Periodicity

Question 1

How are elements arranged in the Periodic Table?

Answer 1

In order of increasing atomic number - the number of protons in the nucleus.

Question 2

Why do chemists find the Periodic Table useful?

Answer 2

They can make predictions about the properties of an element based on its position in the table.

Question 3

Why are elements in the same group similar?

Answer 3

They share the same number of electrons in their outer shell.

Question 4

Definition of periodicity?

Answer 4

A repeating pattern.

Question 5

How are elements in periods organised?

Answer 5

In order of increasing atomic number, showing the increasing number of outer electrons. They also move from metallic to non-metallic characteristics.

Question 6

Name three elements with metallic bonding and structure.

Answer 6

Li, Be, Na, Mg, Al, K, Ca

Question 7

Name three elements with covalent molecular structure.

Answer 7

H2, N2, O2, F2, Cl2, P4, S8 and fullerenes (eg C60)

Question 8

Name three elements with covalent network structure.

Answer 8

B, C (diamond, graphite), Si

Question 9

Name three elements with monatomic structure.

Answer 9

Noble gases - He, Ne, Ar

Question 10

What group are the noble gases?

Answer 10

Group 0/8

Question 11

What does monatomic mean?

Answer 11

The molecules exist as individual atoms.

Question 12

Why are noble gases unreactive?

Answer 12

They have a full outer energy level.

Question 13

What are London dispersion forces (LDFs)?

Answer 13

They are very weak forces of attraction caused by the continual movement of electrons in an atom which causes a temporary uneven distribution of charge known as a temporary dipole. This induces a temporary dipole in a neighbouring atom and forming a force of attraction between them.

Question 14

What is a temporary dipole?

Answer 14

Caused by the continual movement of electrons in an atom which causes a temporary uneven distribution of charge at opposite sides of the atom -this makes one side of an atom slightly negative 𝛿- and one side slightly positive 𝛿+.

Question 15

How do LDFs change going down a group?

Answer 15

Although very weak, the more electrons, the bigger the LDFs, increasing the MPs and BPs as more energy is needed to separate them.

Question 16

What is the difference between intermolecular and intramolecular forces?

Answer 16

INTER = between separate molecules
INTRA = between atoms inside a molecule

Question 17

Name the diatomic molecules.

Answer 17

H2, O2, F2, Br2, I2, N2, Cl2

Question 18

What are diatomic molecules?

Answer 18

They exist as two atoms covalently bonded.

Question 19

What force of attraction holds noble gases together?

Answer 19

LDFs

Question 20

Why are most diatomic molecules gases at room temperature?

Answer 20

The forces between them (LDFs) are so small that there is enough energy at room temperature to overcome the weak force of attraction and separate the molecules from each other.

Question 21

Name the halogens and their group number.

Answer 21

Group 7 - F2, Cl2, Br2, I2, At

Question 22

How do phosphorus and sulfur exist?

Answer 22

P4 and S8. They are solids at room temperature with low melting and boiling points.

Question 23

What bonds hold together P4 and S8?

Answer 23

The atoms which make up the molecules are held together by covalent bonds and the molecules are held together by LDFs.

Question 24

How can carbon exist?

Answer 24

As covalent networks - graphite and diamond, and as covalent molecules - fullerene (C60).

Question 25

Describe fullerene's shape.

Answer 25

A series of carbons arranged in hexagons and pentagons are joined to form a ball shape. It can also exit as tube shapes called nanotubes.

Question 26

How is fullerene held together?

Answer 26

By LDFs but because of its size it requires a lot more energy to separate.

Question 27

What's one thing that a high melting and boiling point indicates?

Answer 27

There are other forces holding the molecules together, apart from just LDFs.

Question 28

Give three points about diamond.

Answer 28

- Each carbon atom is covalently bonded to four other carbons in a tetrahedral shape. - All the outer electrons of each carbon are used to make single covalent bonds with neighbouring atoms, hence the giant covalent network. - No individual molecules or unbonded (delocalised) electrons, so diamond does not conduct electricity.

Question 29

Give three points on graphite.

Answer 29

- Any one carbon is only bonded to three other carbons, making the fourth electron in its outer shell delocalised, and therefore can conduct electricity. - Atoms form hexagonal plates held together by weak LDFs. - The plates are able to slide over each other, explaining why powdered graphite can be used as a lubricant.

Question 30

Why do covalent networks have such high melting and boiling points?

Answer 30

The strong covalent bonds need to be broken requiring a lot of energy.

Question 31

What's unusual about graphite and diamond?

Answer 31

When heated to high temperatures, rather than melt they sublime; meaning they go straight from solid to gas.

Question 32

Describe silicon.

Answer 32

Silicon is similar to diamond in that it forms covalent bonds with four other silicon atoms and makes a covalent network. It has a very high melting point, is extremely hard, and is a poor conductor.

Question 33

Describe boron.

Answer 33

Boron is similar to carbon as it forms a stable covalent network. It is a very hard, black material with a very high melting point.

Question 34

What is the covalent radius?

Answer 34

The measure of the size of an atom. Taken by half the distance between the nuclei of two atoms joined by a single covalent bond.

Question 35

Explain what happens to the covalent radius going down a group.

Answer 35

INCREASES (atoms get bigger) - due to the increasing atomic number, there are more protons in the nucleus and more electrons in each energy level. As the energy levels increase there is shielding of the full pull of the positive nucleus allowing the outermost electrons to move further away.

Question 36

Explain what happens to the covalent radius going across a period.

Answer 36

DECREASES (atoms get smaller) - the atomic number increases by one each time, meaning the atom gains a proton while also gaining an electron in the same energy shell. Thus there is no shielding effect from the inner levels and so the outer electrons are more strongly attracted to the nucleus pulling them closer together.

Question 37

What is ionisation energy?

Answer 37

The energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Question 38

What unit is ionisation energy measured in?

Answer 38

kJ mol-1

Question 39

What does first and second ionisation energy mean?

Answer 39

The first ionisation energy is the energy required to remove one mole. The second is the energy required to remove two moles.

Question 40

Why are second ionisation energy values always higher than the first?

Answer 40

Although electrons are being removed, the number of protons in the nucleus stays the same increasing the pull on the remaining electrons.

Question 41

Describe what happens to ionisation energy as you go down a group.

Answer 41

DECREASES (less energy required) - because the covalent radius has increased and the outer electrons are shielded from the full pull of the nucleus, it is easier to remove an outer electron.

Question 42

Describe what happens to ionisation energy as you go across a period.

Answer 42

INCREASES (more energy required) - although electrons are being added, it is to the same energy level which increases the nuclear charge, holding the electrons closer together. This makes it harder to remove an outer electron.

Question 43

What is electronegativity?

Answer 43

The attraction an atom, involved in a bond, has for the electrons of the bond.

Question 44

What elements are the most and least electronegative?

Answer 44

``` MOST = Fluorine (4.0) LEAST = Caesium (0.8) ```

Question 45

What does electronegativity help indicate?

Answer 45

The difference in electronegativity within a bond helps indicate the type of bonding that primarily exists between them.

Question 46

Describe what happens to electronegativity going down a group.

Answer 46

DECREASES (less attraction) - the covalent radius increases and outer electrons are further from the nucleus and are also shielded from the full pull of the nucleus, combining to result in the nucleus having less of an attraction for the bonding electrons.

Question 47

Describe electronegativity going across a period.

Answer 47

INCREASES (more attraction) - the covalent radius decreases and the number of outer levels stay the same, so the nucleus's attraction for bonding electrons increases.