Chemistry - Periodicity And Reactivity Trends

Front 1

Period

Back 1

A row of elements with the highest energy electron in the same shell

Front 2

What is a group?

Back 2

A column of elements with similar chemical and physical properties and with the same number of electrons in their outer shell

Front 3

Periodicity

Back 3

The repeating trend in properties across each period in the periodic table

Front 4

How are elements arranged in the periodic table?

Back 4

In order of increasing atomic number

Front 5

Where are the different blocks in the periodic table

Back 5

S block to the left, d block in centre, p block to the right

(The width of each block is the same ad the number of electrons that fill the subshell)

Front 6

What is ionisation energy?

Back 6

It's the amount of energy needed to remove electrons

Front 7

What three factors affect ionisation energy?

Back 7

● atomic radius

● nuclear charge

● electron shielding

Front 8

What is the definition of first ionisation energy?

Back 8

It is the energy required to remove one electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.

Mg (g) --> Mg+(g) + e-

Front 9

Why does the value of ionisation energy increase with ionisation number?

Back 9

Electrons are drawn closely to the nucleus when lost because the same number of protons attracts fewer electrons

Front 10

What does a large increase in ionisation energy indicate?

Back 10

The decrease in shell number as electrons closer to the nucleus are harder to remove

Front 11

What happens to ionisation energy down a group, and why?

Back 11

It decreases down a group because increased atomic radius and electron shielding outweigh the effect of increasing nuclear charge.

Front 12

What happens to ionisation energy across a period, and why?

Back 12

It increases because of increasing nuclear charge. Atomic radius decreases across a period and shielding doesn't change as new electrons are added to the same shell

Front 13

Why is there a drop in first ionisation energy between Be and B?

Back 13

Because of a new subshell being filled. The first electron in a new subshell is easier to remove

Front 14

Why is there a drop in first ionisation energy from N to O

Back 14

Because electron pairing begins. And because of electron repulsion between electrons in the pair, the electron from the pair is easier to remove

Front 15

Metallic bonding

Back 15

It is the strong electrostatic attraction between cations and delocalised electrons

Front 16

What are the 3 properties of metals?

Back 16

● strong metallic bonds

● high electrical conductivity

● high melting and boiling points

Front 17

Why are metals conductive?

Back 17

Because when a voltage is applied to a metal, the delocalised electrons can move through the structure and carry charge.

(Ionic compounds cannot conduct electricity as their electrons are 'fixed in the ionic lattice")

Front 18

Explaij why metals have High melting and boiling points

Back 18

Because a high amount of energy is required to overcome the strong electrostatic attraction between the cations and electrons

Front 19

Why don't metals dissolve in water?

Back 19

Because the interactions between the charges in the lattice and the polar water molecules prevent it from doing so

Front 20

What do metals do instead of dissolve?

Back 20

They react in a chemical reaction (eg. Sodium in water)

Front 21

What 3 non metals form a giant covalent lattice?

Back 21

Boron, carbon and silicon

Front 22

State the properties of giant covalent lattices

Back 22

● high mp and bp

● they are insoluble in almost all solvents because the covalent bonds are too strong to be broken by interactions with solvents

● they don't conduct electricity because no electrons are available. (However in graphene and graphite, one free electron for a carbon bonded to 3 other carbons can help it conduct electricity)

●

Front 23

What is graphite? Decribe the structure

Back 23

It is composed of parallel layers of hexagonal arranged carbon atoms. These layers are bonded by weak London forces

Front 24

Describe the trend in melting points across period 2 and 3

Back 24

There is an increase in melting points from group 1 to 4. (Where Si or C is) and then afterwards there is a sharp decrease in melting points. The melting points thus stay low after group 4

Front 25

What is a reducing agent?

Back 25

A substance that adds electrons to other species and loses electrons itself

Front 26

What 3 things distinguish a group 2 metal?

Back 26

● all have 2 electrons in its outer shell in the s block - s2

● all oxidise - lose 2 electrons

● all form 2+ ions

Front 27

What happens to reactivity down group 2?

Back 27

It increases because ionisation decreases

Front 28

Explain the decrease in ionisation energy down group 2

Back 28

● atomic radius increases

● electron shielding increases

● nuclear charge also increases but the effect of this is outweighed

Front 29

Group 2 metal + oxygen?

Back 29

Forms a metal oxide

Front 30

Group 2 metal + water?

Back 30

This forms a metal hydroxide + hydrogen

Front 31

Group 2 metal + dilute acid?

Back 31

Produces a salt (aq) and hydrogen (g)

Front 32

Group 2 oxide + water?

Back 32

Forms an an alkaline solution of group 2 hydroxide

Front 33

Calcium oxide + water?

Back 33

Produces calcium hydroxide

Front 34

How soluble are group 2 hydroxides in water?

Back 34

Not very soluble therefore the hydroxide product would form a precipitate instead an alkaline solution when calcium oxide reacts with water

Front 35

What happens to the solubility of group 2 hydroxides as you move down the group?

Back 35

They become more soluble. And the alkaline solutions become stronger as more OH- ions are added

Front 36

Give 2 examples of group 2 compounds used in industry

Back 36

● calcium hydroxide - used in agriculture to neutralise acidic soils

● magnesium hydroxide/ calcium carbonate - used as 'antacids' to neutralise acids causing indigestion

Front 37

What 3 things distinguish a group 7 halogen?

Back 37

● Contains 7 electrons in its outer shell - s2p5

● all are reduced by gaining 1 electron

● form 1- ions

Front 38

What happens to the bp's of the halogens down the group?

Back 38

It increase as the increased number of electrons down the group increases the strength of the London forces between halogen molecules. Thus, more energy is required to overcome them

Front 39

What is an oxidising agent?

Back 39

A substance that removes electrons from other species and gains electrons itself

Front 40

What happens to rectaivity down group 7 and why?

Back 40

The reactivity decreases as the number of shells increases, causing an increase in atomic radius and electron shielding. Making it harder to gain an electron because nuclear attraction weakens

● nuclear charge increases but this effect is outweighed

Front 41

What organic solvent is added to make the halogen color show up?

Back 41

Cyclohexane

Front 42

When chlorine reacts with potassium bromide what happens?

Back 42

Potassium chloride + bromine is formed and the solution turns orange.

Front 43

What's disproportionation?

Back 43

The simultaneous oxidation and reduction if the same element

Front 44

What are the two disproportionation reactions of chlorine?

Back 44

● chlorine + water - to kill bacteria

● chlorine + sodium hydroxide - to form bleach

Front 45

Chlorine+ water?

Back 45

Produces chloric acid (HClO) and hydrochloric acid (HCl)

Front 46

Chlorine+ sodium hydroxide?

Back 46

Sodium chlorate (NaClO) + sodium chloride (NaCl) + water

Front 47

Carbonate test?

Back 47

Add nitric acid to the sample and see if bubbles are formed

Front 48

Sulfate test?

Back 48

Add barium chloride or barium nitrate. If a white precipitate of barium sulfate is formed... then sulfate is present

Front 49

Halide test?

Back 49

Add silver nitrate. Would form a precipitate depending on the halide

Front 50

Precipitate colours for silver halides?

Back 50

● silver chloride - white

● silver bromide - cream

● silver iodide - yellow

Front 51

To identify silver halides with ammonia?

Back 51

● chloride - dilute

● bromide - concentrated

● iodide - neither

Front 52

How to test for ammonium ions?

Back 52

Add sodium hydroxide to the mixture and heat it. Then test any gas produced with damp litmus paper... it should turn blue if the gas is ammonia

Front 53

Illustrate how calcium hydroxide is used in agriculture with an equation

Back 53

Ca (OH)2 + 2H+ ------> Ca2+ + 2H2O

Front 54

Explain why bromine has a lower boiling point than ICl even though they have the same relative molecular mass

Back 54

Because bromine only has London forces whereas ICl has London forces and permanent dipole dipole interactions. Which are stronger

Front 55

Show the displacement equation reaction between bromine and sodium iodide

Back 55

BR2 + 2I- --> 2Br- + I2

Front 56

Why are the halogens non polar?

Back 56

Because the atoms have the same elecryonegativity value so overall there is no permanent dipole and therefore no charge

Front 57

What are the conditions needed for a reaction between chlorine and sodium hydroxide to form NaClO

Back 57

Cold dilute sodium hydroxide

Front 58

Give the equation for the reaction between chlorine and cold dilute NaOH with state symbols

Back 58

Cl2 (aq) + 2NaOH (aq) --> NaCl (aq) + NaClO (aq) + H2O (l)

Front 59

What is formed in the reaction between chlorine and water

Back 59

Hydrochloric acid (HCl) an acid

And HClO which is a bleach

Front 60

Give the name of the Aqueous reagent that could distinguish between AgBr and AgI

Back 60

Concentrated ammonia sulfate... AgBr dissolves in it, AgI doesn't

Front 61

Give one similarity and one difference between the structures of graphene and graphite

Back 61

Similarity - both have layers of hexagons / delocalised electrons

Difference - graphene is 2d, graphite is 3d / graphite contains layers whereas graphene doesn't. / graphite has intermolecular forces between layers

Front 62

State the C-C-C bond angle of graphene

Back 62

120°

Front 63

Explain why gallium is insoluble in water?

Back 63

Because metallic bonds are very strong and don't break apart by interactions with water dipoles

Front 64

How would successive ionisation energies show that oxygen is in group 6 of the periodic table?

Back 64

Because there'd be a much larger increase between the sixth and seventh ionisation energy values

Front 65

Explain in terms of structure why the melting points of Na, Mg, Al and Si (in the same period) are larger than those of P and S

Back 65

● Na, Mg, Al and all create giant lattices

● P and S create small molecular structures

● in giant structures, strong bonds need to be broken whereas in simple molecular structures, only weak intermolecular forces are broken. Which requires less energy

Front 66

The variation in first ionisation energies across a period of the periodic table provided evidence for what structure within an atom?

Back 66

The sun shells of electrons

Front 67

From Li, Be, and fluorine. Which one will have the largest second ionisation energy and why?

Back 67

Lithium because the second electron is simply in the first shell unlike the second electron in Be and F which is still in the second shell. The outer electron in Li therefore is harder to remove as there is less electron shielding and a stronger pull of the electron