2.4a: explain the meaning of the term electronegativity as applied to atoms in a covalent bond

Electronegativity: the ability of an atom to attract the electrons in a covalent bond

As you go up groups and across periods, electronegativity increases

FONClBrIdge (from highest electronegativity, downwards)

Electronegativity is affected by:

• nuclear charge
• the higher the charge, the more electronegative the atom
• atomic radius
• the smaller the radius, the more electronegative the atom
• shielding
• less shielding = more electronegativity

2.3e: discuss the different structures formed by carbon atoms, including graphite, diamond, fullerenes and carbon nanotubes, and the applications of these, eg the potential to use nanotubes as vehicles to carry drugs into cells

Diamond:

• made up of tetrahedral interlocking carbon atoms, which are all covalently bonded to four other atoms
• a giant covalent structure, extremely hard, chemically inert, poor conductor
• usually colourless

Graphite:

• black/grey, flakey
• made up of layers of trigonal carbon atoms, which are all covalently bonded to three other carbon atoms at 120° bond angles
• London (Van der Waals) forces form between the layers, which allows them to slide away in flakes
• good conductor, chemically inert
• used as a lubricant, in electrodes, in pencils, etc.

Fullerenes:

• buckminsterfullerene = 60 atoms of carbon, 32 sides
• dissolves in petrol to form a dark red solution, can conduct electricity
• has delocalised electrons

Nanotubules:

• elongated cage-like structures, all containing 12 five-membered rings and almost any number of six-membered rings
• make highly complex shapes
• could be used in air and spacecraft, as they have good electrical conductivity as well as enormous strength and lightness
• could also be used to inject drugs into cells

2.3d: demonstrate an understanding of the terms bond length and bond angle and predict approximate bond angles in simple molecules and ions

Bond angle: the angle between two adjacent bonds on the same atom

• linear: 180°
• trigonal planar: 120°
• tetrahedral: 109.5°
• trigonal bipyramidal: 120° and 90°
• octahedral: 90°
• trigonal pyramidal: 107.5°
• bent: 104.5°
• square planar: °

Bond length: the distance between the nuclei of two bonded atoms

2.3c: apply the electron-pair repulsion theory to predict the shapes of molecules and ions analogous to those in 2.3b

The shape of a molecule or ion depends on the number of electron pairs around the central atom

Electron pairs repel each other so that they stay as far away from each other as possible (minimum repulsion/maximum separation)

Lone pairs have greater forces of repulsion than bonded pairs, as they are held closer to the nucleus so have higher charge densities

2.3b: recall and explain the shapes of BeCl2, BCl3, CH4, NH3, NH4 +, H2O, CO2, gaseous PCl5 and SF6 and the simple organic molecules listed in Units 1 and 2

2.3a: demonstrate an understanding of the use of electron-pair repulsion theory to interpret and predict the shapes of simple molecules and ions

LINEAR (180°)

TRIGONAL PLANAR (120°)

TETRAHEDRAL(109.5°)

TRIGONAL BIPYRAMIDAL (120° and 90°)

OCTAHEDRAL (90°)

TRIGONAL PYRAMIDAL (107.5°)

BENT (104.5°)

SQUARE PLANAR (°)

1.6.1b describe the formation of ions in terms of electron loss or gain AND 1.6.1c draw electron configuration diagrams of cations and anions using dots or crosses to represent electrons AND 1.6.1e describe ionic bonding as the result of strong net electrostatic attraction between ions

IONIC BONDING 
- ionic bonds form between a metal and a non-metal
- ionic bonding is the complete transfer of electrons between positive and negative ions. The positive ion transfers electrons to the negative ion to empty its outer shell and become stable, forming a strong electrostatic attraction between the ions.

You should know:
- that cations are ions that have lost electrons and that anions have gained electrons
- how to draw electronic configuration diagrams of cations and anions using dots or crosses to represent electrons
- to describe ionic bonding as the result of strong net electrostatic attraction between ions
- that the melting point of ionic substances is higher the stronger the bond and why
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PRACTICE QUESTIONS
1.i) Name the type of bonding in magnesium chloride.

ii) Draw a diagram (using dots or crosses) to show the bonding in magnesium chloride. Include all of the electrons and the charges present.








iii) Suggest why the melting temperature of magnesium oxide is higher than that of magnesium chloride, even though both are almost 100% ionic.
Magnesium chloride has an equation of MgCl2, as Mg has a charge of 2+ and Cl has a charge of 1-. Magnesium oxide has the equation MgO, as O has a charge of 2-. Therefore, MgCl2 has two individual sets of bonds, while MgO only has one with two pairs within it. This means that MgO has a higher melting point, as more energy is required to break the bond, as it has stronger electrostatic attraction.

2. The bonding in magnesium oxide, MgO, is:
- ionic
- metallic and ionic
- ionic and covalent
- metallic and covalent

3. Element R is in group 1 of the periodic table and element T is in group 6. R and T are not symbols for the elements.
i) The compound of R and T will have the formula:
- RT
- RT6
- RT2
- R2T

ii) The compound of R and T will have bonding which is predominantly:
- ionic
- covalent
- dative covalent
- metallic

iii) In terms of its electrical conductivity, the compound of R and T will:
- conduct when solid and liquid
- conduct when solid but not when liquid
- conduct when liquid but not when solid
- not conduct when solid or liquid