3.4 Atmosphere & Acids Flashcards
The pH scale
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
Methyl Orange indicator
Acidic- Red
Neutral- Yellow
Alkaline- Yellow
Phenolphthalein indicator
Acidic- Colourless
Neutral- Colourless
Alkaline- Pink
Red Litmus paper indicator
Acidic- Red
Neutral- (stays red)
Alkaline- Blue
Blue Litmus paper indicator
Acidic- Red
Neutral- (stays blue)
Alkaline- Blue
Universal indicator
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)
What four gases is the air composed of?
Nitrogen: 78%
Oxygen: 21%
Argon: 0.96%
Carbon dioxide: 0.04%
Measuring the % of oxygen
-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
Element combustion
Element + Oxide –> Element Oxide
Metal combustion
Metal –> Metal Oxide (Solid)
Non-metal combustion
Non-metal –> Non-metal Dioxide (Gas)
Magnesium combustion
-Magnesium + Oxygen → Magnesium oxide
-2Mg(s) + O2(g) → 2MgO(s)
-Observations: bright white light and white powder formed.
Sulphur combustion
-Sulphur combusts according to the following equation:
-Sulphur + Oxygen → Sulphur dioxide
-S(s) + O2(g) → SO2(g)
-Observations: blue flame.
Hydrogen combustion
-Hydrogen combusts according to the following equation:
-Hydrogen + Oxygen → Water
-2H2(g) + O2(g) → 2H2O(l)
-Observations: squeaky pop!
What is the test for oxygen?
-Glowing splint → relights
This is because the oxygen allows the wood in the splint to combust more efficiently.
Metal ions:
Iron(II) = ?
Iron(III) = ?
Copper = ?
Silver = ?
Zinc = ?
Lead = ?
Iron(II) = Fe2+
Iron(III) = Fe3+
Copper = Cu2+
Silver = Ag+
Zinc = Zn2+
Lead = Pb2+
Dissolving oxides
If they dissolve in water:
-Metal oxides form alkaline solutions
-Non-metal dioxides form acidic solutions
How could you determine whether an unknown element was a metal or non-metal?
- Combust it in oxygen
- Dissolve the oxide formed
- Check the pH of the solution with an indicator
Acid
H+ donor
Base
H+ acceptor
Alkali
OH- donor
What is the difference between acid/acidic and alkali/alkaline?
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
What happens when acids dissolve?
-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”
What happens when alkalis dissolve?
-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”
What happens when bases dissolve?
-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”:
Important acid examples?
Hydrochloric acid: HCl
Nitric acid: HNO3
Sulphuric acid: H2SO4
Phosphoric acid: H3PO4
Carbonic acid: H2CO3
Important alkali examples?
Sodium hydroxide: NaOH
Potassium hydroxide: KOH
Ammonium hydroxide: NH4OH
(all soluble metal hydroxides)
Important base examples?
Ammonia: NH3
(all metal oxides)
(all metal hydroxides)
Polyatomic ions
A polyatomic ion is an ion that contains more than one atom.
Acid + Base reaction
Acid + Base –> Salt + Water
e.g. Zinc oxide + Sulphuric acid –> Zinc sulphate + Water
What is an important observation of an acid + base reaction?
-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.
Making salt crystals from an acid + base reaction
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
Acid + Base soluble salt making
We make soluble salts by neutralising acids and crystallising the resulting solution:
Acid(aq) + Base(s)/Alkali(aq) → Salt(aq) + Water(l)
The insoluble base method
- Heat the acid.
Hot acid has more energy so it reacts
faster with the base. - 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. - 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
The titration method
- Perform a titration to measure the volumes of acid and alkali that neutralise each other.
- 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
Crystallisation method
- 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. - 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. - Filter out the crystals.
-Removes them from the excess salt solution in the basin. - Dry the crystals by dabbing them with
filter paper
-Removes the last traces of water from the crystals
Acid + Carbonate reaction
Acid + Carbonate → Salt + Water + Carbon dioxide
e.g. Lead carbonate + Sulphuric acid → Lead sulphate + Water + Carbon dioxide
What are important observations of an acid + carbonate reaction?
-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.
What is the test for carbon dioxide?
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
Precipitate
-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.
If we mixed KOH(aq) and Mg(NO3)2(aq):
Soluble: K+ and NO3-
Insoluble: Mg2+ and OH-
If an insoluble compound is formed, it will appear as a precipitate:
Mg(OH)2 ppt
Making salts by precipitation
- 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
Once the salt has been made by precipitation, we need to separate it from the solution
- Filter out the salt precipitate.
-Residue is the salt we want to keep
-Filtrate is the leftover solution. - Rinse the salt with water.
-washes off traces of the leftover solution.
-doesn’t dissolve our salt because it’s insoluble. - Dry the salt by dabbing it with filter paper.
-Removes the last few traces of water, leaving us with a pure dry salt
Solubility rules
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+)
