2012
Index
Covalent substances
1. Covalent bonds (dot cross, dative, displayed) ...... 2
2. Shapes of covalent molecules ...... 7
3. Electronegativity ...... 12
4. Intermolecular forces ...... 15
5. Explaining properties of substances ...... 19
6. Properties of covalent substances (inc simple & macromolecular) 21
Ionic substances
1. Ionic bonding (dot cross) ...... 22
2. Properties of ionic substances ...... 25
Metallic substances
1. Metallic bonding ...... 26
2. Properties of metallic substances ...... 28
Summary activities ...... 28
Introduction:
A chemical bond is a force which holds particles together in a material. There are three types of chemical bond: ionic, covalent and metallic. The properties of materials depend on the type of bonding present, so there are four types of substance: simple covalent (2 or more non-metal atoms bonded together), giant covalent (diamond, graphite, boron, silicon dioxide), giant ionic (a metal and a non-metal bonded) and giant metallic (the metal elements off the periodic table).
Activity 1 Classify the following as ionic, covalent or metallic substances:
1) NaBr ...... 2) MgBr2 ......
3) S8 ...... 4) CO2 ......
5) Fe ...... 6) Os ......
7) BF3 ...... 8) Mg3N2 ......
9) H2O ...... 10) Rb2S ......
Covalent Substances
Covalent substances have covalent bonds within molecules. The forces between molecules are also important – these are called intermolecular forces.
Definition of a covalent bond: The electrostatic attraction between a positively charged nucleus and a shared pair of electrons.
In covalent bonding each atom donates a valence electron to the bond.
This shared pair helps to fill the outermost energy shell of both atoms to achieve a noble gas configuration. A covalent compound consists of molecules and the bond is formed when electrons are shared.
Typical covalent compounds are formed between two non-metals. Examples of covalent compounds are: -
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Dot and Cross Diagrams for covalent compounds
The cross dot diagrams show the outer valence electrons of the elements as crosses and dots.
Example 1: Hydrogen
Each hydrogen atom has a single valence electron. It is unlikely that the hydrogen atom will lose its electron and become an ion due to the high electrostatic force of attraction for the electron from the nucleus. Instead it donates its electron to a covalent bond.
Diagram:-
The octet rule
Most compounds have eight electrons in their outer shell when involved in bonding. This is seen to be the stable configuration.
Example 2: Hydrogen chloride – HCl
Activity 2 Draw cross dot diagrams for each of the following covalent molecules.
a) Chlorine (Cl2) b) Methane (CH4)
c) Ammonia (NH3) d) Water (H2O)
Example 3: Carbon dioxide (CO2)
Oxygen has six electrons in the outer shell. Carbon has four electrons in its outer shell. Carbon needs to share four electrons to reach the octet and each oxygen atom needs to share two electrons.
Diagram:-
Example 4: Ethene (C2H4)
Dative Covalent Bonds
Dative covalent bonding (sometimes called coordinate bonding) is really a type of covalent bonding.
Definition: A dative bond is a covalent bond where only one of the bonded atoms donates both electrons being shared.
Example: the boron trifluoride/ammonia complex
Draw the cross dot diagram for boron trifluoride (BF3) next to that of ammonia (NH3) below.
Ammonia has a lone pair of electrons on the nitrogen atom, which can be donated to the boron trifluoride to form a co-ordinate or dative bond. This is shown in the dot and cross diagram below: -
Activity 3
1) What feature of the nitrogen (within ammonia) allows it to take part in dative covalent bonding?
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2) Complete the following past paper question (Jan 09 Q2b ii)
Displayed Formulae
Bonds can be drawn as ‘sticks’. This is called a displayed formula. Each stick represents a covalent bond (i.e. a pair of shared electrons). We can also draw lone pairs as two dots if we need to.
A hydrogen molecule, H2 and carbon dioxide, CO2 in dot & cross notation and displayed formula:
Dot & cross diagram Displayed formula
Activity 4
1) Turn the following dot and cross diagrams into displayed formula
a) b)
2) Using your knowledge of valencies, draw displayed formula for the following molecules:
a) Cl2 b) CH4
c) O2 d) C2H6
e) N2 f) CH2O
Shapes of molecules
Electron pair repulsion theory
The shape of a simple molecule or ion is determined by the number of outer electron pairs around the central atom. For example, the shape of a methane molecule (CH4) is dependent on the number of electron pairs around the carbon atom. The electron pairs repel each other so that they are as far apart as possible. The methane molecule has a tetrahedral shape.
Until now we have assumed that all atoms prefer to have 8 electrons in their outer shell (the octet rule). However, there are many exceptions. You can get stable atoms with as many as 12 electrons in their outer shell. These atoms have up to 6 outer electron pairs. They have different shapes because the pairs repel each other.
Complete the table below. Use pencil for all these tables.
Number of electron pairs / Dot and cross diagram of the covalent bonding around the central atom / Bond angles / Name of the shape2
3
4
6
How to determine the shape of a molecule without lone pairs
The pattern shown in the previous table can be used to predict shapes of other molecules without lone pairs. You will only be asked about these shapes listed below.
Formula / Number of electrons in the outer shell of the central atom / Total number of electrons involved in covalent bonding around the central atom / Number of electron pairs / Shape of the moleculeCF4 / 4 / 8 / 4 / tetrahedral
BF3
BeCl2
SiCl4
SCl6
How to determine the shape of a molecule with lone pairs
We have to use our knowledge of covalent bonding to decide whether there are any lone pairs of electrons around the central atom. A lone pair of electrons is held closer to the nucleus of the central atom than a bonding pair. Lone pairs have a greater repulsive effect than bonding pairs. Lone pairs repel each other more than bonding pairs repel each other. A lone pair will repel a bonding pair more than a bonding pair repels a bonding pair.
Activity 5
The three different types of interaction are:
bonding pair-bonding pair; lone pair-lone pair; lone pair-bonding pair
Order them in the table below.
Type of interaction / Repulsionthe most
somewhere in between
the least
We can use this as part of a systematic way of working out the shape of the molecule.
1. Use the periodic table to decide how many outer shell electrons the central atom has (this is the group number).
2. Draw the dot and cross diagram. For a negative molecular ion- add 1 extra electron to the central atom. For a positive molecular ion- take away 1 electron from the central atom.
3. Decide how many lone pairs are present and how many bonding pairs.
4. For molecules that have four pairs of electrons apply the following rules:
· 0 lone pairs, all the bond angles will be 109.5o;
· 1 lone pair, the bond angles will be 107o;
· 2 lone pairs, the bond angles will be 104.5o.
We can draw a 3-dimensional representation of methane as below.
Formula / Outer shell electrons of the central atom / Dot and cross diagram / Number of bonding pairs / Number of lone pairs / ‘3-dimensional’ representation of the shapeH2O
NH3
NH2-
NH4+
Molecules with multiple bonds
They repel each other in the same way that single bonds do. Here are some examples:
Name Formula Shape
Hydrogen cyanide ______
Carbon dioxide ______
Methanal ______
Activity 6
Draw dot and cross diagrams for the following molecules, predict their shapes and draw their shapes using the 3D representation.
Molecule / Dot and cross diagram / Name of shape and drawinga) PH3
b) SH2
c) H3O+
Activity 7
1) Exam style question: Use a dot and cross diagram to draw a displayed structure of NCl3. Predict with reasons, the bond angle and name of the shape. (hint: use phrases from previous pages).
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......
......
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2) Complete the following past paper question (Jan 10 Q5d ii)
Electronegativity and bond polarity
In covalent molecules where the atoms are the same (Br2, Cl2) the electrons in the bond are equally distributed between the bonding atoms. In covalent bonds between unlike atoms, one atom has a stronger attraction for the shared electrons than the other does.
Examples
Cl Cl F F
H Cl O
H H
The electron attracting ability of an atom is called electronegativity.
Definition of electronegativity: The power of attraction between a bonded atom and a pair of electrons in a covalent bond.
The scientist Pauling assigned numbers to different elements where the higher the number, the greater the attraction. Electronegativity increases across a period and decreases down a group.
Which element has the highest electronegativity? …………………
Which element has the lowest electronegativity? …………………
Example:
The bond in hydrogen chloride is predominantly covalent with one pair of electrons shared between the two atoms.
Notice the two electrons lie closer to the chlorine as the chlorine has a greater attraction for the electrons
Hd+ Cl
The HCl molecule is polarised with a small positive charge on the hydrogen atom and a small negative charge - on the chlorine atom. The HCl molecule is polar with a permanent dipole.
Symmetrical and unsymmetrical molecules
In symmetrical molecules dipoles cancel and there is no overall permanent dipole.
CCl4 is non polar. Although the four bonds are polar the dipoles act in opposite directions. The overall effect is that the dipoles cancel each other out. H2O is polar as the dipoles do not cancel each other out.
Bonds within a molecule can have dipoles (are polar bonds), but if the dipoles cancel out, the molecule has no overall dipole.
Activity 8
Label the following molecules as polar (overall dipole) or non-polar.
Intermolecular forces: forces acting between molecules
What are the different types of intermolecular force?
In molecules, atoms are held together by shared pairs of electrons called covalent bonds, which are very strong bonds. Molecules are also attracted to each other by weaker forces called intermolecular forces. There are three types of intermolecular force:
1. hydrogen bonding
2. permanent dipole-dipole forces
3. Van der Waal’s forces
Van der Waal’s forces
Van der Waal’s forces are temporary, induced dipole-dipole forces. They occur between non-polar molecules. The electrons in a molecule are constantly moving. At one instant, the electron distribution may be unsymmetrical. This produces a temporary dipole.
If another molecule gets close to the molecule with the temporary dipole, this affects its electron distribution. An opposite dipole is induced into the adjacent molecule. The two molecules are therefore attracted to each other; this is known as a Van der Waal’s force.
Example- Chlorine
Van der Waal’s forces only act for a short time, as the electron density is constantly changing. They are the weakest intermolecular force; typically about 1% of the strength of a covalent bond. Symmetrical molecules with no overall dipole interact through Van Der Waal’s forces.
The strength of Van der Waal’s interactions increases as:
a) The size of the molecule or atom increases;
b) The surface area of the molecule or atom increases.
Activity 9
Write strongest and weakest (Van der Waals forces) next to the respective particles, for each pair.
Permanent dipole-dipole forces
Polar molecules occur when atoms with different electronegativities bond, have permanent dipoles. Electrostatic forces of attraction act between the opposite charges of the adjacent molecules.
Example- hydrogen chloride
Very small differences in electronegativity do not produce significant dipoles; hydrocarbons are not polar.
Activity 10
Using the following Pauling electronegativity values, draw dipoles on the molecule of PF3 below:
N3.0 / O
3.5 / F
4.0
P
2.1 / S
2.5 / Cl
3.0
Hydrogen bonding
Definition: The interaction between the lone pair of electrons on a nitrogen, oxygen or fluorine atom and a hydrogen atom which is bonded to an N, O or F atom.
N, O and F are the three most electronegative atoms. The large difference in electronegativity between a hydrogen atom and either an N, O or F atom means that very polar bonds are formed, with the hydrogen having a strong partial charge (δ+).
Hydrogen bonding is a special case of dipole-dipole forces and are the strongest type of intermolecular force; typically about one tenth of the strength of a covalent bond.
Example: Water
Activity 11
1) Write underneath the substances if the intermolecular force is hydrogen bonding or permanent dipole-dipole.
2) Past paper question (Jan 10 Q3b ii)
Anomalous properties of water
There are three anomalous properties of water that you need to be able to explain.
Activity 12
Carefully match the explanation to the abnormal property of water.
Property / ExplanationHigher melting/boiling than expected / High strength of hydrogen bonding across the surface of the liquid
Solid is less dense than the liquid (ice floats on water) / Strong hydrogen bonds holding the water molecules together so a lot of energy is needed to break the intermolecular forces
High surface tension / Hydrogen bonding holds the molecules apart in an open lattice structure
Common cause of the anomalous properties is ......