Topic 4 Inorganic Chemistry and the Periodic Table

Question 1

Group 2: Trend in ionisation energy

Answer 1

First and second ionisation energy decreases going down the group:
-Nuclear charge increases, so does force of attractions for the electron being removed (this causes an increase going down).
-As each quantum shell is added, energy of the outermost electron increases.
-As the number of filled electrons increases, the force of repulsion on electron being removed increases.
The two decrease down the group outweighs the one increase meaning an overall decrease.,

Question 2

Group 2: Reactions with oxygen

Answer 2

Similar reactions then when magnesium burns in air (very bright flame and the formation of a white solid).
However calcium, strontium and barium react with more vigour (barium is the most reactive-often stored in oil to stop reactions with air).
–>2M(s) + O2(g) –> 2MO(s)
–> The element needs to be heated for the reaction to start
(Reacts with increasing vigour going down the group).

Question 3

Group 2: Reactions with chlorides

Answer 3

Reacts with increasing vigour going down the group
–> M(s) + Cl2(g) –> MCl2(s)

Question 4

Group 2: Reactions with water

Answer 4

-The reaction between magnesium and water is very slow and does not proceed completely.
-Calcium, strontium and barium react with increasing vigour (which can be see by the increase of effervescence)
–> M(s) + 2H2O(l) –> M(OH)2(aq) + H2(g)
-Calcium hydroxide is only slightly soluble in water, so the liquid in this experiment goes cloudy as a precipitate of calcium hydroxide forms

Question 5

Group 2: Magnesium and steam

Answer 5

Magnesium reacts differently when heated with steam- it rapidly forms magnesium oxide (a white solid) and hydrogen gas is a vigorous reaction.
(The hydrogen formed is burned as it leaves the tube. This is for safety reasons, to prevent highly flammable gas into the lab).

Question 6

Group 2: Reactions of beryllium and radium

Answer 6

(Don’t need to know specific reactions)
Trend is increasing reactivity down the group:
-So beryllium is less reactive than magnesium
-And radium is more reactive than barium

Question 7

Group 2: Reactions of the oxides with water

Answer 7

-They react with water to form alkalis (form colourless solution)
–> MO(s) + H2O(l) –> M(OH)2(aq)

Question 8

Group 2: Trends in solubility of the hydroxides

Answer 8

Solubility of hydroxides increases going down group 2.

Question 9

Group 2: Testing for carbon dioxide (hydroxides)

Answer 9

(Limewater goes cloudy)
Limewater is the name for saturated aqueous solution of calcium hydroxide; carbon dioxide reacts to form calcium carbonate which is insoluble in water and is the white precipitate.
–> CO2 + Ca(OH)2 –> CaCO3 + H2O

Question 10

Group 2: Milk of magnesia (hydroxides)

Answer 10

Taking milk of magnesia neutralises some of the hydrochloric acid in the stomach reducing symptoms of indigestion.
–> Mg(OH)2 + 2HCl –> MgCl2 +2H2O

Question 11

Group 2: Reactions of the oxides and hydroxides with acids

Answer 11

All react with acid to form salts and water (neutralisation reactions).
Observations made: a white solid reacts to form a colourless solution.
–> MgO + H2SO4 –> MgSO4 + H2O
–> Ba(OH)2 + 2HCl –> BaCl2 + 2H2O

Question 12

Group 2: Use in agriculture

Answer 12

Farmers use lime to control soil acidity so a greater yield of crops can be obtained.
Lime is mostly calcium hydroxide.
–> Ca(OH)2 + 2HNO3 –> Ca(NO3)2 + 2H2O (nitric acid is used to represent the acid in the soil).

Question 13

Group 2: Trends in solubility of the group 2 sulphates

Answer 13

Solubility of sulfates decreases going down group 2.
-Magnesium sulphate is classed as soluble
-Calcium sulfate is slightly soluble
-Strontium and barium sulfate
(The very low solubility of barium sulfate is used to test for sulfate ions in solution).

Question 14

Group 2: Testing for sulfate ions

Answer 14

-Add dilute hydrochloric acid (or nitric acid)
-Add a few drops of barium chloride solution
-A white precipitate of barium sulfate confirms presence of sulfate ion
–> Ba(NO3)2(aq) + Na2SO4(aq) –> BaSO4(s) + 2NaNO3(aq)

Question 15

Group 2: Barium meals

Answer 15

Barium meals contain barium sulfate (which is not poisonous to humans as it insoluble).
They are used in hospitals to allow soft tissues to show up in an x-ray.

Question 16

Group 2: Thermal stability

Answer 16

Thermal stability indicates how stable a compound is when it is heated: it doesn’t decompose (thermally stable), it decomposes as much as possible (not at all thermally stable).

Question 17

Group 2: Factors affecting thermal stability

Answer 17

-The charge on a group 2 cation (is double that of a group 1 cation)
-The size (ionic radius) of group 2 cation (is smaller that that of the group 1 cation)
-The nitrate (NO3-) and carbonate (CO3^2-) anions are more complex than the Cl- ions.
These differences mean that when group 2 nitrates and carbonates are heated, they do not melt.

Question 18

Group 2: Thermal stability of nitrates

Answer 18

All of group 1&2 elements are white solids. When they are heated, they all decompose to nitrates or oxides and give off nitrogen dioxide (brown fumes) and/or oxygen.
If the nitrate contains water of crystallisation, then steam will be observed.
—> If no browns fumes are observed, this indicates a lesser decomposition: metal nitrate –> metal nitrate + oxygen
—> If brown fumes are observed, this indicates a greater decomposition: metal nitrate –> metal oxide + nitrogen dioxide + oxygen

Question 19

Group 2: Thermal stability of carbonates

Answer 19

All of group 1&2 elements are white solids. When they are heated they either don’t decompose, or decompose to oxides and give off carbon dioxide.
Gas given off is colourless so no observations can be made.
-Lithium carbonate decomposes at lower temperatures than the other group 1 carbonates: Li2CO3 –> Li2O + CO2.
Other group 1 carbonates do not decompose on heating, except at very high temperatures.
All group 2 carbonates decompose in the same way, but with increasing difficulty down the group.
-A typical equation for one of these decompositions: CaCO3 –> CaO + CO2.

Question 20

How to do a flame test

Answer 20

-Add a few drops of concentrated hydrochloric acid to the solid metal and mix together so the metal begins to dissolve
-Dip a clean metal wire (platinum or nichrome) or a silica rod into the mixture
-Hold the end of the wire in the flame and observe the colour

Question 21

Colours of flame tests

Answer 21

-Lithium (Li+): red
-Sodium (Na+): yellow/orange
-Potassium (K+): lilac
-Rubidium (Rb+): red/purple
-Caesium (Cs+): blue/violet
-Beryllium (Be2+): no colour
-Magnesium (Mg2+): no colour
-Calcium (Ca2+): brick red
-Strontium (Sr2+): crimson red
-Barium (Ba2+): apple green

Question 22

What causes the colours in flame tests?

Answer 22

Electron transitions.
-Electrons absorb energy and move to higher energy levels from ‘ground state’
-The electron in a higher energy level is in an ‘excited state’
-The movement is immediately followed by the return of the electron to its ground state, which releases energy
-If this energy corresponds to radiation in the visible light spectrum, then a colour appears.
-This is the flame test colour

Question 23

The test for ammonium ions

Answer 23

-Add sodium hydroxide solution
-Warm the mixture
NH4^+ + OH- –> NH3 + H2O
Warming releases ammonia gas. Damp red litmus paper turns blue if ammonia gas is present.
(Hydrogen chloride gas reacts with ammonia to form white fumes of ammonium chloride: NH3 + HCl –> NH4Cl).

Question 24

Intro to group 7 elements

Answer 24

States at room temperature:
Fluorine: gas
Chlorine: gas
Bromine: liquid
Iodine: solid
Astatine: solid

Question 25

Group 7: Dipoles (mpt & bpt)

Answer 25

-Halogen molecules are non-polar (covalent bonds between diatomic molecule is shared equally)
-The force of attraction is described as an instantaneous dipole-induced dipole attractions, and these are the intermolecular forces of attraction that exist between halogen molecules.
–> This explains the increase in both melting and boiling temperatures going down group 7.

Question 26

Group 7: Equations for changes of state

Answer 26

-When bromine is left at room temperature, it gives off a brown vapour, as its boiling point (59°C) is not much higher than room temperature: Br2(l) –> Br2(g).
-When iodine is warmed, most of it changes directly into a vapour without melting (sublimation): I2(s) –> I2(g).

Question 27

Group 7: Trend in electronegativity

Answer 27

The electronegativity of an atom depends on:
-Its nuclear charge; the bigger the nuclear charge, the higher the electronegativity
-The distance between the nucleus and the bonding pair of electrons; the shorter the distance, the higher the electronegativity
-The shielding effect of electrons in inner energy levels; the fewer energy levels, the higher the electronegativity.
The electronegativity of the group 7 elements is the highest of any group in the periodic table (the electronegativity of fluorine is the highest of all elements).

Question 28

Group 7: Trend in reactivity

Answer 28

Reactivity decreases going down group 7.
–> Fluorine is an extremely reactive element.

Question 29

Group 7: Reactions with metals in group 1&2

Answer 29

-Reactions are most vigorous between elements at the bottom of group 1&2 and elements at the top of group 7 (so the most vigorous reaction should be between caesium and fluorine, and the least vigorous between beryllium and iodine).
-The products of these reactions are salts (ionic salts that are usually white).
-Reactions involve electron transfer to the halogen, so they are redox reactions in which the halogen acts as an oxidising agent.
-The oxidation number of the halogen decreases from 0 to -1 and the oxidation number of the metal increases from 0 to +1 or +2

Question 30

Group 7: Halogen/halide displacement reactions

Answer 30

A more reactive halogen can displace a less reactive halogen from one of its compounds:
-Chlorine displaces bromine and iodine
-Bromine displaces iodine but not chlorine
-Iodine does not replace either chlorine or bromine
These reactions occur in aqueous solution, so any reaction that takes place is indicated by a colour change: add cyclohexane and shake to distinguish between each one.

Question 31

Disproportionation reactions of chlorine: Chlorine with water

Answer 31

When chlorine is added to water, it dissolves to form a solution called ‘chlorine water’.
(Chloric(I) acid- HClO).
Disproportionation reaction:
Cl2 + H2O –> HCl + HClO
0 -1 chlorine is reduced
0 +1 chlorine is oxidised

Question 32

Disproportionation reactions of chlorine: Chlorine with cold alkali

Answer 32

When chlorine is added to cold dilute aqueous sodium hydroxide, it reacts to form the salts of the acids; sodium chloride and sodium chlorate(I) (aka sodium hypochlorite).
Disproportionation reaction:
Cl2 + 2NaOH –> NaCl + NaClO + H2O
0 -1 chlorine is reduced
0 +1 chlorine is oxidised
The sodium(I) chlorate formed is also a disinfectant, but is mainly known for its bleaching action.

Question 33

Disproportionation reactions of chlorine: Chlorine with hot alkali

Answer 33

When chlorine is added to hot concentration sodium hydroxide solution, it reacts to for sodium chloride and sodium chlorate(V).
Disproportionation reaction:
3Cl2 + 6NaOH –> 5MNaCl + NaCLO3 + 3H2O
0 -1 chlorine is reduced
0 +5 chlorine is oxidised
The sodium chlorate(V) formed is also used in bleaching, and as a weed killer.
Bromine and iodine react in a similar way.

Question 34

Group 7: Redox reactions

Answer 34

-Oxidising power of halide ions decrease going down the group.
-Reducing power of halide ions increase going down the group: 2X- –> X2 + 2e-

Question 35

Group 7: Sodium bromide

Answer 35

The formation of misty fumes of hydrogen bromide can be represented by the equation:
NaBr + H2SO4 –> NaHSO4 + HBr
Full half equation:
2HBr + H2SO4 –> 2H2O + SO2 + Br2 (represents the oxidation of misty fumes of hydrogen bromide).

Question 36

Group 7: Sodium iodide

Answer 36

The formation of misty fumes of hydrogen iodide can be represented by the equation:
NaI + H2SO4 + HI
Full half equation: 6HI + H2SO4 –> 4H2O + S + 3I2 (represents the oxidation of the misty fumes of hydrogen iodide)

Question 37

Observations and products: NaCl (1)

Answer 37

Observations: misty fumes
Products: hydrogen chloride

Question 38

Observations and products: NaBr (3)

Answer 38

Observations: misty fumes, brown fumes, colourless gas with choking smell
Products: hydrogen bromide, bromine, sulfure dioxide

Question 39

Observations and products: NaI (5)

Answer 39

Observations: misty fumes, purple fumes or black solid, colourless gas with choking smell, yellow solid, colourless gas with rotten egg smell
Products: hydrogen iodide, iodine, sulfur dioxide, sulfur, hydrogen sulfide

Question 40

Group 7: sodium chloride

Answer 40

Misty fumes
NaCl + H2SO4 –> NaHSO4 + HCl

Question 41

Testing for halide ions in solution

Answer 41

-Add dilute nitric acid to halide solution (so other anions don’t form precipitates eg. carbonates)
-Add silver nitrate
-A precipitate is obtained:
Chloride: white ppt.
Bromide: cream ppt.
Iodide: yellow ppt.
-Ammonium is added to distinguish between each ion
Chloride: ppt. dissolves in dilute ammonia (/both)
Bromide: ppt. dissolves in only conc. ammonia
Iodide: ppt. does not dissolve in either

Question 42

Group 7: Hydrogen halides acting as acids

Answer 42

All of the hydrogen halides are colourless gases and exits as polar diatomic molecules.

Question 43

Hydrogen halides acting as acids: Reactions with water

Answer 43

They readily react with water to form acidic solutions, all of which are colourless.

Question 44

Hydrogen halides acting as acids: Reactions with ammonia

Answer 44

Hydrogen halides all react with ammonia gas to form salts (all of which are white solids).
Ammonia and hydrogen chloride gases are given off from the cotton wool pieces soaked in conc. aqueous ammonia and conc. hydrochloric acid. These colourless gases move through the tube until they meet and react to form ammonium chloride. Ammonia gas molecules move more quickly and therefore travel further in the same time as hydrogen chloride molecules.