Applying your knowledge and understanding of
electronegativity (Topic 3), explain why metals and nonmetals react to form ionic compounds.
High difference in EN values
• Metals – low EN, weaker attraction to
electrons
• Non-metals – high EN, stronger attraction to
electrons
• Very different tendencies to lose or gain
electrons.
EN difference of 1.8 or higher
• General rule, when the difference in
electronegativity values of two elements is 1.8
or higher, the compound they formed is
predominantly ionic.
Which fluoride is the most ionic?
1.
2.
3.
4.
NaF
CsF
MgF2
BaF2
Bonding experience!
• Take a piece of paper and draw a line down the
middle
• Place 20 Skittles in the center
• Play 10 rounds of Rock-paper-scissors
• Winner of each round moves two Skittles to
his/her side of the paper
• Count the number of Skittles on either side of the
paper at the end of 10 rounds
• Report result to Ms Tsui
• Don’t eat your Skittles just yet!
Results
•
•
•
•
•
•
•
Ms Tsui 6, Sakura 14
Radina 14, Sarah 6
Klaudia 10, Vera 10
Michael 12, August 8
Fintan 8, Isabel 12
Leo 12, Dominik 8
Chang 12, Nick 8
How can we relate this to chemistry in
terms of…
• Electronegativity?
• Nature of bonding?
Bonding continuum
1. Determine the position of your bond on the
bonding continuum
2. “Covalent”
– “Co” = together
– “Valent” = valence electrons
Limitation of this activity
• In reality, when atoms share electrons to form
covalent bond, they all begin with different
number of valence electrons.
• Atoms share valence electrons to achieve
octet structure, meaning to form a stable
arrangement of eight electrons in their outer
shell.
4.2 Covalent bonding
4.3 Covalent structures
When electrons are shared
Essential ideas
• Covalent compounds form by the sharing of
electrons.
• Lewis (electron dot) structures show the
electron domains in the valence shell and are
used to predict molecular shape.
4.2 Understandings
1. A covalent bond is formed by the electrostatic
attraction between a shared pair of electrons
and the positively charged nuclei.
2. Single, double and triple covalent bonds involve
one, two and three shared pairs of electrons
respectively.
3. Bond length decreases and bond strength
increases as the number of shared electrons
increases.
4. Bond polarity results from the difference in
electronegativities of the bonded atoms.
4.2 Applications and skills
• Deduction of the polar nature of a covalent
bond from electronegativity values.
4.3 Understandings
1.
2.
3.
4.
5.
6.
Lewis (electron dot) structures show all the valence
electrons in a covalently bonded species.
The “octet rule” refers to the tendency of atoms to gain a
valence shell with a total of 8 electrons.
Some atoms, like Be and B, might form stable compounds
with incomplete octets of electrons.
Resonance structures occur when there is more than one
possible position for a double bond in a molecule.
Shapes of species are determined by the repulsion of
electron pairs according to VSEPR theory.
Carbon and silicon form giant covalent/network covalent
structures.
4.3 Applications and skills
1.
2.
3.
4.
5.
6.
Deduction of Lewis (electron dot) structure of molecules and ions
showing all valence electrons for up to four electron pairs on each
atom.
The use of VSEPR theory to predict the electron domain geometry
and the molecular geometry for species with two, three and four
electron domains.
Prediction of bond angles from molecular geometry and presence
of non- bonding pairs of electrons.
Prediction of molecular polarity from bond polarity and molecular
geometry.
Deduction of resonance structures, examples include but are not
limited to C6H6, CO32- and O3.
Explanation of the properties of giant covalent compounds in
terms of their structures.
What are these bonds
between atoms like?
We’ll use the most simple
example, H2, to explain
•
First, look at the structure of one hydrogen atom
•
The positively charged nucleus and the negatively
charged electron attract each other.
•
•
Then two hydrogen atoms approach each other . . .
. . . the electron of one atom and the nucleus of the
other atom start to attract each other.
•
•
•
When the atoms are close, the attractions are so strong
that a molecule of hydrogen (H2) is formed.
The atoms are held together by the electrostatic
attractions between the two nuclei and the shared pair of
electrons.
This is a single covalent bond.
Definitions
• A covalent bond is the electrostatic attraction between a shared
pair of electrons and positively charged nuclei
• A double covalent bond consists of 2 pairs of shared electrons
• A triple covalent bond consists of 3 pairs of shared electrons
• Simple covalent structures are small molecules held together by
covalent bonds e.g. O2, N2, CO2, HCN, C2H4 (ethene) and C2H2 (ethyne)
O=O
N≡N
O=C=O
H-C≡N
H-C≡C-H
Lewis Structures
• Diagrams that show all the
valence electrons in a
molecule are Lewis structures
• Dot-cross diagrams show
origins of shared electrons
• Non-bonding pairs of electrons
are called lone-pairs e.g. water
Drawing Lewis structure instruction
• https://www.youtube.com/watch?v=nw3xVV
mEAU8
Worksheet, worksheet, worksheet..
• Practice drawing Lewis structures
Naming covalent compounds
• Different rules from naming ionic compounds
CARBON DIOXIDE
Chemical formula?
CARBON DIOXIDE
CO2
What connection can you make between the name
“carbon dioxide” and the chemical formula CO2
CARBON DIOXIDE
CO2
NITROGEN DIOXIDE
Chemical formula?
NITROGEN DIOXIDE
NO2
SULPHUR DIOXIDE
Chemical formula?
SULPHUR DIOXIDE
SO2
CARBON MONOXIDE
CO
SULPHUR TRIOXIDE
SO3
Between Nitrogen and Oxyen…
Nitrogen monoxide
Nitrogen dioxide
Dinitrogen monoxide
Dinitrogen trioxide
Dinitrogen tetroxide
Dinitrogen pentoxide
NO
NO2
N2O
N2O3
N2O4
N2O5
N2O3
N2O4
N2O5
Covalent compounds and their naming
• Two or more nonmetals bonded together
• Many can combine in more than one way
• Prefixes are used to indicate the ratio in which
nonmetal atoms are combined together
Naming covalent compounds
• The least electronegative element is named
first
• “Mono” is never used on the first element. If
there is only one atom of the first element,
this is shown by the absence of a prefix
• The ending of the last element named is
replaced with the suffix “-ide”
Naming covalent compounds
• With oxygen, if the prefix ends in an “a” or “o”,
the vowel of the prefix is dropped.
– E.g.
• carbon pentoxide ✔
• Carbon pentaoxide ✖
Practice examples
1.
2.
3.
4.
5.
6.
CO2
CO
CCl4
N2O5
SF6
CS2
7. Diphosphorus trinitride
8. Oxygen difluoride
9. Silicon tetrabromide
10. Phosphorus pentachloride
11. Iodine trichloride
12. Dinitrogen tetroxide
Practice examples
1.
2.
3.
4.
5.
6.
CO2
CO
CCl4
N2O5
SF6
CS2
Carbon dioxide
Carbon monoxide
Carbon tetrachloride
Dinitrogen pentoxide
Sulphur hexafluoride
Carbon disulphide
Practice examples
7. Diphosphorus trinitride
8. Oxygen difluoride
9. Silicon tetrabromide
10. Phosphorus pentachloride
11. Iodine trichloride
12. Dinitrogen tetroxide
P2N3
OF2
SiBr4
PCl5
ICl3
N2O4
Coordinate or dative bonds
• Sometimes, both the shared electrons come from the same atom.
• A coordinate or dative bond forms when a lone-pair of electrons, donates its
electrons to other atoms or ions that are deficient in electrons.
Coordinate or dative bonds
Carbon Monoxide
Coordinate or dative bonds
A beaker of water contains H2O and also H+ ions (protons), and
OH- ions and H3O + hydronium ions
• The charge is now delocalised across the polyatomic ion and is
shown by the square brackets [ ]
The 3 examples of coordinate bond
you need to know
Incomplete octets
• Some atoms, like beryllium Be and boron B,
might form stable compounds with
incomplete octets of electrons
1. Write the formula
2. Name this molecule
Beryllium chloride
• NOT Beryllium dichloride
• Follows ionic naming rule
• But draw as a covalent
molecule