structure and bonding

Trigonal planar

120

Bent

104.5 2 lone pairs

Linear

180

Tetrahedral

109.5

Pyramidal

107 one lone pair

Octahedral

90

Trigonal bipyramidal

120 90

Name the shape of a phosgene molecule and explain why it has this shape

• trigonal planar

• 3 bonding regions/3 bonded pairs and no lone pairs

• electron pair repulsion — bonded pairs repel each other equally

Name the shape of this molecule and explain why the molecule has this shape.

• pyramidal

• 3 bonded pairs and 1 lone pair

• the lone pair repels the bonded pairs more greatly than they repel each other so will give a pyramidal shape overall

Use a labelled diagram to explain why methanol is soluble in water.

State the shape and bond angle around a carbon atom in the alkyl group of propanoic acid. Explain the shape.

• tetrahedral

• 109.5

• four bonded pairs repel each other equally

Suggest a value for the C–O–H bond angle in propanoic acid.

104.5

State what is meant by the term ionic bond.

electrostatic attraction between positive and negative ions

Explain why the boiling points are different.

• 2-methylpropan-1-ol has less surface contact

• weaker london forces

• therefore less energy required to break london forces

This question is about halogens. Solid chlorine and solid bromine have a similar structure. Name this structure.

• simple molecular lattice

The shape around the oxygen atom in butan-2-ol is non-linear. Predict the C−O−H bond angle and explain this shape

• 104.5

• 2 bonded pairs and 2 lone pairs

• the lone pairs repels the bonded pairs more greatly than they repel each other

Explain the differences in the melting points of phosphorus and chlorine

• phosphorous has more electrons

• stronger london forces

• more energy required to break the london forces

Explain the different boiling points of NH3, F2 and Br2.

• NH3 has hydrogen bonding

• F2 and Br2 have london forces

• there are forces between molecules in ammonia and fluorine and bromine

• the london force in br2 are greater than in f2 because bromine has more electrons than flourine

• the london forces in br2 are greater than hydrogen bonding in NH3 but the hydrogen bonding in NH3 is stronger than the london forces in F2

Explain what is meant by the term electronegativity.

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

Draw a 3-D diagram of a molecule of CH2Cl2. Use partial charges to indicate polar bonds.

Explain why a CH2Cl2 molecule is polar.

• the dipoles do not cancel out because the molecule is asymmetrical

Describe what is meant by the term ionic lattice, in terms of the type and arrangement of particles present.

• repeating pattern of oppositely charged ions

What is meant by the term covalent bond?

the electrostatic attraction between 2 nuclei and a shared pair of electrons

Complete the diagram below to show hydrogen bonding between the H2O molecule shown and one other H2O molecule.

State and explain two anomalous properties of ice caused by hydrogen bonding.

• ice is less dense than water

• the molecules in ice are held apart by hydrogen bonds and ice has an open lattice

• ice has a relatively high melting point

• hydrogen bonds are relatively strong and more energy needed to overcome it

Predict the shape of a molecule of SbCl3.

• pyramidal

• 3 bonded pairs and one lone pair of electrons

• the lone pair repels the bonded pairs more greatly than the bonded pairs repel each other

SbCl3 molecules are polar. explain why

• there is a difference in electronegativities and so bonds are polar

• the molecule is not symmetrical and the dipoles do not cancel

The variation in boiling point can be explained by intermolecular bonding. Explain why H2S has a lower boiling point than H2O and H2Se.

• h2o has hydrogen bonding

• hydrogen bonding is stronger and requires more energy to overcome

• london forces in h2s are weaker as it has fewer electrons

• less energy the overcome london forces

Predict the type of structure and bonding of SO2 and MgO and explain the difference in their melting points.

• MgO = giant ionic

• SO2 = simple molecular

• ionic bonds in mgo are stronger than intermolecular bonds

• ionic bonds need more energy to overcome

Explain the physical properties shown in Table 16.1 using your knowledge of structure and bonding.

MAGNESIUM

• giant lattice

• metallic bonding

• delocalised electrons

bromine

• simple molecular

• london forces between molecules

magnesium bromide

• giant lattice

• ionic bonding between oppositely charged ions

• metallic and ionic bonds are stronger than london forces

• magnesium conducts due to delocalised electrons that cna move

• magnesium bromide = solids ions cannot move but in solution ions can move

• bromine does nto conduct as no mobile charge carriers

Use your knowledge of structure and bonding to explain the properties in the table

• calcium = metallic bonding and is a giant lattice

  • electrons are delocalised

  • metallic bonds are strong need large amount of energy

• br2 = simple molecular and has london forces between molecules

  • charge carriers are not mobile

  • london forces are weak and need little energy

Explain these physical properties of strontium, in terms of bonding and structure. Include a labelled diagram in your answer.

• melting point high

• very good electrical conductivity

• metallic bond strong and require a lot of energy to break — high mp

• strong attraction between electrons and positive ions

• delocalised electrons can move — good conductivity

Solid SiO2 melts at 2156 °C. Solid CO2 melts at −56 °C. Suggest the type of lattice structure in solid SiO2 and in solid CO2 and explain the difference in melting points in terms of the types of force within each lattice structure.

• sio2 = giant covalent

  • covalent bonds

• co2 = simple molecular

  • london forces

• covalent bonds are stronger than london forces so more energy to break bonds in sio2 than london forces in co2

Explain the properties shown in Table 1.1 in terms of bonding and structure.

• mgcl2 = giant ionic lattice

  • ions are mobile in liquid state

• sicl4/pcl3/scl2 = simple molecular

  • london forces

• ionic bonds are stronger than london forces

Explain the differences in the melting points of sodium and magnesium, using the model of metallic bonding.

• magnesium has more outer electrons

• magnesium ions have a greater positive charge

• magnesium has a greater attraction between ions and delocalised electrons

Magnesium and silicon have different types of giant structures. Describe the bonding in magnesium and in silicon

• magnesium has metallic bonds, positive ions and delocalised electrons

• silicon has covalent bonds between atoms

Describe and explain the electrical conductivity of sodium oxide, Na2O, and sodium in their solid and molten states.

• sodium conducts in solid and molten states

  • sodium has delocalised electrons

• na2o conducts when molten and not when solid

• molten na2o has ions which are mobile

• sold na2o has ions which are fixed in an ionic lattice

Explain the trend in melting point from Si to Cl across Period 3. • Comment, with reasons, on the similarities and differences in the trends across Period 3 and Period 4.

• si = giant covalent, covalent bonding and atoms

• p,s,cl = simple molecular, london forces, molecules

• covalent bonds in si are much stronger than london forces

• london forces greater with larger molecules as there are more electrons

• the stronger the force the higher the melting point

• period 4 = ge, se, br similar trend

  • as has much higher mp suggesting giant structure

  • ge has low mp suggesting weaker covalent bonds