Comprehensive GCSE Chemistry: Chemical Changes, Atmosphere, and Bonding Overview

Acids, Bases, and the pH Scale

• The pH Scale Overview:

  • The pH scale is used to measure how acidic or alkaline a solution is.

  • The range of the scale is from $0$ to $14$.

  • Lower pH values indicate higher acidity; higher pH values indicate higher alkalinity.

  • A neutral substance, such as pure water, has a pH of exactly $7$.

• Measuring pH:

  • Indicators: Dyes that change color based on the pH of the solution. Some are "wide range indicators" which contain a mixture of dyes to provide gradual color changes over the entire pH scale.

  • Universal Indicator: A common wide range indicator that follows a red-to-green-to-purple scale.

  • pH Probes and Meters: These tools measure pH electronically by attaching a probe to a meter, providing numerical readings (e.g., $10.25$).

• pH Examples from Industry and Nature:

  • pH $0$: Car battery acid.

  • pH $1$: Stomach acid.

  • pH $3$: Vinegar, lemon juice.

  • pH $4$: Acid rain.

  • pH $5$: Normal rain. (Note: Human skin is slightly acidic at pH $5.5$).

  • pH $7$: Pure water.

  • pH $8$: Pancreatic juice.

  • pH $9$: Soap powder.

  • pH $10$: Washing-up liquid.

  • pH $12$: Bleach.

  • pH $14$: Caustic soda (drain cleaner).

• Chemical Definitions:

  • Acids: Form aqueous solutions containing hydrogen ions ($H^+$). Their pH is less than $7$.

  • Bases: Substances that react with acids to neutralize them.

  • Alkalis: Bases that are soluble in water and form solutions containing hydroxide ions ($OH^-$). Their pH is greater than $7$.

• Neutralization Reactions:

  • The general reaction: $\text{acid} + \text{base} \rightarrow \text{salt} + \text{water}$.

  • The products of neutralization (pure salt and water) have a pH of $7$.

  • The ionic equation for neutralization between an acid and an alkali is:
            $H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)}$

The Reactivity Series and Metal Reactions

• Reactivity Principles:

  • The reactivity series lists metals in order of their reactivity toward other substances.

  • Metals react by losing electrons to form positive ions. The more easily a metal does this, the higher its position in the series.

• The Reactivity Series Order (High to Low):

  • Potassium ($K$) - Very reactive.

  • Sodium ($Na$) - Very reactive.

  • Lithium ($Li$) - Very reactive.

  • Calcium ($Ca$) - Very reactive.

  • Magnesium ($Mg$) - Fairly reactive.

  • Carbon ($C$) - Included for comparison (non-metal).

  • Zinc ($Zn$) - Fairly reactive.

  • Iron ($Fe$) - Fairly reactive.

  • Hydrogen ($H$) - Included for comparison (non-metal).

  • Copper ($Cu$) - Not very reactive.

• Reaction with Acids:

  • General equation: $\text{Acid} + \text{Metal} \rightarrow \text{Salt} + Hydrogen$

  • The speed of reaction is shown by the rate of hydrogen bubble production.

  • Explosive reactions: Potassium, Sodium, Lithium, and Calcium.

  • Vigorous reactions: Magnesium reacts quickly with cold dilute $HCl$ or $H_2SO_4$.
            $Mg_{(s)} + 2HCl_{(aq)} \rightarrow MgCl_{2(aq)} + H_{2(g)}$
            $Mg_{(s)} + H_2SO_{4(aq)} \rightarrow MgSO_{4(aq)} + H_{2(g)}$

  • Slow reactions: Zinc and Iron react slowly with cold dilute acid but more strongly if heated.

  • No reaction: Copper generally does not react with cold dilute acids.

• Reaction with Water:

  • General equation: $\text{Metal} + \text{Water} \rightarrow \text{Metal Hydroxide} + \text{Hydrogen}$.

  • More reactive metals ($K$, $Na$, $Li$, $Ca$) react with water at room temperature.

  • Example: $Ca_{(s)} + 2H_2O_{(l)} \rightarrow Ca(OH)_{2(aq)} + H_{2(g)}$.

  • Less reactive metals ($Zn$, $Fe$, $Cu$) will not react with water.

Oxidation, Reduction, and Metal Extraction

• Key Definitions (Oxygen-based):

  • Oxidation: Gain of oxygen (e.g., $2Mg + O_2 \rightarrow 2MgO$).

  • Reduction: Loss of oxygen.

• Metal Ores:

  • Metals like iron and aluminum react with oxygen in the ground to form oxides (ores).

  • Extraction requires separating the metal from oxygen (reduction).

• Extraction by Reduction with Carbon:

  • Metal oxides are reacted with carbon. Carbon takes the oxygen away to form carbon dioxide, leaving pure metal.

  • Example (Iron extraction): $2Fe_2O_3 + 3C \rightarrow 4Fe + 3CO_2$.

  • Rule: Carbon can only extract metals below it in the reactivity series. Metals above carbon require more expensive methods like electrolysis.

  • Unreactive Metals: Gold is so unreactive that it is found in the earth as its elemental form and does not need chemical extraction.

Redox Reactions and Equations

• Electron Transfer (OIL RIG):

  • Oxidation Is Loss (of electrons).

  • Reduction Is Gain (of electrons).

  • These occur simultaneously in "Redox" reactions.

• Half Equations:

  • Show only one part of the reaction (either oxidation or reduction).

  • Example: For the reaction between sodium and acid:

    • Oxidation: $Na \rightarrow Na^+ + e^-$

    • Reduction: $2H^+ + 2e^- \rightarrow H_2$

  • To combine them, the number of electrons must be the same: $2Na + 2H^+ \rightarrow 2Na^+ + H_2$.

• Displacement Reactions:

  • A more reactive metal displaces a less reactive metal from its compound.

  • Example: Iron added to copper sulfate solution produces iron sulfate and copper metal.         $Fe_{(s)} + CuSO_{4(aq)} \rightarrow FeSO_{4(aq)} + Cu_{(s)}$

  • In displacement, the metal ion is reduced (gains electrons), and the metal atom is oxidized (loses electrons).

• Ionic Equations:

  • Only show the particles that change during a reaction.

  • Spectator ions (those that do not change, like $SO_4^{2-}$ in the above example) are crossed out.

  • The ionic equation for iron displacing copper is: $Fe_{(s)} + Cu^{2+}_{(aq)} \rightarrow Fe^{2+}_{(aq)} + Cu_{(s)}$.

Electrolysis

• The Process:

  • An electric current is passed through an electrolyte (a molten or dissolved ionic compound).

  • Electrodes: Inert solid rods submerged in the electrolyte.

  • Cahode: Negative electrode; attracts positive ions; reduction occurs.

  • Anode: Positive electrode; attracts negative ions; oxidation occurs.

• Electrolysis of Molten Ionic Compounds:

  • Used to extract metals more reactive than carbon.

  • Example (Aluminum from Bauxite): Aluminum oxide ($Al_2O_3$) is mixed with cryolite to lower its extremely high melting point.

  • Cathode: $Al^{3+} + 3e^- \rightarrow Al_{(l)}$.

  • Anode: $2O^{2-} \rightarrow O_2 + 4e^-$. Carbon anodes must be replaced regularly as they react with oxygen to form $CO_2$.

  • Overall: $2Al_2O_{3(l)} \rightarrow 4Al_{(l)} + 3O_{2(g)}$.

• Electrolysis of Aqueous Solutions:

  • Solutions contain $H^+$ and $OH^-$ ions from water beside the salt ions.

  • Cathode rule: If the metal is more reactive than hydrogen, $H_2$ gas is produced ($2H^+ + 2e^- \rightarrow H_{2(g)}$). If less reactive (e.g., Copper), the metal forms.

  • Anode rule: If halide ions ($Cl^-$, $Br^-$, $I^-$) are present, the halogen is produced. If not, oxygen is produced from hydroxide ions.

Chemical Analysis and Gas Tests

• Chromatography:

  • Used to separate mixtures (like dyes).

  • Mobile Phase: The solvent (liquid or gas) where molecules move.

  • Stationary Phase: The paper or solid where molecules move slowly.

  • Substances separate based on how they distribute between phases. More soluble substances travel further.

• Tests for Common Gases:

  • Chlorine: Damp litmus paper bleaches white.

  • Oxygen: A glowing splint relights.

  • Carbon Dioxide: Bubbling through limewater turns the solution cloudy.

  • Hydrogen: A lit splint creates a "squeaky pop" sound.

Evolution of the Atmosphere

• Phase 1 (Volcanoes): The early atmosphere (4.6 billion years ago) was mostly $CO_2$ with little oxygen, similar to Mars/Venus. Volcanoes released nitrogen, water vapor, methane, and ammonia.

• Phase 2 (Oceans and Life): Water vapor condensed to form oceans. $CO_2$ dissolved and formed carbonate precipitates/sediments. Algae (2.7 billion years ago) and plants evolved, absorbing $CO_2$ for photosynthesis.

• Locked-up Carbon: Carbon became trapped in fossil fuels (coal, crude oil, gas) and sedimentary rocks (limestone from shells/skeletons).

• Phase 3 (Oxygen build-up): Plants produced oxygen via photosynthesis:
      $6CO_2 + 6H_2O \xrightarrow{\text{light}} C_6H_{12}O_6 + 6O_2$

• Modern Composition (200 million years ago to present):
  

  • $\sim 80\%$ Nitrogen.

  • $\sim 20\%$ Oxygen.

  • $\lt 1\%$ Other gases (Noble gases, $CO_2$, water vapor).

Atomic Structure and the Periodic Table

• Electronic Structure Rules:

  • Electrons occupy shells (energy levels).

  • 1st shell: Max $2$ electrons.

  • 2nd shell: Max $8$ electrons.

  • 3rd shell: Max $8$ electrons.

  • Example: Sodium (Atomic number $11$) is $2, 8, 1$.

• Development of the Periodic Table:

  • Early 1800s: Arranged by physical properties and atomic weight.

  • Dmitri Mendeleev (1869): Arranged 50 elements by atomic weight but switched order to group similar properties. He left gaps for undiscovered elements (e.g., ekasilicon/germanium) and predicted their properties accurately.

• Group Characteristics:

  • Group 1 (Alkali Metals): One outer electron. Highly reactive and soft. Reactivity increases as you move down the group because the outer electron is further from the nucleus (weaker attraction).

  • Group 7 (Halogens): Seven outer electrons. Exist as diatomic molecules (e.g., $Cl_2$). Reactivity decreases down the group. Melting/Boiling points increase down the group.

  • Group 0 (Noble Gases): Full outer shells. Unreactive and non-flammable. Boiling points increase as you move down the group.

Bonding and Matter

• Ionic Bonding:

  • Occurs between metals (lose electrons to form cations) and non-metals (gain electrons to form anions).

  • Structure: Giant ionic lattice with strong electrostatic forces in all directions.

  • Properties: High melting points; conduct electricity only when molten or dissolved.

• Covalent Bonding:

  • Non-metals share pairs of electrons.

  • Simple Molecular: Strong covalent bonds within molecules but weak intermolecular forces between them. Low melting points.

  • Giant Covalent (Macromolecules): All atoms linked by strong covalent bonds. High melting points. Examples: Diamond (each Carbon forms 4 bonds), Graphite (3 bonds, layers, delocalised electrons), Silicon Dioxide (sand).

• Allotropes of Carbon:

  • Diamond: Very hard, no free electrons.

  • Graphite: Soft, slippery (layers slide), conducts electricity (delocalised electrons).

  • Graphene: Single layer of graphite; strong and light.

  • Fullerenes: Spheres (e.g., Buckminsterfullerene $C_{60}$) or tubes (nanotubes). Used for drug delivery, catalysts, and strengthening materials.

• States of Matter:

  • Solids: Strong attraction, fixed regular lattice, particles vibrate.

  • Liquids: Weak attraction, random arrangement, particles move past each other.

  • Gases: Very weak attraction, far apart, move randomly in straight lines.

  • State Symbols: $(s)$ solid, $(l)$ liquid, $(g)$ gas, $(aq)$ aqueous (dissolved in water).

Questions & Discussion

• Q: What color would universal indicator turn in lemon juice?

  • A: Red or orange, as it is acidic (pH $3$).

• Q: Is a solution with pH $8$ acidic or alkaline?

  • A: Alkaline.

• Q: Why can't calcium be extracted from calcium oxide by reduction using carbon?

  • A: Because calcium is more reactive than carbon; carbon cannot displace oxygen from a more reactive metal.

• Q: Why does graphite conduct electricity and diamond does not?

  • A: Graphite has one delocalised electron per carbon atom that is free to move. In diamond, all four outer electrons are used in rigid covalent bonds.

• Q: How do you test for Carbon Dioxide?

  • A: Bubble it through limewater; it turns cloudy.