d. Why does the ionisation energy drop between beryllium and boron?
The slightly lower ionisation energy of boron occurs because the extra electron has filled the 2p1 sub-shell and is being shielded by the 1s and 2s orbital’s, reducing the effective nuclear charge and lowering the energy required to remove the electron.
e. Why does the ionisation energy drop between nitrogen and oxygen?
The additional electron in oxygen has filled the 2pX2 sub-orbital which is already occupied. Because electrons exert a repulsive force on each other less energy is required to remove this electron. (it was this drop in ionisation energy which provided some proof of the existence of sub-orbital’s).
Task 2: Bonding
1. What type of bonding is present in sulphur dioxide? Draw a dot-cross diagram to show the electron arrangement within the molecule.
As both species are non-metals and have similar levels of electronegativity, they form covalent bonds, ‘sharing’ valence electrons, rather than transferring them. The negative electrons are held between the sulphur and oxygen atoms by the attraction of both (positively charged) nuclei. For the sulphur and oxygen atoms to satisfy the octet rule, there must be two types of covalent bond in the molecule; a dative bond with sulphur supplying both electrons to the bonding pair, and a double bond with an oxygen and the sulphur atom providing a pair each.
2. What type of bonding is present in potassium fluoride? Draw diagrams to show the electron arrangements of the atoms. Also draw a diagram to show how the atoms are arranged in 3D.
Potassium fluoride is formed by ionic bonds. The two species are electrostatically attracted to each other due to opposing charges, caused by the transfer of an electron from potassium to sulphur. As potassium fluoride is electrically neutral it must be composed of each element in a 1:1 ratio; and as ionic bonds are non-directional this facilitates the formation of the cuboid ionic lattice.
3. What type of bonding is present in magnesium metal? Draw a diagram to show how the atoms are and electrons are arranged.
Metallic bonds, when magnesium atoms bond, they off-load there two valence electrons to a delocalised electron sea. The difference in electrical charge between the magnesium nuclei and the delocalised electrons is what holds the metal together.
4. Name and draw the shapes of the following molecules. Explain your working.
a. NCl3 – Nitrogen trichloride (Trichloramine).
The Lewis diagram of nitrogen trichloride shows that there are four pairs of electrons around the central nitrogen atom. Three bonding pairs and one loan pairs, in three dimensions these electron pairs form a tetrahedral arrangement as this keeps them furthest apart. As only the bonded pairs are included in the molecules shape and as loan pairs exert a greater repulsive force than bonded pairs. We can predict that the molecular geometry of NCl3 will be trigonal planar.
b. H2S – Hydrogen sulphide (Sulphane).
From the Lewis structure of hydrogen sulphide we can see that again, there are four pairs of electrons around the central sulphur atom, only this time there are two loan pairs present. This means that the hydrogen-sulphur bonds assume two points of a tetrahedron but are even further compressed because of the two loan pair, resulting angular in a (bent) molecular geometry.
c. CCl4 – Carbon tetrachloride (Tetrachloromethane).
Here the central carbon is using all four of its valence electrons to bond with chlorine atoms. As there are no loan pair electrons carbon tetrachloride will have a tetrahedral shape with bond angles of 109.5o.
5. Use your knowledge of bonding and intermolecular forces to explain the following observations:
a. Lithium iodide is a solid at room temperature.
The bonds which form between lithium and iodine are mostly ionic, arising from the electrostatic attraction between the positive lithium cation and the negative iodine anion. As lithium’s charge density (charge to volume ratio) is high the bonds do exhibit some covalent characteristics, such as dissolving in organic solvents. But this does not affect the high bond energy between the two species, meaning that as the chemical bonds are so strong separating them requires more energy than is available at room temperature (23oC).
b. Aluminium is a solid at room temperature.
The aluminium atoms in aluminium metal are again held together by the electrostatic attraction that results from the transfer of electrons; this time to a delocalised electron sea which results in metallic bonding. Because each aluminium atom donates three electrons to the delocalised sea, the electron density which holds the atoms together is high. Meaning a large amount of energy is needed to separate the individual atoms, and that the melting point of aluminium is well above room temperature.
c. Methane is a gas at room temperature.
Methane is a discreet covalent molecule, formed of only a carbon atom and four hydrogen atoms which are held together by strong single covalent bonds. However the attractive force between the discreet molecules is a weak inter-molecular force, van der Waals force. Because van der Waals force is only weak methane has a exceptionally low boiling point, far below room temperature. Although it is only the inter-molecular force which is being broken not the covalent bonds.
d. Water is a liquid at room temperature.
Although water is also a discreet covalent molecule, the inter-molecular force of attraction between the molecules is different. Because the oxygen atom has a strong electronegativity the water molecule has a net polarity. Meaning that the positively charge hydrogen atoms for electrostatic bonds with the loan pair electron on surrounding oxygen atom. These hydrogen bonds require a high energy to break them. Because of this water has a high boiling point than room temperature.
6. Use your knowledge of bonding to explain the following observations
a. Aluminium conducts.
The metallic bonds between the aluminium atoms are formed by the valence shell orbitals overlapping to form a single molecular orbital. Because of this, the three valence electrons which each atom off-loads are free to move throughout the structure of the metal; and can act as mobile charge carriers.
b. Sodium chloride conducts electricity only when molten or dissolved.
When sodium and chlorine are bonded together in an ionic lattice both the ions and electrons are held rigidly in place by the electrostatic bonds. This means there are no free particles available to act as charge carriers. Once these bonds have been broken, either through melting or dissolving in a polar solvent the ions become available to carry a charge, allowing the compound to conduct electricity.
c. Methane does not conduct in any state.
Because there are no free charged particles in methane, with all the electrons involved in strong directional covalent bonds and no ions present in any state.
Task 3: bonding
Using the examples of PCl5, methane (CH4), ethene (C2H4), water and the ammonium ion (NH4+), explain how the shape of the molecule may be predicted. A dot-cross diagram should be included for each molecule as well as a diagram showing and naming the shape.
The simplest method of predicting the shapes (molecular geometry) of small covalent molecules and polyatomic ions is VSEPR (valence-shell electron-pair repulsion) theory. This uses Lewis structure’s to find the numbers of bonded and non-bonded electron pairs in the valence shell of a molecules central atom. The process works due to the repulsive force which electrons exert on each other because of their negative charge.
For example to find the approximate shape of the molecule PCl5 (phosphorous pentachloride) we start with a Lewis diagram:
While this is unusual, as the central phosphorous atom has 10 valence electrons, it is explained by the occurrence of hypervalence in elements from and above the 3rd period, due to the presence of d-orbitals.
As there are no loan pairs, the 5 bonds will repel each other to be as far apart in space as possible. Molecules with five bonds and no loan pairs have trigonal bipyramidal geometry with equatorial bond angles of 120o and 90o between the axial and equatorial bond. So this is the molecular geometry of phosphorus pentachloride.
The same process works with all molecules with a central atom and no loan pairs.
With methane the, which has four electrons pairs about the central carbon atom, the molecular geometry will be tetrahedral with bond angles of 109.5o. As this minimises the repulsive force between the bonds.
Ethane which is more complex should be considered as two trigonal planar molecules joined double bond in the middle. This gives bond angles of 120o around each carbon atom.
With molecules such as H2O where the central oxygen atom has two loan pairs in its valence shell, the shape of the molecule is affected. Instead of a linear geometry like CO2, the molecular orbital forms a tetrahedron, including the bond electrons and loan pairs. As such the hydrogen-oxygen bonds take and angular (bent) geometry. However due to loan pair electrons exerting a greater repulsive force than bonded ones, the bond angle is not 109.5o, but 104.5o as a result of the compression this causes.
VSEPR theory also works on polyatomic ions, such as ammonium (NH4+). This is because the bonds between the nitrogen and hydrogen atoms are covalent, thus have a fixed position and angle.
In the ammonium ion the central nitrogen has a full octet, with all four electron pairs involved in a bond. Even though one of the covalent bonds is dative the repulsion it exerts on the other bond is the same as a single or double bond; so has no effect on the geometry. This means four equal pairs, giving a tetrahedral geometry with 109.5o bond angels.