Topic 2

Question 1

What is the definition of ionisation energy?

Answer 1

The measure of energy required to remove 1 electrons from 1 mole of an atom at a gaseous state.

Question 2

What effects ionisation energy?

Answer 2

-nuclear charge: the number of protons in nucleus’s:
As the number of protons increase more ionisation is needed, this is because it’s harder to pull the electron away due to the attraction of the electrons to the proton.

-distance of electron to the nucleus:
As the distance of the electron from the nucleus increase, less ionisation energy is needed.

Shielding: as the number of shells increases, there will be more repulsion between the electrons, which makes it easier to remove the electrons. So less ionisation energy is required.

Question 3

What happens to ionisation energy down the group:

Answer 3

-ionisation energy decreases down the group as the electron is further away from the nucleus.

Question 4

What happens to ionisation energy across the period:

Answer 4

• ionisation energy increases, because the number of protons increase, and the nuclear charge increase.
  The distance between the electron and nucleus decrease, which increase the ionisation energy.

Question 5

What is the quantum shell?

Answer 5

The whole shell

Question 6

What is the quantum shell?

Answer 6

The whole shell

Question 7

What is sub shell?

Answer 7

The s, p, d sub shells

Question 8

What is an orbital?

Answer 8

The box which contains 2 electrons:
1s2—> has 1 orbital
2p3—> has 3 orbitals

Question 9

Why do we use log on the graph rather than the ionisation energy?

Answer 9

Because ionisation energy is way too big to plot on the graph

Question 10

What does isoelectronic mean?

Answer 10

Elements which have the same electronic configuration

Question 11

What are the stages of the spectrometer?

Answer 11

1. Vaporisation, happens before entering the spectrometer
2. Ionisation
3. Acceleration
4. Deflection
5. Detection

Question 12

Why does the element has to be in a gaseous state?

Answer 12

If element is not in a gaseous state, it wont be accelerated or detected

Question 13

What is the relative molecular mass?

Answer 13

The relative molecular mass is the average mass of a molecule compared to 1/12 of a carbon atom.

Question 14

What can the spectrometer can tell us?

Answer 14

It can tell us the relative isotopic masses and abundance of different elements

Question 15

What does m/z mean?

Answer 15

It means mass/charge

Question 16

How can you predict the mass spectra for diatomic molecules?
Diatomic: molecule containing two atoms.

Answer 16

1.Draw a table showing all different molecules
2.multiply the abundance for each isotope to find the relative abundance for each one.
3.find the 2 isotopes which have the same abundance and add them up.
4.divide all relative abundance with the smallest relative abundance.

Question 17

How to find a molecular mass from spectrometry:

Answer 17

o find the molecular mass of a compound, you look at the molecular ion peak.
o the mass/charge value of the molecular ion peak is the molecular mass value

Question 18

What happens to the atomic radius across a period?

Answer 18

Due to the number of protons increasing, the nuclear charge increases.
This means that electrons are pulled closer to the nucleus, making the atomic radius smaller.

Question 19

How does the drop from groups 2 and 3 shows subshell structure?

Answer 19

Group 2 elements require more ionisation energy than group 3.
This is because elements in group 3 have a 3p orbital, rather than 3s orbital like the elements in group 2.
3p orbital has more energy than 3s, so less ionisation energy is required

Question 20

Explain the drop between group 5 and 6?

Answer 20

If the elements are in the same period they have the same amount of shielding.
Group 6 elements have paired electrons in the p orbital.
Group 5 laments have singly occupied electrons.
So group 6 orbital has more repulsion, so less ionisation energy is needed to remove the electron.

Question 21

In terms of structure and bonding explain why silicon has a high melting point:

Answer 21

Silicon has a giant lattice structure.
Silicon has a giant macro molecular structure with very strong covalent bond which requires a lot of energy to be broken.

Question 22

Why is the melting point of sulfur more than that of phosphorus?

Answer 22

Sulfur (s8) has more electrons than phosphorus (p4), so sulfur has stronger forces of attraction between the molecule.
And
Phosphures has weak intermolecular forces between the molecule so less energy is needed to break the bond.

Question 23

What is ionic bonding?

Answer 23

An ionic bond is the electrostatic attraction between oppositely charged ions.

Question 24

What two things that affect the strength of an ionic bond?

Answer 24

The ionic charge:
-The greater the charge of an ion the greater the strength of the ion
E.g Ca+2 is stronger than Na+1, which means that Ca+2 has a greater melting and boiling point.
Also O-2 is stronger than F-1

The ionic Radii:
- the smaller the ionic radius the stronger the ionic bond.
- electrostatic attraction between the molecule get weaker with distance.

Question 25

Does the ionic radius increase or decrease down a group?

Answer 25

The radius of the ion increases because there are more number of shells and they have the same charge.

Question 26

What are isoelectronic ions?

Answer 26

Ions which have the same electronic configuration like: N-2, O-2, F-1, Na+1, Mg+2, Al+3 They all have 10 electrons So their electronic configuration is 2,8

Question 27

What type of structure does sodium chloride have?

Answer 27

Sodium chloride has a giant ionic lattice structure. Sodium chloride would have a high melting point, because a lot of energy is required to overcome the strong electrostatic attraction between the opposite charged ions.

Question 28

Does sodium bromide have a higher or lower melting point compared to sodium chloride

Answer 28

Sodium bromide would have lower melting point. Bromide ion has an extra quantum shell, 1 more shell than chloride. So the radius of sodium bromide is larger than the sodium chloride. As the distance of the radius of the ionic bond increases, the strength of the ionic bond decreases. Because ions in the sodium bromide aren’t as closely packed as the ions in the sodium chloride. So less energy is required to break the bond in odium bromide.

Question 29

Why does sodium calcium oxide does not conduct electricity, but molten calcium oxide does?

Answer 29

Because sodium calcium oxide is a solid, and in a solid ions at eheld tightly together due to the ionic bond. But when molten ions are mobile/free to move around and carry electricity

Question 30

What is a covalent bond?

Answer 30

When two non-metals share electrons to have a full outer shell. It’s the electrostatic attraction between the two positive nuclei and the shared electrons in the bond.

Question 31

How are giant covalent bonds different to simple covalent bond?

Answer 31

The giant covalent bonds have stronger electrostatic attractions holding the atoms together, the have much stronger electrostatic attraction than simple covalent bonds.

Question 32

How can graphite conduct electricity ?

Answer 32

Graphite is a covalent bond of carbon atoms. It has free electrons which can move and conduct electricity

Question 33

What is graphene?

Answer 33

Graphene is one layer of graphite which is a sheet of carbon atoms joined together in hexagons. It’s useful because it conducts electricity due to the lose electron. It is also incredibly strong, transparent and really light.

Question 34

Describe the metallic giant metallic lattice structure :

Answer 34

In a metallic lattice electrons in the outer most shell are free to move. This leaves the a positive metal ion—-> Na+1, Mg+2 The positive metal ion is electrostatically attracted to the free electrons in the outer most shell. They form a metallic lattice ion, where all positive ions are packed together.

Question 35

Why do a giant metallic lattice have high melting point:

Answer 35

They have a high melting point due to the strong metallic ionic bond because the more electrons there are the stronger the bond. Mg+2 has 2 delocalised electrons per atom Where’s Na+1 has 1 free electron per atom So Mg+2 has a higher meting point.

Question 36

Explain and give the characteristics of metals with giant metallic lattice structure :

Answer 36

1. Malleable: Because there are no bonds holding specific ions together, the layer of positive ion are separated by a layer of electrons. 2. Good heat conductor: The delocalised electrons can pass kinetic energy to teach other. 3. Good conductor of electricity: The delocalised electrons can move and conduct electricity

Question 37

Explain why graphite is able to conduct electricity:

Answer 37

Graphite consists of carbon sheets, where each carbon atom is bonded with 3 carbon atom. So each atom has 1 free electron which are free to conduct electricity.