OCR (A) Chemistry A-Level Module 2 Quiz

OCR (A) Chemistry A-Level Module 2: Foundations in Chemistry

• This module covers fundamental chemistry concepts, including atomic structure, formulae, amount of substance, and moles.

2.1 Atomic Structure and Isotopes

• Isotopes are atoms of the same element with different numbers of neutrons.
• Mass number (A) = number of protons (p) + number of neutrons (n)
• Atomic number (Z) = number of protons (p)
• Different isotopes react similarly because their chemical behavior depends on the number of electrons, which is the same for all isotopes of an element.
• Heavy water (D₂O) is an example of water containing deuterium (²H) isotopes, impacting physical properties like melting and boiling points.

2.2 Relative Masses

• Mass defect: a small amount of mass lost due to strong nuclear forces holding protons and neutrons together.
• Relative isotopic mass: mass of an isotope relative to 1/12 the mass of carbon-12.
• Relative atomic mass: weighted average mass of an element's isotopes relative to 1/12 the mass of carbon-12. This reflects the abundance of each isotope.
• Mass spectrometry is used to determine isotope abundances.

2.3 Formulae and Equations

• Binary compounds contain two elements.
• Polyatomic ions are ions with more than one element, e.g., ammonium (NH₄⁺), hydroxide (OH⁻), nitrate (NO₃⁻), etc.
• Know how to name and write the formulae for common ions and compounds (including those with polyatomic ions).
• Diatomic molecules consist of two atoms, e.g. H₂, N₂, O₂, etc.

3.1 Amount of Substance and the Mole

• Avogadro constant: 6.02 x 10²³ particles per mole of a substance (e.g., carbon-12).
• Mole (mol): a unit for amount of substance.
• Molar mass: mass in grams of one mole of a substance (units: g mol⁻¹).
• The relationship n=m/Mr is a key way to calculate amount of substance.

3.2 Determination of Formulae

• Molecular formula: shows the actual number of atoms of each element in a molecule.
• Empirical formula: simplest whole-number ratio of atoms of each element in a compound. Useful for ionic compounds but not for molecular compounds.

3.3 Moles and Volume

• Standard solution: a solution of known concentration.
• Molar gas volume (Vm): volume occupied by one mole of a gas at a specified temperature and pressure (RTP).
• Ideal gas equation (PV = nRT) is crucial for calculations involving gases and moles.

3.4 Reacting Quantities

• Stoichiometry is the quantitative relationship between reactants and products in a chemical reaction.
• Limiting reactant is calculated and used in calculations.
• Atom economy: measures how many atoms from a reactant are used to produce a desired product in a reaction. A higher atom economy is more efficient and environmentally friendly.

4.1 Acids, Bases, and Neutralisation

• Strong acids completely dissociate in water, weak acids partially dissociate.
• Alkalis are bases that dissolve in water, releasing hydroxide ions (OH⁻) into solution.
• Neutralisation involves the reaction of H⁺ ions (from acid) with OH⁻ ions (from base) to form water (H₂O)

4.2 Acid-Base Titrations

• Titration is used to accurately measure the volume of one solution needed to react completely with another solution.
• Preparation of standard solutions is essential in titration experiments.
• Volumetric flasks and burettes are important equipment in titrations.

4.3 Redox Reactions

• Oxidation is loss of electrons and increase in oxidation number.
• Reduction is gain of electrons and decrease in oxidation number.
• Oxidation numbers: rules for assigning oxidation numbers to elements in compounds/ions.

5.1 Electron Structure

• Electron shells/energy levels (n): Electrons exist around the nucleus in energy levels.
• Atomic orbitals (s, p): Regions where electrons are most likely to be found.
• Orbital filling rules: Rules for predicting which atomic orbitals are filled, often summarized in terms of electronic configuration.

5.2 lonic Bonding and Structure

• Ionic bonding: electrostatic attraction between oppositely charged ions.
• lonic compounds are composed of positive cations and negative anions.

5.3 Covalent Bonding

• Covalent bonding: sharing of electron pairs between atoms to achieve a stable electron configuration (often noble gas configuration).
• Types of covalent bonds: single, double, triple.

6.1 Shapes of Molecules and lons

• VSEPR (Valence Shell Electron Pair Repulsion) theory: predicts the shapes of molecules/ions.
• How to draw Lewis dot structure and molecular shape.

6.2 Electronegativity and Polarity

• Electronegativity: tendency of an atom to attract shared electrons in a covalent bond.
• Polar bonds: Bonds between atoms with significantly different electronegativities.

6.3 Intermolecular Forces

• London dispersion forces: temporary, weak attractions between molecules due to constantly changing electron distributions.
• Permanent dipole–dipole forces: attractions between polar molecules where the positive end of one molecule attracts the negative end of another.
• Hydrogen bonding: a special type of dipole–dipole force that occurs when hydrogen is bonded to a highly electronegative atom (O, N, or F). Strong attractive forces amongst molecules.

6.4 Hydrogen Bonding

• Hydrogen bonding: a special dipole-dipole force that occurs when hydrogen is covalently bonded to a highly electronegative atom.
• Its influence on properties of substances.

This quiz covers essential concepts in Chemistry A-Level, focusing on atomic structure, isotopes, and relative masses. Understand how isotopes differ and learn about mass defects in nuclear chemistry with engaging questions designed for learners at this level.