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wow woah insane its the best science flash cards in the observable universe confirmed > Inorganic chemistry > Flashcards

Inorganic chemistry Flashcards

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1
Q

Why ionisation energy decreases down group 2

A
  • more electron shielding
  • larger atomic radius
  • less attraction between nucleus and outer shell
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2
Q

Why reactivity increases down group 2

A
  • metals lose electrons in reactions
  • shielding and atomic radius increase
  • it becomes easier to lose electrons
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3
Q

Group 2 reaction with oxygen

A

Group 2 + oxygen&raquo_space; group 2 oxide

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4
Q

Group 2 reaction with chlorine

A

Group 2 + chlorine&raquo_space; group 2 chloride

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5
Q

Group 2 reaction with water

A

Group 2 + water&raquo_space; group 2 hydroxide + hydrogen

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6
Q

Group 2 oxide reaction with water

A

Group 2 oxide + water&raquo_space; group 2 hydroxide

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7
Q

Group 2 oxide / hydroxide reaction with dilute acid

A

Group 2 oxide / hydroxide + acid&raquo_space; group 2 salt + water

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8
Q

Solubility of group 2 sulfates down the group

A

Decreases (remember insoluble barium sulfate precipitate from the test for sulfates)

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9
Q

Solubility of group 2 hydroxides down the group

A

Increases (magnesium hydroxide can be used as medicine because of its low solubility)

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10
Q

Explain the trend in thermal stability of metal carbonates and nitrates

A
  • thermal stability decreases as polarising power of the cation increases (size decrease, charge increase)
  • as the anion’s electrons are distorted, its covalent bonds become weaker
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11
Q

How flame tests work

A
  • electrons absorb energy from the flame
  • they move into a higher energy level
  • as they move back down, they release the energy as specific wavelengths of light
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12
Q

How to do a flame test

A
  • dip nichrome wire into concentrated HCl then into the solid sample
  • hold the wire over a flame and observe the colour
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13
Q

How to test thermal stability of nitrates

A
  • heat the nitrate
  • normal decomposition gives no observation
  • greater decomposition (so a less stable nitrate) produces brown fumes (nitrogen dioxide)
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14
Q

How to test thermal stability of carbonates

A
  • weigh the sample of carbonate
  • heat it
  • weigh the sample again
  • if mass decreased, decomposition occurred (so a less stable carbonate)
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15
Q

Lithium flame

A

Red

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16
Q

Sodium flame

A

Yellow

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17
Q

Potassium flame

A

Lilac

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18
Q

Rubidium flame

A

Purple

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19
Q

Caesium flame

A

Blue

20
Q

Calcium flame

A

Orange

21
Q

Strontium flame

A

Crimson

22
Q

Barium flame

A

Apple green

23
Q

Why boiling point increases down group 7

A
  • amount of electrons increases

- london forces become stronger

24
Q

Why electronegativity decreases down group 7

A
  • electron shielding increases
  • atomic radius increases
  • less attraction between nucleus and bonding electron pairs
25
Q

Why reactivity decreases down group 7

A
  • halogens want to gain electrons in a reaction
  • shielding and atomic radius increase
  • it becomes harder to gain electrons
26
Q

Displacement redox reactions of halogens

A

A more reactive halogen displaces and oxidises a less reactive halide ion

27
Q

Why is cyclohexane added to identify halogens?

A
  • it is hard to tell the difference between halogens in aqueous solution
  • halogens are more soluble in organic solvent and the colour difference is clearer
28
Q

Chlorine in cyclohexane

A

Pale green (same as in water)

29
Q

Bromine in cyclohexane

A

Orange (same as in water)

30
Q

Iodine in cyclohexane

A

Purple (changed from brown in water)

31
Q

Disproportionation of chlorine with water

A
  • chlorine + water&raquo_space; hydrochloric acid + chloric(i) acid
  • Cl reduced from 0 to -1 in HCl
  • Cl oxidised from 0 to +1 in HClO
  • reaction is used to disinfect water
32
Q

Reduction of halogens with metal

A
  • metal + halogen&raquo_space; metal halide

- halogen reduced from 0 to -1 in the metal halide

33
Q

Disproportionation of chlorine with cold, dilute sodium hydroxide

A
  • chlorine + sodium hydroxide&raquo_space; sodium chloride + sodium chlorate(i) + water
  • Cl reduced from 0 to -1 in NaCl
  • Cl oxidised from 0 to +1 in NaClO
  • reaction is used to make bleach (NaClO)
34
Q

Disproportionation of chlorine with hot, concentrated sodium hydroxide

A
  • chlorine + sodium hydroxide&raquo_space; sodium chloride + sodium chlorate(v) + water
  • Cl reduced from 0 to -1 in NaCl
  • Cl oxidised from 0 to +5 in NaClO3
  • reaction is used to make weed killer (NaClO3)
35
Q

Oxidation number of sulfur in sulfuric acid

A

+6

36
Q

Oxidation number of sulfur in sulfur dioxide

A

+4

37
Q

Oxidation number of sulfur in hydrogen sulfide

A

-2

38
Q

Reactions of halide ions with sulfuric acid

A

Halide ions act as reducing agents for sulfur:

Chloride - sulfur dioxide produced
Bromide - sulfur dioxide and sulfur
Iodide - sulfur dioxide, sulfur, and hydrogen sulfide

Reducing power increases down the group

39
Q

Test for chloride ions

A

Add silver nitrate - white precipitate

Add dilute ammonia - solid dissolves

40
Q

Test for bromide ions

A

Add silver nitrate - cream precipitate
Add dilute ammonia - solid does not dissolve
Add concentrated ammonia - dissolves

41
Q

Test for iodide ions

A

Add silver nitrate - yellow precipitate
Add dilute ammonia - solid does not dissolve
Add concentrated ammonia - still does not dissolve

42
Q

Reactions of hydrogen halides

A
  • hydrogen halide + water&raquo_space; acidic solution

- hydrogen halide + ammonia&raquo_space; ammonium halide salt

43
Q

Test for carbonate / hydrogencarbonate ions

A
  • react with dilute acid
  • colourless gas (carbon dioxide) is produced
  • white precipitate when gas is bubbled through limewater
44
Q

Test for sulfate ions

A
  • react with acidified barium chloride

- white precipitate (barium sulfate) produced

45
Q

Test for ammonium ions

A
  • heat with dilute sodium hydroxide
  • colourless gas (ammonia) is produced
  • gas causes damp red litmus paper to turn blue

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