IGCSE Double Award Chemistry Study Guide (Edexcel)
States of Matter
Changes of State:
Melting:
Process: The transition of a substance from a solid state to a liquid state upon heating.
Mechanism: Particles within the solid gain thermal energy and consequently vibrate more vigorously. This increase in kinetic energy allows them to overcome the intermolecular forces that hold them in a rigid structure, resulting in a more disordered liquid phase.
Boiling:
Process: The phase change of a substance from liquid to gas that occurs at its boiling point.
Mechanism: In this process, the particles in the liquid gain sufficient energy to break free from the liquid’s surface, forming vapor bubbles that escape into the air, causing the liquid to transition into a gaseous state. The boiling point is the temperature at which the vapor pressure of the liquid equals the atmospheric pressure.
Sublimation:
Process: The direct transition of a substance from a solid to a gas without passing through the liquid phase.
Examples: Common substances undergoing sublimation include iodine crystals and dry ice (solid carbon dioxide), which can transition to gas at room temperature or under low-pressure conditions.
Diffusion:
In gases: The process by which gas particles move from areas of high concentration to areas of low concentration, facilitated by their relatively high kinetic energy which allows for rapid movement. The rate of diffusion in gases is influenced by temperature, as increased thermal energy enhances particle movement.
In liquids: Although diffusion occurs in liquids as well, it happens at a slower rate due to the closer proximity of particles, which restricts their movement. Factors such as temperature and agitation can influence the rate of diffusion.
Particle Theory:
Solids: Characterized by tightly packed particles that are fixed in a definite arrangement, allowing only vibrational movements. This results in a defined shape and volume as the particles maintain their positions due to strong intermolecular forces.
Liquids: Consist of particles that are closely packed but not in fixed positions, allowing them to slide past one another; this gives liquids a defined volume but not a fixed shape, adapting to the contours of their container.
Gases: Composed of widely spaced particles that move freely and rapidly; gases possess neither a fixed shape nor a defined volume, expanding to fill their container completely.
Atomic Structure and the Periodic Table
Subatomic Particles:
Proton: Positively charged particle found in the nucleus, with a mass of approximately 1 atomic mass unit (amu).
Neutron: Neutral particle also located in the nucleus, with a mass similar to that of a proton, contributes to the atomic mass.
Electron: Negatively charged particle with a negligible mass compared to protons and neutrons, orbits the nucleus in defined energy levels.
Isotopes:
Definition: Atoms of the same element with identical numbers of protons but differing numbers of neutrons, leading to variations in atomic mass.
Relative Atomic Mass (Ar):
Formula: Ar = \frac{\text{Sum of isotope abundance} \times \text{mass number}}{100}, this calculation takes into account the natural isotopic abundances of an element in a sample.
Electronic Configuration:
Electrons fill energy shells in the order of 2, 8, 8, etc., following Hund's rule and the Pauli exclusion principle. Each shell can hold a maximum number of electrons determined by the formula 2n² where n is the shell level.
Example: Oxygen (atomic number 8) has the configuration 2,6, indicating it has two electrons in the first shell and six in the second.
Periodic Trends:
Groups: Vertical columns in the periodic table that contain elements with the same number of outer electrons, which confer similar chemical properties.
Periods: Horizontal rows indicating the number of energy shells occupied by electrons.
Reactivity Trends:
Group 1 (Alkali Metals): Reactivity increases down the group due to the increase in atomic size and decreasing ionization energy, making it easier for these metals to lose their outermost electron.
Group 7 (Halogens): Reactivity decreases down the group, largely due to increasing atomic radius and decreasing electronegativity, which makes it harder for these elements to gain an electron.
Chemical Bonding
Ionic Bonding:
Involves a transfer of electrons from a metal to a non-metal, resulting in the formation of charged ions (cations and anions).
Mechanism: The loss of one or more electrons by the metal atom creates a positively charged ion, while the gain of these electrons by the non-metal creates a negatively charged ion; these oppositely charged ions attract each other to form an ionic compound.
Covalent Bonding:
Involves the sharing of electron pairs between two non-metal atoms.
Mechanism: Atoms share valence electrons in order to achieve full outer electron shells, leading to the formation of molecular compounds.
Properties:
Ionic Compounds: Typically characterized by high melting points, electrical conductivity in molten or dissolved state, and solid crystalline structure.
Covalent Compounds: Generally possess lower melting points, are non-conductive in solid and liquid states, and may exist as gases, liquids, or low-melting solids.
Metallic Bonding:
Description: Consists of positively charged metal ions immersed in a 'sea' of delocalized electrons that move freely throughout the structure.
Properties: This structure enables metals to conduct electricity, be malleable (able to be hammered into shapes), and ductile (able to be drawn into wires).
Stoichiometry and Chemical Calculations
Balancing Equations:
Aim: To ensure the chemical equation reflects the law of conservation of mass, achieving an equal number of atoms for each element on both sides of the equation.
Mole Calculations:
Formula: \text{Moles} = \frac{\text{Mass}}{\text{Molar Mass}}, allowing for the conversion between mass and the quantity of substance.
For gases: \text{Volume} = \text{Moles} \times 24 \text{ dm}^3, applicable at standard temperature and pressure (STP).
Concentration: \text{Concentration} = \frac{\text{Moles}}{\text{Volume (dm}^3)}, indicating the amount of solute in a given volume of solution.
Empirical/Molecular Formula:
Empirical: Represents the simplest whole-number ratio of elements in a compound.
Molecular: Indicates the actual number of atoms of each element in the molecule, which may be a whole-number multiple of the empirical formula.
Chemical Reactions
Types of Reactions:
Combustion: Reaction of a fuel with oxygen to produce carbon dioxide and water, often releases energy in the form of heat and light.
Neutralisation: The reaction between an acid and a base to produce a salt and water, resulting in a pH balance.
Displacement: Involves a more reactive element replacing a less reactive one in a compound, often observed in halogen reactions.
Energy Changes:
Exothermic: A type of reaction that releases energy, leading to an increase in temperature of the surroundings (e.g., combustion).
Endothermic: A reaction that absorbs energy from surroundings, resulting in a decrease in temperature (e.g., photosynthesis).
Rates of Reaction:
Factors: Various factors including temperature, concentration of reactants, surface area of solid reactants, and the presence of catalysts influence the speed of a reaction.
Measurement Methods: Rates of reaction can be measured by observing changes in gas volume, mass change, and observable color changes.
Acids, Bases, and Salts
Properties:
Acids: Substances with a pH less than 7, possessing a sour taste, turn litmus paper red, and react with metals to produce hydrogen gas.
Bases: Substances with a pH greater than 7, characterized by a bitter taste, feel slippery, and turn litmus paper blue.
pH and Indicators:
Universal indicator:
A scale ranging from red (strong acid) through green (neutral) to purple (strong base) indicating the pH level of a solution.
Salt Preparation:
Methods:
For soluble salts: Crystallization is typically used to purify and isolate the salt from solutions.
For insoluble salts: Precipitation reactions are employed where solutions containing ions react to form an insoluble compound.
The Periodic Table
Group 1: Alkali Metals:
Highly reactive metals that react vigorously with water to produce hydroxide and hydrogen gas, known for their reactivity increasing down the group as atomic size increases.
Group 7: Halogens:
Diatomic molecules (e.g., Cl₂, F₂) that exhibit reactivity decreasing down the group; they undergo displacement reactions with halide salts, forming different halogen compounds.
Group 0: Noble Gases:
Inert gases that have a full outer shell of electrons, making them largely unreactive; these gases are utilized in applications such as lighting and welding due to their stability.
Transition Metals:
Characterized by high melting points, variable oxidation states, and the ability to form colorful compounds; they often serve as catalysts in various chemical reactions.
Metals and Reactivity Series
Extraction of Metals:
Based on reactivity:
Metals above carbon in the reactivity series require extraction through electrolysis as they cannot be reduced by carbon.
Metals below carbon can be extracted by reduction with carbon, as carbon can displace them from their ores.
Reactivity Series:
Order: Potassium > Sodium > Calcium > Magnesium > Aluminum > Carbon > Zinc > Iron > Lead > Copper > Silver > Gold, indicating the relative tendency of metals to react and displace each other.
Corrosion:
Rusting: A common type of corrosion occurring with iron, represented by the reaction of iron with oxygen and water to form iron oxide (rust).
Prevention methods: Include galvanising (coating with zinc), painting to create a protective layer, or using sacrificial protection methods that involve more reactive metals protecting the iron surface.
Organic Chemistry
Alkanes:
Definition: Saturated hydrocarbons characterized by single C-C bonds, leading to relatively stable structures.
General formula: CnH_{2n+2}, where n is the number of carbon atoms.
Alkenes:
Definition: Unsaturated hydrocarbons distinguished by at least one C=C double bond, resulting in more reactive compounds.
General formula: CnH_{2n}, allowing for the possibility of addition reactions with other compounds.
Crude Oil:
Fractional Distillation: A separation technique that divides crude oil into various hydrocarbons based on their boiling points, producing fractions such as petrol, diesel, and bitumen for different uses.
Polymers:
Formed from the process of addition polymerisation of alkenes, creating long-chain molecules; while useful in many applications, they are often non-biodegradable and contribute to environmental pollution.
Environmental Chemistry
Greenhouse Gases:
Examples: Carbon dioxide (CO₂), methane (CH₄), and water vapor (H₂O). These gases trap heat in the atmosphere, playing a significant role in contributing to global warming and climate change.
Air Pollutants:
Examples: Carbon monoxide (CO) from incomplete combustion, sulfur dioxide (SO₂), and nitrogen oxides (NOₓ). Their presence in the atmosphere can lead to adverse effects such as acid rain and respiratory problems.
Sustainable Development:
Strategies: Include practices such as recycling materials, utilizing renewable resources, and minimizing carbon footprints to promote environmentally sound development that meets current needs without compromising future generations' ability to meet theirs.