3.1 periodicity

Question 1

How are the elements arranged in the periodic table?

Answer 1

In the modern periodic table elements are arranged in the order of increasing atomic number.

Question 2

What is meant by periodicity?

Answer 2

It is the trend in properties across the periodic table.
Across each period elements change from metals to non-metals.
Proton number increases
Ionisation energy generally increases.
Atomic radius decreases

Question 3

Define first ionisation energy.

Answer 3

Energy required to remove one electron from each atom in one mole of gaseous atoms
Li(g)– Li+ +e-

Question 4

Define second ionisation energy.

Answer 4

Energy required to remove a second mole of electrons from a gaseous atom
Li+ (g)—- Li2+ + e-

Question 5

What factors affect ionisation energy?

Answer 5

Atomic radius
Nuclear charge
Electron shielding

Question 6

How does atomic radius affect ionisation energy?

Answer 6

The larger the atomic radius the smaller the nuclear attraction experienced by the outer electron.

Question 7

How does nuclear charge affect ionisation energy?

Answer 7

The higher the nuclear charge the larger is the attraction force on the outer electron.

Question 8

How does shielding affect ionisation energy?

Answer 8

The inner shell electrons repel the outer shell electrons and create a smaller nuclear attraction on the outer electron

Question 9

Why does first ionisation energy decrease between group 2 and 3?

Answer 9

Group 3 elements have their outermost electron in the p-orbital, whereas group 2 have theirs in the s.
P- orbitals have a slightly higher energy level than S-orbitals and so are marginally further away from the nucleus do they are easier to remove.

Question 10

Why does first ionisation energy decrease between group 5 and 6?

Answer 10

In both groups the outer electrons are in the P-orbital. However, group 6 outer electron is paired up in the Px orbital so it experiences electron-electron repulsions making it slightly easier to remove.

Question 11

What is the trend of first ionisation energies down a group?

Answer 11

Shielding increases and so does the atomic radius causing it to decrease.

Question 12

How does the trend in ionisation energy look across a group?

Answer 12

Number of protons increases (nuclear charge), but the shielding remains the same so atomic radius decreases. So, first ionisation energy increases

Question 13

Describe the properties of giant metallic lattices.

Answer 13

High melting and boiling points
Electrical conductivity
Malleability and ductility

Question 14

Why do metallic lattices have high melting and boiling points?

Answer 14

Strong attraction between positive ions and negative, delocalised electrons.

Question 15

Explain why giant metallic lattices have electrical conductivity?

Answer 15

celocalised electrons are free to mave and carry charge.

Question 16

Explain why giant metallic lattices are malleable and ductile?

Answer 16

The delocalised electrons can move giving the structure a certain degree of ‘give’ allowing atoms or layers to slide past each other.

Question 17

Describe the structure, forces and bonding of the elements across period 2.

Answer 17

Li,Be— giant metallic, strong forces between cations and negative electrons, metallic bonding.

B,C– giant covalent, strong covalent bonds between atoms

N2,O2,F2,Ne — simple molecular, weak van der waals between molecules

Question 18

Describe the structure, forces and bonding of the elements across period 3.

Answer 18

Na, Mg, Al — giant metallic lattices, metallic bonding
Si – giant covalent, strong forces between atoms
P4,S8,Cl2,Ar — simple molecular with weak van der Waals between molecules

Question 19

Describe melting points across period 3

Answer 19

g1,2,3 melting points increase steadily due to their giant structure and increasing nuclear charge.

G4 has a large increase due to its giant covalent structure and strong covalent bonds

G5,6,7,8, is relatively low due to weak van der Waals forces between molecules (S8 is slightly higher due to it being a larger molecule than P4)

Question 20

What are the physical properties of group 2 elements?

Answer 20

Reasonably high melting points
Light metals with low densities.
Form colourless(white ) compounds

Question 21

Describe and explain the configuration down group 2.

Answer 21

Mg - 1s2 2s2 2p6 3s2 [Ne] 3s2

Question 22

Describe and explain the ionisation energies down group 2.

Answer 22

1st and 2nd ionisation energies decrease

They are strong reducing agents so they are oxidised in reactions and form 2+ ions.

Question 23

Describe the reaction between group 2 elements and oxygen.

Answer 23

They react vigorously with oxygen.
2Ca +02 — 2CaO
OXIDISED - Ca – Ca2+ + 2e-
REDUCED - 0 +2e- — O2-

Question 24

Describe the reactions between group 2 and dilute acids.

Answer 24

All group two elements except beryllium react with acids to form salts and hydrogen. (compound with ionic ionic assembly of anions and cations)

Ca+2HCl — CaCl2 +H2
OXIDISED- Ca— Ca2+ +2e-
REDUCED - 2H+ + 2e- —- H2

Question 25

Describe the reaction between group 2 elements and water.

Answer 25

All of them except beryllium react with water to form hydroxides x(OH)2 and they react more vigorously as you go down the group.

Ca + 2H2O = Ca(OH)2 +H2
OXIDISED- Ca= Ca2+ +2e-
REDUCED - 2H2O +2e- = 2OH- + H2

Question 26

Describe the equation of a reaction between a metal oxide and water.

Answer 26

MgO + H20 = Mg(OH)2

Question 27

Describe the solubility of group 2 metal hydroxides.

Answer 27

The solubility of hydroxides in water increases down a group.
When a hydroxide is more soluble than another it will release more OH- ions and will make a more alkaline solution with a higher pH.

Question 28

What are some uses of group 2 compounds?

Answer 28

Calcium hydroxide is used by farmers to neutralise soils.
Magnesium hydroxide is used as an antacid.
Calcium carbonate is a useful building material.

Question 29

What is the halogen trend in boiling points?

Answer 29

Moving down the group the boiling point increases and their physical states change.
This is because each successive element has an extra shell of electrons leading to a higher level of London forces.

F2- gas
Cl2-gas
Br2-liquid
I2-solid
At2 -solid

Question 30

Describe the reactivity of the halogens

Answer 30

They are very reactive and electronegative so they are good at attracting and capturing electrons so they are strong oxidising agents and form 1- halide ions.

The reactivity decreases because atomic radius increases as well as electron shielding so the ability to gain an electron in the P-subshell and form 1- ions decreases

Question 31

What is a disproportionation reaction?

Answer 31

Reaction where the same element is both oxidised and reduced.

Question 32

Give some examples of disproportionation reactions.

Answer 32

Reaction of chlorine with water to make bleach and hydrochloric acid.
Cl2+ H20 = HClO + HCl
Chlorine is both reduced and oxidised

Reaction of chlorine with cold dilute aqueous sodium hydroxide to form sodium hypochlorate
Cl2 + 2NaOH = NaClO + NaCl+ H2O
chlorine is both reduced and oxidised

Question 33

Describe the redox reactions that occur between the halogen group.

Answer 33

They occur between aqueous solutions of halides and aqueous halogens. A more reactive halogen will oxidise and displace the halide

E.G.
Cl2+ 2Br- = 2Cl- + Br2

Question 34

What colour does the solution go when it has been displaced by a halogen

Answer 34

water cyclohexane
Cl2 pale green pale green
Br2 orange orange
I2 brown violet

Question 35

why is cyclohexane used in the displacement pag?

Answer 35

To help distinguish between bromine and iodine

Question 36

What is the test and positive result for carbonate ions?

Answer 36

Add a dilute strong acid to suspected carbonate and pass collected gas through lime water
This causes the limewater to go cloudy and fizzes

Question 37

What is the test and positive result for sulphate ions?

Answer 37

Add dilute hydrochloric acid and barium chloride
A white precipitate called barium sulphate is formed

Question 38

What is the test and positive result for halide ions?

Answer 38

Dissolve the suspected halide in water and add aqueous silver nitrate and then ammonia.
Silver chloride forms a white precipitate and is soluble in dilute ammonia
silver bromide forms a cream precipitate and is soluble in concentrated ammonia
silver iodide forms a yellow precipitate and is insoluble in ammonia

Question 39

What is the test and positive result for ammonium ions?

Answer 39

Add sodium hydroxide to the suspected ammonium and heat gently.
The evolved gas will turn red litmus paper blue and will have a distinct smell.

Question 40

What is the order for the ion tests and why?

Answer 40

carbonate
sulphate
halide
Barium ions form insoluble white precipitates in carbonates and sulphates.
Silver ions form insoluble silver sulphate so you wouldn’t be able to distinguish between it and a silver halide.

Question 41

Explain why the test for carbonates work.

Answer 41

The CO3 2- ions react with acids

CO3 2- + 2H+ = H20+CO2

Question 42

Explain why the test for sulphate ions works.

Answer 42

Sulphate ions react with barium ions to form an insoluble salt.

Ba2+ +SO4 2- = BaSO4

Question 43

Explain why the test for halide ions works.

Answer 43

Halide ions react with silver ions to form different coloured halide precipitates.
Ag+ +Cl- = AgCl
Ag+ +Br- = AgBr
Ag+ +I- =AgI

Question 44

Explain why the test for ammonium ions works.

Answer 44

The ammonium NH4= ions react with OH- ions to form ammonia and water