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Chemistry > 3.4 Atmosphere & Acids > Flashcards

3.4 Atmosphere & Acids Flashcards

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1
Q

The pH scale

A

The pH scale is used to describe how acidic or alkaline a solution is:
In general:
* pH < 7 is acidic
* pH = 7 is neutral
* pH > 7 is alkaline

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2
Q

Methyl Orange indicator

A

Acidic- Red
Neutral- Yellow
Alkaline- Yellow

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3
Q

Phenolphthalein indicator

A

Acidic- Colourless
Neutral- Colourless
Alkaline- Pink

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4
Q

Red Litmus paper indicator

A

Acidic- Red
Neutral- (stays red)
Alkaline- Blue

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5
Q

Blue Litmus paper indicator

A

Acidic- Red
Neutral- (stays blue)
Alkaline- Blue

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6
Q

Universal indicator

A

pH 0-2= Red (Strong acid)
pH 3-4= Orange (Weak acid)
pH 5-6= Yellow (Weak acid)
pH 7= Neutral
pH 8-10= Light Blue (Weak alkali)
pH 11-13= Dark Blue (11- Weak alkali) (12-14 Strong alkali)
pH 14= Purple (Strong alkali)

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7
Q

What four gases is the air composed of?

A

Nitrogen: 78%
Oxygen: 21%
Argon: 0.96%
Carbon dioxide: 0.04%

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8
Q

Measuring the % of oxygen

A

-Reacting it with another element, often a metal
-Air gradually decreases in volume until all of the oxygen has been added to the metal
Results for this experiment might look like this:
-Initial volume of air: 100cm3
-Final volume of air: 79cm3
-Decrease in volume 100 − 79 = 21
So the percentage of O2 in the original air is:
% O2 = 21 cm3/100 cm3 × 100 = 21%
-Air nearly always contains 21% O2: so whatever the initial volume, it should decrease by 21%.
-For example: 147 cm3 of air should decrease (by 30.87 cm3) to 116.13 cm3
-If the air in one of these experiments didn’t decrease by 21%, it could be because:
* The metal wasn’t heated for long enough for all the oxygen to be reacted
* There’s a leak in the apparatus

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9
Q

Element combustion

A

Element + Oxide –> Element Oxide

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10
Q

Metal combustion

A

Metal –> Metal Oxide (Solid)

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11
Q

Non-metal combustion

A

Non-metal –> Non-metal Dioxide (Gas)

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12
Q

Magnesium combustion

A

-Magnesium + Oxygen → Magnesium oxide
-2Mg(s) + O2(g) → 2MgO(s)
-Observations: bright white light and white powder formed.

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13
Q

Sulphur combustion

A

-Sulphur combusts according to the following equation:
-Sulphur + Oxygen → Sulphur dioxide
-S(s) + O2(g) → SO2(g)
-Observations: blue flame.

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14
Q

Hydrogen combustion

A

-Hydrogen combusts according to the following equation:
-Hydrogen + Oxygen → Water
-2H2(g) + O2(g) → 2H2O(l)
-Observations: squeaky pop!

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15
Q

What is the test for oxygen?

A

-Glowing splint → relights
This is because the oxygen allows the wood in the splint to combust more efficiently.

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16
Q

Metal ions:
Iron(II) = ?
Iron(III) = ?
Copper = ?
Silver = ?
Zinc = ?
Lead = ?

A

Iron(II) = Fe2+
Iron(III) = Fe3+
Copper = Cu2+
Silver = Ag+
Zinc = Zn2+
Lead = Pb2+

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17
Q

Dissolving oxides

A

If they dissolve in water:
-Metal oxides form alkaline solutions
-Non-metal dioxides form acidic solutions

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18
Q

How could you determine whether an unknown element was a metal or non-metal?

A
  1. Combust it in oxygen
  2. Dissolve the oxide formed
  3. Check the pH of the solution with an indicator
19
Q

Acid

A

H+ donor

20
Q

Base

A

H+ acceptor

21
Q

Alkali

A

OH- donor

22
Q

What is the difference between acid/acidic and alkali/alkaline?

A

It’s best to think of the words as:
* Acid or alkali (or base) describe a substance in terms of ions
* Acidic or alkaline describe a solution in terms of pH

23
Q

What happens when acids dissolve?

A

-Acids dissolve
-Then donate (release) H+ ions into the water.
-This is what causes the solution to have a pH < 7 and for it to be labelled “acidic”

24
Q

What happens when alkalis dissolve?

A

-Alkalis dissolve
-Then donate (release) OH− ions into the water.
-This is what causes the solution to have a pH > 7 and for it to be labelled “alkaline”

25
Q

What happens when bases dissolve?

A

-Bases dissolve
-Then accept (take) H+ ions from the H2O molecules.
-But when an H2O molecule loses an H+
ion, what’s leftover is an OH− ion.
-It’s these OH− ions that cause the solution to have a pH > 7 and for it to be labelled “alkaline”:

26
Q

Important acid examples?

A

Hydrochloric acid: HCl
Nitric acid: HNO3
Sulphuric acid: H2SO4
Phosphoric acid: H3PO4
Carbonic acid: H2CO3

27
Q

Important alkali examples?

A

Sodium hydroxide: NaOH
Potassium hydroxide: KOH
Ammonium hydroxide: NH4OH
(all soluble metal hydroxides)

28
Q

Important base examples?

A

Ammonia: NH3
(all metal oxides)
(all metal hydroxides)

29
Q

Polyatomic ions

A

A polyatomic ion is an ion that contains more than one atom.

30
Q

Acid + Base reaction

A

Acid + Base –> Salt + Water
e.g. Zinc oxide + Sulphuric acid –> Zinc sulphate + Water

31
Q

What is an important observation of an acid + base reaction?

A

-The base disappears
-This is because it is used up in the reaction.
-If the salt is soluble, it dissolves into the water to form a salt solution.

32
Q

Making salt crystals from an acid + base reaction

A

e.g. Zinc sulphate crystals
-Add excess base (ZnO) to hot acid (H2SO4), makes the salt solution (ZnSO4) quickly, ensures all the acid is used up.
-Filter out the excess base (ZnO), salt solution (ZnSO4) is the filtrate and is
collected in an evaporating basin.
-Heat the salt solution (ZnSO4) until crystals
start to form around the edges, saturates the solution so that crystals
form most efficiently.
-Leave the salt solution to crystallise, evaporation of water allows crystals to form

33
Q

Acid + Base soluble salt making

A

We make soluble salts by neutralising acids and crystallising the resulting solution:
Acid(aq) + Base(s)/Alkali(aq) → Salt(aq) + Water(l)

34
Q

The insoluble base method

A
  1. Heat the acid.
    Hot acid has more energy so it reacts
    faster with the base.
  2. While stirring, add base until no more
    will dissolve.
    This guarantees the base is in excess so
    that we know all the acid has been used
    up.
  3. Filter out the excess base.
    (Because you don’t want it.)
    -ACID + BASE –> SALT + WATER
    e.g. CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l
35
Q

The titration method

A
  1. Perform a titration to measure the volumes of acid and alkali that neutralise each other.
  2. Repeat the titration, but this time:
    a. Don’t put the indicator in.
    b. Use the burette to add exactly the right volume (e.g. 21.05 cm3) for neutralisation.
    -ACID + ALKALI –> SALT + WATER
    e.g. HCl(aq) + KOH(aq) → KCl(aq) + H2O(l
36
Q

Crystallisation method

A
  1. Heat the salt solution until crystals just
    start forming.
    -You can tell by looking at the edges, or by occasionally dipping in a glass rod.
    -Causes the solution to become saturated.
  2. Allow the solution to cool in an evaporating basin for a few days.
    -Lowers the solubility of the salt so that lots of solid crystals form.
  3. Filter out the crystals.
    -Removes them from the excess salt solution in the basin.
  4. Dry the crystals by dabbing them with
    filter paper
    -Removes the last traces of water from the crystals
37
Q

Acid + Carbonate reaction

A

Acid + Carbonate → Salt + Water + Carbon dioxide
e.g. Lead carbonate + Sulphuric acid → Lead sulphate + Water + Carbon dioxide

38
Q

What are important observations of an acid + carbonate reaction?

A

-The carbonate disappears because it is used up in the reaction.
-There is fizzing because a gas (carbon dioxide) is produced.
-If the salt is soluble, it dissolves into the water to form a salt solution.

39
Q

What is the test for carbon dioxide?

A

The chemical test for carbon dioxide (CO2) is:
* Limewater → turns cloudy
This is because the CO2 reacts with the limewater to make an insoluble white solid

40
Q

Precipitate

A

-A precipitate (ppt) is an insoluble solid that forms inside a solution.
-Ions in the solution collide with each other, a ppt forms and (slowly) sinks to the bottom.
-e.g. when Ag+ and Cl− ions collide in solution
-When mixing two solutions, you can tell if a precipitate will form by mixing the ions up and
seeing if either of the resulting compounds is insoluble.

41
Q

If we mixed KOH(aq) and Mg(NO3)2(aq):

A

Soluble: K+ and NO3-
Insoluble: Mg2+ and OH-
If an insoluble compound is formed, it will appear as a precipitate:
Mg(OH)2 ppt

42
Q

Making salts by precipitation

A
  • Insoluble salts made by mixing solutions to create precipitate
  • Each solution contains one of the ions needed
  • Example: Making BaSO₄ precipitate with BaCl₂(aq) (Ba²⁺) and Na₂SO₄(aq) (SO₄²⁻)
  • Fully balanced equation: BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
  • Cl⁻ ions (from BaCl₂) and Na⁺ ions (from Na₂SO₄) are spectator ions
  • Spectator ions remain in solution as “leftovers”
  • Simpler equation involving only the precipitate: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)
  • This is termed an ionic equation
43
Q

Once the salt has been made by precipitation, we need to separate it from the solution

A
  1. Filter out the salt precipitate.
    -Residue is the salt we want to keep
    -Filtrate is the leftover solution.
  2. Rinse the salt with water.
    -washes off traces of the leftover solution.
    -doesn’t dissolve our salt because it’s insoluble.
  3. Dry the salt by dabbing it with filter paper.
    -Removes the last few traces of water, leaving us with a pure dry salt
44
Q

Solubility rules

A

SOLUBLE: Sodium, Potassium, Ammonium, Nitrate
Chloride (except with Ag+ and Pb2+)
Sulphate (except with Ca2+, Ba2+ and Pb2+)
INSOLUBLE: Carbonate (except with Na+, K+, and NH4+)
Hydroxide (except with Na+, K+ and Ca2+)
Oxide (except with Na+, K+, and NH4+)

Chemistry (12 decks)

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