Topic 1.3: Structure and bonding

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1.3(a) Understand the nature of ionic, covalent and dative covalent bonds and the simple charge cloud representation of s and p bonds.

• MgO, CaF₂  and Li₂O are other examples of compounds that use ionic bonding.
  Task 1.3a.1 Draw dot and cross diagrams for these examples. LiF, KF, LiCl, NaF, MgCl₂, AlF₃, MgO, MgS, Na₂O, CaF₂  and Li₂O, Al₂O₃.
• Ionic size (radii) depends on: Principal quantum number of outer shell. Higher number, bigger ion. K⁺  bigger than Na⁺.
• Nuclear charge. Higher charge, smaller ion. Mg (+12 charge) smaller than Na (+11 charge)
• Strength of ionic bonds: Ionic radii - small ions, strong bonds.  Ionic charge - high charge, strong bonds.
• Direction: Ionic bonds are non-directional and have equal influence in all directions

• Covalent bonding involves atoms of non-metals completing their outer shells by sharing pairs of electrons. E.g. the hydrogen atom has one electron in its outer shell. Within the hydrogen molecule, (H₂) each atom contributes one electron to the bond. Consequently each hydrogen atom now has control of 2 electrons, one of its own and a second from the other atom, giving it the configuration of an inert gas (He).

Dative covalent bonds.

• Dative covalent bonding. Normally each atom contributes 1 electron to a covalent bond. In a dative covalent bond, 1 atom contributes both electrons to the covalent bond. For a dative covalent bond, an arrow replaces the straight line.

Charge Cloud Representation
Bonding orbitals are molecular orbitals formed by the bonding of two or more atoms.  Common bonds are between s, p, sp³ hybrid or sp² hybrid orbitals.
There are two types:- 1) Sigma  s and 2) Pi  p

1) A Sigma bond or a Sigma molecular orbital is formed by the single overlap of 2 s-orbitals, 2 p-orbitals or s and p orbitals overlapping, these are shown in the diagrams below.
=(Sigma) s  bond.

= (Pi) p bond.

Task 3.1a.4 Draw lines to show bonds and identify the sigma and pi bonds in the following: CHCl₃, O₂,CO₂, C₂H₄, C₂H₂.

1.3(b)Understand the intermediate nature of most bonds in terms of (i)Electronegativity differences leading to polarity in bonds. 
(ii)Polarising power of cations and polarisibilty of anions and the factors affecting these.

• Large anion : distant electrons less firmly held.
• Small cation : high charge density
• Highly charged cation : high charge density.
• easily polarised anions : Chloride^-, Bromide^-, Iodide^-.
• highly polarising cations : Lithium⁺, Berylium²⁺, Aluminium ³⁺.

Task 1.3bii.1 Select the most easily polarised anion and give a reason for the following pairs: bromide and chloride, chloride and iodide, bromide and iodide.
Task 1.3ii.2  Select the cation with the best polarising power and give a reason for: sodium and magnesium, aluminium and lithium, lithium and sodium, potassium and sodium.

1.3 (c)Understand that polar bonds may give rise to polar molecules.

1.3 (d) Understand the nature of intermolecular forces, resulting from interactions between permanent dipoles and induced dipoles and from the formation of hydrogen bonds.

• Electrostatic attractions between two molecules with permanent dipoles.
• Induced dipole interactions, when the dipole of one molecule polarises a neighbouring molecule.
• Dispersion forces arising because of small temporary dipoles in atoms.

Hydrogen bond
This is a very strong dipole-dipole force as hydrogen has a very high charge density (whenever hydrogen is joined to an electronegative element (E.g. F, N, and O)). (It is a bare proton). E.g hydrogen bonds are found in H-F.

Task 1.3d Name the bonding between the following molecule pairs. CH₄ and C₂H₆, H₂O and H₂O, CO and CH₄, CO₂ and NH₃, HF and H₂O, CO and Cl₂O.
 
(e) Show how these various types of bond give rise to giant atomic structures (eg diamond and graphite), hydrogen bonded molecular structures (eg ice), ionic structures (eg sodium chloride), simple molecular structures (e.g. iodine) and polymers (eg poly(ethene)) and how the properties of solids are related to their structure and bonding.
Simple molecular structures.

substance	density/gcm-3	melting point/K	electrical conductivity
O2	1.15 at 90K	55	very low
H2	0.07 at 20K	14	very low
S2	2.0	390	very low
CH4	0.46	91	very low
I2	4.93	387	very low

• Strong covalent bonds between atoms in molecules.

• Weak Van der Waals’ forces between molecules.

• Low melting points and boiling points because of weak forces between molecules. They require very little energy to be broken; therefore we have melting points and boiling points at low temperatures.

• Brittle (elements) because the weak forces between the molecules are unable to withstand a large external force.

• Electrical insulators because there are no charged particles to carry the charge.

• Often insoluble in water, especially the non-polar molecules. There is little interaction between polar water molecules and non-polar molecules. The exceptions are polar molecules.
  
• Low density because molecules are not pulled strongly together in solids or liquids and many are gases in which particles are far apart.

Task 1.3e.1

Ionic structures.
Caesium chloride

Sodium Chloride 6:6 co-ordination:  each chloride ion has 6 sodium ions as it’s nearest neighbours. Each sodium ion is also surrounded by six chloride ions as it’s nearest neighbours. The electrostatic forces holding the ions in place are not directional. mp = 1081K, density = 2.17gcm^-3
All ionic compounds:

• Have high melting points and boiling points and are solids at room temperature because each ion is held firmly in place by strong ionic electrostatic forces.
•  Have densities higher than water but much lower than typical metals.  Although efficiently packed there is some empty space between ions.
• Are often soluble in water because polar H₂O molecules are attracted to ions and the attractive forces of many H₂O molecules can pull ions away from their crystal structures.
  
• Do not conduct electricity while solid as no particles are free to carry the charge. ‚
• Conduct electricity when molten or in a solution as the ions are free to carry the charge.

The structure of Diamond.
• It is hard and has a very high melting point and boiling point because each atom is held firmly in place by 4 strong, short, covalent bonds and a lot of energy is required to break these strong bonds.
• Doesn’t conduct electricity even when molten as no charged particles to carry charge.
• Insoluble in water as forces between solvent and carbon atoms are too weak.
• Thermal conductor as rigid structure allows heat to be passed through vibrations.
• High density (3.51gcm^-3) because atoms are packed tightly together.

Task 1.3e.2 

Hydrogen bonded molecular structures
H₂O, NH₃,  Hydrogen bonds between molecules are stronger than other intermolecular forces.  So water has a higher melting temperature than expected and is a liquid and not a gas.  These are all electrical insulators as there are no charged particles to carry current.

The Structure of ice.

Ice has a molecular structure but unusual properties, e.g. it solid has a lower density than its liquid. Water is at its maximum density at 4^oC. Four hydrogen bonds, two through its hydrogen atoms and two through its oxygen atoms hold each water molecule in place. On melting, many hydrogen bonds are broken and the structure collapses to give a lower volume liquid.  Water expands on freezing so ice floats as it has a lower density than water.
 

The structure of polymers.
Low electrical conductivity as there are no charged particles to carry current.

Branched chain.
Low density.  Low melting point.  Chains further apart.  Van der Waal’s forces between the tips of the branches are weaker because fewer electrons are involved.

Linear chain.

1.3 (f)Understand the existence of interparticle forces in the liquid state and hence explain the trends in the boiling temperatures of the noble gases and the hydrides of the elements of Groups 4, 5, 6 and 7.

Trends in boiling points of  group 5,6 and 7 hydrides

The boiling temperatures of the hydrides of these groups also increase as you go down the group. There is one exception. The element at the top of the group has a higher boiling temperature and doesn’t seem to fit the pattern.  The increased boiling tempertures of the first hydride in each group is because the first element in each group is highly electronegative and therefore they form hydrogen bonds that are difficult to break. For this reason the boiling temperatures are higher.
 

(g)To interpret changes of state in terms of the types, motion and arrangement of particles (atoms, molecules and ions) present and explain associated energy changes.

Liquid:  Shape of container but otherwise spherical.  Finite volume.  Less compressible than a gas.
In a solid particles may only vibrate about a fixed position but in a liquid, they may also rotate and translation is also possible. In a liquid, particles also move from place to place although not as freely as in a gas.  There is greater disorder in the liquid but some order is retained and particles have higher energies.

Gas: Takes shape of container.  Volume greatly affected by pressure and temperature.  The energy of gas particles is much greater than for liquid particles and the separation is also much greater.  The kinetic theory describes the behaviour of particles in a gas.  Much energy is needed to vaporise a liquid.

• The spherical particles are in constant random motion.
• The particles are constantly colliding. Collisions are perfectly elastic, so energy through collisions is all given off (no energy is lost).
• There are no forces between the particles, except during collisions.
• The particles are tiny compared to the distance of separation.
• The time taken for the collision is very short compared to the time between each collision.

Melting - endothermic (energy taken in)
simple molecular structures: Weak Van der Waals forces between atoms or molecules so little energy needed to melt hence low melting temperatures.
hydrogen bonded molecular structures: Hydrogen bonds much stronger so more energy and higher melting temperatures needed.
Giant ionic structures: Strong electrostatic attractions between each ion and many other ions of opposite charge require much energy to melt hence high melting temperatures.
Giant atomic structures: Each atom held firmly by strong covalent bonds with no simple molecules.  Much energy needed to melt these structures and hence high melting temperatures.
Boiling - endothermic
With much of the bonding disrupted the difference between melting temperatures and boling temperatures is relatively small.
freezing - exothermic (energy taken in)
condensation - exothermic

(h) Recall the shapes of the following molecules and ions; BeCl₂ CO₂ HCl SO₂ H₂O BCl₃ NH₄⁺ CH₄ NH₃ SO₃^2- CO₃^2- NO₃^- PCl₅ SF₆

Task 1.3h.1 Draw diagrams to show the bonding in all of the above species.
Task 1.3h.2 Draw diagrams to show the shapes of the above species.

Story 1.3h Trainer SF6 Climate Change Danger

1.3 (i) Interpret the shapes in (h) in terms of the valence shell electron pair repulsion theory, and predict shapes of related molecules and ions (eg from a knowledge of NH₃ the shape of PH₃ can be predicted

2 pairs linear  angle = 180^o

3 pairs trigonal planar  BCl₃ 

4 pairs tetrahedral    angle = 109.5^o  CH₄ NH₃ NH₄⁺ (or bent) 

5 pairs trigonal bipyramidal

6 pairs octahedral (or square planar)

In a molecule, pairs of electrons create negative centres that surround the central atom. These negative centres repel each other and arrange themselves to be as far apart as possible, thus minimising repulsion. Lone pairs repel more strongly than bonding pairs of electrons therefore a lone pair distorts the shape of a molecule.  
Phosporous  has 5 outer electrons just like nitrogen so PH₃ will have four electron pairs and be tetrahedral with respect to electron pairs like NH₃.  Also like ammonia, PH₃ has a lone pair which generates more repulsion than bonding pairs of electrons.  The shape of PH₃ is therefore distorted tetrahedral.
Task 1.3i Predict and draw the shape for the following molecules or ions: BeF₂, SeO₂, SCl₆, H₂S, AlCl₃, SiCl₄, PF₅, PO₃^-, SiO₃^2-, Cl₂O.

1.3 (j)Describe metallic bonding and explain (a) the electrical conductivity of metals and graphite in terms of mobility of electrons (b) the melting points and melting points of metals

Graphite has layers of strongly bonded carbon atoms. The layers are stacked and joined together by weak bonds between the layers. Between the stacked layers there are delocalised electrons which can transfer electricity. This allows electricity to be conducted along the length of carbon planes but electricity can not be conducted across the layers because the electrons are localised inside each layers.