Bonding and Structure

Question 1

What is meant by an ionic bond?

Answer 1

A strong electrostatic attraction between the oppositely charged ions

Question 2

Explain why ionic compounds conduct electricity only when molten/dissolved in water?

Answer 2

There are free moving ions when dissolved in water/molten

Question 3

Explain why ionic compounds have high melting points:

Answer 3

Because in order to melt an ionic compound, you have to break many strong ionic bonds (between oppositely charged ions), this takes a lot of energy

Question 4

Explain why magnesium oxide has a higher melting point than lithium chloride:

Answer 4

The charge of an ion is related to the strength of the ionic bond. MgO has greater charge, greater attraction, stronger forces of attraction, .˙. stronger ionic bonding .˙. more energy required to break the bonds

Question 5

Why do larger ions with greater ionic radius have weaker attractions?

Answer 5

Larger ions that have a greater ionic radius will have a greater attraction to the oppositely charged ion because the attractive forces have to act over a greater distance

Question 6

Why do larger ions with greater ionic radius have weaker attractions?

Answer 6

Larger ions that have a greater ionic radius will have a greater attraction to the oppositely charged ion because the attractive forces have to act over a greater distance

Question 7

Explain why converting a lattice into gaseous ions is an endothermic process:

Answer 7

We have to put energy in, to break the strong ionic bonds between oppositely charged ions

Question 8

What is the trend in lattice energy as you go down the group?

Answer 8

Gets less exothermic as you go down the group
- Lattice enthalpy becomes less exothermic and less negative as the size of the negative ions increases. This indicates weaker attraction between ions and hence weaker ionic bonding

Question 9

What is the trend in lattice energy as you go across the period?

Answer 9

Gets more exothermic across a period
- Charge increases and produces a greater attraction between these positive ions and negative ions. The ionic radius decreases resulting in the ions in the lattice being closer together producing more attraction

Question 10

What happens to the ionic radius as you go down the group?

Answer 10

Increases as the atomic number increases. This is because extra electron shells are added

Question 11

What happens to the ionic radius of a set of isoelectronic?

Answer 11

The ionic radius of a set of isoelectronic ions decreases as the atomic number increases
- As you go through this series of ions the number of electrons stays the same, but the number of protons increases. This means that the electrons are attracted to the nucleus more strongly, pulling them in a little, so the ionic radius decreases

Question 12

Define covalent bond:

Answer 12

When two atoms share one or more pairs of electrons. Strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond

Question 13

What is a pure (non-polar) covalent bond:

Answer 13

Bonding electrons shared equally between two atoms
No charges on atoms

Question 14

What is a polar covalent bond?

Answer 14

Bonding electrons shared unequally between two atoms
Partial charges on atoms
Electrons shared unequally, if electron is pulled to one end, that end will be negative. Electronegativity of oxygen: 3.5, hydrogen: 2.1

Question 15

What is a dative (coordinate) bond:

Answer 15

Shared pair of electrons that came from the same atom
- In AlCl3, Al has 3 electron pairs around it. Hence, Al has an empty orbital

Question 16

What are the positive nuclei in covalent bonds attracted to?

Answer 16

• The positive nuclei are attracted to the area of electron density between the two nuclei. But there’s also a repulsion.
• The two positively charged nuclei repel each other, as do the electrons.

Question 17

What is bond length?

Answer 17

The distance between the two nuclei is the distance where the attractive and repulsive forces balance each other.

Question 18

What is the trend in bond lengths? Why?

Answer 18

• Increases down a group
• Decreases across a period
  This is because the smaller the atomic radius the closer the two atoms can get and the shorter the bond will be. This also tends to make the bond stronger

Question 19

What is the trend in bond lengths? Why?

Answer 19

• Increases down a group
• Decreases across a period
  This is because the smaller the atomic radius the closer the two atoms can get and the shorter the bond will be. This also tends to make the bond stronger

Question 20

Explain the bond lengths in NH3 vs PH3

Answer 20

Bond length will be greater in PH3 than NH3. This is because P is lower down the group, so has more shells of electrons, and therefore a bigger atomic radius. The nuclei of the two atoms are further from the outer shell electrons so therefore the attraction for the shared pair of electrons will be weaker and the bond will be longer.

Question 21

What happens when electron density increases?

Answer 21

The higher the electron density between the nuclei, the stronger the attraction between the atoms, the higher the bond enthalpy and the shorter the bond length.

Question 22

What does molecular shape depend on?

Answer 22

Depends on electron pairs around the central atom
- Shape depends on the number of pairs of electrons in the outer shell of the central atom

Question 23

What is a pi bond?

Answer 23

The two P orbitals overlap sideways (laterally) to produce regions of electro density above and below the axis joining the two nuclear centres -pi bonds

Question 23

What is a pi bond?

Answer 23

The two P orbitals overlap sideways (laterally) to produce regions of electro density above and below the axis joining the two nuclear centres -pi bonds

Question 24

What is the order of the VSEPR shapes?

Answer 24

2 bonded pairs - 180° linear
3 bonded pairs - 120° trigonal planar
2 bonded pairs, one lone pair - 119° bent
4 bonded pairs - 109.5° tetrahedral
3 bonded pairs, one lone pair - 107° trigonal pyramidal
2 bonded pairs, two lone pairs - 104.5° bent
5 bonded pairs - 90/120° trigonal bipyramidal

Question 25

What is the order of strengths of repulsion:

Answer 25

Lone pair/lone pair > bonded pair/lone pair > bonded pair/bonded pair

Question 26

What is the lone pair repulsion?

Answer 26

Each lone pair reduces the bond angle by 2.5 degrees

Question 27

Explain the bond angles in NH3 vs NH4+

Answer 27

• In NH3 all electron pairs repel to achieve maximum separation, but lone pairs repel more than bonding pairs so the shape is pyramidal and the bond angle is 107
• NH3 - three bonding pairs and one lone pair
• In NH4+ all electron pairs are bonding pairs, therefore they all repel equally. Therefore the shape is tetrahedral and has a bond angle of 109.5
• NH4+ - has four bonding pairs (one is dative covalent)

Question 28

What are London Forces?

Answer 28

• Even in molecules with no polar bonds, there are temporary dipoles due to uneven electron distribution due to the constant movement of electrons
• This induces a temporary/instantaneous dipole in a neighbouring molecule, producing a temporary induced dipole-dipole attraction

Question 29

What happens if the molecules get bigger?

Answer 29

The bigger the molecule, the greater the London forces

Question 30

What are permanent dipole-dipole attraction?

Answer 30

Some molecules with polar bonds have an overall dipole (eg/HCl)
- But individual dipoles cancel each other out as the molecule is symmetrical in every direction

Question 31

What are hydrogen bonds?

Answer 31

• Where an H atom is bonded to a very electronegative atom (i.e. F,O,N)
• The polar bond leaves the H nucleus exposed as H only has one electron
• Therefore there is a strong attraction from the lone pair on the N,O or F of one molecule to the exposed H nucleus of another molecule
  Eg/ NH3, H2O.HF

Question 32

What is the order of the strength of intermolecular forces?

Answer 32

Hydrogen bonding > permanent dipole-dipole > London forces

Question 33

Explain the boiling point trend across the hydrides of group 7:

Answer 33

HF: highest boiling point as it has hydrogen bonding which is the strongest IMF
HCl, HBr and HI all have dipole-dipole interactions, which are less strong, so have a lower bpt

From HCl → HI, Mr increases, so stronger London forces

Question 34

What is metallic bonding:

Answer 34

Array of cations in a sea of delocalised electrons

Question 35

Explain why molecular substances have low melting points

Answer 35

They have weak intermolecular forces in between the molecules therefore require less energy to overcome the forces

Question 36

Explain why molecular substances do not conduct electricity

Answer 36

They have no charge (no delocalised electrons/ions) .˙. cannot conduct

Question 37

Why do giant covalent substances have high melting points?

Answer 37

Many and strong covalent bonds, which require a lot of energy to break

Question 38

Why is diamond hard?

Answer 38

Giant covalent structure therefore many covalent bonds

Question 39

Why does graphite conduct electricity?

Answer 39

Layers with delocalised electron, therefore have free moving electrons and can carry current

Question 40

Why is graphite soft and slippery?

Answer 40

Weak forces of attraction between layers therefore can slide

Question 41

Use of fullerenes and carbon nanotubes:

Answer 41

Delivery tube for cancer drugs

Question 42

Why do metals conduct electricity?

Answer 42

Delocalised electrons can move and when a current is applied (eg/circuit) they all move in the same direction

Question 43

Why do metals have high melting points?

Answer 43

There are very strong electrostatic forces of attraction between the oppositely charged particles. Requires a lot of energy to break the forces