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chemistry > Chapter 8 (Approved) > Flashcards

Chapter 8 (Approved) Flashcards

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AP® Chemistry flashcards Periodic Table flashcards
1
Q

Where are alkali metals found

A

In stable compounds as they don’t occur in their elemental form naturally

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2
Q

Most common reaction for group 2 elements

A

Redox reaction

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3
Q

Group 2 elements: oxidising or reducing agents, and why

A

Reducing agents, as they cause reduction (gain of electrons) in the other species

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4
Q

How do group 2 elements react with oxygen?

A

Forms a metal oxide with general formula MO, made up of M2+ and O2- ions

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5
Q

How do group 2 elements react with water?

A

To form an alkaline hydroxide with the general formula M(OH)2, and hydrogen gas
Water reacts very slowly with magnesium but the reaction becomes more vigorous with metals further down the group

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6
Q

Group 2 elements trend in reactivity

A

Increases down the group

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7
Q

Group 2 strontium and water reaction balanced equation

A

Sr(s) + 2H20(l) ——> Sr(OH)2(aq) + H2(g)

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8
Q

What does the reaction between metals and dilute acids form and what type of reaction is it?

A

Metal + acid —-> salt + hydrogen
Redox reaction

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9
Q

Group 2 Magnesium + dilute hydrochloric acid reaction balanced equation

A

Mg(s) + 2HCl(aq) —-> MgCl2(aq) + H2 (g)

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10
Q

Why does reactivity increase down group 2

A

As you go down the group, the atomic radius increases and electron shielding increases due to more inner shells, which reduces the electrostatic attraction between nuclei and outer shell electrons, decreasing ionisation energies as less energy is needed to remove the outer electron, making it more reactive

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11
Q

What does the reaction between group 2 metals and water form

A

Alkaline metal + water ——> Alkaline hydroxide and hydrogen gas

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12
Q

What do Group 2 oxides and water form? (reaction)

A

Group 2 oxide + water —-> group 2 cations and hydroxide ions in solution

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13
Q

What happens when group 2 oxides react with water? (explanation)

A

Hydroxide ions and group 2 metal ions are released in solution, but group 2 hydroxides are only slightly soluble in water, so the solution becomes saturated, and further metal and hydroxide ions will react to form a solid precipitate of group 2 hydroxide

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14
Q

Reaction between group 2 calcium ion and hydroxide ions balanced equation

A

Ca2+(aq) + 2OH-(aq) —-> Ca(OH)2(s)

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15
Q

Reaction between group 2 calcium oxide and water balanced equation

A

CaO(s) + H2O(l) ——> Ca2+(aq) + 2OH-(aq)

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16
Q

What is the trend in pH for group 2 metal hydroxides and why?

A

pH increases down the group, as the hydroxides get more soluble down the group, their solutions contain more OH- ions and are therefore more alkaline

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17
Q

Experiment to show solubility, pH and alkalinity of group 2 hydroxides increases down the group (3)

A
  1. Add a spatula of each group 2 oxide to water in a test tube
  2. Shake the mixture, there won’t be enough water to dissolve all the metal hydroxide that forms, so you will have a saturated solution of each metal hydroxide with some white solid undissolved at the bottom of each test tube.
  3. Measure the pH of each solution, the alkalinity will increase down the group.
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18
Q

Uses of group 2 compounds

A

In agriculture: calcium hydroxide Ca(OH)2 is added to fields to increase the pH of acidic soils. It neutralisés the acid, forming neutral water

In medicine: Antacids for treating acid indigestion. Magnesium hydroxide or calcium carbonate are often used to neutralise stomach acids

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19
Q

Reaction of calcium hydroxide in agriculture ionic equation

A

Ca(OH)2(s) + 2H+ ——> Ca2+(aq) + 2H2O(l)

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20
Q

Magnesium hydroxide and calcium carbonate neutralising stomach acid balanced equations

A

Mg(OH)2(s) +2HCl(aq) —> MgCl2(aq) + 2H20(l)

CaCo3(s) + 2HCl(aq) —-> CaCl2(aq) + H20(l) + CO2(g)

21
Q

Fluorine, Chlorine, bromine, iodine appearance and state at RTP

A

F2: Pale yellow gas
Cl2: Pale green gas
Br2: Red-brown liquid
I2: Shiny grey-black solid

22
Q

Why does boiling point increase down group 7

A

More electrons down the group, so more London forces, so more energy is required to break them

23
Q

Halogen most common reaction type

A

Redox reaction

24
Q

Are halogens reducing or oxidising agents

A

They are oxidising agents as they cause the other species to lose electrons

25
Q

Halogen halide displacement reactions (3)

A

Halogens can displace less reactive halides from their compounds.

If displacement fails, there is no colour change and the halogen was less reactive.

If the displacement happens, then the halogen was more reactive and the solution will take on the colour of the displaced halide.

26
Q

What is the colour of chlorine, bromine and iodine in water, and how can you tell them apart?

A

In water chlorine is pale green, bromine is orange and iodine is brown, but depending on the concentration iodine and bromine may be difficult to tell apart.

To tell the apart, cyclohexane can be added and the mixture shaken. In cyclohexane, chlorine is pale green, bromine is orange and iodine violet.

27
Q

Reaction of chlorine with bromide ions (in sodium compound) balanced equation

A

Cl2(aq) + 2NaBr(aq) —> 2NaCl(aq) + Br2(aq)

28
Q

Halogen trend in reactivity and why?

A

Halogen reactivity decreases down the group as atomic radius increases and electron shielding increases, weakening electrostatic attraction between nucleus and outer shell electrons

29
Q

Disproportionation

A

A redox reaction in which the same element gets oxidised and reduced

30
Q

Chlorine and water balanced equation

A

Cl2(aq) + H20(l) —-> HClO(aq) + HCl(aq)

31
Q

What is the use of chloric acid and chlorate ions?

A

Chloric acid and chlorate ions kill bacteria, and chloric acid acts a a weak bleach

32
Q

How can you demonstrate that chloric acid acts as a weak bleach?

A

Add some indicator solution to a solution of chlorine in water, the indicator first turns red due to the acids then the colour disappears due to the bleach

33
Q

Reaction of chlorine with cold dilute aqueous sodium hydroxide balanced equation

A

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

34
Q

Benefits of chlorine use (1)

A

Ensuring water is fit to drink by killing bacteria that causes diseases like typhoid and cholera

35
Q

Risks of chlorine use (2)

A

Chlorine is an extremely toxic gas.
In drinking water it can react with organic hydrocarbons and form chlorinated hydrocarbons which are suspected of causing cancer

36
Q

Carbonate test

A
  1. In a test tube, add dilute nitric acid to the solid or solution to be tested
  2. If you see bubbles , the unknown compound could be a carbonate
  3. To prove that the gas formed is carbon dioxide, bubble the gas through limewater / a saturated solution of calcium hydroxide
  4. Carbon dioxide would react to form a white precipitate, calcium carbonate, which would turn the solution cloudy
37
Q

Why can barium sulfate be used for the sulfate test but not other sulfates?

A

Because unlike other sulfates, barium sulfate is insoluble in water

38
Q

Barium ion and sulfate ion ionic equation

A

Ba2+(aq) + SO4(2-)(aq) —-> BaSO4(s)

39
Q

Why should barium nitrate be used instead of barium chloride of the Halide test is being carried out afterwards?

A

Because you are introducing chloride ions to the solution

40
Q

Why are silver halides used for the halide test?

A

Because unlike most halides, silver halides are insoluble

41
Q

Halide test

A
  1. Add aqueous silver nitrate AgNO3 to an aqueous solution fo a halide
  2. The silver halide precipitates are different colours silver chloride is white, silver bromide is cream and silver iodide is yellow
  3. Add aqueous ammonia to tell the three similar colours apart
42
Q

How to tell the difference between the silver halides with aqueous ammonia

A

Silver chloride dissolves in dilute ammonia, silver bromide dissolves in concentrated ammonia and silver iodide is insoluble in concentrated ammonia

43
Q

Why is the carbonate test done first?

A

Because it also produces a white precipitate, which is the marker for the sulfate test, so you must ensure the unknown compound isn’t a carbonate first

44
Q

Why must the halide test be done last?

A

Because silver carbonate and silver sulfate are both insoluble in water and will form as white precipitates, so the other tests must be carried out before to rule out carbonates and sulfates as possibilities

45
Q

Carrying out the tests on the same solution (3)

A

If you see bubbles after the carbonate tests, keep adding dilute nitric acid (sulfuric or hydrochloric will show up in next tests) until bubbling stops so there will be no more carbonate to react in next tests

Add barium nitrate (barium chloride will show up in halide test) to the solution left, any sulfate ions will become a barium sulfate precipitate which you then filter out

Add aqueous silver nitrate to the precipitate formed. Since carbonate or sulfates ions have been removed, any precipitate formed must contain halide ions. Add aqueous ammonia to confirm which halide you have.

46
Q

What happens when aqueous ammonium ions and aqueous hydroxide ions are heated together?

A

They react to form ammonia gas and water

47
Q

aqueous ammonium ions and aqueous hydroxide ions balanced equation

A

NH4+(aq) + OH-(aq) —> NH3(g) +H20(l)

48
Q

Ammonium ion test

A
  1. Aqueous sodium hydroxide, NaOH, is added to a solution of an ammonium ion
  2. Ammonia gas is produced, but the gas bubbles are unlikely to be seen as ammonia is very soluble in water
  3. The mixture is warmed and ammonia gas is released
  4. You might be able to smell the ammonia, or test it with a moist pH indicator, which should turn blue

chemistry (12 decks)

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