Summary Notes

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• Introduction to A-level Chemistry: Resources for OCR (A) Chemistry A-level, Module 2: Foundations in Chemistry

  • Provided by PMT Education, authored by Amie Campbell.

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2.1 Atomic Structure and Isotopes

• Isotopes: Atoms of the same element with different numbers of neutrons and different masses.

• Definitions:

  • Mass number = protons (p) + neutrons (n)

  • Atomic number = number of protons

• Chemical Behavior: Isotopes react similarly as they have the same number of electrons, which participate in chemical reactions. Neutrons do not affect chemical reactions.

• Heavy Water:

  • Composed of deuterium (2H), with the formula D2O.

  • Higher melting point (m.p.), boiling point (b.p.), and density; leads to more frequent ice formation if all water were heavy water.

• Ions:

  • Cations: positively charged ions (e.g., Na+).

  • Anions: negatively charged ions (e.g., Cl-).

2.2 Relative Masses

• Mass Defect: Small mass lost due to the strong nuclear force binding protons and neutrons.

• Relative Isotopic Mass: Mass of an isotope relative to 1/12 mass of a carbon-12 atom.

• Relative Atomic Mass: Weighted mean mass of an atom of an element relative to carbon-12.

• Isotope Abundance Measurement with Mass Spectrometry:

  • Sample is vaporized and ionized to form positive ions.

  • Ions are accelerated, heavier ions move slower, allowing separation.

  • Ions detected as mass-to-charge ratio (m/z).

2.3 Formulae and Equations

• Binary Compounds:

  • Compounds that consist of two different elements.

  • Naming involves changing the second element's ending to -ide; metals come first in ionic compound names.

• Polyatomic Ions:

  • Ions with more than one element are bonded together (e.g., NH4+, OH-, NO3-, etc.).

• Diatomic Molecules:

  • Molecules composed of two atoms (e.g., H2, N2, O2, etc.).

• Other Small Molecules:

  • Commonly noted as P4 or S8.

3.1 Amount of Substance and the Mole

• Avogadro Constant: The number of particles in one mole of carbon-12 (6.02 x 10^23 particles).

• Molar Mass: The mass in grams of one mole of a substance (g mol-1).

• Formula Calculation:

  • n = m / Mr, where n = number of moles, m = mass, Mr = molar mass.

3.2 Determination of Formulae

• Molecular Formula vs. Empirical Formula:

  • Molecular formula: The actual number of atoms of each element in a molecule.

  • Empirical formula: Simplest whole number ratio of atoms in a compound.

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Relative Molecular Mass

• Compares mass of a molecule to mass of carbon-12.

• Relative Formula Mass: Compares mass of formula unit to mass of carbon-12.

• Chemical Analysis: Investigating the composition of substances.

• Water of Crystallization: The water molecules integrated into crystalline structures.

• Heating Hydrated Crystals:

  • Heating causes bonds to break and water to evaporate (e.g., CuSO4·5H2O (s) goes to CuSO4 (s) + 5H2O (l)).

• Finding Formula for Hydrated Salts:

  • Weigh the empty crucible, add hydrated salt, and reweigh.

  • Heat for 1 min strong, 3 mins gentle and cool before weighing again.

3.3 Moles and Volume

• Concentration: 1 mol/dm3 contains 1 mol of solute per 1 dm³ solution.

• Standard Solution: A solution with known concentration prepared by dissolving a precise solute mass in a solvent to a specific volume.

• Molar Gas Volume (Vm): The volume per mole of gas at standard temperature and pressure (approximately 24 dm³/mol at RTP).

• Gas Laws: pV = nRT

  • p = pressure (Pa), V = volume (m³), n = moles, R = ideal gas constant (8.314 J mol-1 K-1), T = temperature (K).

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3.4 Reacting Quantities

• Stoichiometry: The ratio of amounts of substances in a balanced equation.

• Theoretical Yield: Maximum product possible if all reactants convert to products.

• Causes for Yield Limitations:

  • Reactions may not complete, side reactions occur, purification loses product.

• Percentage Yield Calculation: (Actual yield / Theoretical yield) x 100.

• Limiting Reagent: The reactant that is consumed first, which stops the reaction.

• Calculating Atom Economy: Atom economy = (sum molar masses of desired product / sum molar masses of all products) x 100.

• Relevance of High Atom Economy: Indicates efficient use of materials; critical for sustainability.

4.1 Acids, Bases, and Neutralization

• Strong vs. Weak Acids:

  • Strong acids fully dissociate in solution; weak acids partially dissociate.

• Neutralization Reaction: H+ ions from acids react with OH- ions from bases, forming water and salts.

4.2 Acid-Base Titrations

• Purpose: Accurate measurement of volumes for chemical reactions.

• Procedure for Preparing Standard Solutions:

  • Solid is dissolved in water, transferred to a volumetric flask, and adjusted to the mark with distilled water.

• Titration Process: Add a solution to a conical flask, record initial readings, and swirl until reaching the endpoint.

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4.3 Redox Reactions

• Oxidation Numbers: Defined for common elements (e.g. O = -2, H = +1).

• Electron Transfer:

  • Reduction: Gain of electrons; oxidation: loss of electrons.

5.1 Electron Structure

• Energy Levels: Shells represent energy levels, with higher levels corresponding to increased energy.

• Atomic Orbitals: Regions around the nucleus for 2 electrons (e.g., s-orbitals are spherical; p-orbitals are dumbbell-shaped).

5.2 Ionic Bonding and Structure

• Ionic Bonding: Attraction between positive and negative ions, yielding high melting and boiling points.

• Solubility Factors: Depends on strengths of attraction within the lattice and interactions with solvents.

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Conductivity:

• Solid state: Ions are fixed and do not conduct electricity.

• Liquid/molten: Ions are free to move, allowing conduction.

5.3 Covalent Bonding

• Definition: Strong attraction between shared electron pairs and the nuclei of participating atoms.

• Bond Types:

  • Single, double, triple, dative bonds are distinguished by the number of shared pairs of electrons.

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6.1 Shapes of Molecules and Ions

• Electron Pair Repulsion Theory: Shapes determined by positioning to minimize electron repulsion.

• Molecular Geometry: Characterized by the arrangement of bonding pairs and lone pairs.

6.2 Electronegativity and Polarity

• Electronegativity: Measure of the attraction of bonded atoms for shared electrons, increasing across periods and decreasing down groups.

• Bond Polarity:

  • Non-polar bonds share electrons equally; polar bonds share unequally, creating dipoles.

  • Permanent Dipoles: Exist in polar covalent bonds, creating partial charges.

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6.3 Intermolecular Forces

• Types of Forces:

  • London forces, permanent dipole-dipole interactions, and hydrogen bonding explain behavior of molecules under various conditions.

• Solubility Trends: Polar molecules typically dissolve in polar solvents; nonpolar molecules dissolve in nonpolar solvents.

6.4 Hydrogen Bonding

• Hydrogen Bonds: Occur between molecules containing electronegative atoms with lone pairs attached to hydrogen.

• Unique Water Properties: Hydrogen bonds cause unusual characteristics like lower density of ice compared to liquid water.