Chapter 7 - Periodicity

Question 1

Who was responsible for a more accurate periodic table?

Answer 1

Mendeleev

Question 2

What did Mendeleev do that was different?

Answer 2

He left gaps for undiscovered elements so that elements fir in patterns

Question 3

What are the vertical columns in the periodic table called?

Answer 3

Groups

Question 4

What are the horizontal rows in the periodic table called?

Answer 4

A period

Question 5

How are the elements in the periodic table ordered?

Answer 5

By proton (atomic) number

Question 6

4 Properties which vary across a period

Answer 6

Electron configuration
Ionisation energy
Structure
Melting point

Question 7

How can the periodic table be used to work out electron configuration?

Answer 7

Each period signals the start of a new subshell, so you can read along the period and work out how many electrons are in each sub shell

Question 8

What is the name for the elements in group 1?

Answer 8

Alkali metals

Question 9

What is the name for elements in group 2?

Answer 9

Alkaline earth metals

Question 10

What is the name for elements in groups 3-12?

Answer 10

Transition elements

Question 11

What is the first ionisation energy of an element?

Answer 11

The energy needed to remove one mole of electrons from one mole of gaseous atoms

Question 12

Why is ionisation an endothermic process?

Answer 12

It requires energy input

Question 13

Equation for first ionisation energy

Answer 13

X (g) -> X+ +e-

Question 14

What does a low ionisation energy indicate?

Answer 14

It is easier to form an ion

Question 15

3 factors affecting ionisation energy

Answer 15

Nuclear charge
Atomic radius
Shielding

Question 16

What does a high ionisation energy indicate?

Answer 16

There is strong attraction between the electron and the nucleus so more energy is needed to overcome it and remove the electron

Question 17

How does nuclear charge impact ionisation energy?

Answer 17

The more protons there are, the stronger the attraction between the nucleus and electron will be

Question 18

How does atomic radius impact ionisation energy?

Answer 18

The greater the distance between electron and nucleus, the weaker the force of attraction

Question 19

How does shielding impact ionisation energy?

Answer 19

The more electrons shielding the electron from the nucleus, the weaker the force of attraction

Question 20

Why does ionisation energy decrease down a group?

Answer 20

There is more electron shielding and the atomic radius is greater

Question 21

Why does ionisation energy increase across a period?

Answer 21

There is a greater nuclear charge, so a stronger force of attraction

Question 22

Why does ionisation energy drop between groups two and three?

Answer 22

The outer shell in group 3 is in a p-orbital, rather than an s-orbital. This means it is further from the nucleus and there is additional shielding, so ionisation energy drops

Question 23

Why does ionisation energy drop between groups 5 and 6?

Answer 23

In group 5 elements, the electron is being removed from a singly occupied orbital, but in group 6, the orbital contains two electrons. The repulsion between the two electrons makes it easier to remove

Question 24

What are successive ionisation energies?

Answer 24

Removing additional electrons

Question 25

What can a graph of successive ionisation energies tell us?

Answer 25

When different shells are having electrons removed and how many electrons are in each sub shell

Question 26

How are metals held together?

Answer 26

Metallic bonding

Question 27

What is the name for the structure of most metals?

Answer 27

Giant metallic lattice structure

Question 28

What is metallic bonding?

Answer 28

Delocalised electrons are electrostatically attracted to the positively charged nuclei

Question 29

What happens when the number of delocalised electrons increases?

Answer 29

Melting point increases

Question 30

Why are metals ductile and malleable?

Answer 30

There are no bonds holding the layers together so they can slide past each other

Question 31

Why are metals good thermal and electrical conductors?

Answer 31

The delocalised electrons can carry kinetic energy or electrical charge

Question 32

Why are metals insoluble?

Answer 32

The metallic bonds are very strong

Question 33

What are giant covalent lattices?

Answer 33

Huge networks of covalently bonded atoms

Question 34

What are different forms of the same element called?

Answer 34

Allotropes

Question 35

Properties of diamond

Answer 35

Very high melting point Very hard Insoluble Can't conduct electricity

Question 36

Why does diamond have such specific properties?

Answer 36

All four of its outer carbons are bonded to another carbon

Question 37

Why is graphite used as a lubricant?

Answer 37

Weak forces between layers

Question 38

Why can graphite and graphene conduct electricity?

Answer 38

Delocalised electrons to carry charge

Question 39

Why is graphite less dense than diamond?

Answer 39

The layers are far apart

Question 40

Why does graphite have a high melting point?

Answer 40

It has strong covlent bonds

Question 41

Why is graphite insoluble?

Answer 41

The covalent bonds in the sheets are very strong

Question 42

Why is graphene so strong?

Answer 42

Delocalised electrons strengthen covalent bonds

Question 43

What is a simple molecular structure?

Answer 43

A structure containing only a few atoms

Question 44

Why do simple molecules have low melting points?

Answer 44

Weak induced dipole-dipole forces to overcome

Question 45

Why do noble gases have low melting points?

Answer 45

They are monatomic, meaning weak dipole-dipole forces

Question 46

Why does melting point change across a period?

Answer 46

The type of bond formed changes

Question 47

What will happen to the melting point of metals across a period?

Answer 47

It will increase - the number of delocalised electrons increases so the bonds are harder to overcome

Question 48

Why so carbon and silicone have high melting points?

Answer 48

They form giant covalent lattices, which require lots of energy to overcome

Question 49

Why are giant covalent lattices insoluble?

Answer 49

The covalent bonds holding the atoms in the lattice are far too strong to be broken by solvents

Question 50

Across period 2 and 3, what happens to melting points between groups 1 and 14?

Answer 50

There is an increase

Question 51

What happens to melting points of elements in groups 15-18 in periods 2 and 3?

Answer 51

They decrease

Question 52

In periods 2 and 3, why does melting point decrease between group 14 and 15?

Answer 52

The elements in group 15 onwards form simple molecular structures, not giant, which require less energy to overcome

Question 53

How does the bonding in a simple molecular lattice differ to that in a giant covalent lattice?

Answer 53

A simple molecular lattice only has London forces between atoms, whereas a giant covalent lattice has covalent bonds

Question 54

Why does the melting point of metals increase with the number of electrons?

Answer 54

The number of delocalised electrons and the charge of the ion increase, which makes the metallic bonds stronger and harder to break