∴ the greater the
electronegativity
 The larger the nuclear charge, the greater the
electronegativity
Trends in
electronegativity
Small atoms
Less shielding
Increasing electronegativity
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
Br
2.8
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Polar
 This leads to a charge difference across the H-Cl bond
Shows a small
positive charge
on the
hydrogen
Shows a small
negative
charge on the
chlorine
 This charge difference leads to a permanent dipole
and we represent it with the δ+ or δ- symbol
 We now have a polar covalent bond
Questions
 Identify which of these covalent bonds is/are polar, and
explain your answer.
 H–H
 H–F
 F–F
 Add dipoles to the polar covalent bonds
 Arrange the following covalent bonds in order of
increasing polarity: H – O, H – F, H – N
 Explain your answer
Polar molecules
 Poles many cancel each other out.
 E.g. HCl is linear – it has polar bonds and is a polar
molecule. The molecule is non-symmetrical and so there
is a charge across the whole molecule.
 Symmetrical molecules, the dipoles of any bond may
cancel out. E.g. Tetrachloromethane
Non polar polar
Each C-Cl
bond is polar
= a non-polar
molecule with polar
bonds!
Dipoles act in different
directions and cancel
each other out
Symmetrical molecule
(tetrahedral)
Patterns
Trends in
electronegativity
 Which atoms form the most electronegative atoms?
 Which atoms form the least electronegative atoms?
Electronegativity and
bonding type
 What does this diagram describe to you?
Questions
Complete the exam style questions
Homework
 Complete the relevant section of the webquest