Chemistry EOY

Metals, chemical tests, halogens and the atmosphere

Alkali metal group

Found in group 1 as they only have 1 electron in their outermost shell

Physical properties of alkali metals

• Good conductors of heat and electricity

• Soft compared to other metals

• Low density compared to other metals

• Low melting/boiling points compared to other metals

Chemical properties of alkali metals

• Form ionic compounds with non-metals

• React violently with chlorine

• Burst into flames when heated with oxygen

• Produce soluble white compounds

• React with cold water

Pattern of reactivity in alkali metals as you go down

As outer electrons get further from the nucleus, there is less force of attraction making the electron more likely to bond with another atom, increasing reactivity

Patterns in alkali metals (aside from reactivity) as you go down

• Melting/boiling point decrease

• Density increases

• Softness increases

General equation for alkali metal reaction to water

Metal + water —> metal hydroxide + hydrogen

Observation of potassium and cold water + equation

• Reacts very violently; explodes as it glides across the water

• Hydrogen gas is produced and ignited by heat

• Water becomes alkali

• 2K₍ₛ₎ + 2H₂O₍ₗ₎ —>2KOH₍ₐq₎ + H₂₍₉₎

Observation for sodium and cold water + balanced equation

• Reacts violently, gliding across the water

• Hydrogen gas produced and MAY ignite and explode

• Water becomes alkali

• 2N + 2H₂O —>2NaOH + H₂

General equation: group 1 metal with chlorine

group 1 metal + chlorine —>metal chloride

Balanced equation for lithium and chlorine

2Li +Cl₂—>2LiCl

General equation: metal and heated oxygen (fire)

Metal + oxygen—> metal oxide

Balanced equation for potassium and oxygen

2K + O₂—>K₂O₂

Observation of magnesium and oxygen (fire)

• Starts silver shiny metal

• Bright white light of sparks during reaction

• Produces white powder of MgO

Observation of iron and oxygen (fire)

• Starts a dull and grey metal

• Iron glows hot, sparks, during reaction

• Produces rust-coloured FeO

General equation: metal with hydrochloric acid (HCl)

metal + hydrochloric acid—>salt + hydrogen gas

Observation of copper with HCl

No reaction at all

Observation of iron with HCl

• Very slow fizzing

• Hydrogen present (squeaky pop)

Observation of zinc with HCl

• Slow fizzing, few bubbles

• Hydrogen present (squeaky pop)

Observation of magnesium with HCl

• Exothermic reaction

• Fizzing

• Hydrogen produced (squeaky pop)

Displacement reactions

A more reactive metal will displace a less reactive metal from it’s aqueous solution of one of its salts

Displacement reactions of magnesium with metal sulphates

• Displaces CuSO₄ (black powder formed), ZnSO₄ (turned black), FeSO₄ (turned slightly black and dissolves slightly)

• The most reactive

Displacement reactions of copper with metal sulphates

• Displaces no metal sulphates

• The least reactive

Displacement reactions of zinc with metal sulphates

• Displaces CuSO₄ (turned black), FeSO₄ (solution turns colourless)

• The second most reactive

Displacement reactions of iron with metal sulphates

• Displaces CuSO₄ (rust formed)

• The third most reactive

Define an anion

An atom that has gained an extra electron and thus becomes more negatively charged

Define a cation

An atom that has lost an electron and thus becomes positively charged

Method for flame test

1. Dip platinum loop into solid sample and place on the edge of blue flames (must be a metal that has low reactivity and high melting point)

2. Record result

3. Clean loop with HCl

4. Repeated process

Flame test for cations: Li⁺

Flame becomes brick-red

Flame test for cations: Na⁺

Flame becomes yellow

Flame test for cations: K⁺

Flame turns lilac

Flame test for cations: Ca²⁺

Flame becomes orange-red

Flame test for cations: Cu²⁺

Flame turns green-blue

Sodium hydroxide test for cations method

1. Add NaOH (sodium hydroxide) into the solution

2. Record result (typically precipitate)

Sodium hydroxide test for cations: Cu²⁺

Blue precipitate of Cu(OH)₂ is formed

Sodium hydroxide test for cations: Fe²⁺

Green precipitate of Fe(OH)₂ is formed

Sodium hydroxide test for cations: Fe³⁺

Brown precipitate of Fe(OH)₃ is formed

Test for ammonium ions (NH⁴⁺) method

1. Add sodium hydroxide (NaOH) to ammonium chloride

2. Warm gently with Bunsen burner

3. Test gas produced using damp, red litmus paper

4. If ammonia gas is present, red litmus paper turns blue

Equation for ammonium chloride and sodium hydroxide (symbol)

NH₄Cl + NaOH—>NaCl +NH₃ + H₂O

Equation for ammonium chloride and sodium hydroxide (word)

Ammonium chloride + sodium hydroxide—>sodium chloride + ammonia gas + water

Method for testing halide ions

1. Remove carbonate ions with dilute nitric acid

2. Add silver nitrate (AgNO₃)

3. Observe precipitate formed

Test for halide ions: iodide (I⁻)

Yellow precipitate of AgI formed

Test for halide ions: bromide (Br⁻)

Cream precipitate of AgBr formed

Test for halide ions: chloride (Cl⁻)

White precipitate of AgCl formed

Method for testing sulphate ions (SO₄²-)

1. Add hydrochloric acid to remove carbonate ions

2. Add barium chloride (BaCl₂) to the solution

3. Sulphate ions present—>white precipitate formed

Method for testing carbonate ions (CO₃²⁻)

1. Add hydrochloric acid (HCl) to solution; if effervescence present, test for CO₂

2. To test for CO₂, run the gas through lime water

3. Carbonate ion—>solution turns cloudy/milky

Test for carbon dioxide (CO₂)

• CO₂ is a colourless, odourless gas

• If you bubble it through limewater and it turns cloudy/milky, CO₂ is present

Test for hydrogen (H₂)

• Hydrogen is a colourless, odourless gas

• When you place a lit splint near the mouth of the test tube and it makes a squeaky pop sound (small explosion), hydrogen is present

Testing for oxygen (O₂)

• Oxygen is a colourless and odourless gas

• By placing a glowing splint (barely lit) at the mouth of a test tube and it lights back up, oxygen is present

• Oxygen is needed to light fire

Testing for chlorine (Cl₂)

• Chlorine gas has a pale green colour and a choking smell

• Place damp blue litmus paper near the mouth of a test tube and if the paper turns white, chlorine gas is present

• When the litmus paper turns white, we call the paper “bleached”

Test for ammonia gas

• Ammonia is a colourless gas with a pungent smell

• Place damp, red litmus paper near the mouth of the test tube and if the litmus paper turns blue, ammonia gas is present

Test for water (chemical way)

• Pure copper sulphate is white when there is no water

• When copper (II) sulphate turns blue when the solution is poured on it, water is present

Define halogens

Elements in group 7 that have seven electrons in their outer shell. They are non-metals and consist of diatomic molecules

Physical observation of Fluorine (F)

• Gas at room temperature

• Yellow in colour

Physical observation of chlorine (Cl)

• Gas at room temperature

• Pale green in colour

• Green-blue in solution

Physical observation of bromine (Br)

• Liquid at room temperature

• Red-brown in colour

• Orange in solution

Physical observation of iodine (I)

• Solid at room temperature

• Black in colour

• Dark brown in solution

Patterns of halogens as you go down

• Melting/boiling point increases

• Colour becomes darker

• Reactivity decreases

Reactivity of halogens

As you go down group 7, the halogens become less reactive. This is because as the outer shell grows further away from the nucleus, it is harder to attract another electron

Gases in the atmosphere

• Nitrogen 78%

• Oxygen 21%

• Argon 0.9%

• Carbon dioxide 0.04%

Method for calculating oxygen using iron

1. Place iron fillings at the end of a burette

2. Use clamp to hold burette vertically in trough of water

3. Measure the initial height of water in the burette

4. Leave apparatus for weeks

5. Measure and note the final height of water level in the burette

Method of calculating oxygen using phosphorous

1. Add Phosphorus onto an evaporating dish and place it on a trough of water, making sure it is floating

2. Ignite Phosphorus using a candle

3. Cover this with a bell jar

4. Measure and note the starting height of the water level in the bell jar

5. Leave apparatus for several days

6. Measure and note the final height of the water level in the bell jar

Calculating the amount of oxygen used in rusting

• Volume of air at start = total burette volume - initial burette reading

• Volume of oxygen used = initial reading - final reading

• Percentage of oxygen = (volume of oxygen used ÷ volume of air at start) × 100

Define combustion

A chemical reaction between a substance and oxygen, producing an exothermic reaction with the substance, also known as burning

Observation of combusting magnesium

• Intense white flame

• White powder produced (magnesium oxide)

• 2Mg + O₂—>

Observation of combusting hydrogen

• Exothermic

• Water is produced

• 2H₂ + O₂—>2H₂O

Observation of combusting sulfur

• Blue flame

• Forms colourless, poisonous gas (sulphur dioxide: can lead to acid rain)

• S + O₂—>SO₂

Define thermal decomposition

The process in which heat (thermal energy) is being used to break chemical substances down

General equation for thermal decomposition

Metal carbonate—>metal oxide + carbon dioxide

Symbol equation for copper thermal decomposition

CuCO₃—>CuO + CO₂

The difference between global warming and climate change

• Global warming is the result of human activities warming the earth; it is the cause of climate change

• Climate change is the effect of global warming including change of weather, winds, etc.

Effect of greenhouse gases

Greenhouse gases absorbs infrared radiation (sun rays and warms up the atmosphere. At regular amounts, they are essential for life but if the concentration is too high, it leads to global warming.

Examples of greenhouse gases

• Methane

• Nitrous oxide

• Chlorofluorocarbons

• Carbon dioxide

Problems resulting from global warming

• Ice sheets melting—>sea levels rise, increase flood risk for coastal countries

• Warmer atmosphere—>extreme weather

• Change in climate—>crops and animals are harder to raise

• Increase in smog—>heart disease, lung cancer, asthma

Define rusting

When iron/steel corrodes. The rust is essentially just hydrated iron oxide (see equations below)

• Iron + oxygen + water—>hydrated iron (III) oxide

• 4Fe + 3O₂ + 𝑥H₂O—>2Fe₂O₃ • 𝑥H₂O

Rust prevention: barrier

• By coating iron with a barrier, it prevents iron from coming into contact of water and oxygen, avoiding rust

• Common barrier methods include paint, oil, grease, and plastic

• However, if the coating is washed away, the iron will rust

Rust prevention: galvanising/sacrificial

• Iron is coated in a layer of zinc, ZnCO₃

• Since zinc is more reactive than iron, it will react with the oxygen before it can reach the iron, thus preventing rust

• If the coating of zinc is scratched, iron is still protected from rusting by the sacrificial method