Topic 4: Inorganic Chemistry and the Periodic Table

Question 1

Explain reasons for the trend in reactivity of the Group 2 elements down the group

Answer 1

Reactivity increases as you go down the group because atomic radius increases due to more shells.
So less energy is required down the group to remove the outermost 2 e

Question 2

What are the reactions of elements in Group 2 with oxygen?

Answer 2

You see a bright white flame and the formation of a white solid, with more vigorous reactions down the group.
If the burning metal is placed in a glass jar of 02 then the same reaction occurs but more vigorously.
For the reaction to start the element must be heated, but even w/o heat hay slow reaction entre the air and the metal, and an oxide coating prevents further reaction.
Ba is so reactive that it must be stored in oil.
general equation: 2M+ O2——-> 2MO

Question 3

What are the reactions of the oxides of Group 2 elements with water and dilute acid?

Answer 3

G2 oxides= basic oxides, meaning they react w/ water to form alkalies. You observe solids reacting to form colourless solutions
G2 oxides react w/ acids to form salt+water (neutralisation). You observe a white solid reacting to form a colourless solution, and the reaction is exothermic.

Question 4

What are the reactions of the hydroxides of Group 2 elements with dilute acid?

Answer 4

SAME AS G2 OXIDES: G2 hydroxides react w/ acids to form salt+water (neutralisation).
You observe a white solid reacting to form a colourless solution, and the reaction is exothermic.

Question 5

Describe the trends in solubility of the hydroxides and sulfates of Group 2 elements

Answer 5

solubility of G2 hydroxides increases down the group.

Solubility of G2 sulphates increases up the group

Question 6

What are the flame colours for Group 1 and 2 compounds?

Answer 6

K= lilac
Na= orange 
Li= orange/red 
Cu= blue/green
Ba= pale green
Ca= brick red
Sr= crimson
Mg= no colour

Question 7

Devise an experiment to show patterns in thermal decomposition of Group 1 and 2 carbonates

Answer 7

-Weigh an empty boiling tube and place 1cm depth of metal carbonate into it.
-reweigh the boiling tube, and clamp it at a 45 degree angle.
-fit a bung w/ a delivery tube at the top of the boiling tube. The delivery tube should contain a test tube of limewater.
-Strongly heat the boiling tube for 5 mins and record when the limewater is cloudy.
-After 5 mins remove the the delivery tube from the limewater, leave it to cool and reweigh it.
-Repeat with other metal carbonates.
-Heat causes air in the test tube to expand, so the first bubbles you observe in the tube are expanded air, NOT CO2.
CaCO3 —–> CaO + CO2

Question 8

Devise an experiment to show flame colours in compounds of Group 1 and 2 elements

Answer 8

Dip a nichrome wire in HCL to avoid any impurities. Dip wire into the pure solid metal. Hold it to the flame and observe the colour change

Question 9

Explain reasons for the trend in reactivity of Group 7 elements down the group

Answer 9

Group 7 elements must gain an electron.
As atomic radius increases down the group it becomes harder to attract an electron as the positive attraction of the nucleus is weakened by additional shielding.
Therefore reactivity decreases down the group
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Question 10

Explain The reactions of halogens with metals and their ability as oxidising agents

Answer 10

Most halogens react with metals to produce metal halides. All halogens react with Na: 2Na + X2 ———> 2NaX. You observe a bright yellow flame and white solid forming of the sodium halide.
All halogens except iodine react with Fe to produce solid Fe (iii) halides:
2Fe + 3X2 ———> 2FeX3.
Iodine isn’t a powerful enough oxidising agent to oxidise Fe from 0 to +3 but it can oxidise it to +2. Fe + I2 ———> FeI2
Overall halogens are stronger oxidising agents at the top of the group

Question 11

Describe and explain the precipitation reactions of the aqueous anions Cl–, Br– and I– with silver nitrate solution, followed by aqueous ammonia solution

Answer 11

Add silver nitrate. Add nitric acid to prevent other precipitates forming. Add the halide sample, and then ammonia.
Chloride ions = white precipitate that dissolves in dilute NH3
Bromide ions= Cream precipitate that dissolves only in concentrated NH3
Iodide ions= Yellow precipitate that does not dissolve in ammonia

Question 12

Describe and explain the reactions of:

iii hydrogen halides with ammonium hydroxide, with water and with ammonia

Answer 12

Hydrogen halide + ammonium hydroxide –> salt + water
HX + NH4OH ———> NH4X + H2O

Hydrogen halide + water —> X⁻ + H3O⁺

Hydrogen halide + ammonia –> NH4X
HX + NH3 —> NH4X

Question 13

What are the reactions of elements in Group 2 with chlorine?

Answer 13

M+ Cl2——-> MCL2

Reactions get more vigorous down the group

Question 14

What are the reactions of elements in Group 2 with water and steam?

Answer 14

Magnesium + water is a very slow reaction, but down the group there is an increase in fizzing. Metal hydroxides form
Mg+ steam rapidly forms MgO (white solid) and H gas. The H gas is burned when it leaves the test tube for safety to prevent a highly flammable gas entering a lab

Question 15

Give a test to identify solid halide ions

Answer 15

Add concentrated H2SO4 to a solid sample. Also soak a splint in ammonia solution to test for the hydrogen halide gas that will be produced:
Floride and chloride ions: white misty fumes of HF and HCl
Bromide: misty acidic fumes contaminated with brown bromine vapour
Iodide: some steamy fumes of HI but lots of red and purple vapour

Question 16

Why are all the halogens potential oxidising agents?

Answer 16

Each of the halogens could take electrons from other species to make their ions, so they can be oxidising agents.

Whenever one of these halogens oxidise something IN SOLUTION, the halogen ends up as halide ions with water molecules attached

Question 17

What kind of oxidising agent is fluorine?

Answer 17

It is such a powerful oxidising agent. F oxidises water to oxygen so it’s impossible to do simple solution reactions with it.

Question 18

What happens to the oxidising ability of the halogens down the group?

Answer 18

Oxidising ability falls down the group.

So a halogen higher in the group can oxidise all those lower down

Question 19

Give the equation for chlorine oxidising iodide to iodine. what is observed?

Answer 19

Cl2 + 2I- ———> 2Cl- +I2

The iodine appears as a red solution if you use little chlorine, or a dark grey precipitate if chlorine is used in excess

Question 20

Give the equation and observations for bromine oxidising iodide ions into iodine. what is observed?

Answer 20

Br2 + 2I- ——-> 2Br +I2

A red solution of iodine is formed until the bromine is in excess, the u get a dark grey precipitate

Question 21

When halogens oxidise something in solution, how do the halogens end up as halide ions w water molecules attached?

Answer 21

The halogen starts as diatomic molecules which may be a solid liquid or gas.
These must be split apart to make individual atoms.
Those atoms gain an electron, and isolated ions become wrapped in water molecules to form hydrated ions

Question 22

What is atomisation energy?

Answer 22

The energy needed to produce 1 mole of isolated gaseous atoms starting from an element in its standard state, (eg gas for Cl2, liquid for Br2 etc)

Question 23

What would be the atomisation energy for the halogens?

Answer 23

For a gas like Cl2, it’s half the bond enthalpy because breaking the Cl-Cl bond produces 2 atoms, not 1.
For a liquid or solid like Br2 or I2, it also includes the energy needed to convert them into gases

Question 24

State the definition for electron affinity

Answer 24

The energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions

X(g) + e- –> X- (g)

Question 25

Devise an experiment to show patterns in thermal decomposition of Group 1 and 2 nitrates.

Answer 25

Put the substance in the boiling tube, heat it and see how long it takes for orange gas to form, or rekindle a glowing splint to check for oxygen gas.
NO2 is orange/brown fumes

Question 26

What is the effect of heat on G2 carbonates?

Answer 26

All G2 carbonates undergo thermal decomposition to give the metal oxide and CO2.
Carbonates get more stable to heat as you go down G2.
Therefore Mg should decompose most easily.

Question 27

What is the effect of heat on G2 nitrates?

Answer 27

All G2 nitrates undergo thermal decomposition to give the metal oxide, nitrogen dioxide and O2.
Nitrates get more stable to heat as you go down G2. Nitrogen dioxide = brown/orange fumes

Question 28

What is the effect of heat on G1 nitrates? Give the equations

Answer 28

G1 compounds are more stable to heat than G2 compounds, but Li compounds behave similarly to G2.
Lithium nitrate when heated:
4LiNO3 (s) ———> 2Li2 (s) + 4NO2 (g) + O2 (g)
The rest of the group don’t decompose so completley at bunsen burner temps. The rest of the group produce the metal nitrite and O2, but no brown fumes of NO2:
2XNO3 (s) ——-> 2XNO2 (s) + O2 (g)

Question 29

What is the effect of heat on G1 carbonates?

Answer 29

Li carbonate behaves like G2- producing LiO and CO2:
Li2CO3 (s) ——> Li2O (s) + CO2 (g)
The rest of G1 carbonates dont decompose at bunsen temps, but will at higher temps.
Again, decomposition temps increase down G1 because thermal stability increases

Question 30

Explain why thermal stability increases down the group for the metal carbonates and nitrates

Answer 30

G1 and G2 metal ions get larger down the group. This means the positive charge of the ion is spread over a larger area, meaning a lower charge density. The polarising power of the cation decreases. Polarisation of the carbonate or nitrate ion thus decreases, so thermal stability increases.

Question 31

Why does magnesium nitrate decompose more readily on heating than potassium nitrate?

Answer 31

Potassium has a greater ion radius than magnesium because it has more shells. Therefore potassium has a lower charge density than magnesium. The polarising power of the magnesium ion increases. Polarisation of the nitrate ion increases.

Question 32

Explain the solubility of the halogens

Answer 32

Halogens are diatomic so the only intermolecular forces between the molecules are London forces. This dictates their solubility.
All halogens are more soluble in organic solvents compared to water because the main intermolecular force in organic solvents are also London forces. Solubility in water DEcreases down the group
Solubility in hexane INcreases down the group

Question 33

What happens when you add a halogen to a halide salt of KCl or NaCl? Describe all colour changes

Answer 33

Chlorine + KCl gives no reaction, it remains colourless.
Bromine + KCL gives no reaction, it remains red/brown.
Iodine + KCL gives no reaction, it remains brown

Question 33

What happens when you add a halogen to a halide salt of KBr or NaBr? Describe all colour changes

Answer 33

Cl2 + KBr: solution goes from colourless to red/brown as chlorine displaces bromide ions.
Br2 + KBr: No reaction, the solution remains brown/red
I2 + KBr: No reaction, the solution remains brown

Question 33

What happens when you add a halogen to a halide salt of KI or NaI? Describe all colour changes

Answer 33

Cl + KI: Solution goes from colourless to brown as chlorine displaces iodine.
Br + KI: Solution goes from red to brown as bromine displaces iodine.
I + KI: solution remains brown

Question 33

What happens when you add cyclohexane to the halogens?

Answer 33

Chlorine turns into a very pale yellow.
Bromine turns orange/yellow.
Iodine turns purple

Question 34

Give reasons for the trends in electronegativity for Group 7 elements

Answer 34

Electronegativity falls down the group. Atoms become worse at attracting bonding pairs of electrons bc nuclear attraction decreases.
In fluorine, solo hay the 1S and 2S shell shielding the nuclear charge. The larger pull from the closer fluorine nucleus gives it more electronegativity

Question 38

Give reasons for the trends in melting and boiling temperatures for Group 7 elements

Answer 38

Melting and boiling points increase as you go down the group.
As molecules get bigger there are more electrons which can move around and set up temporary dipoles which create more London forces. These need more energy to break.

Question 39

Describe the use of chlorine in water treatment

Answer 39

Chlorine is used in small quantities to kill bacteria. Chlorine can be toxic but the benefits of clean water outweigh the risks.

Question 40

Give the general equation for G2 metals reacting w/ Oxygen

Answer 40

general equation: 2M+ O2——-> 2MO

Question 41

Give the general equation for the thermal decomposition of G2 carbonates

Answer 41

XCO3 (s) ——-> XO (s) + CO2 (g)

Question 42

Give the ionic equation for carbonate ions once you test for them by adding aqueous acid

Answer 42

Acid + metal carbonate ———> salt + water + carbon dioxide.

2H+(aq) + CO3 2-(aq) ———> H2O(l) + CO2 (g)

Question 43

Give the equation for F oxidising water into oxygen.

Answer 43

2F2 + 2H2O ——> 4HF + O2

Question 44

Give the general equation for the thermal decomposition of G2 nitrates

Answer 44

2X(NO3)2 (s) ——–> 2XO (s) + 4NO2 (g) + O2 (g)

NO2 is orange gas

Question 45

Describe the concept of charge density and polarising ability

Answer 45

smaller ions at the top of a group, and larger charges increase the charge density.
High charge density ions distort/polarise carbonate and nitrate ions more. Low charge density ions polarise nitrate and carbonate ions less.
Less distortion means a more stable carbonate!!

Question 46

Describe and explain ALL of the disproportionation reactions of chlorine, giving equations

Answer 46

Disproportionation of chlorine and water:
Cl2 + H20 ————> ClO- + Cl- + 2H+

Chlorine with cold, dilute aqueous sodium hydroxide forms bleach, aka sodium chlorate
2NaOH + Cl2 ————> NaClO + NaCl + H2O.

Disproportionation of chlorine with hot alkali:
3NaOCl(aq) ———> 2NaCl(aq) + NaClO3(aq)

If chlorine is bubbled directly into hot alkali:
3Cl2(g) + 6NaOH(aq) –> 5NaCl(aq) + NaClO3(aq) + 3H20
The oxidation state goes from 0 to -1 to 5