Unit 1

Question 1

What was the name of the scientist who developed the modern version of the periodic table?

Answer 1

Dimitri Mendeleev

Question 2

What was similar about the elements in each group of the periodic table?

Answer 2

He placed elements with similar chemical properties in the same group.

Question 3

What key did Mendeleev use to arrange the elements?

How did this differ from previous attempts by other chemists?

Answer 3

Previously elements had only been ordered using atomic mass, Mendeleev combined this with chemical properties.

Question 4

What did Mendeleev do at points in the table where his system did not work?

Answer 4

He left gaps and made predictions about elements still to be discovered

Question 5

What group from the modern periodic table was missing from Mendeleev version?

Why is this likely to be?

Answer 5

Nobel gases

Due to them not reacting and can’t be seen or smelled so had not been discovered

Question 6

What happens to the covalent radius as you go across a period?

Answer 6

It decreases.

The increased nuclear charge means increased attraction between the nucleus and the outer electrons making the atoms smaller

Question 7

What happens to the covalent radius of an element as you go down a group?

Answer 7

Increases

There are more electron shells so the outer electrons are further away. The nuclear charge is screened by the inner shells so the attraction between the nucleus and the outer electrons are weaker.

Question 8

What is first ionisation energy?

Answer 8

The energy required to remove 1 mole of electrons from 1 mole of free atoms in the gas state.

Question 9

What happens to the ionisation energy as you go across a period?

Answer 9

The ionisation energy increases because as you go across the period there are more protons and the attraction between the nucleus and the outer electrons becomes greater, so it become harder to remove moles of electrons.

Question 10

What happens to the ionisation energy of an element as you go down a group?

Answer 10

The ionisation energy going down a group decreases because there are more outer shells so the attraction is weaker and screened so it is easier to remove moles of electrons.

Question 11

What is the definition of second ionisation energy?

Answer 11

The energy required to remove mole of electrons from a mole of positively charged ions in the gas state.

Question 12

Formula for first ionisation?

Answer 12

E(g) —> E+(g) + e-

Question 13

What is the formula for second ionisation energy?

Answer 13

E+(g) —> E2+(g) + e-

Question 14

Why does successive ionisation increase?

Answer 14

Successive ionisation energies increase because the nuclear charge remains the same but as electrons are removed, the attraction of the nucleus on the remaining electrons increases making the, harder to remove and therefore increasing the ionisation energies.

Question 15

What is the formula for the third/successive ionisation energy?

Answer 15

E2+(g) —> E3+(g) + e-

Question 16

Why are there no electronegativity values recorded for noble gases?

Answer 16

Nobel gases do not form covalent bonds, they were not really discovered because they are very un-reactive

Question 17

What is electronegativity?

Answer 17

A measure of attraction an atom involved in a covalent bond has for the shared electrons of that bond

Question 18

What happens to the electronegativity as you go across a period?

Answer 18

Increase because there are more protons and the attraction between the protons and the electrons are stronger.

Question 19

What happens to the electronegativity of an element as you go down a group?

Answer 19

Decrease because there are more electron shells which also screen the attractions so the electronegativity is lower.

Question 20

What structural feature of metals allow them to conduct electricity?

Answer 20

The electrons are delocalised allowing them to move through the structure

Question 21

Explain the change in terms of delocalised electrons down a group?

Answer 21

The metallic bond gets weaker because there are more shells so the nuclear charge is screened by the inner shells. Making the attraction weaker.

Question 22

Explain the change in terms of delocalised ions across a period?

Answer 22

The metallic bonds get stronger because there is higher nuclear charge giving a stronger attraction and there are more delocalised electrons meaning there are more metallic bonds.

Question 23

How do London dispersion forces arise between two none bonded atoms?

Answer 23

Electrons clouds wobble causing partial positive and negative charges on each atom.

LDF is the weak attraction between the slight positive on one atom and the slight negative on another

Question 24

What happens to the boiling point of the noble gases as you go down the group?

What does this tell us about the strength of the London dispersion forces between the atoms?

Answer 24

The boiling increases, this tells us the LDF are getting stronger.

And that is is due to more wobbling if the atoms

Question 25

How does the size of an atom affect the strength of the London dispersion forces?

Answer 25

Larger atoms wobble more giving stronger dipoles and therefore stronger LDF between atoms

Question 26

What are uses of fullerenes?

Answer 26

Drug delivery systems in the body

In lubricant and as a catalyst
Abrasive powder
Medicine, cosmetics
Superconductor, solar cell, organic materials

Question 27

What are uses of carbon nanotubes?

Answer 27

Antennas
Used in brushes for commercial electric motors
Cleaning polluted waters
Boosting solar energy

Question 28

How does the use of a diamond relate to its structure and properties?

Answer 28

Jewellery- long lasting

Drill bits- hard and strong

Question 29

What are the three crystalline forms of carbon?

Answer 29

Diamond
Graphite
Fullerenes (recently discovered)

Question 30

How does the use of graphite relate to its structure and properties?

Answer 30

Electrodes- conduct electricity

Pencil lead- a layer can be easily left behind on the paper