IGCSE Cards for pnwa

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UPDATED TO 2023-2025 SYLLABUS

EDEXCEL IGCSE

CHEMISTRY

SUMMARIZED NOTES ON THE THEORY SYLLABUS

Prepared for Ananthiya for personal use only.

EDEXCEL IGCSE CHEMISTRY

1. State of Matter

1.1. Solid, Liquid, and Gases

Properties	Solid	Liquid	Gas
Particle arrangement
Forces between particles	Strongest	Weaker than solids	Weakest
Motion of particles	Vibrate in fixed positions and regular arrangement	Slip and slide over other particles	Move randomly at high speed in all directions
Shape	Fixed shape	Take the shape of the container	Takes the shape of the container
Expansion when heated	Expands slightly	Expands more than solids	Expands the most
Can it be compressed?	No	Slightly	Yes

1.2. Changes in the State of Matter

• Melting point: when a solid melts

• Boiling point: when a liquid turns to a gas

• Freezing point: when the liquid changes to a solid

• Boiling: When a liquid is heated so strongly that all the particles move fast and can overcome the forces of attraction between them.

• Evaporation: When some fast-moving particles at the surface of the liquid have enough energy to change into a gas.

• Sublimation: Solid to gas

• Deposition: Gas to solid

1.3. Diffusion

• Diffusion: The spreading out of particles from where they are at a higher concentration to a lower concentration.

• Diffusion in liquids is very slow if the liquid is very still

1.4. Solutes, Solvents, and Solutions

• Solute: The substance that dissolves in the liquid

• Solvent: The liquid the solid dissolves in

• Solution: The liquid formed

• Saturated Solution: A solution that contains as much dissolved solid as possible at a particular temperature

• Solubility: The mass of solid that must dissolve in 100g of solution at that temperature to form a saturated solution

2. Elements, Compounds and Mixtures

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• Elements: Substances that can’t be split into anything simpler by chemical means. It contains only 1 type of atom

• Compounds: Contains Two or more elements that are chemically combined together

• Mixture: Various substances mixed with no chemical reaction that occurs

• Pure substances: Melt and boil at fixed temperatures

• Mixtures: boil over a range of temperatures

2.1. Methods to Separate Mixtures

3. Atomic Structure

An Atom is the smallest piece of an element that can be still known as that element

Molecule - Two or more atoms chemically combined together

• Atomic Number - Number of protons in the nucleus of an atom

• Mass Number - The number of protons and neutrons in the nucleus of an atom

• Isotopes - Atoms of the same element that have the same atomic number but a different mass number. (same number of protons, different number of neutrons)

• Relative atomic mass - Average mass of an atom. Takes into account the isotopes naturally occurring in a sample present.

• Relative atomic mass = [(% of atom naturally occurring x atomic number) + (% of atom naturally occurring x atomic number) / 100 ]

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4. The Periodic Table

• Vertical Columns are called groups

• Horizontal rows are called periods

• The elements are arranged in increasing order of atomic number in the periodic table

• Electrons are arranged around the nucleus in shellsand lower ones are filled before the higher ones

• In the first shell, a max of 2 electrons can fit, in the second a max of 8 electrons can fit and in the third shell a max of 8 electrons can fit (only for the first 20 elements - till Calcium)

• The arrangement of electrons in its shell is called its electronic configuration

• Elements in the same group of the periodic table have the same number of electrons in their outermost shell

• The period number gives the number of occupied shells

• Groups in the periodic table have similar chemical properties

• Noble Gases: These occupy group 0, also called group 8, as they contain 8 electrons in their outermost shell so are extremely inert and unreactive

4.1. Properties of the Elements

5. Chemical Formulae, Equations and Calculations - Part 1

“2H22​O” The larger “2” Shows 2 moles of “H2OH2​O”, whereas the smaller subscript 2 shows 2 atoms of Hydrogen

• Relative Atomic Mass (Ar)(Ar​) : The relative atomic mass is the weighted average of the isotopes of the element and is measure on a Carbon-12 scale where each carbon-12 atom has a mass of exactly 12.

• Relative Formula Mass (Mr)(Mr​) : Measure the masses of compounds using the Carbon-12 scale.

• To find the percentage of an element in a compound, find the Mr of the compound and the Ar of the element and divide them and multiply the answer by 100

  (Ar/Mr)∗100(Ar​/Mr​)∗100

1 Mol = 6.022 * 10231023 atoms

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• Empirical Formula: Shows the simplest whole number ratio of atoms present in a compound

• Molecular Formula: Shows the actual number of atoms of each element present in the compound

The sample table above shows the table method that can be used to calculate the empirical formula

Percentage yield: (Actual Yield) / (Theoretical Yield) * 100

Calculating excess: Find the ratio of the compounds by looking at the formula for example - Na2CO3(s)+2HClNa2​CO3​(s)+2HCl(l) shows 2 mols of HCl and 1 mol of Na2CO3Na2​CO3​ so the molar ratio is 2:1. Using the data given we can calculate the moles of both and whichever is in excess can then be easily identified.

5.1. Chemical Formulae, Equations and Calculations - Part 2

Avogadro’s Law - Equal Volumes of all gasses at the same temperature and pressure will have the same number of molecules.

IMPORTANT FORMULAS -

• For a solid: n=m/Mrn=m/Mr

• For aqueous solutions: n=cVn=cV

• For a gas: n=V/Mn=V/M

1 litre = 1dm31dm3 = 1000cm31000cm3

6. Ionic Bonding

Ionic Bonding is the electrostatic force of attraction between the oppositely charged ions (positive and negative ions)

Ionic compounds usually contain a metal

• Positive ions are called Cations

• Negative ions are called Anions

Elements in groups 1,2,3 will form 1+, 2+, 3+ ions. Whereas elements in groups 5,6,7 will form 3-,2-,1- ions as they gain electrons to become stable

6.1. Finding the Formula of the Compound

Just cross the charges of the elements over and the Charge for Chlorine (Cl−Cl−) is 1 itself so there is no extra number added.

Whereas if the charges are the same, there is no need to cross the charges as they are the same so become balanced.

6.2. Giant Ionic Structures

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• A lattice is a regular array of particles and each lattice is held together by the electrostatic force of attraction between the oppositely charged ions.

• In Magnesium Oxide, the 2+ ions and 2- ions are being attracted, whereas the 1+ and 1- charge in Sodium Chloride and so due to the 2 ions being transferred, magnesium oxide has stronger electrostatic forces of attraction.

   

Ionic Compounds:

• Have high melting and boiling points because of the strong electrostatic forces of attraction holding the lattice together so a lot of thermal energy has to be supplied to break these strong bonds.

• Tend to be crystalline

• Tend to be brittle as any distortion of the structure can cause like charges to come together and as like charges repel, the crystal splits itself apart.

• Tend to be soluble in water

• Tend to be insoluble in organic solvents

• Don’t conduct electricity when in the solid state as the ions aren’t free to move but can conduct electricity in molten and aqueous.

7. Covalent Bonding

A covalent bond is the electrostatic force of attraction between the nuclei of the atom and shared pair of electrons.

7.1. Diatomic Molecules -

• A Hydrogen Molecule consists of diatomic atoms hence hydrogen has the formula H2H2​

• Hydrogen forms molecules as the H2H2​ molecule is more stable than two separate hydrogen atoms

Examples of diatomic molecules would be Chlorine, Fluorine and the Halogens, Hydrogen, Oxygen, Nitrogen

7.2. Inorganic Molecules -

• Methane -

Carbon has 4 electrons in its outermost shell so would share 1 each with the 4 hydrogen atoms and form Methane - CH4CH4​

• Ammonia -

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Nitrogen has 5 electrons in its outermost shell so would share 1 each amongst the 3 hydrogen atoms and form ammonia - NH3NH3​

7.3. Organic Molecules -

The structure is for Ethane and consists of Carbon-Hydrogen and Carbon-Carbon bonds and is called a hydrocarbon. (These structures will be discussed in more detail in Unit 4.)

• Oxygen has a double Covalent bond as 4 electrons are shared

• Nitrogen has a triple covalent bond as 6 electrons are shared

7.4. Simple Molecular Structures

• Contain Intermolecular forces that are weaker than the covalent bonds

• These substances tend to be solids with low melting and boiling points as not a lot of thermal energy is needed to break these weak intermolecular forces of attraction

• Intermolecular forces of attraction increase as the relative molecular mass increases

• Example - H2OH2​O, CO2CO2​, CH4CH4​, NH3NH3​, C2H4C2​H4​

7.5. Giant Covalent Structures

Diamond -

• Pure Carbon form

• Tetrahedral Structure

• Each carbon bonds strongly to 4 other carbons

• High melting and boiling points as the strong covalent bonds require a lot of thermal energy to break

• Very Hard as a lot of energy has to be applied of break the strong covalent bonds

• Doesn’t conduct electricity as the electrons in the outer shells are held tightly and not free to move around

Graphite -

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• Form of Carbon

• Layer structure

• Soft material as even though it has covalent bonds, only the weak forces between the layers have to be broken so can easily be separated and flaked off

• Has high melting and boiling points as to melt or boil the structure all bonds have to be broken including the covalent bonds which are very strong and require large amounts of thermal energy to break

• Conducts electricity as each carbon atom is only joined to 3 others so the 4th electron of the atom in each shell are free to move and can carry the charge (delocalised electrons) allowing graphite to conduct electricity

   

Fullerene -

• Lower melting and boiling points than diamond and graphite as only weak intermolecular forces need to be broken

• Not as hard as diamond as it doesn’t take a lot of energy to break the weak intermolecular forces

• Doesn’t conduct electricity as even though the 4 electron is free to move it can only move around within the atom and can’t jump from atom to atom so can’t carry the charge or conduct electricity

• Non-toxic so used in the body for medicines

   

Allotropes: Different forms of the same element

8. Metallic Bonding

Electrostatic force of attraction between the lattice of positive ions and sea of delocalised electrons \n

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• Metals have giant structures, so don’t contain individual molecules

• Metals are hard and have high melting and boiling points - Due to the strong electrostatic forces of attraction between the lattice of positive ions and sea of delocalised electrons

• Conduct electricity as the delocalised electrons are free to move throughout the structure

• Malleable and Ductile as when a force is applied the layers of positive ions slide over each other

9. Electrolysis

NOTE - PAPER 2 CHAPTER

9.1. Electrolysis of Ionic Compounds

Molten compounds undergo electrolysis and always produce their respective elements.

Lead (II) Bromide is an ionic compound and consists of lead (II) ions and Bromine ions packed together. When heated they melt and the ions become free to move.

The lead ions are attracted to the cathode and pick up two ions and become lead atoms and fall to the bottom forming molten lead.

The bromine ions are attracted to the anode and lose an electron forming a bromine atom.

Half-ionic equations are used to show what happens at each electrode.

At the cathode: Pb2+Pb2+ + 2e−2e−   ⟹  ⟹ Pb

At the anode:  2Br−2Br−   ⟹  ⟹ Br2Br2​ + 2e−2e−  (don’t forget that Bromine is di-atomic)

The discharge of ions means that ions are losing their charge.

9.2. Electrolysis of Aqueous Solutions

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Aqueous solutions will always have water present and some split up into hydrogen and hydroxide ions, H+ and OH- and participate in electrolysis reactions.

• At the positive electrode, the OH- ions or the non-metal ions are discharged and lose electrons or gain oxygen (oxidised)

• At the negative electrode, the H+ ions or the metallic ions are discharged but only one would gain electrons or lose oxygen (reduction)

• Aqueous Sodium Chloride solution contains Na+ and H+ ions which are attracted to the cathode.

• Hydrogen is less reactive than sodium so we can more easily add an electron to hydrogen ions to form a hydrogen molecule.

   

Hydroxide ions present in the solution due to the water splitting make the solution alkaline \n 2H+2H+(aq) + 2e−2e−   ⟹  ⟹H2H2​ (g)

At the anode, Cl- and OH- ions are present but as there are many more chloride ions these are oxidised and form Cl2Cl2​

2Cl−2Cl−(aq) - 2e−2e−   ⟹  ⟹ Cl22​(g)

The remaining solution now consists of Na+Na+ and OH−OH−ions. So we are left with Sodium Hydroxide (NaOH)

We get less chlorine than we expect as chlorine is more soluble in water, reducing its yield

During the electrolysis of Dilute Sulfuric acid, twice as much hydrogen is produced

2H+2H+ (aq) + 2e−2e−   ⟹  ⟹ H2(g)

4OH−4OH− (aq)  ⟹  ⟹ 2H2O2H2​O (l) + O2O2​ + 4e−4e−

If you look at the equations above you would see that 2 electrons produce 1 mol of hydrogen meaning 4 electrons would produce 2 mols of hydrogen.

In the second, 4 electrons produce 1 mol of oxygen

This means that twice as much hydrogen is producedcompared with oxygen

9.3. Oxidation and Reduction

• Reduction is the gain of electrons or the loss of oxygen.

• Oxidation is the loss of electrons or the gain of oxygen.

OILRIG: Oxidation is loss (of electrons), Reduction is gain (of electrons)

• Reduction always occurs at the cathode while oxidation occurs at the anode.

• Not all ionic compounds can be electrolysed as they break up into similar chemical compounds before their boiling point making it impossible to melt.

CCRG: Cathode, Cation, Reduction (is) Gain of electrons, Reducing Agent

AAOL: Anode, Anion, Oxidation (is) Loss of electrons, Oxidising Agent

9.4. Reactivity Series

• If the metal is above hydrogen in the reactivity series, you would get Hydrogen produced at the cathode eg: Potassium would produce Hydrogen at the cathode

• If the metal is below hydrogen, you would get the metal at the cathode, eg: Gold would produce gold itself at the cathode.

• If you have halides, you get that at the anode but any other negative ions would produce oxygen at the anode.

IMPORTANT IONS AND CHARGES -

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9.5. Experiments

Electrolysis of aqueous NaCl (Sodium Chloride) solution

• Take a glass tube, close it with a rubber bung, place the electrodes through it and connect it to the battery.

• Pour the concentrated NaCl into the glass tube

• Invert a test tube and place it over the electrodes and ensure that the electrodes aren’t completely covered, or the ions won't be able to flow

• Connect the electrodes to the battery

• Conduct the experiment in a fume cupboard to ensure that the poisonous fumes produced from chlorine gas are ventilated away

• Hydrogen forms at the cathode and Chlorine gas forms at the anode

   

Quantitative electrolysis

• Take a glass tube, close it with a rubber bung, place the electrodes through it, connect it to the battery and add a variable resistor and ammeter to the circuit.

• Place an inverted test tube over the electrodes to collect the Chlorine and Hydrogen gas

• Pour 50cm350cm3 of concentrated NaCl into the glass tube

• Place a glass burette with NaCl solution over the cathode

• Take the initial reading on the glass burette

• Set the current to 0.2A using the variable resistor and connect the battery

• Start the timer

• Repeat for different currents (0.4A, 0.6A, 0.8A and 1A)

• Repeat the experiment to get accurate and reliable results

• Collect the data in a table and plot a graph

ZNOTES.ORG

UPDATED TO 2023-2025 SYLLABUS

EDEXCEL IGCSE

CHEMISTRY

SUMMARIZED NOTES ON THE THEORY SYLLABUS

Prepared for Ananthiya for personal use only.

EDEXCEL IGCSE CHEMISTRY

1. Alkali Metals

Elements in group 1 are the alkali metals

Elements -

• Li - Lithium

• Na - Sodium

• K - Potassium

• Rb - Rubidium

• Cs - Caesium

• Fr - Francium

   

All the alkali metal elements have 1 electron in their outermost shell which is easily lost. As we move down the group, the atoms get bigger with more shells, the electron in the outermost shell is less strongly attracted to the nucleus and so more easily lost.

1.1. Physical and Chemical Properties

• The melting and boiling points decrease as you move down the group

• Their densities increase as we move down the group

• The metals are soft and can easily be cut with a knife

• They are shiny and silver when freshly cut but can tarnish quickly when exposed to air

• Become more reactive as you move down the group

• React with oxygen forming oxides and with water forming alkaline solutions

• Stored under oil to prevent them from reacting with oxygen or water

• Group 1 metal Ions are colourless or white, unless they are with a coloured negative ion

1.2. Reactions with water

alkali metal + water   ⟹  ⟹ alkali metal hydroxide + hydrogen

Reaction of water with Sodium -

• Sodium Floats

• Melts into a ball

• Fizzing as hydrogen gas is produced

• Moves around on the surface of the water

The piece of sodium gets smaller and disappears

The reaction of water with Lithium is less vigorous and so the Lithium used doesn’t melt but the rest of the observations are similar

The reaction of water with Potassium is more vigorous and produces a lilac flame causing it to spit around and explode

The reaction of water with Rubidium or Caesium is extremely reactive and violent.

1.3. Reactions with Air

alkali metal + oxygen   ⟹  ⟹alkali metal Oxide

Freshly cut pieces of Alkali metals tarnish very quickly when exposed to the air

A freshly cut piece of Potassium tarnishes extremely quickly, which is faster than that of Sodium, which is faster than that of Lithium

If heated with a Bunsen burner all Alkali metals react vigorously

• Lithium burns with a red flame

• Sodium burns with a yellow flame

• Potassium burns with a lilac flame

1.4. Predicting Properties

We can predict properties of Francium as we know the trend in reactivity of elements above it.

Hence we know that Francium will be -

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• Very soft

• Density over 2g/cm32g/cm3

• Melting point around room temperature

• More reactive than Caesium

• Tarnish almost instantly in air

• Form colourless compounds and colourless solutions

• Form Francium Hydroxide when dissolved in water and be alkaline in nature due to the presence of OH−OH− ions

2. The Halogens

The halogens are part of Group 7

They React with metals and produce a wide range of salts

All elements are non-metallic and contain diatomic molecules

The reactivity decreases as you move down the group as the electron that’s gained is less strongly attracted to the nucleus do to a larger shielding effect. The electron for Chlorine will be closer to the nucleus

2.1. Physical and Chemical Properties

• The melting and boiling points increase as we move down the group as there is an increase in their relative molecular mass. This increase in relative molecular mass increases the intermolecular forces of attraction requiring more thermal energy to break these stronger bonds.

• Halogens poor conductors of heat and electricity as they are non-metals

• Extremely corrosive and dangerous so have to be used in a fume cupboard (Fluorine and Chlorine are extremely toxic and reactive) and liquid bromine is corrosive so great care has to be taken to keep of off the skin

2.2. Reactions

Hydrogen + Halogens

• Halogens react with hydrogen forming hydrogen halides

• Hydrogen halides are all acidic, poisonous gases

• When dissolved in water they form acids

• HCl (g)  ⟹  ⟹(dissolved in water)  ⟹  ⟹HCl (aq)

Halogens react with Alkali metals to form salts

• 2Na (s) + Cl22​   ⟹  ⟹NaCl (s)

In the reaction above Sodium burns with Chlorine with a yellow flame to produce white Sodium Chloride

Displacement Reactions

• Chlorine is more reactive than Bromine, which is more reactive than Iodine

• Chlorine would displace Bromine as it’s more reactive. A Halogen more reactive than the other Halogen would displace it

• 2KBr (aq) + Cl2Cl2​ (aq)  ⟹  ⟹2KCl (aq) + Br2Br2​ (aq)

• If something is more reactive it is more likely to form a compound while the other is more likely to go back to being stable in its element form

2.3. Ionic Equations

Key Points:

• Solids don’t have Ions so it stays as it is in the original equation

• If the substance is an ionic compound in the aqueous state, it splits into Ions

• If it’s a strong acid or alkali split it into Ions

• Don’t make changes to other compounds if not needed

Original Equation: Ca (s) + Cu(NO3NO3​)22​ (aq)  ⟹  ⟹Ca(NO33​)22​(aq) + Cu (s)

Split the compounds into Ions: Ca + Cu2+2+​ + 2NO33​   ⟹  ⟹Ca2+2+​+ 2NO3−3−​ + Cu

Remove the spectator Ions: Ca + Cu2+2+​ + 2NO3−3−​  ⟹  ⟹Ca2+2+​ + 2NO3−3−​ + Cu

Final Net Ionic Equation (Doesn’t include spectator Ions):Ca (s)+ Cu2+2+​ (aq)  ⟹  ⟹Ca2+2+​ (aq) + Cu (s)

3. The Composition of Air

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3.1. Combustion of elements with Oxygen

Magnesium:

• Magnesium burns in oxygen to form a white ash(Magnesium Oxide) which even though isnt very soluble, dissolves to form an alkaline solution

• 2Mg (s) + O22​ (g)   ⟹  ⟹ 2MgO (s)

Sulfur:

• Sulfur burns with oxygen producing a blue flame and poisonous Hydrogen sulfide gas. When dissolved in water it forms Sulfurous Acid.

• S (s) + O22​ (g)   ⟹  ⟹ SO22​ (g)

Hydrogen:

• Hydrogen burns with oxygen producing a pale blue flame and water.

• H22​ + O22​   ⟹  ⟹2H22​O

3.2. Carbon Dioxide and Global Warming

• Carbon dioxide is a greenhouse gas that causes the enhanced greenhouse effect and global warming

• Carbon dioxide is produced due to the combustion of vast quantities of fossil fuels like coal, oil and natural gas

• The greenhouse effect occurs when UV light from our sun passes through the atmosphere and the radiation is reflected by the Earth and gets absorbed by the CO22​ in the atmosphere and radiated back to Earth heating it up

• Global Warming can cause Climate change leading to the melting of Ice caps and glaciers, causing a rise in sea levels leading to flooding and other extreme weather conditions

4. Reactivity series

Carbon (even though it’s a non-metal) is included in the reactivity series as a marker and used to show that if the metal is less reactive than carbon it can be easily and cheaply extracted by heating with carbon.

• Reducing agent: Substance that reduces something else and gets oxidised in the process

• Oxidising agent: Substance that oxidises something else and gets reduced in the process

• Oxidation: Loss of electrons, Gain of oxygen

• Reduction: Gain of electrons, Loss of oxygen

• REDOX: When oxidation and reduction happen simultaneously in a reaction

4.1. Reactions with water

Metals above hydrogen in the reactivity series react with water (l or g form) producing Hydrogen

• Metal + Cold water   ⟹  ⟹Metal Hydroxide + Hydrogen

• Metal + Steam  ⟹  ⟹Metal Oxide + Hydrogen

Metals below hydrogen in the reactivity series don’t react with water or steam

Calcium:

• Reacts gently with cold water and the mixture heats up as heat is produced.

• Calcium Hydroxide forms which isn’t very soluble in water

• Most of CaOH is left as a white insoluble solid, even though some dissolves and forms a colourless solution

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Magnesium:

Cold Water:

• Magnesium doesn’t react with cold water as the magnesium gets coated with insoluble magnesium hydroxide which prevents any more magnesium from reacting with water

  Steam:

  In this method, magnesium isn’t heated directly but the mineral wool containing water evaporates and moves left towards the magnesium ribbon.

• Magnesium burns with a bright white flame

• White Magnesium oxide is formed

Zinc:

With steam, Zinc oxide is formed which is yellow when hot and white when cool

Iron:

With steam, iron becomes darker forming tri-iron tetroxide

4.2. Reactions with Dilute Acids

• Metal + Acid -→ Salt + Hydrogen

• Metals such as Copper, Gold and Silver (below hydrogen in the reactivity series) don’t react with dilute acids

• Potassium, Sodium, Lithium, Calcium are all very reactive and violent causing it to be unsafe to react with

Magnesium:

• Reacts vigorously with cold dilute acids causing heat to be produced, making the mixture hot

Mg (s) + H22​SO44​ (aq)   ⟹  ⟹ MgSO44​ (aq) + H22​ (g)

Aluminium:

• Reacts slowly when cold but when heated the acid is able to overcome the strong layer of aluminium oxide on the surface allowing aluminium to react properly with the dilute acid

2Al (s) + 6HCl (aq)   ⟹  ⟹ 2AlCl33​ (aq) + 3H22​ (g)

Zinc & Iron:

• Slow to react with cold acids but when heated, it reacts more rapidly.

• The vigour of these 2 reactions is less than that of Aluminium as it’s higher on the reactivity series than these two elements.

Zinc: Zn(s) + H22​SO44​ (aq)   ⟹  ⟹ZnSO44​ (aq) + H22​ (g)

Iron: Fe (s) + 2HCl (aq)   ⟹  ⟹ FeCl22​ (aq) + H22​ (g)

4.3. Rusting of Iron

• Iron rusts in the presence of oxygen and water

• Metals corrode, but only Iron rusts

• Rusting occurs with Iron (3+) ions

• Rust: Hydrated Iron(III) oxide (Fe22​O33​ .xH22​O)

Ways of preventing rusting:

• Barrier method: Coating the iron with a metal below it in the reactivity series

  • Cheap ways of preventing rusting

  • Iron will rust if any part of the coating is broken

  • Example: Painting, coating with Oil or grease, covering with plastic

• Galvanised iron: Iron that’s coated with a layer of Zinc

  • Iron doesn’t rust even if it’s exposed to oxygen as Zinc reacts and forms a layer of zinc oxide protecting the iron beneath it

  • Example: Cars, Iron nails and tools

• Sacrificial protection: Use a metal more reactive than Iron that reacts and protects the iron beneath it

  • Have to be replaced occasionally when all the metal has reacted and formed an oxide

  • Used on large structures

  • Example: Ships and tankers, underground pipelines

5. Extraction of metals

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• Ore: A sample of mineral with enough of a mineral to be beneficial to extract

• Native: Element exists naturally in its uncombined state

• Roasting: When sulfides can be easily converted to oxides by heating it in air

5.1. Metals below Carbon in the reactivity series

• Cheapest method is to heat the ore with carbon

• Carbon would be higher on the reactivity series than the ore so would take oxygen away from it getting oxidised producing carbon dioxide and the required element

• Extraction of iron carried out in a blast furnace

Fe22​O33​ (s) + 3C (s)   ⟹  ⟹2Fe (l) +3CO (g)

(No need to learn details of the blast furnace)

5.2. Metals above Carbon in the reactivity series

• Ores of metals higher than carbon in the reactivity series have to be separated by electrolysis which is expensive (requires large amounts of electricity)

• Aluminium is extracted from Aluminium oxide by electrolysis

Cathode: Al3+3+ + 3e−−   ⟹  ⟹Al

Anode: 2O2−2−   ⟹  ⟹ O22​ + 4e−−

5.3. Alloys

• Mixture of metals with other metals

• Alloys are harder than individual pure metals as in alloys the atoms are differently sized and breaks the regular arrangement of atoms in the lattice, making it more difficult for the layers of atoms to slide over each other

• Example of alloys: Bronze, stainless steel, Cupronickel

5.4. Properties and uses of some metals and alloys

• Pure aluminum isn’t very strong so alloys are used

• Aluminium resists corrosion as a layer of magnesium oxide forms, preventing anything reaching the surface and reacting with the aluminium

• It has a low density

Mild Steel:

• Contains 0.25% of carbon as its an alloy of Iron

• Can be easily hammered into various shapes and drawn into wires (malleable and ductile)

• Will rust when exposed to water and oxygen

• 3x denser than aluminium

• Examples: Car bodies, Ship building, girders, nails

High-Carbon Steel:

• Contains 0.6%-1.2% carbon

• Usually contains small amounts of manganese

• Harder and less resistant to wear than Mild Steel

• Less malleable and ductile than Mild Steel

• Examples: Cutting tools, Masonry nails

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Stainless Steel:

• Contains Chromium and Nickel

• Resistant to corrosion

• More expensive than mild steel

• Examples: Cutlery, cooking utensils, kitchen sinks

Copper:

• Good conductor of heat and electricity

• Ductile and malleable

• Unreactive

• Anti-microbial properties

• Examples: Water pipes, Electric wires, Pots and pans, Hospital surfaces

6. pH and Indicators

The pH scale tells you how acidic or alkaline a solution is

Acid-Alkali indicators -

6.1. Acids and Bases

Acids:

• Acids: act as a source of Hydrogen ions in a solution

• All acids contain Hydrogen ions. (H++)

• All acids have a replaceable H (not all “H” are replaceable though) in the compounds for example -

  HCl (aq) + NaOH (aq)   ⟹  ⟹NaCL (aq) + H22​O (l)

• When acids are in water, they dissociate to form H++ ions

• We measure the acidity by measuring the concentration of these H++ ions in the solution

Bases:

• Alkalis: Source of Hydroxide ions (OH−−)in the solution

• When in solution -

  NaOH (aq)   ⟹  ⟹ Na++ (aq)+ OH−− (aq)

6.2. Neutralisation Reactions

• Acid react with an Alkali or a base in a neutalisation reaction

   

Sodium Hydroxide reacts with Hydrochloric acid, forming Sodium chloride and water (neutralisation reaction)

NaOH (aq) + HCl (aq)   ⟹  ⟹NaCl (aq) + H22​O (l)

If you write an ionic equation for this reaction, removing the spectator ions you get:

OH−− (aq) + H22​ (aq)   ⟹  ⟹H22​O (l)

All neutralisations reactions will have the same ionic equation given above

6.3. Titration

Method -

“Using HCl and NaOH solutions with phenolphthalein indicator”

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1. Measure 25cm33 of Hydrochloric acid using a pipette

2. Transfer it to a conical flask

3. Fill the burette with Sodium Hydroxide solution

4. Take initial reading on the burette and record to 2dp

5. Add phenolphthalein indicator to the conical flask (acid)

6. Add Sodium hydroxide until the indicator changes colour

7. Take the reading on the burette when the indicator turns a pale pink colour

8. Find the final volume of alkali added to the acid to neutralise it

7. Acids, Bases and Salt Preparations

An acid is a proton donor

A base is a proton acceptor

7.1. Reactions with Acids

• Metals above hydrogen react to produce hydrogen gas (Metal + Acid -→ Salt + Hydrogen)

• Metals below hydrogen don’t react with the dilute acids

• The higher the metal in the reactivity series the more violent and vigorous the reaction

Mg + H22​SO44​   ⟹  ⟹ MgSO44​ + H22​

In this reaction, there is a rapid fizzing and a rapid gas is produced. Heat is produced which makes the reaction warm. A colourless solution of magnesium sulfate is produced.

Mg + 2HCl   ⟹  ⟹ MgCl22​ + H22​

This reaction is similar to the previous reaction above only here a mixture of Magnesium Chloride is formed.

Reactions between Zinc and the acids are slower as Zinc is lower down the periodic table than Magnesium

7.2. Reactions with Bases

Metals oxides are bases

Copper(II) oxide reacts with dilute Hydrochloric acid to produce a blue solution of copper(II) sulfate

Metal oxide + Acid -→ Salt and Water

This is a neutralisation reaction

NaOH (aq) + HCl (aq)   ⟹  ⟹ NaCl (aq) + H22​O (l)

A colourless solution is formed and gets warmer.

7.3. Making Soluble salts

Making Copper(II) Sulfate crystals (Non-SAP, Crystallization method)

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• Measure 50 cm33 of dilute Sulfuric Acid and heat it on a tripod and gauze using a Bunsen Burner

• Add Copper(II) Oxide and continue heating until the Copper(ll) Oxide is in excess

• Filter off the excess Copper(ll) Oxide and transfer the solution (filtrate) to an evaporating basin

• Heat the filtrate over a Bunsen Burner to boil off some of the water and concentrate the solution

• Keep heating until a saturated solution is formed which can be tested by dipping a glass rod and seeing if crystals form on the surface when removed

• Stop heating and allow the solution to cool to room temperature

• Remove the blue crystals by filtration

• Dry the crystals by blotting it with a paper towel or left in a warm place

Making Sodium Sulfate crystals (SAP, Titration method)

• Measure 25 cm33 of Sodium hydroxide solution into a conical flask using a pipette

• Add methyl orange as the indicator

• Add dilute sulfuric acid from the burette until the indicator changes from yellow to orange

• Note the volume of acid added

• Mix the same volumes of the acid and alkali noted without the indicator in a clean flask

• Heat the filtrate over a Bunsen Burner to evaporate off some of the water and concentrate the solution

• Stop heating and allow the solution to cool to room temperature

• Remove the crystals by filtration

• Dry the crystals by blotting it with a paper towel or left in a warm place

7.4. Making Insoluble salts

Reactants:

Nitrate of the metal part

Sodium/Potassium of the non-metal part

Making Lead(II) Sulfate crystals (Precipitation method)

• Measure 25 cm33 of Lead(II) nitrate solution into a beaker

• Measure 25 cm33 of sodium sulfate solution into a beaker

• A white precipitate of Lead(II) Sulfate forms

• Remove the crystals by filtration

• Wash the residue with distilled water

• Dry the crystals by blotting it with a paper towel or left in a warm place

7.5. Testing of Gases

Hydrogen:

• A lighted splint is held to the end of the test tube causing the hydrogen to explode with a squeaky pop sound

Oxygen:

• A lighted splint is held to the end of the tube. Oxygen would relight the glowing splint

Carbon Dioxide:

• Carbon dioxide gas is bubbled through limewater causing it to turn the limewater cloudy/milky

Chlorine:

• A damp blue litmus paper kept to the end of the test tube would turn red and then bleach

Ammonia:

• A damp red litmus paper kept to the end of the test tube would turn blue

7.6. Physical and Chemical Test for Water

Chemical:

• Water turns anhydrous copper sulfate blue as anhydrous copper sulfate lacks the water of crystallisation, so adding water helps turn it blue

Physical:

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• The melting and boiling points of water can be used to show if it is pure or not.

• Water boils at 100oo Celsius

• Water freezes at 0oo Celsius

7.7. Flame Test

A platinum or nichrome wire is dipped in conc. Hydrochloric acid to remove any impurities left behind and then dipped into the salt and held in the luminous Bunsen burner flame

7.8. Testing for Cations with NaOH

Dissolve the salt in water and add 1cm33 of this to a test tube and add the same volume of NaOH to this. A precipitate forms

7.9. Testing for Carbonates

Add dilute HCl and then look for fizzing/effervescence as a gas is given off. We can test if this gas is carbon dioxide and contains carbonate ions by bubbling the gas through limewater causing the limewater to turn cloudy/milky.

7.10. Testing for Sulfates

Add Dilute Hydrochloric acid and then Barium Chloride. This would produce a white precipitate of Barium sulfate if sulfate ions are present.

• The acid is added to remove any other ions present in the solution

• Never add sulfuric acid to this solution for the test as the sulfuric acid would contain sulfate ions so you are bound to get a positive test for this test

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7.11. Test for Halide ions

Make a solution of your suspected halide and add enough nitric acid to make it acidic and then proceed to add silver nitrate solution