Electrons, Bonding and Structure

Question 1

Why do ionic compounds have relatively high melting and boiling points?

Answer 1

Ionic compounds have high melting and boiling points as a large amount of energy is required to overcome the strong electrostatic forces of attraction between oppositely charged ions

Question 2

Why does MgO have a higher boiling point than NaCl?

Answer 2

As the charges on the Mg2+ and O2- are greater than that of the Na+ and the Cl- ions, the greater the charge on the ions, the stronger the electrostatic forces of attraction between oppositely charged ions

Question 3

Can ionic compounds conduct electricity when solid?

Answer 3

No, ions are fixed in place when the compound is solid meaning it can’t conduct electricity

Question 4

Can ionic compounds conduct electricity when aqueous?

Answer 4

Yes, ions are free to move throughout the solution as the solid lattice breaks down, the ionic compound is now a conductor of electricity

Question 5

Explain why ionic compounds such as NaCl are soluble in water

Answer 5

• Water molecules attract the Na+ and Cl- ions
• Na+ attracts water molecules, the delta negative oxygen is attracted to the Na+
• Cl- attracts the delta positive hydrogen in the water molecule
• This disrupts the lattice and ions are pulled out of the lattice, this causes the ionic compound to dissolve

Question 6

How does an ammonium ion form?

Answer 6

NH3 has a lone pair, this lone pair is used to form a dative covalent bond with a H+ ion, this forms an NH4+ ion

Question 7

How is an hydronium ion formed?

Answer 7

H2O has two lone pairs, one of these lone pairs is used to form a dative covalent bond with an H+ ion, this forms an H3O+ ion

Question 8

What is the structure of BF3?

Answer 8

The central boron atom is covalently bonded to 3 fluorine atoms, this means that the boron atom has 6 electrons in its outer shell, it does this so that all of its outer electrons are paired

Question 9

What is the structure of SF6?

Answer 9

The central sulfur atom forms 6 covalent bonds with 6 fluorine atoms, this means that sulfur has 12 electrons in its outer shell, it has expanded it’s octet

Question 10

How are simple molecules held together?

Answer 10

Atoms between molecules are held together by strong covalent bonds, different molecules are held together by weak intermolecular forces such as dipole-dipole interactions and London forces

Question 11

Can simple molecular structures conduct electricity?

Answer 11

No, there are no charged particles free to move

Question 12

Are simple molecular structures soluble in polar solvents?

Answer 12

Weak London forces are able to form between the covalent molecules and the polar solvents, this breaks down the molecular lattice and the substance dissolves

Question 13

Why do simple molecular structures have relatively low melting and boiling points?

Answer 13

Not much energy is required to overcome the weak intermolecular forces holding the molecules together, this means they have relatively low melting and boiling points

Question 14

Give some examples of giant covalent structures

Answer 14

Diamond, graphite and SiO2

Question 15

Why do giant covalent structures have high melting and boiling points?

Answer 15

As the covalent bonds holding the lattice together are very strong, this means that a lot of energy is required to overcome these forces and therefore giant covalent lattices have high melting and boiling points

Question 16

Can giant covalent structures conduct electricity?

Answer 16

• No, they have no charged particles that are free to move
• Graphite is the only giant covalent structure that can conduct electricity as it has delocalised electrons between layers that are able to move freely when a voltage is applied

Question 17

Are giant covalent structures soluble in any solvents?

Answer 17

No, the covalent bonds are too strong to be broken by either polar or non-polar solvents

Question 18

What is the shape and bond angle of a molecule with a central atom with 2 bonded regions and no lone pairs?

Answer 18

Linear, 180º

Question 19

What is the shape and bond angle of a molecule with 3 bonded regions and no lone pairs?

Answer 19

Trigonal planar, 120º

Question 20

What is the shape and bond angle of a molecule with 4 bonded regions and no lone pairs?

Answer 20

Tetrahedral, 109.5º

Question 21

What is the shape and bond angle of a molecule with 5 bonded regions and no lone pairs?

Answer 21

Trigonal bipyramid, 90º and 120º

Question 22

What is the shape and bond angle of a molecule with 6 bonded regions and no lone pairs?

Answer 22

Octahedral, 90º

Question 23

How much does each lone pair reduce the bond angle by?

Answer 23

2.5º, lone pairs repel more than bonded pairs as they’re slightly more electron dense meaning that bonded pairs move closer together to minimise repulsion from the lone pair, this reduces the bond angle

Question 24

What is the shape and bond angle of a molecule with 3 bonded regions and 1 lone pair?

Answer 24

Pyramidal, 107º, it has a lone pair in place of a bonded pair meaning the bond angle is reduced from 109.5º to 107º

Question 25

What is the shape and bond angle of a molecule with 2 bonded regions and 2 lone pairs?

Answer 25

Non-linear, 104.5º, it has 2 lone pairs in place of 2 bonded pairs meaning the bond angle is reduced by 2x2.5º (5º), this means the bond angle is reduced from 109.5º to 104.5º

Question 26

What is the shape and bond angle of an ammonium ion and how is the charge shown?

Answer 26

Tetrahedral as there are 4 bonded pairs and no lone pairs around the central atom, 109.5º, square brackets around the ion with a superscript +

Question 27

When is a bond described as polar?

Answer 27

• When there is a covalent bond between two different bonding atoms, then it is likely that the electrons will lie closer to one of the bonding atoms
• This is because one of the bonding atoms is more electronegative than the other atoms meaning it attracts the bonding pair of electrons more than the other atom
• The bonding electrons lie slightly closer to the atom that is more electronegative
• This creates a charge difference between the two atoms, a small positive charge on the less electronegative atom is shown by a delta + and the small negative charge on the more electronegative atom is shown by a delta -
• for example H𝛿+—– Cl𝛿-

Question 28

Is methane a polar molecule?

Answer 28

Methane has many C-Cl bonds which are polar, however, the molecule is symmetrical so the dipoles act in different directions and cancel each other out

Question 29

Why is CO2 non-polar but water is highly polar?

Answer 29

• There are two polar C=O bonds in CO2 but the molecule is symmetrical meaning the dipoles are cancelled out
• In water, there are 2 polar H-O bonds and the molecule is non linear meaning that there is an overall dipole

Question 30

If the electronegativity difference between two bonding atoms is very large, what may happen?

Answer 30

An ionic bond could be formed

Question 31

What is a permanent dipole-induced dipole interaction?

Answer 31

If a polar molecule moves close to a non-polar molecule, it is able to cause electrons in the the shells of the nearby non-polar molecule to shift slightly towards the 𝛿+ side of the polar molecule, this causes the non-polar molecule to become slightly polar and then an attraction occurs between the two molecules

Question 32

What is a permanent dipole-permanent dipole interaction?

Answer 32

Molecules with permanent dipoles are attracted to other molecules with permanent dipoles, if there are two H-Cl molecules close to each other, there will be an attraction between the H𝛿+ and the Cl𝛿-:
- H𝛿+——Cl𝛿- - - - - - - - H𝛿+——Cl𝛿-

Question 33

What are London forces?

Answer 33

• London forces are attractions between molecules caused by the constant random movement of electrons in atom’s shells, this movement unbalances the distribution of charge within the electron shells
• At any moment there will be an instantaneous dipole across the molecule
• The instantaneous dipole induces a dipole in neighbouring molecules, which in turn induce dipoles in neighbouring molecules
• Small induced dipoles attract one another, this causes weak intermolecular forces between molecules known as London forces

Question 34

What is the trend in boiling point going down a group?

Answer 34

The number of electrons as you go down the group., as the number of electrons increases, so does the strength of the London forces. This means that more energy is required to overcome these intermolecular forces and therefore the boiling point increases as the number of electrons increases and as you go down the group

Question 35

Which molecules can hydrogen bonding occur between?

Answer 35

H and NOF

Question 36

Why is ice less dense than water?

Answer 36

• When ice forms, the water molecules arrange themselves into an orderly pattern and hydrogen bonds form between the molecules
• Ice has an open lattice with hydrogen bonds holding the water molecules apart
• When ice melts the rigid hydrogen bonds collapse allowing the H2O molecules to move closer together, hence ice is less dense than water and therefore ice floats on water
• There are still hydrogen bonds in water but they don’t occur as often as molecules move past each other and hence overcoming these forces

Question 37

Why does water have a higher than expected melting and boiling point?

Answer 37

• Hydrogen bonds are much stronger than the other intermolecular forces
• The extra strength of these forces has to be overcome to melt or boil H2O, this results in H2O having a higher melting or boiling point than if hydrogen bonds weren’t present

Question 38

Why does water have a relatively high surface tension and high viscosity?

Answer 38

Due to the hydrogen bonds that between water molecules