4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei.
4.2.2 Describe how the covalent bond is formed as a result of electron sharing.
4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom.
4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength.
4.2.5 Predict whether a compound of two elements in would be covalent from the position of the elements in the Periodic Table or from their electronegativity values.
4.2.6 Predict the relative polarity of bonds from electronegativity values.
4.2.7 Predict the shape of the shape and bond angles for species with four, three, and two negative charge centres on the central atom using the valence shell electron pair repulsion theory (VSPER).
4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities.
4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C60 fullerene).
Graphite
Diamond
Fullerene C60
4.2.10 Describe the structure and bonding in silicon and silicon dioxide.
Silicon
Silicon dioxide (silica/quartz)
- Covalent bond: electrostatic attraction between a pair of electrons and positively charged nuclei
- Molecule: group of atoms held together by covalent bonds
4.2.2 Describe how the covalent bond is formed as a result of electron sharing.
- The shared pair of electrons is concentrated in the region between the two nuclei and is attracted to them both.
4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom.
- Calculate the total number of valence electrons in the molecule by multiplying the group number of each element by the number of atoms of the element in the formula and totalling these.
- If ion: calculate the valence electrons and add for each negative charge or subtract for each positive charge.
- B and Be can have incomplete octets
- P, Cl, S, Xe, Br, I can have expanded octets
- Dative bonds: both electrons in pair originate from the same atom (H3O+, NH4, CO)
4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength.
- Bond length: measure of distance between the two bonded nuclei
- Bond strength: bond enthalpy
- Multiple bonds have a greater number of shared electrons and therefore have a stronger electrostatic attraction between bonded nuclei. Greater pulling power on the nuclei brings them closer resulting in a bond shorter and stronger than a single bond.
4.2.5 Predict whether a compound of two elements in would be covalent from the position of the elements in the Periodic Table or from their electronegativity values.
- 0 to 0.5 is pure covalent
4.2.6 Predict the relative polarity of bonds from electronegativity values.
- 0.5 to 1.8 is polar covalent
4.2.7 Predict the shape of the shape and bond angles for species with four, three, and two negative charge centres on the central atom using the valence shell electron pair repulsion theory (VSPER).
- Electron pairs found in the outer energy level of atoms repel each other and thus position themselves as far apart as possible
- The repulsion applies to both bonding and non-bonding pairs of electrons
- Double and triple bonded electron pairs are orientated together and so behave in terms of repulsion as a single unit known as a negative charge centre
- The total number of charge centres around the central atom determines the geometrical arrangement of the electrons
- The shape of the molecule is determined by the angles between the bonded atoms
- Non-bonding pairs of electrons (lone pairs) have a higher concentration of charge than a bonding pair because they are not shared between two atoms and so they cause more repulsion than bonding pairs
- The repulsion decreases in the following order: Lone pair–lone pair --> lone pair–bonding pair --> bonding pair–bonding pair
4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities.
- Dipole: bond has two separated opposite electric charges
- Polar: unsymmetrical with respect to electron distribution
- 0.5 to 1.8 and shape looks unsymmetrical
4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C60 fullerene).
Graphite
- sp2 hybridization
- covalently bonded to 3 others forming hexagons in parallel layers with bond angles of 120
- Layers held together with weak VDW forces so they can slide over each other
- 2.26 g/cm3
- conducts electricity due to the mobility of one non-bonded delocalized electron per atom
- non-lustrous, grey solid
- used as lubricant and in pencils
Diamond
- sp3 hybridization
- covalently bonded to 4 others tetrahedrally arranged with bond angles 109.5
- hardest know natural substance
- 3.51 g/cm3
- non-conductor because all electrons are bonded so there are no mobile electrons
- lustrous crystal
- polished for jewelry and ornamentation. Used in tools and machinery for grinding and cutting glass
Fullerene C60
- sp2 hybridization
- covalently bonded to 3 others in a sphere with 60 carbons consisting of 12 pentagons and 20 hexagons
- 1.72 g/cm3
- semiconductor as it easily accepts electrons to form negative ions so has some electron mobility
- yellow crystalline solid, soluable in benzene
- reacts with K to make superconducting crystalline material
- related forms are used to make nanotubes for the electronics industry, catalysts, and lubricants
4.2.10 Describe the structure and bonding in silicon and silicon dioxide.
Silicon
- Four valence shell electrons
- Covalently bonded to 4 others forming a tetrahedral arrangement
- Giant lattice structure like diamond
Silicon dioxide (silica/quartz)
- Bonds between Si and O
- Each Si bonded to 4 O and each O bonded to 2 Si forming a tetrahedral arrangement
- Strong, insoluble in water, high melting point, cannot conduct electricity or heat