5) Shapes and Intermolecular Forces Flashcards
Describe the repulsion between lone pairs and bonded pairs.
Lone pairs repel more than bonded pairs.
For each lone pair in a molecule, how much is the bond angle reduced by?
2.5°.
How many bonded pairs and lone pairs are there in a tetrahedral molecule? Give the bond angle.
- Bonded pairs: 4.
- Lone pairs: 0.
- Bond angle: 109.5°.
How many bonded pairs and lone pairs are there in a pyramidal molecule? Give the bond angle.
- Bonded pairs: 3.
- Lone pairs: 1.
- Bond angle: 107°.
How many bonded pairs and lone pairs are there in a non-linear molecule? Give the bond angle.
- Bonded pairs: 2.
- Lone pairs: 2.
- Bond angle: 104.5°.
How many electron pair regions are there in a linear molecule? Give the bond angle.
- Electron regions: 2.
- Bond angle: 180°.
How many electron pair regions are there in a trigonal planar molecule? Give the bond angle.
- Electron regions: 3.
- Bond angle: 120°.
How many electron pair regions are there in an octahedral molecule? Give the bond angle.
- Electron regions: 6.
- Bond angle: 90°.
Define electronegativity.
The ability of an atom to attract the bonding electrons in a covalent bond.
What does electronegativity increase towards in the periodic table?
F.
What is the bond type for something that has an electronegativity difference of 0?
Covalent.
What is the bond type for something that has an electronegativity difference of 0-1.8?
Polar covalent.
What is the bond type for something that has an electronegativity difference greater than 1.8?
Ionic.
Predict the bond angle in an F2O molecule. Explain your answer.
- 104° – 105°.
- There are 2 bonded pairs and 2 lone pairs.
- Lone pairs repel more than bonded pairs.
Describe and explain two anomalous properties of water which results from hydrogen bonding.
- Liquid H2O is denser than solid H2O. In a solid state, ice has an open lattice therefore, H2O molecules are held apart by hydrogen bonds.
- H2O has a relatively high boiling and melting point as the relatively strong hydrogen bonds need to be broken and a lot of energy is needed to overcome hydrogen bonds.
Why does water have a high surface tension?
Strong hydrogen bonds on the surface.
Explain, with the aid of a diagram, the intermolecular forces in H2O that lead to the relatively high boiling point of H2O. Use ‘Shapes & Imf’ card to test knowledge.
Rate knowledge 1-5.
Suggest why H2S has a much lower boiling point than H2O.
No hydrogen bonding, therefore, weaker intermolecular forces.
Explain why silicon has a much higher boiling point than phosphorus.
- Silicon has strong covalent bonds between atoms.
- Phosphorus has weak forces between molecules meaning that the London forces are easily broken.
Explain why the boiling point increases from sodium to aluminium.
- From Sodium to Aluminium, the number of delocalised electrons increases.
- Therefore, the metallic bonding gets stronger.
Name the shape of an NCl3 molecule.
Pyramidal.
The O–H bonds in water and the N–H bonds in ammonia have dipoles. Why do these bonds have dipoles?
Oxygen/nitrogen is more electronegative which attracts a bonded electron pair more.
Describe and explain the density of ice compared with water.
- Ice is less dense than water.
- Because hydrogen bonds hold H2O molecules apart in ice which causes an open lattice structure.
Water and carbon dioxide both have polar bonds. Explain why water has polar molecules but carbon dioxide has non-polar molecules.
- CO2 is symmetrical, meaning the dipoles cancel.
- H2O is not symmetrical, meaning the dipoles don’t cancel.
Explain Sodium Chloride’s solubility in water in terms of bonding and structure.
- Good solubility in water.
- Water is polar and the ions interact with water molecules.
Explain graphite’s solubility in water in terms of bonding and structure.
- Insoluble.
- Graphite is non-polar, therefore no interaction with water.
When given values for electronegativity, how are the bonds labelled?
- The atom with the larger electronegativity is given the 𝛿- symbol.
- The atom with the smaller electronegativity is given the 𝛿+ symbol.
Describe how London forces arise.
-Uneven distribution of electrons.
-Creates an instantaneous dipole
-Which causes induced dipoles in neighbouring
molecules.
