Chemistry Paper 1 overview

🧪 Chemical Substances

🧱 Atoms and Elements

• All substances are made of atoms.

• Different types of atoms are called elements.

• Elements are represented by symbols in the periodic table.

⚗ Compounds

• A compound contains two or more different types of atoms chemically bonded.

  Example: Water (H2OH2​O) - two hydrogen atoms and one oxygen atom.

• Absence of a number after a symbol implies one atom.

⚛ Chemical Reactions

• Atoms change their bonding through chemical reactions.

• Reactions can be represented by:

  • Word equations

  • Chemical equations (using symbols)

• Atoms are not created or destroyed in chemical reactions.

  • The number of each type of atom must be the same on both sides of the equation.

Balancing Equations

1. Start balancing atoms that are only in one compound.

2. Do not change the small numbers within a compound's formula.

3. Add numbers in front of elements or compounds to multiply them.

4. Always finish balancing elements that appear on their own last (e.g., O2O2​).

   Example: CH4+O2→CO2+H2OCH4​+O2​→CO2​+H2​O

   Balanced: CH4+2O2→CO2+2H2OCH4​+2O2​→CO2​+2H2​O

🫗 Mixtures

• A mixture is a combination of different elements and compounds that are not chemically bonded.

  Example: Air (oxygen, nitrogen, etc.)

• Solutions are mixtures.

  Example: Saltwater (water and sodium chloride)

Separation Techniques

• Filtration: Separates large, insoluble particles from a liquid.

  Example: Sand from water

• Crystallization: Separates a solute from a solvent by evaporating the solvent.

  Example: Salt from water

• Distillation: Heats a solution, cools the gas, and condenses it back into a liquid.

  Fractional distillation separates different liquids in a mixture based on boiling points.

  These are physical processes, not chemical reactions, because no new substances are made.

🧊 States of Matter

• Three main states of matter: solid, liquid, and gas.

• To change states (melt or evaporate), supply energy to overcome electrostatic forces.

• These are physical changes.

• State symbols in chemical equations:

  • (s) for solid

  • (l) for liquid

  • (g) for gas

  • (aq) for aqueous (dissolved in solution)

⚛ Atomic Structure

• JJ Thompson: Atoms contain positive and negative charges (plum pudding model).

• Ernest Rutherford: Positive charge is concentrated in a tiny nucleus; electrons orbit far away. Atoms are mostly empty space.

• Niels Bohr: Electrons exist in shells or orbitals.

• James Chadwick: Nucleus contains neutral particles called neutrons.

📊 The Periodic Table

• The atomic number (bottom number) is the number of protons, determining the element.

• Neutral atoms have equal numbers of protons and electrons.

• Ions are formed when atoms gain or lose electrons.

• The mass number (top number) is the number of protons and neutrons.

☢ Isotopes

• Isotopes are atoms of the same element with different numbers of neutrons.

• Periodic tables show average mass, considering the relative abundance of isotopes.

  Example: If you have some chlorine gas it turns out that 75% of the atoms will have a mass of 35 while 25% of the atoms will be 37 these are what we call their relative abundance to find the average we just pretend that we have 100 atoms we add up the total masses of all the Isotopes then just divide by 100 that's why chlorine average relative atomic mass is 35.5

📜 History of the Periodic Table

• Early tables ordered elements by atomic weights, grouping similar properties.

• Dimitri Mendeleev grouped elements by properties, even if it meant deviating from atomic weight order, predicting undiscovered elements.

orbital Shells

• Electron shells fill from the inside:

  • First shell: max 2 electrons

  • Second and third shells: max 8 electrons

  • Fourth shell: max 2 electrons (up to Calcium)

• Electron configuration:

  Magnesium (12 electrons): 2,8,2

🔩 Metals vs. Nonmetals

• Metals (left of staircase) donate electrons.

• Nonmetals (right of staircase) accept electrons.

• Group indicates the number of outer shell electrons.

Alkali Metals

• Group 1 atoms are alkali metals with one outer shell electron.

• They have similar properties when reacting with water.

• Reactivity increases as you go down the group.## ⚛ Periodic Table Trends

Group 1: Alkali Metals

• As you move down Group 1 (alkali metals), elements become more reactive.

• This is because the outermost electron is further from the nucleus.

• The electrostatic attraction between the negative electron and the positive nucleus is weaker.

• The electron is more readily donated.

Group 7: Halogens

• Halogens are essentially the opposite of alkali metals.

• They have seven electrons in their outer shell, needing one more to complete it.

• As you move down the group, elements become less reactive.

• Electrons are less readily accepted because the shell is further from the nucleus.

• Boiling points increase as you go down the group.

Group 0 (8): Noble Gases

• Noble gases have a full outer shell.

• They are very unreactive, though they can react under special conditions.

• The term "Group 8" is becoming outdated to include Helium, which only has two electrons in its outer shell.

➕ Ion Formation

Metals and Positive Ions

• Metals become positively charged when they lose electrons.

• Group 1 metals lose one electron, forming ions with a 1+ charge.

• Group 2 metals lose two electrons, forming ions with a 2+ charge.

Non-metals and Negative Ions

• Group 7 elements gain one electron, forming ions with a 1- charge.

• Group 6 elements gain two electrons, forming ions with a 2- charge.

Elements Not Forming Ions

• Atoms in Groups 3, 4, and 5 generally do not form ions, except for aluminum (Al) which forms Al3+.

Transition Metals

• Transition metals can donate different numbers of electrons.

  • For example, iron can form Fe2+Fe2+ or Fe3+Fe3+ ions.

  • These are distinguished as Iron (II) and Iron (III) respectively.

• Generally harder and less reactive than alkaline metals.

• They often form colored compounds.

🔗 Chemical Bonding

Metallic Bonding

• Metal atoms bond to each other through metallic bonding.

• A lattice of ions is formed with a "sea" of delocalized electrons around them.

• These free-moving electrons make metals good conductors of electricity and heat.

Ionic Bonding

• Metals bond to non-metals through ionic bonding.

• A metal atom loses an electron, while a non-metal atom gains one.

• Example: Lithium (Li) donates its outer electron to Chlorine (Cl).

  • Dot and cross diagrams can visually represent electron transfer.

  • Include brackets and charges of the ions.

  • The charges of all ions in an ionic compound must add up to zero.

Examples of Ionic Compounds

Properties of Ionic Compounds

• Ionic compounds consist of repeating units of ions in a lattice to form a crystal.

• High melting points and boiling points due to strong electrostatic forces.

• Conduct electricity only in liquid form (molten or dissolved in solution) when ions are free to move and carry charge.

Molecular Ions

• Molecular ions consist of multiple atoms.

  • Example: Hydroxide ion (OH−OH−).

• Example: Magnesium Hydroxide requires two hydroxide ions (OH−OH−) to balance the Mg2+Mg2+ charge.

Naming Conventions for Ionic Compounds (Salts)

• Any ionic compound can be called a salt, not just sodium chloride.

• The name is always the metal ion (positive ion/cation), followed by the non-metal ion (anion).

• Anion names are altered:

  • Not sodium chlorine, but sodium chloride.

Covalent Bonding

• Non-metals bond to each other with covalent bonding to form molecules.

• They achieve full outer shells by sharing electrons.

• Example: Chlorine gas (Cl2Cl2​)

  • Each chlorine atom shares an electron with the other.

Dot and Cross Diagrams and Structural Formulas

• Dot and cross diagrams illustrate shared electrons.

• Structural formulas use symbols and lines to represent bonds.

  • Each line represents a shared electron pair.

Multiple Bonds

• Oxygen (O2O2​): Each oxygen atom shares two electrons, forming a double bond.

• Nitrogen (N2N2​): Forms a triple bond.

Bonding Capacity

• The number of electrons an atom needs is the same as the number of bonds it makes.

  • Hydrogen can only make one bond.

  • Carbon makes four bonds.

Simple Molecular (Covalent) Structures

• Individual molecules that can mix together.

• Relatively low boiling points due to weak intermolecular forces between molecules.

  • Heating overcomes intermolecular forces, not covalent bonds.

• Cannot conduct electricity, even as liquids.

Giant Covalent Bonding

• Atoms form covalent bonds to other atoms, creating a giant molecule.

• Example: Diamond

  • A crystal of carbon atoms bonded to each other.

  • Extremely hard with a high melting point because covalent bonds need to be broken.

Allotropes of Carbon

• Allotropes are different forms of the same element.

• Example: Graphite

  • Consists of layers of carbon atoms with three bonds each in a hexagonal structure.

  • Spare delocalized electrons form weak bonds between layers.

  • Can conduct electricity as electrons move between layers.

  • Layers can slide over each other easily, making it useful in pencils.

Metal Alloys

• Metal alloys are stronger than pure metals.

• Mixtures of metals with different-sized atoms disrupt the regular lattice, preventing layers from sliding easily.

More Carbon Allotropes

• Graphene: A single layer of graphite.

• Fullerenes: 3D structures of carbon atoms.

  • Buckminsterfullerene: Spherical structure with 60 carbon atoms.

• Nanotubes: Tube-shaped fullerenes.

Nanoparticles and Surface Area

• Surface area to volume ratio is just one divided by the other, s/vs/v.

• If the length of a side of a cube doubles, the ratio halves.

• Nanoparticles have huge surface area to volume ratios, requiring fewer to fulfill a purpose.

🧪 Quantitative Chemistry

Conservation of Mass

• Total mass of all substances is conserved in a chemical reaction.

• Atoms that go in must come out.

• Equations must be balanced to reflect this.

Relative Formula Mass

• Add up the individual relative atomic masses (Rams) of atoms in a compound.

• Example: CO2CO2​ is 12+(2×16)=4412+(2×16)=44.

• Some reactions produce gas, which, if it leaves the reaction vessel, results in a seeming decrease in mass.

Moles

• A way of comparing amounts of substances.

• If you have as many grams of a substance as its relative atomic or formula mass, you have one mole.

• One mole of carbon has a mass of 12g.

🧪 Stoichiometry: Mass Relationships in Chemical Reactions

Calculating Moles

The relationship between moles, grams, and relative atomic mass (RAM) or relative formula mass is given by:

moles=gRAMmoles=RAMg​

Balancing Chemical Reactions and Mole Ratios

In a balanced chemical reaction, the coefficients indicate the mole ratios of reactants and products. For example, in the combustion of methane:

CH4+2O2⟶CO2+2H2OCH4​+2O2​⟶CO2​+2H2​O

This equation tells us that for every 1 mole of methane (CH4CH4​) reacted, we need 2 moles of oxygen (O2O2​) and we'll produce 2 moles of water (H2OH2​O). This relationship is essential for stoichiometric calculations.

Mass-to-Mass Conversions: The Mole Middleman

To convert from the mass of one substance to the mass of another in a chemical reaction:

1. Convert the mass of the first substance to moles.

2. Use the stoichiometry of the balanced equation to find the moles of the second substance.

3. Convert the moles of the second substance back to mass.

Example: How many grams of water would be made if 64 g of methane reacted completely with oxygen?

1. Moles of methane: moles=64 g16 g/mol=4 molesmoles=16 g/mol64 g​=4 moles

2. From the balanced equation, the mole ratio of CH4CH4​ to H2OH2​O is 1:2. So, moles of water = 2×4=8 moles2×4=8 moles

3. Mass of water: mass=8 moles×18 g/mol=144 gmass=8 moles×18 g/mol=144 g

Working with Different Mass Units

You can use kilograms or tons in place of grams, so long as you stick with the same unit throughout the entire calculation.

Limiting Reactants

The limiting reactant is the reactant that runs out first in a chemical reaction, thus determining the maximum amount of product that can be formed.

Solution Concentration

Concentration can be expressed in grams per cubic decimeter (g/dm3g/dm3), but it's often more useful to convert this into moles per cubic decimeter (mol/dm3mol/dm3), also known as Molarity (M). 1 mol/dm3mol/dm3 is the same as saying 1 Moler.

Example: If one mole of HCl is dissolved in 1 dm3dm3 of water, the concentration is 1 mol/dm3mol/dm3.

⚗ Chemical Changes: Yield and Atom Economy

Percentage Yield

Percentage Yield=(Actual YieldTheoretical Yield)×100Percentage Yield=(Theoretical YieldActual Yield​)×100

You must be given the actual masses involved in percentage yield questions.

Example: If you start with 20 g of reactants and obtain 10 g of ammonia, the percentage yield is 1020×100=50%2010​×100=50%

Atom Economy

Atom Economy=(RAM of desired productTotal RAM of reactants)×100Atom Economy=(Total RAM of reactantsRAM of desired product​)×100

Example: For the methane reaction CH4+2O2⟶CO2+2H2OCH4​+2O2​⟶CO2​+2H2​O, if CO2CO2​ is the desired product:

Atom Economy=(4444+(2×18))×100=(4480)×100=55%Atom Economy=(44+(2×18)44​)×100=(8044​)×100=55%

This can also be calculated by dividing the RAM of the desired product by the RAM of all products, due to conservation of mass.

🎈 Molar Volume of Gases

One mole of any gas occupies a volume of 24 dm3dm3 at Room Temperature and Pressure (RTP), which is 20∘C20∘C and one atmosphere.

Volume=Moles×24Volume=Moles×24

To convert between moles and volume, multiply or divide by 24.

🔩 Reactivity Series and Displacement Reactions

Reactivity Series

Metals vary in their reactivity, with some donating electrons more readily than others. The reactivity series lists metals in order of their reactivity, with hydrogen and carbon included for comparison.

Displacement Reactions

A more reactive metal will displace a less reactive metal from a compound. For example, zinc will displace copper from copper sulfate solution:

Zn+CuSO4⟶ZnSO4+CuZn+CuSO4​⟶ZnSO4​+Cu

Metal Extraction

Metals less reactive than carbon can be extracted from their ores by reduction with carbon, a process called smelting. For example, iron can be displaced from iron oxide with carbon.

Oxidation and Reduction (Redox)

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

The mnemonic "OIL RIG" helps remember this: Oxidation Is Loss, Reduction Is Gain.

Example: In the reduction of iron oxide (Fe2O3Fe2​O3​) with carbon:

Fe3++3e−⟶FeFe3++3e−⟶Fe

The iron ions gain electrons and are reduced.

Reactions with Acids and Alkalis

• Metals more reactive than hydrogen can displace it from an acid, producing a salt.

• Alkalis (pH > 7) react with acids (pH < 7) to produce a salt and water.

• If quantities are correct, neutralization occurs, leaving no unused reactants.

Salts

• Hydrochloric acid produces metal chlorides.

• Sulfuric acid produces metal sulfates.

• Nitric acid produces metal nitrates.

These salts are left in solution (dissolved in water).

Obtaining Salt Crystals

Solid crystals of a dissolved salt can be obtained by gently warming the solution to evaporate the water.

⚖ The pH Scale

The pH scale is logarithmic (base 10), not linear.

• An acid with a pH of 3 has 10 times the concentration of H+H+ ions compared to an acid with a pH of 4.

• Alkalies work similarly with OH−OH− ions.

Strong vs. Weak Acids

• Strong acids: Dissociate (ionize) completely in solution (e.g., hydrochloric, nitric, and sulfuric acids).

• Weak acids: Only partially dissociate (e.g., ethanoic, citric, and carbonic acids).

For acids of the same concentration, a stronger acid will have a lower pH.

🧪 Titrations (Triple Science Only)

1. Measure a known volume of alkali using a glass pipette and put it in a conical flask.

2. Add a few drops of an indicator (e.g., methyl orange).

3. Put the acid of unknown concentration in a burette above the flask.

4. Slowly drip the acid into the flask while swirling.

5. When the indicator changes color (e.g., to pink), close the tap.

Titration Calculations 🧪

After neutralization, indicated by a color change (e.g., pink after swirling), we can perform titration to find the concentration of an unknown solution. A rough titration helps estimate the required volume, followed by a more precise titration, adding the titrant drop by drop near the end point.

Let's work through an example with sodium hydroxide (NaOH) and sulfuric acid (H₂SO₄):

• Balanced Equation: 2NaOH+H2SO4→Na2SO4+2H2O2NaOH+H2​SO4​→Na2​SO4​+2H2​O

• Given: 50 cm³ of 0.2 moles/dm³ NaOH

1. Convert volume to dm³: 50 cm3÷1000=0.05 dm350cm3÷1000=0.05dm3

2. Calculate moles of NaOH: 0.05 dm3×0.2 moles/dm3=0.01 moles0.05dm3×0.2moles/dm3=0.01moles

3. Using stoichiometry (1:2 acid:alkali ratio), calculate moles of acid needed: 0.01 moles NaOH÷2=0.005 moles H2SO40.01molesNaOH÷2=0.005molesH2​SO4​

4. If the measured volume of acid is 0.0125 dm³ (12.5 cm³), calculate the concentration: 0.005 moles÷0.0125 dm3=0.4 moles/dm30.005moles÷0.0125dm3=0.4moles/dm3

Electrolysis Explained⚡

Electrolysis involves passing an electric current through a molten or dissolved ionic compound, using inert electrodes (e.g., carbon) that do not react.

Process Breakdown

• Molten Ionic Compound: An ionic compound that has been heated to a liquid state. This allows ions to move freely and conduct electricity.

• Cations: Positive metal ions (e.g., Al³⁺) move to the cathode (negative electrode).

  • At the cathode, cations are reduced (gain electrons) and become neutral atoms.

  • Example: Al3++3e−→Al(s)Al3++3e−→Al(s)

• Anions: Negative ions (e.g., O²⁻) move to the anode (positive electrode).

  • At the anode, anions are oxidized (lose electrons) and form neutral molecules.

  • Example: 2O2−→O2(g)+4e−2O2−→O2​(g)+4e−

Applications

Electrolysis is used for:

• Purifying metals

• Extracting metals from compounds when carbon displacement is not viable due to reactivity.

Example: Aluminium Oxide (Al₂O₃)

• Oxygen produced at the graphite anode reacts with it, requiring frequent replacement.

• Aluminium oxide is mixed with cryolite to lower its melting point, reducing extraction costs.

Electrolysis in Solutions

In aqueous solutions, multiple ions are present (e.g., Na⁺, Cl⁻, H⁺, OH⁻ in NaCl solution). The ion that gets reduced or oxidized depends on its reactivity:

• Cathode: The less reactive ion is reduced.

  • Example: In NaCl solution, H⁺ is reduced instead of Na⁺, producing hydrogen gas.

    • 2H++2e−→H2(g)2H++2e−→H2​(g)

• If the metal is less reactive than hydrogen (e.g., copper in copper sulfate), the metal forms on the cathode.

• Anode: If a halide ion (e.g., Cl⁻) is present, it is oxidized. Otherwise, oxygen from OH⁻ is oxidized, producing oxygen gas.

  • 2Cl−→Cl2(g)+2e−2Cl−→Cl2​(g)+2e−

Energy Changes in Reactions 🔥

Chemical reactions involve energy transfers related to bond breaking and bond formation.

• Bond Breaking: Requires energy

• Bond Formation: Releases energy

Exothermic Reactions

• More energy is released from bond formation than is required for bond breaking.

• Net energy release results in a temperature increase.

• Example: Combustion

  • "X" means "out," so heat exits

Endothermic Reactions

• More energy is required for bond breaking than is released from bond formation.

• Net energy input results in a temperature decrease.

Practical Example: Neutralization in a Cup

1. Carry out a neutralization reaction in an insulated polystyrene cup with a lid and thermometer.

2. Measure the maximum temperature reached.

3. Repeat with increasing volumes of alkali.

4. The maximum temperature plateaus when all acid has reacted.

5. Plot temperature vs. volume of alkali to determine the neutralization point.

Energy Profile Diagrams

• Y-axis: Potential Energy

• Inverse Relationship: Potential energy and kinetic energy often balance. Lower potential energy usually corresponds to higher kinetic energy and thus a higher temperature.

• Exothermic Reaction:

  • Products have lower potential energy than reactants (higher kinetic energy/temperature).

  • Requires activation energy (the "bump") to initiate the reaction.

• Endothermic Reaction:

  • Products have higher potential energy than reactants.

Bond Energies

• Every bond has a specific bond energy (energy needed to break one mole of that bond).

• The same amount of energy is released when the bond is formed.

Calculating Net Energy Change

Let's revisit the combustion of methane (CH4+2O2→CO2+2H2OCH4​+2O2​→CO2​+2H2​O)

1. Draw out the structures to visualize all bonds.

2. Calculate total energy needed to break bonds in reactants:

   • 4×C−H4×C−H bonds (4×413 kJ/mol4×413kJ/mol)

   • 2×O=O2×O=O bonds (2×495 kJ/mol2×495kJ/mol)

   • Total energy in: 2642 kJ/mol2642kJ/mol

3. Calculate total energy released from forming bonds in products:

   • 2×C=O2×C=O bonds (2×799 kJ/mol2×799kJ/mol)

   • 4×O−H4×O−H bonds (4×467 kJ/mol4×467kJ/mol)

   • Total energy out: 3466 kJ/mol3466kJ/mol

4. Net Energy Change:

   • Energy released - Energy input

   • 3466 kJ/mol−2642 kJ/mol=824 kJ/mol3466kJ/mol−2642kJ/mol=824kJ/mol (exothermic)

Batteries and Fuel Cells 🔋

Batteries

• Contain chemicals that produce a potential difference (voltage).

• Composed of two different metals in contact with an electrolyte.

• Non-renewable batteries: Reactants are consumed and the battery stops working.

• Rechargeable batteries: A supplied current reverses the reaction, recharging the battery.

Hydrogen Fuel Cells

• Water is split into hydrogen and oxygen via electrolysis.

• Recombining hydrogen and oxygen generates a voltage.