Periodicity

Question 1

What are metalloids

Answer 1

• Elements such as silicon that touch the ‘staircase line’ on the periodic table that separates the metals from non-metals.
• Metalloids have a combination of metallic and non-metallic properties, for example silicon is a non-metal but looks shiny and conducts electricity. (Just not as well as actual metals)

Question 2

What do the s,p,d and f blocks on the periodic table tell us

Answer 2

Which type of orbital (s,p,d or f) the elements outer electron is situated in

Question 3

What is the issue with labelling the transition metals and d block as exactly the same

Answer 3

• There are some exceptions
• Scandium and zinc are not transition metals because they do not form any compounds in which they have partly filled d-orbitals, but they are located within the d-block on the periodic table.

Question 4

What is the key characteristic of a transition metal

Answer 4

That they form compounds with a partly filled d-orbital

Question 5

Describe the patterns in reactivity of different groups on the periodic table

Answer 5

• In the S block elements (metals) get more reactive as you go down a group
• to the right (non-metals) tend to get more reactive going up a group
• Transition metals are pretty unreactive

Question 6

What are the lanthanides

Answer 6

• A group of metals not often encountered.
• They all tend to form 3+ ions in their compounds and have broadly similiar reactivity

Question 7

What are the actinides

Answer 7

• Radioactive metals
• Only thorium uranium occur naturally in the earths crust in anything less than trace quantities.

Question 8

Where is helium placed in the periodic table and why is this unusual

Answer 8

• Helium is placed above the noble gases because of its properties.
• it is unusual as it is not a p-block element: its electron configuration is 1s2

Question 9

Describe how the electron arrangements of the elements across period three explains how trends in properties occur

Answer 9

1) The elements in group 1,2 and 3 are metals. They have giant structures. They lose their outer electrons to form ionic compounds.
2) Silicon is group 4 (so it has four electrons in its outer shell) and forms 4 covalent bonds. The element has some metallic properties and is classed as a semi-metal.
3) The elements in groups 5,6 and 7 are non-metals. They either accept electrons and form ionic compounds or share electrons and form covalent compounds.
4) Argon is a group 0 noble gas- it has a full outer shell and is unreactive

Question 10

Why do the elements on the left of period three have higher melting and boiling points than those on the right

Answer 10

The elements on the left have higher melting and boiling points than the elements on the right because they form giant structures whereas the elements on the left form molecular or atomic structures (the one exception to this is sodium)

Question 11

Describe the shape of the graph which shows the melting and boiling points of the elements across period three

Answer 11

• The melting and boiling points of the metals increase from sodium to aluminium.
• Silicon also has high melting and boiling points due to its giant structure
• The melting and boiling point then decreases dramatically from silicon to phosphorus
• Then increases from phosphorus to sulphur
• They then decrease to chlorine and decrease again to argon.

Question 12

Why do the melting of boiling points of the metals in period three increase from sodium to aluminium

Answer 12

• The strength of the metallic bonding increases
• As you go from left to right the charge on the ion increases so more electrons join the delocalised electron ‘sea’
• So the electrostatic attraction that holds the metal together becomes stronger.

Question 13

Explain why silicon has a high melting point despite not being a metal

Answer 13

It has a giant structure so there are many strong covalent bonds which must be broken

Question 14

Explain the trend in the melting and boiling points across period three of phosphorus, sulphur and chlorine

Answer 14

• The melting points of these non-metals with molecular structures depend on the sizes of the Van Der Waals forces between the molecules
• This in turn depends on the number of electrons in the molecule and how closely the molecules can pack together.
• As a result the melting points are ordered S8> P4 > Cl2

Question 15

Why cant you measure the radius of an isolated atom easily

Answer 15

Because there is no clear point at which the electron cloud density around it drops to zero.

Question 16

How do we measure atomic radii

Answer 16

We halve the distance between the centres of a pair of atoms.

Question 17

Why can the atomic radius of an element differ

Answer 17

It is a general term and depends on the type of bond that is forming- covalent, ionic, metallic, Van Der Waals etc.

Question 18

What do we most commonly use as a measure of the size of an atom

Answer 18

The covalent radius

Question 19

What does the graph that plots covalent radius against atomic number show us

Answer 19

• Atomic radius is a period property because it decreases across each period and there is a jump when starting the next period.
• Atoms get larger down any group

Question 20

Explain why the radii of atoms decrease across a period

Answer 20

• As you move across a period, there are more protons and neutrons in the nucleus and so the nuclear charge increases.
• this increased charge pulls electrons closer to the nucleus.
• there is the same level of shielding due to the outer electrons being in the same main level.
• therefore the size of the atoms decreases

Question 21

Explain the radii of atoms increase down a group

Answer 21

Going down a group on the periodic table, the atoms of each element have one extra complete main level of electrons compared with the one before so the atomic radius increases.

Question 22

Define first ionisation energy

Answer 22

First ionisation energy is the energy needed to remove one electron from every atom in one mole of gaseous atoms to form one mole of singly positively charged ions.

Question 23

What is the general symbol equation for first ionisation energy

Answer 23

X (g) —> X+(g) + E-

Question 24

What is the general pattern in first ionisation energies as you go across a period

Answer 24

The first ionisation energy generally increases as you go across a period

Question 25

What is the pattern in first ionisation energies as you go down a group

Answer 25

The first ionisation energies decrease as you go down any group

Question 26

Explain why, generally, the first ionisation energies increase across a period

Answer 26

• As you go across a period from left to right, the number of protons in the nucleus increases but the electrons enter the same main level.
• The increased nuclear charge means it gets increasingly difficult to remove an electron due to a stronger force of attraction from the nucleus.

Question 27

Explain why first ionisation energies decrease going down a group

Answer 27

• The number of filled inner levels increases going down the group.
• This results in an increase in shielding so the outer electron experiences a weaker force of attraction from the nucleus and is therefore easier to remove.

Question 28

Why is there a drop in first ionisation energy from one period to the next.

Answer 28

• Moving from neon in period 0 to Sodium in period 1 there is a sharp drop in first ionisation energy.
• This is because at sodium a new main level starts and so there is an increase in atomic radius, the outer electron is further from the nucleus, less strongly attracted and easier to remove.

Question 29

Describe the shape of the graph which shows the first ionisation energies across period three

Answer 29

• The first ionisation energies increase as you go across the period.
• there are two exceptions
• the first ionisation energy of aluminium is lower than that of magnesium.
• the first ionisation energy of sulphur is lower than that of phosphorus

Question 30

Explain why the first ionisation energy of aluminium is lower than that of magnesium

Answer 30

• The outer electron that aluminium (group 3) has to lose is in the 3P sub shell
• The outer electron that magnesium (group 2) has to lose is in the 3s sub shell
• the 3P sub-level sits at a higher energy level than the 3S sub-level does and therefore the electron is easier to lose.

Question 31

Explain why the first ionisation energy of sulphur is lower than that of phosphorus

Answer 31

• The outer electron in phosphorus is in the 3p3 position and is unpaired as the sub shell is half full
• The outer electron in sulphur is a 3p4 electron and therefore is a paired electron and is easier to remove than the unpaired electron in phosphorus due to the repulsion of the paired electrons.

Question 32

What is the trend in successive ionisation energies within an element

Answer 32

• there is an increase in successive ionisation energies.
• This increase is a bigger jump when the next electron is being removed from a completely different main level.