Chemistry Paper 1 Key Concepts

Atomic Structure and the Evolution of the Atom

• Atomic Components:

  • Nucleus: Contains protons and neutrons. It is described as tiny compared to the overall size of the atom.
  • Electrons: Orbit the nucleus in fixed shells or energy levels.
  • Electrical Neutrality: Atoms are neutral because they possess an identical number of protons and electrons.

• Properties of Sub-atomic Particles:

  • Protons: Mass = $1$, Charge = $+1$.
  • Neutrons: Mass = $1$, Charge = $0$.
  • Electrons: Mass = $\frac{1}{2000}$, Charge = $-1$.

• Historical Model of the Atom:

  • Dalton's Theory: Originally suggested the atom was the smallest particle of matter and indivisible.
  • Modern Discovery: The identification of sub-atomic particles (electrons, protons, and neutrons) proved atoms could be broken down.
  • Mass Distribution: It is now known that the vast majority of an atom's mass is concentrated within the nucleus.

• Atomic Identity and Isotopes:

  • Atomic Number: The number of protons in an atom. This number is unique to each element.
  • Mass Number: The total number of protons plus neutrons ($\text{Mass number} = \text{no. protons} + \text{no. neutrons}$).
  • Isotopes: Atoms of the same element that have the same number of protons but a different number of neutrons, resulting in different mass numbers.
  • Relative Atomic Mass (RAM) - HIGHER ONLY: The mean relative mass of the isotopes in an element. This is often not a whole number.

The Periodic Table and Electronic Configuration

• Mendeleev’s Periodic Table:

  • Arranged elements by increasing atomic mass.
  • Encountered exceptions to mass order due to isotopes (which were unknown at the time).
  • Grouped elements based on similar properties.
  • Predicted unknown elements and their properties by leaving gaps in the table.

• The Modern Periodic Table:

  • Arranged by increasing atomic number.
  • Groups: Vertical columns containing elements with similar properties.
  • Periods: Horizontal rows. The period number corresponds to the number of electron shells.
  • Layout: Metals are located to the left of the zig-zag line, and non-metals are to the right.

• Electronic Configuration:

  • States the number of electrons in each shell, filling from the first shell outwards.
  • First Shell: Maximum of $2$ electrons.
  • Second Shell: Maximum of $8$ electrons.
  • Third Shell: Maximum of $8$ electrons.
  • Example: Sodium ($Na$) has $11$ electrons: configuration = $2.8.1$. It is in Group 1 (due to 1 outer electron) and Period 3 (due to 3 shells).

Ions and Chemical Bonding

• Ions:

  • Atoms or groups of atoms with a positive or negative charge.
  • Cations: Positive ions, usually metals (e.g., $Li^+$, $Mg^{2+}$, $Ca^{2+}$). Formed by losing electrons.
  • Anions: Negative ions, usually non-metals (e.g., $Cl^-$, $O^{2-}$, $SO_4^{2-}$). Formed by gaining electrons.
  • Stability: Atoms become ions to achieve a full outer shell.

• Ionic Bonding:

  • Occurs when electrons are transferred between atoms.
  • Example: Sodium Chloride ($NaCl$). $Na (2.8.1) + Cl (2.8.7) \rightarrow Na^+ (2.8) + Cl^- (2.8.8)$.
  • Deducing Formulas: Magnesium oxide ($MgO$), Magnesium chloride ($MgCl_2$), Potassium oxide ($K_2O$), Copper nitrate ($Cu(NO_3)_2$).
  • Terminology: Compounds ending in –ate contain oxygen (e.g., Magnesium sulphate $MgSO_4$ vs. Magnesium sulphide $MgS$).

• Covalent Bonding:

  • A shared pair of electrons between two atoms, resulting in the formation of molecules.
  • Scale: Typical size of an atom and small molecules is $10^{-10}\,m$.
  • Double Bonds: Involve four shared electrons (two pairs), such as in Oxygen ($O_2$) and Carbon Dioxide ($CO_2$).

• Detailed Comparison of Bonding Types:

  • Ionic:
    • Structure: Lattice with regular arrangement of positive and negative ions.
    • Attraction: Strong electrostatic forces between oppositely charged ions.
    • Properties: High melting/boiling points; most are soluble; conduct electricity only when liquid or dissolved (not as a solid).
  • Simple Molecular (Covalent):
    • Structure: Individual molecules held together by covalent bonds.
    • Attraction: Weak intermolecular forces between molecules.
    • Properties: Low melting/boiling points; few are soluble; poor electrical conductors.
  • Metallic:
    • Structure: Lattice of positive ions in a ‘sea’ of delocalised electrons.
    • Attraction: Strong electrostatic forces between positive ions and delocalised electrons.
    • Properties: High melting/boiling points; insoluble in water; good electrical conductors (delocalised electrons carry current).
    • Malleability: Metals can be hammered into shape because layers of ions slide over each other while electrons hold them together.

Carbon Structures and Polymers

• Allotropes of Carbon:

  • Diamond: Each carbon atom is bonded to $4$ others in a giant lattice. Very hard; high melting point; non-conductive.
  • Graphite: Each carbon atom is bonded to $3$ others in layers. One free electron per atom allows electrical conductivity. Used as a lubricant because layers slide. Insoluble; high melting point.
  • Fullerenes (e.g., $C_{60}$): Molecules like hollow balls. Insoluble; low melting point due to weak forces between molecules; non-conductive.
  • Graphene: Single sheets of carbon. Light, strong, and conducts electricity. Potential future technology use.

• Polymers: Large molecules containing long chains of carbon atoms, such as poly(ethene).

• Atomic Model Limitations:

  • Dot and Cross: Does not show 3D shape.
  • Displayed Formula: Does not show atom size or how bonds are actually formed.
  • 3D Model: Does not show details of bond formation or specific numbers of bonds clearly.

Calculations Involving Masses

• Relative Formula Mass (RFM):

  • The sum of relative atomic masses in a compound.
  • Example (Na2O): $(2 \times 23) + 16 = 62$.
  • Example (Mg(OH)2): $24 + [(16 + 1) \times 2] = 24 + 34 = 58$.

• The Mole (HIGHER ONLY):

  • One mole = Avogadro’s number ($6.02 \times 10^{23}$ particles).
  • One mole has a mass (in grams) equal to the atomic mass or RFM.
  • Formula: $\text{No. moles} = \frac{\text{mass (g)}}{\text{Relative formula mass}}$.

• Empirical Formula:

  • The simplest whole number ratio of atoms in a substance.
  • Example: Molecular formula $C_2H_6$ has an empirical formula of $CH_3$.
  • Calculating Empirical Formula (Mn example): $5.5\,g$ of Mn reacts with $3.2\,g$ of O.
    • Moles of Mn: $5.5 / 55 = 0.1$.
    • Moles of O: $3.2 / 16 = 0.2$.
    • Ratio: $1:2$. Formula = $MnO_2$.
  • Practical Finding (MgO): Burn magnesium in a crucible; weigh product; subtract Mg mass to find Oxygen mass; calculate ratio.

• Conservation of Mass:

  • Mass is never created or destroyed.
  • Closed System: Mass before equals mass after.
  • Non-enclosed System: Mass might appear to decrease if a gas is released, but the missing mass is simply in the escaped gas.

• Concentration:

  • $\text{Concentration (g/dm}^3\text{)} = \text{mass (g)} \div \text{volume (dm}^3\text{)}$.
  • $1\,dm^3 = 1000\,cm^3$.

• Stoichiometry and Limiting Reactants (HIGHER ONLY):

  • The mass of product formed is limited by the reactant that runs out first (the limiting reactant).
  • Example: Moles in $8\,g$ of $Na_2O = 8 / 62 = 0.13\,mol$.

States of Matter and Pure Substances

• Three States: Solid, liquid, and gas.

  • State changes are physical: arrangement, movement, and energy change, but particles stay the same.
  • Melting: Particles gain energy, bonds break, arrangement becomes random and free.

• Heating Curves:

  • Temperature does not increase during state change (melting/boiling points).
  • Energy is used to break bonds rather than raise temperature.

• Pure Substances vs. Mixtures:

  • Pure Substance: Fixed composition (e.g., pure gold, pure sodium chloride). Has a sharp melting point.
  • Mixture: Not chemically joined; contains various elements/compounds. Melts over a range of temperatures.

Separation and Purification

• Separation Methods:

  • Filtration: Separates insoluble solids from liquids (e.g., chalk and water).
  • Crystallisation: Separates a solute from a solvent by evaporating the liquid (e.g., salty water).
  • Distillation: Collects a liquid from a mixture by evaporating and then condensing it (e.g., collecting solvent from ink).
  • Fractional Distillation: Separates multiple liquids based on different boiling points (e.g., crude oil).
  • Chromatography: Separates soluble substances. Uses a stationary phase (paper) and a mobile phase (solvent). A pencil baseline is used because pencil lead is insoluble.

• Providing Potable Water:

  • Ground Water: Purified via sedimentation (impure particles settle), filtration, and chlorination (kills microorganisms).
  • Sea Water: Made potable via distillation.
  • Chemical Analysis: Requires pure water because dissolved salts would interfere with test results.

Chemical Change: Acids, Bases, and Salts

• Acids and Alkalis:

  • Acids: Source of hydrogen ions ($H^+$); pH < 7.
  • Alkalis: Source of hydroxide ions ($OH^-$); pH > 7. Soluble bases.

• Indicators:

  • Litmus: Red (acid), Blue (alkali).
  • Universal: Red/Yellow (acid), Blue/Purple (alkali).
  • Methyl Orange: Red (acid), Yellow (alkali).
  • Phenolphthalein: Colourless (acid), Pink (alkali).

• Concentration and pH (HIGHER ONLY):

  • An increase in $H^+$ concentration by a factor of $10$ decreases pH by $1$.
  • Strong Acids (e.g., HCl): Dissociate completely in water, releasing many $H^+$ ions.
  • Weak Acids (e.g., Ethanoic acid): Do not dissociate completely.

• Chemical Equations for Salts:

  • $\text{Metal} + \text{Acid} \rightarrow \text{Salt} + \text{Hydrogen}$
  • $\text{Metal oxide/hydroxide (Base)} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$
  • $\text{Metal carbonate} + \text{Acid} \rightarrow \text{Salt} + \text{Water} + \text{Carbon dioxide}$
  • Naming: HCl makes chlorides, $H_2SO_4$ makes sulphates, $HNO_3$ makes nitrates.

• Gas Tests:

  • Hydrogen: Lighted splint makes a ‘squeaky pop’.
  • Carbon Dioxide: Bubble through limewater; turns from colourless to milky/cloudy.

• Salt Preparation:

  • Neutralisation (Soluble Salt): Warm acid, add excess base (e.g., Copper Oxide), filter to remove excess, evaporate filtrate to crystallise.
  • Titration: For acid + alkali. Use pipette for acid and burette for alkali. Find the end point with indicator, then repeat without indicator to get pure salt and water.
  • Precipitation (Insoluble Salt): Mix two soluble salts (e.g., Silver nitrate + Sodium chloride), filter precipitate, rinse with distilled water, and dry.

Solubility Rules

• Soluble:

  • All Sodium, Potassium, and Ammonium salts.
  • All Nitrates.
  • Most Chlorides (except Silver and Lead).
  • Most Sulphates (except Lead, Barium, and Calcium).
  • Sodium, Potassium, and Ammonium carbonates and hydroxides.

• Insoluble:

  • Silver chloride, Lead chloride.
  • Lead sulphate, Barium sulphate, Calcium sulphate.
  • Most carbonates.
  • Most hydroxides.

Electrolysis

• Key Principles:

  • Electrolyte: An ionic substance with moving ions that conducts electricity (liquid or dissolved).
  • Process: Electrical energy from a d.c. supply decomposes electrolytes.
  • Cations (+): Move to the negative cathode.
  • Anions (-): Move to the positive anode.

• Electrolysis Products Table:

  • Copper chloride (aq): Cathode = Copper; Anode = Chlorine gas.
  • Sodium chloride (aq): Cathode = Hydrogen gas; Anode = Chlorine gas.
  • Sodium sulphate (aq): Cathode = Hydrogen gas; Anode = Oxygen gas.
  • Water (acidified): Cathode = Hydrogen gas; Anode = Oxygen gas.
  • Molten lead bromide ($PbBr_2$): Cathode = Lead (liquid); Anode = Bromine gas (orange).

• Redox in Electrolysis (HIGHER ONLY):

  • Oxidation: Loss of electrons (at the anode).
  • Reduction: Gain of electrons (at the cathode).
  • OIL RIG: Oxidation Is Loss, Reduction Is Gain.
  • Half Equations:
    • Cathode: $Cu^{2+} + 2e^- \rightarrow Cu$ or $2H^+ + 2e^- \rightarrow H_2$
    • Anode: $2Cl^- \rightarrow Cl_2 + 2e^-$ or $2O^{2-} \rightarrow O_2 + 4e^-$

• Purifying Copper:

  • If using copper electrodes, the anode (impure copper) loses mass and the cathode (pure copper) increases in mass.

Extracting Metals and Equilibria

• Reactivity and Redox:

  • Reactivity Series: Lists metals from most reactive (Potassium) to least.
  • Displacement: A more reactive metal displaces a less reactive one.
  • Oxidation: Gain of oxygen or loss of electrons.
  • Reduction: Loss of oxygen or gain of electrons.

• Extraction Methods:

  • Unreactive Metals: Found as pure elements (e.g., Gold).
  • Metals less reactive than Carbon: Extracted by heating with Carbon (e.g., Iron). Carbon reduces the metal oxide.
  • Metals more reactive than Carbon: Extracted via electrolysis (expensive, high energy).

• Alternative Extraction (HIGHER ONLY):

  • Bioleaching: Bacteria create toxic substances but extract metal.
  • Phytoextraction: Plants absorb metal; expensive and slow.

• Recycling and Life Cycle Assessment (LCA):

  • Pros: Resources last longer, less mining/landfill, less pollution.
  • LCA Stages: Raw materials, manufacturing, use, and disposal.

• Reversible Reactions and Equilibrium:

  • Symbol: $\rightleftharpoons$. Both directions happen simultaneously.
  • Dynamic Equilibrium: Forward and backward reactions occur at the same rate.
  • Haber Process: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$. Conditions: $450\,^\circ C$, $200\,atm$, Iron catalyst.
  • Influencing Equilibrium (HIGHER ONLY):
    • Temperature: Increasing it shifts in the endothermic direction.
    • Gas Pressure: Increasing it shifts toward the side with fewer gas molecules.
    • Concentration: Increasing it shifts toward the side that uses up the added substance.