Chemistry Notes

States of Matter

• 1.1 States of Matter:

  • Three states of matter are understood based on:
    • Arrangement of particles.
    • Movement of particles.
    • Energy of particles.

• 1.2 Interconversions:

  • Interconversions between the three states of matter involve:
    • Names of the interconversions (e.g., melting, boiling, freezing, condensation, sublimation, deposition).
    • How interconversions are achieved (e.g., heating, cooling).
    • Changes in arrangement, movement, and energy of particles during these interconversions.

• 1.3 Dilution and Diffusion:

  • Experimental results from:
    • Dilution of colored solutions.
    • Diffusion of gases.
  • These experiments can be explained by particle theory.

• 1.4 Definitions:

  • Key terms:
    • Solvent: A substance that dissolves a solute.
    • Solute: A substance that is dissolved in a solvent.
    • Solution: A homogeneous mixture of a solute and a solvent.
    • Saturated Solution: A solution containing the maximum amount of solute that can dissolve at a given temperature.

Elements, Compounds, and Mixtures

• 1.5C Solubility (Higher Tier):

  • Solubility is the amount of solute that dissolves in a given amount of solvent.
  • Units: grams per 100 grams of solvent (g/100g).

• 1.6C Solubility Curves (Higher Tier):

  • Plotting and interpreting solubility curves.

• 1.7C Practical (Higher Tier):

  • Investigating the solubility of a solid in water at a specific temperature.

• 1.8 Classification:

  • Classifying substances as:
    • Element: A substance made of only one type of atom.
    • Compound: A substance made of two or more types of atoms chemically bonded together.
    • Mixture: A combination of two or more substances that are physically combined.

• 1.9 Pure Substances vs. Mixtures:

  • Pure substance:
    • Fixed melting and boiling points.
  • Mixture:
    • Melts or boils over a range of temperatures.

• 1.10 Separation Techniques:

  • Experimental techniques for separating mixtures:
    • Simple distillation: separates a liquid from a solution.
    • Fractional distillation: separates multiple liquids with different boiling points.
    • Filtration: separates an insoluble solid from a liquid.
    • Crystallization: obtains pure solid crystals from a solution.
    • Paper chromatography: separates substances based on their differing solubilities and adsorption to the paper.

• 1.11 Chromatography Information:

  • A chromatogram provides information about the composition of a mixture.

• 1.12 Rf Values:

  • Calculating R_f values to identify the components of a mixture.
  • R_f = \frac{Distance travelled by the substance}{Distance travelled by the solvent}

• 1.13 Practical:

  • Investigating paper chromatography using inks/food colorings.

• 1.14 Atoms and Molecules:

  • Atom: The smallest particle of an element that can exist.
  • Molecule: Two or more atoms held together by chemical bonds.

Atomic Structure

• 1.15 Atomic Structure (Subatomic Particles):

  • Structure of an atom:
    • Protons: located in the nucleus, relative mass of 1, relative charge of +1
    • Neutrons: located in the nucleus, relative mass of 1, relative charge of 0
    • Electrons: located in shells around the nucleus, relative mass of 1/1836 (negligible), relative charge of -1

• 1.16 Definitions:

  • Atomic Number: The number of protons in the nucleus of an atom.
  • Mass Number: The total number of protons and neutrons in the nucleus of an atom.
  • Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons.
  • Relative Atomic Mass (A_r): The weighted average mass of the isotopes of an element compared to 1/12th the mass of a carbon-12 atom.

• 1.17 Calculating Relative Atomic Mass:

  • Calculating A_r from isotopic abundances:
    • A_r = \frac{(\% { of isotope 1} {\times} {mass of isotope 1}) + (\% { of isotope 2} {\times} {mass of isotope 2}) + …}{100}

The Periodic Table

• 1.18 Arrangement of Elements:

  • In order of atomic number.
  • Arranged in groups (vertical columns) and periods (horizontal rows).

• 1.19 Electronic Configurations:

  • Deducing the electronic configurations of the first 20 elements from their positions in the Periodic Table.

• 1.20 Metals vs. Non-metals:

  • Using electrical conductivity and the acid-base character of oxides to classify elements as metals or non-metals.

• 1.21 Identifying Metals and Non-metals:

  • Identifying an element as a metal or non-metal according to its position in the Periodic Table.

• 1.22 Group Relationship:

  • Understanding how the electronic configuration of a main group element is related to its position in the Periodic Table.

• 1.23 Similar Chemical Properties:

  • Understanding why elements in the same group have similar chemical properties (same number of valence electrons).

• 1.24 Noble Gases:

  • Understanding why the noble gases (Group 0/18) do not readily react (full outer electron shell).

Ionic Bonding

• 1.37 Ion Formation:

  • Ions are formed by the loss or gain of electrons.
    • Metals lose electrons to form positive ions (cations).
    • Non-metals gain electrons to form negative ions (anions).

• 1.38 Charges of Ions:

  • Common ion charges:
    • Metals in Groups 1, 2, and 3: +1, +2, and +3 respectively.
    • Non-metals in Groups 5, 6, and 7: -3, -2, and -1 respectively.
    • Silver (Ag+), Copper(II) (Cu2+), Iron(II) (Fe2+), Iron(III) (Fe3+), Lead(II) (Pb2+), Zinc (Zn2+)
    • Hydrogen (H+), Hydroxide (OH-), Ammonium (NH4+), Carbonate (CO32-), Nitrate (NO3-), Sulfate (SO42-)

• 1.39 Formulae of Compounds:

  • Writing formulae for compounds formed between the ions listed above (ensuring the overall charge is zero).

• 1.40 Dot-and-Cross Diagrams:

  • Drawing dot-and-cross diagrams to show the formation of ionic compounds by electron transfer.
  • Limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7.
  • Only outer electrons need to be shown.

• 1.41 Ionic Bonding:

  • Understanding ionic bonding in terms of electrostatic attractions between oppositely charged ions.

• 1.42 High Melting and Boiling Points:

  • Understanding why compounds with giant ionic lattices have high melting and boiling points (strong electrostatic forces).

• 1.43 Electrical Conductivity:

  • Ionic compounds do not conduct electricity when solid (ions are fixed in position).
  • Ionic compounds do conduct electricity when molten or in aqueous solution (ions are free to move).

Covalent Bonding

• 1.44 Covalent Bond Formation:

  • A covalent bond is formed between atoms by the sharing of a pair of electrons.

• 1.45 Covalent Bonds (Electrostatic Attractions):

  • Understanding covalent bonds in terms of electrostatic attractions between the positive nuclei and the shared negative electrons.

• 1.46 Dot-and-Cross Diagrams:

  • Using dot-and-cross diagrams to represent covalent bonds in:
    • Diatomic molecules: hydrogen, oxygen, nitrogen, halogens, and hydrogen halides.
    • Inorganic molecules: water, ammonia, and carbon dioxide.
    • Organic molecules (up to two carbon atoms): methane, ethane, ethene, and those containing halogen atoms.

• 1.47 Simple Molecular Structures:

  • Explaining why substances with simple molecular structures are gases or liquids, or solids with low melting and boiling points (weak intermolecular forces).
  • The term 'intermolecular forces of attraction' represents all forces between molecules.

• 1.48 Melting/Boiling Points and Molecular Mass:

  • Explaining why the melting and boiling points of substances with simple molecular structures increase, in general, with increasing relative molecular mass (stronger intermolecular forces).

• 1.49 Giant Covalent Structures:

  • Explaining why substances with giant covalent structures are solids with high melting and boiling points (strong covalent bonds throughout the structure).

• 1.50 Properties of Allotropes:

  • Explaining how the structures of diamond, graphite, and C_{60} fullerene influence their physical properties, including electrical conductivity and hardness.

Metallic Bonding

• 1.52C Metallic Lattice (Higher Tier):

  • Representing a metallic lattice by a 2-D diagram (atoms arranged in regular layers).

• 1.53C Metallic Bonding (Higher Tier):

  • Understanding metallic bonding in terms of electrostatic attractions between positive metal ions and delocalized electrons.

• 1.54C Properties of Metals (Higher Tier):

  • Explaining typical physical properties of metals, including electrical conductivity and malleability, due to the delocalized electrons.

Chemical Formulae, Equations, and Calculations

• 1.32 Empirical and Molecular Formulae:

  • Empirical Formula: The simplest whole number ratio of atoms of each element in a compound.
  • Molecular Formula: The actual number of atoms of each element in a molecule.

• 1.33 Formula Calculations:

  • Calculating empirical and molecular formulae from experimental data.

• 1.36 Practical:

  • Determining the formula of a metal oxide by combustion (e.g., magnesium oxide) or by reduction (e.g., copper(II) oxide).

• 1.26 Relative Formula Mass:

  • Calculating relative formula masses (including relative molecular masses) (Mr) from relative atomic masses (Ar).

• 1.25 Chemical Equations:

  • Writing word equations and balanced chemical equations (including state symbols):
    • For reactions studied in this specification
    • For unfamiliar reactions where suitable information is provided

Group 1 (Alkali Metals)

• 2.1 Reactions with Water:

  • Understanding how the similarities in the reactions of these elements with water provide evidence for their recognition as a family of elements.

• 2.2 Reactivity Trends:

  • Understanding how the differences between the reactions of these elements with air and water provide evidence for the trend in reactivity in Group 1.

• 2.3 Property Prediction:

  • Using knowledge of trends in Group 1 to predict the properties of other alkali metals.

• 2.4C Explanation of Reactivity (Higher Tier):

  • Explaining the trend in reactivity in Group 1 in terms of electronic configurations (ease of losing the outer electron).

Group 7 (Halogens)

• 2.5 Physical Properties:

  • Knowing the colors, physical states (at room temperature), and trends in physical properties (e.g., boiling point) of these elements.
    • Fluorine (F2) - pale yellow gas
    • Chlorine (Cl2) - green gas
    • Bromine (Br2) - red-brown liquid
    • Iodine (I2) - grey solid

• 2.6 Property Prediction:

  • Using knowledge of trends in Group 7 to predict the properties of other halogens.

• 2.7 Displacement Reactions:

  • Understanding how displacement reactions involving halogens and halides provide evidence for the trend in reactivity in Group 7.

• 2.8C Explanation of Reactivity (Higher Tier):

  • Explaining the trend in reactivity in Group 7 in terms of electronic configurations (ease of gaining an electron).

Gases in the Atmosphere

• 2.9 Composition of Air:

  • Knowing the approximate percentages by volume of the four most abundant gases in dry air:
    • Nitrogen (~78%)
    • Oxygen (~21%)
    • Argon (~0.9%)
    • Carbon Dioxide (~0.04%)

• 2.10 Oxygen Percentage Determination:

  • Understanding how to determine the percentage by volume of oxygen in air using experiments involving the reactions of metals (e.g., iron) and non-metals (e.g., phosphorus) with air.

• 2.11 Combustion in Oxygen:

  • Describing the combustion of elements in oxygen, including magnesium, hydrogen, and sulphur.
    • 2Mg(s) + O_2(g) \rightarrow 2MgO(s)
    • 2H2(g) + O2(g) \rightarrow 2H_2O(g)
    • S(s) + O2(g) \rightarrow SO2(g)

• 2.14 Practical:

  • Determining the approximate percentage by volume of oxygen in air using a metal or a non-metal.

• 2.12 Carbon Dioxide Formation:

  • Describing the formation of carbon dioxide from the thermal decomposition of metal carbonates, including copper(II) carbonate.
    • CuCO3(s) \rightarrow CuO(s) + CO2(g)

• 2.13 Greenhouse Gas:

  • Carbon dioxide is a greenhouse gas, and increasing amounts in the atmosphere may contribute to climate change.

Reactivity Series

• 2.15 Reactivity Series from Reactions with Acids and Water:

  • Understanding how metals can be arranged in a reactivity series based on their reactions with:
    • Water
    • Dilute hydrochloric or sulfuric acid

• 2.16 Reactivity Series from Displacement Reactions:

  • Understanding how metals can be arranged in a reactivity series based on their displacement reactions between:
    • Metals and metal oxides
    • Metals and aqueous solutions of metal salts

• 2.17 Order of Reactivity:

  • Knowing the order of reactivity of these metals: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver, gold.

• 2.21 Practical:

  • Investigating reactions between dilute hydrochloric and sulfuric acids and metals (e.g., magnesium, zinc, and iron).

• 2.18 Rusting Conditions:

  • Knowing the conditions under which iron rusts: presence of both oxygen and water.

• 2.19 Rust Prevention:

  • Understanding how the rusting of iron may be prevented by:
    • Barrier Methods: painting, coating with plastic, or greasing.
    • Galvanizing: coating with zinc.
    • Sacrificial Protection: using a more reactive metal (e.g., magnesium) to corrode instead of the iron.

• 2.20 Redox Reactions:

  • In terms of gain or loss of oxygen:
    • Oxidation: Gain of oxygen.
    • Reduction: Loss of oxygen.
  • In terms of loss or gain of electrons:
    • Oxidation: Loss of electrons (OIL).
    • Reduction: Gain of electrons (RIG).
  • Redox: A reaction involving both oxidation and reduction.
  • Oxidizing agent: A substance that causes oxidation (and is itself reduced).
  • Reducing agent: A substance that causes reduction (and is itself oxidized).

Extraction and Uses of Metals

• 2.22C Metal Extraction (Higher Tier):

  • Most metals are extracted from ores found in the Earth's crust.
  • Unreactive metals are often found as the uncombined element.

• 2.23C Extraction Methods (Higher Tier):

  • Explaining how the method of extraction of a metal is related to its position in the reactivity series, illustrated by carbon extraction for iron and electrolysis for aluminium.
    • Carbon Extraction: Used for metals less reactive than carbon.
    • Electrolysis: Used for highly reactive metals (e.g., aluminium).

• 2.24C Metal Extraction Commentary (Higher Tier):

  • Being able to comment on a metal extraction process, given appropriate information.

• 2.25C Uses of Metals (Higher Tier):

  • Explaining the uses of aluminium, copper, iron, and steel in terms of their properties.
    • Aluminium: low density, corrosion resistance (aircraft, drink cans)
    • Copper: high electrical conductivity (electrical wiring)
    • Iron: strong but rusts easily (construction)
    • Steel: Alloys of iron with improved properties.
      • Low-carbon (mild) steel: Malleable and ductile.
      • High-carbon steel: Hard but brittle.
      • Stainless steel: Corrosion-resistant.

• 2.26C Alloys (Higher Tier):

  • An alloy is a mixture of a metal and one or more elements, usually other metals or carbon.

• 2.27C Alloy Hardness (Higher Tier):

  • Explaining why alloys are harder than pure metals (different sized atoms disrupt the regular arrangement).

Acids, Alkalis, and Titrations

• 2.28 Indicators:

  • Describing the use of litmus, phenolphthalein, and methyl orange to distinguish between acidic and alkaline solutions.
    • Litmus: Red in acid, blue in alkali.
    • Phenolphthalein: Colorless in acid, pink in alkali.
    • Methyl Orange: Red in acid, yellow in alkali.

• 2.29 pH Scale:

  • Understanding how to use the pH scale, from 0-14, to classify solutions as:
    • Strongly acidic (0-3)
    • Weakly acidic (4-6)
    • Neutral (7)
    • Weakly alkaline (8-10)
    • Strongly alkaline (11-14)

• 2.30 Universal Indicator:

  • Describing the use of universal indicator to measure the approximate pH value of an aqueous solution (shows a range of colors).

• 2.31 Ions in Solutions:

  • Acids in aqueous solution are a source of hydrogen ions (H^+).
  • Alkalis in aqueous solution are a source of hydroxide ions (OH^-).

• 2.32 Neutralization:

  • Alkalis can neutralize acids.

• 2.33C Acid-Alkali Titration (Higher Tier):

  • Describing how to carry out an acid-alkali titration (to determine the concentration of an acid or alkali).

Acids, Bases, and Salt Preparations

• 2.34 Solubility Rules:

  • General rules for predicting the solubility of ionic compounds in water:
    • All common sodium, potassium, and ammonium compounds are soluble.
    • All nitrates are soluble.
    • Common chlorides are soluble, except those of silver (AgCl) and lead(II) (PbCl_2).
    • Common sulfates are soluble, except for those of barium (BaSO4), calcium (CaSO4), and lead(II) (PbSO_4).
    • Common carbonates are insoluble, except for those of sodium, potassium, and ammonium.
    • Common hydroxides are insoluble, except for those of sodium, potassium, and calcium (calcium hydroxide is slightly soluble).

• 2.35 Acid/Base Definition:

  • Understanding acids and bases in terms of proton transfer.

• 2.36 Proton Donor/Acceptor:

  • An acid is a proton donor.
  • A base is a proton acceptor.

• 2.37 Reactions of Acids:

  • Describing the reactions of hydrochloric acid (HCl), sulfuric acid (H2SO4), and nitric acid (HNO_3) with metals, bases, and metal carbonates (excluding the reactions between nitric acid and metals) to form salts.

• 2.38 Types of Bases:

  • Metal oxides, metal hydroxides, and ammonia can act as bases.
  • Alkalis are bases that are soluble in water.

• 2.39 Salt Preparation:

  • Describing an experiment to prepare a pure, dry sample of a soluble salt, starting from an insoluble reactant (e.g., reacting an insoluble metal oxide with an acid).