Chemistry Module 3: Reactive Chemistry
Module 3 Notes: Reactive chemistry (Parts 1–3)
The module provides a comprehensive exploration of fundamental chemical principles, focusing on the classification and prediction of chemical reactions, the factors governing reaction rates, and the core concepts of redox chemistry. It integrates theoretical understanding with practical applications, including the principles of galvanic cells and real-world examples such as metal reactivity and traditional detoxification methods.
Part 1: Types of Reactions and Signs of Chemical Reactions
Synthesis reaction
Definition: Two or more simpler substances (elements or compounds) combine chemically to form a single, more complex product.
General form: A + B -> AB
Example: The formation of water from hydrogen and oxygen gases.
2H2(g) + O2(g) -> 2H2O(l)
Combustion reaction
Definition: A rapid reaction between an element or compound (often a hydrocarbon) and an oxidant, usually oxygen, to produce oxides and release energy vigorously, typically as heat (exothermic) and sometimes light.
Complete combustion with plentiful O2: Occurs when there is an adequate supply of oxygen, resulting in the formation of carbon dioxide and water.
General form for hydrocarbons: Hydrocarbon + O2 -> CO2 + H2O
Example: Methane combustion
CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l)
Example: Octane combustion (gasoline, often balanced with fractional coefficients for simplicity, which are then cleared)
C8H18(g) + 12.5O2(g) -> 8CO2(g) + 9H2O(l)
Incomplete combustion (limited O2): Occurs when oxygen supply is insufficient, leading to the formation of carbon monoxide (CO) and/or solid carbon (soot, C) along with carbon dioxide and water. This is less efficient and produces toxic byproducts.
Decomposition reaction
Definition: A single, complex substance breaks down into two or more simpler substances.
General form: AB -> A + B
Examples:
Decomposition of hydrogen peroxide:
2H2O2(l) -> 2H2O(l) + O2(g)
Thermal decomposition of copper(II) carbonate:
CuCO3(s) -> CuO(s) + CO2(g)
Photodecomposition of silver chloride (e.g., in photographic film):
2AgCl(s) light-> 2Ag(s) + Cl2(g)
Electrolysis of water (electrolytic decomposition):
2H2O(l) electricity-> 2H2(g) + O2(g)
Precipitation reaction
Definition: A reaction in which two soluble ionic compounds in aqueous solution combine to form an insoluble solid product called a precipitate. This solid separates from the solution.
Example: The formation of lead iodide precipitate from potassium iodide and lead(II) nitrate solutions.
2KI(aq) + Pb(NO3)2(aq) -> PbI2(s) + 2KNO3(aq)
Micro vs. macro changes in reactions
Micro level: Involves the breaking and forming of chemical bonds, leading to the rearrangement of atoms or ions. For example, in a double displacement reaction: AB + CD -> AC + BD. Atoms are conserved but connected differently.
Macro level indicators: Observable signs that a chemical reaction has occurred within a system.
Common indicators: Colour change (e.g., phenolphthaleinredpink), temperature change (exothermic/endothermic), gas production (bubbles), formation of a solid precipitate (turbidity), light emission, or odour change.
Chemical vs. physical change
Chemical change: Results in the formation of one or more new substances with different chemical properties from the original substances. These changes are often difficult to reverse.
Signs often include: Colour change, temperature change (heating or cooling), gas evolution, light emission, and precipitate formation.
Physical change: Alters the form or appearance of a substance but does not create a new substance. The chemical composition remains the same, and these changes are often reversible.
Examples: Phase changes (melting, boiling, freezing, condensation), dissolving a solid in a liquid, cutting, grinding.
Part 1: Solubility, Precipitation, and Salt Formation
Precipitation and solubility rules
Concept: Solubility rules are a set of general statements used to predict whether an ionic compound will be soluble (dissolve) or insoluble (precipitate) in water. These rules are crucial for predicting the products of precipitation reactions.
General Guidelines (simplified):
Soluble: Most nitrates (NO3-), acetates (CH3COO-), group 1 metal ions (Li+, Na+, K+, etc.), ammonium (NH4+), most chlorides (Cl-), bromides (Br-), and iodides (I-) (except with Ag+, Pb2+, Hg22+), most sulfates (SO42-) (except with Ba2+, Pb2+, Ca2+, Sr2+).
Insoluble: Most carbonates (CO32-), phosphates (PO43-), sulfides (S2-), and hydroxides (OH-) (except with Group 1 metals and NH4+, and Ca2+, Sr2+, Ba2+ hydroxides are slightly soluble).
Example precipitate: Lead iodide (PbI2) is an insoluble yellow solid that forms when potassium iodide and lead(II) nitrate solutions are mixed.
Net ionic equations (concept)
Definition: A chemical equation that shows only the species (ions, atoms, molecules) that participate directly in a chemical reaction. Spectator ions, which remain unchanged on both sides of the equation, are omitted.
Importance: They provide a clearer representation of the actual chemical change occurring.
Steps:
Write the balanced molecular equation.
Write the complete ionic equation by separating all soluble ionic compounds into their constituent ions.
Cancel out spectator ions (ions appearing identically on both sides of the equation).
Write the remaining species as the net ionic equation.
Example (for PbI2 precipitation):
Molecular: 2KI(aq) + Pb(NO3)2(aq) -> PbI2(s) + 2KNO3(aq)
Complete ionic: 2K+(aq) + 2I-(aq) + Pb2+(aq) + 2NO3-(aq) -> PbI2(s) + 2K+(aq) + 2NO3-(aq)
Net ionic: Pb2+(aq) + 2I-(aq) -> PbI2(s)
Part 1: Neutralisation and Salt Naming
Acid–base reactions (neutralisation)
Definition: A specific type of acid–base reaction where an acid reacts with a base to form a salt and water. It neutralizes the acidic and basic properties.
Arrhenius definition: Acids produce H+ ions in solution; bases produce OH- ions in solution.
General equation:
Acid + Base -> Salt + H2O
Example: Hydrochloric acid reacting with sodium hydroxide:
HCl(aq) + NaOH(aq) -> NaCl(aq) + H2O(l)
Salt naming
Definition: A salt is an ionic compound formed from the cation of a base and the anion of an acid as a product of a neutralisation reaction.
Naming convention: Combine the name of the metal/polyatomic cation from the base with the name of the anion from the acid.
Example from HCl + NaOH: The cation is Na+ (from NaOH) and the anion is Cl- (from HCl), forming sodium chloride (NaCl).
Other examples: Sulphuric acid (H2SO4) + Potassium hydroxide (KOH) yields potassium sulfate (K2SO4).
Nitric acid (HNO3) + Ammonia (NH3) yields ammonium nitrate (NH4NO3).
Part 1: Acid–Base Reactions and Test Concepts
Acid is a substance with pH < 7; base has pH > 7. A neutral solution has pH = 7. The pH scale is a logarithmic scale used to specify the acidity or basicity of an aqueous solution.
Indicators of acid–base reactions: Besides salt and water formation, in solution, these reactions involve the combination of H+ (or H3O+) ions from the acid and OH- ions from the base to form H2O. pH indicators (e.g., litmus, phenolphthalein) show distinct color changes over specific pH ranges.
Part 1: Example Reactions and Tests
Tests and indicators commonly used in chemistry to identify products or types of reactions:
Solubility tests: Used to determine if a product (e.g., hydroxides or carbonates) is soluble or precipitates out, often by observation of turbidity or solid formation.
Conductivity tests: To check for the presence of ions in solution (e.g., MgO or Ca(OH)2 solutions can conduct electricity).
Gas tests:
Glowing splint test: A glowing splint relights in the presence of oxygen (O2).
Lit splint test: A lit splint produces a 'pop' sound in the presence of hydrogen gas (H2).
Limewater test: Carbon dioxide (CO2) turns limewater (calcium hydroxide solution) cloudy.
Colour changes: Often directly observed in decomposition or oxidation reactions (e.g., AgCl turning dark upon exposure to light, indicating Ag formation).
Universal indicator/pH paper: For estimating the pH of a solution.
Part 2: Predicting Reactions of Metals
This section extends the concepts of synthesis and decomposition to reactions involving metals with common substances like water, acids, and oxygen, based on their inherent reactivity.
Metals with water
Mechanism: Metals can react with water to form metal hydroxides (or oxides, depending on conditions) and hydrogen gas.
Activity of metals: The reactivity of a metal dictates how readily and vigorously it reacts with cold water, hot water, or steam. More reactive metals react more energetically.
With cold water: Highly reactive metals (e.g., Group 1 alkali metals like K, Na; some Group 2 alkaline earth metals like Ca) react vigorously to produce metal hydroxide and hydrogen gas.
Example: 2Na(s) + 2H2O(l) -> 2NaOH(aq) + H2(g)
With hot water: Moderately reactive metals (e.g., Mg) may react with hot water, but not cold, forming hydroxide and hydrogen.
Example: Mg(s) + 2H2O(l) -> Mg(OH)2(s) + H2(g)
With steam (red heat): Less reactive metals (e.g., Fe, Zn) that do not react with cold or hot water may react only with steam at high temperatures, forming metal oxide and hydrogen gas.
Example: 3Fe(s) + 4H2O(g) -> Fe3O4(s) + 4H2(g)
Indicators of reactivity: Observation of gas bubbles, heat release, or disappearing metal piece.
Examples: Potassium (K) and Sodium (Na) react extremely vigorously with cold water; Magnesium (Mg) reacts slowly with cold water but faster with steam; Copper (Cu), Silver (Ag), and Gold (Au) show no reaction with water under typical conditions.
General pattern (conceptual):
Metal(s) + H2O(l) -> M(OH)x(aq/s) + H2(g) (for reactive metals with cold/hot water)
Metal(s) + H2O(g) -> MOx(s) + H2(g) (for less reactive metals with steam)
Metals with acids (dilute acid)
Mechanism: Reactive metals react with dilute acids to produce hydrogen gas and a metal salt. This is essentially a single displacement reaction where the metal displaces hydrogen from the acid.
Observations during experiments: The rate and extent of reaction vary widely.
Rapid fizzing: Highly reactive metals (e.g., Na, K, Ca, Mg) react vigorously with cold dilute acids.
Slow fizzing: Moderately reactive metals (e.g., Zn, Fe) react more slowly, possibly requiring heating.
No reaction: Less reactive metals (e.g., Cu, Ag, Au) do not react with dilute acids to produce hydrogen gas, as they are below hydrogen in the activity series.
Typical trend: More reactive metals (those above hydrogen in the activity series) react readily with acids. Reactivity increases with temperature.
Important notes:
Passivation: Some metals (e.g., lead, aluminium) may form an insoluble protective coating (e.g., PbCl2 or Al2O3) on their surface when reacting with certain acids, which prevents further reaction.
Activity Series: The observed reactivities of metals with acids provide a practical basis for constructing a metal activity series, ranking metals by their tendency to lose electrons (get oxidized).
Metals with oxygen
Mechanism: Metals react with oxygen (oxidation) to form metal oxides. The speed and conditions required for this reaction depend on the metal's reactivity.
Reported trends: Some metals oxidize very rapidly (e.g., Group 1 metals tarnish quickly in air), while others react very slowly or only at high temperatures.
Examples of observed oxide products (with colour hints):
Magnesium (Mg): Reacts readily to form white magnesium oxide (MgO).
Aluminium (Al): Forms a thin, protective layer of white aluminium oxide (Al2O3) that prevents further corrosion.
Zinc (Zn): Forms zinc oxide (ZnO), which is yellow when hot and white when cool.
Iron (Fe): Rusts (oxidizes slowly in the presence of oxygen and water) to form reddish-brown iron(III) oxide (Fe2O3) or other iron oxides like FeO.
Tin (Sn): Can form tin(IV) oxide (SnO2), which is a white powder, though some conditions may yield black oxides.
Lead (Pb): Can form lead(II) oxide (PbO), typically a yellow or reddish-brown solid, or lead(IV) oxide (PbO2), a dark brown solid, depending on conditions.
Copper (Cu): Oxidizes very slowly, forming a black copper(II) oxide (CuO) layer or a green patina over long periods.
Summary: Transportable ideas
Metal activity series: Metals can be experimentally ranked according to their chemical reactivity, forming an activity series. This series generally places the most reactive metals (e.g., K, Na, Ca) at the top and the least reactive metals (e.g., Ag, Au, Pt) at the bottom.
Basis of the series: The activity series is derived from observed reactions with water, acids, and oxygen, as well as displacement reactions where a more reactive metal displaces a less reactive metal from its compound, and thermal decomposition studies of metal compounds.
Predictive power: This series allows for the prediction of whether a metal will react with specific reagents and the relative vigour of such reactions.
Part 3: Redox Chemistry and Oxidation–Reduction (Oxidation States, Electron Transfer)
Redox reactions: A class of chemical reactions that involve the transfer of electrons between two species.
Oxidation and reduction definitions
Oxidation: The process of losing electrons, resulting in an increase in oxidation number.
Example (half-reaction): In the context of a zinc-copper cell, zinc metal undergoes oxidation.
Zn(s) -> Zn2+(aq) + 2e- (Zinc loses two electrons)
Reduction: The process of gaining electrons, resulting in a decrease in oxidation number.
Example (half-reaction): In the context of a zinc-copper cell, copper(II) ions undergo reduction.
Cu2+(aq) + 2e- -> Cu(s) (Copper(II) ions gain two electrons)
Mnemonic: OIL RIG (Oxidation Is Losing, Reduction Is Gaining) or LEO the lion says GER (Loss of Electrons is Oxidation, Gain of Electrons is Reduction).
Oxidizing agent (oxidant): The species that causes oxidation by accepting electrons (itself getting reduced).
Reducing agent (reductant): The species that causes reduction by donating electrons (itself getting oxidized).
Oxidation number (or oxidation state): A hypothetical charge assigned to an atom in a molecule or ion based on a set of rules. It is a bookkeeping tool used to track electron transfer and identify which species are oxidized and reduced during a reaction.
Important note: Oxidation numbers are a useful concept for identifying redox couples and balancing complex redox equations, even though they may not represent actual charges in covalent compounds.
Spontaneity of redox reactions and galvanic cells
Spontaneous redox reaction: A redox reaction that occurs naturally without any external energy input. These reactions release free energy.
Example: When a zinc metal strip is submerged in a copper(II) sulfate solution, a spontaneous redox reaction occurs.
Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s)
In this reaction, zinc metal (Zn) is oxidized to Zn2+ ions, while copper(II) ions (Cu2+) are reduced to copper metal (Cu).
Electrons flow directly from Zn atoms to Cu2+ ions, releasing energy as heat.
Galvanic (voltaic) cell: An electrochemical cell designed to harness the energy released from a spontaneous redox reaction to produce electrical current. It separates the oxidation and reduction half-reactions, forcing electrons to flow through an external circuit.
Electrode: A conductor (typically metallic) through which electrons enter or leave a non-metallic medium (electrolyte).
Anode: The electrode where oxidation occurs. In a galvanic cell, the anode is the negative terminal because it is the source of electrons.
Cathode: The electrode where reduction occurs. In a galvanic cell, the cathode is the positive terminal because electrons flow towards it.
Electrolyte: An ionic solution that allows for the movement of ions to maintain charge neutrality in the half-cells and complete the circuit.
Salt bridge: Connects the two half-cells, allowing the migration of ions to balance the charges that build up in the compartments, completing the circuit.
Mnemonic to recall roles: An Oil Rig Cat (Anode: Oxidation; Oil: Oxidation Is Lost; Rig: Reduction Is Gain; Cat: Cathode: Reduction).
Standard electrode potential (E°):
Definition: The potential difference (voltage) of an electrode in its standard state (1 M concentration for solutions, 1 atm pressure for gases, 25°C) relative to the Standard Hydrogen Electrode (SHE), which is arbitrarily assigned an E° of 0.00 V.
Electromotive force (emf) or cell potential (E°cell)**: The potential difference between the cathode and the anode in a galvanic cell, representing the maximum electrical energy that can be obtained from the reaction under standard conditions. A positive E°cell indicates a spontaneous reaction. For a spontaneous reaction, electrons flow from the anode (lower E°) to the cathode (higher E°).
Examples and half-reactions
In the Zn–Cu galvanic cell (Daniell cell) example:
Overall reaction: Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s)
Oxidation half-reaction at the anode (negative terminal):
Zn(s) -> Zn2+(aq) + 2e- (E°Zn2+/Zn = -0.76 V)
Reduction half-reaction at the cathode (positive terminal):
Cu2+(aq) + 2e- -> Cu(s) (E°Cu2+/Cu = +0.34 V)
Cell potential calculation: E°cell = E°cathode - E°anode = (+0.34 V) - (-0.76 V) = +1.10 V
Part 3: Rates of Reactions
Rate of reaction: The speed at which a chemical reaction proceeds. It quantifies how quickly reactants are consumed or products are formed.
Definition: The change in concentration of a reactant or product per unit time.
Units: Typically expressed in units of mol * L-1 * s-1 or M * s-1.
Measurement: Can be determined by monitoring changes in concentration, pressure, volume, pH, or conductivity over time.
Collision theory and activation energy
Collision theory: States that for a chemical reaction to occur, reactant particles must collide with each other with:
Correct orientation: The particles must be aligned in a specific way for the reacting parts of the molecules to come into contact and allow bonds to break and form.
Sufficient kinetic energy: The collision must have kinetic energy equal to or greater than the activation energy (Ea) to overcome the energy barrier and break existing bonds.
Activation energy (Ea): The minimum amount of energy that colliding reactant particles must possess for a reaction to occur. It represents an energy barrier that must be surmounted.
Effective collisions: Only collisions that meet both criteria (correct orientation and sufficient energy) lead to a chemical reaction.
Rate dependence: The overall rate of reaction is directly proportional to the number of effective collisions per unit time.
Maxwell–Boltzmann distribution: A statistical distribution that describes the range of kinetic energies of particles at a given temperature. Only a fraction of particles, those with kinetic energy greater than or equal to Ea, are capable of reacting upon collision.
Important implication: Increasing temperature increases the fraction of particles that possess energy >= Ea (by shifting the Maxwell–Boltzmann distribution), thus leading to more successful collisions and a faster reaction rate. However, temperature **does not alter the Ea of a reaction.
Factors influencing rate of reaction
Temperature:
Effect: Higher temperature generally leads to a faster reaction rate.
Explanation: Increased temperature increases the average kinetic energy of reactant particles, causing them to move faster. This results in:
More frequent collisions.
A significantly larger proportion of collisions having energy equal to or greater than the activation energy (Ea).
Surface area (for solid reactants):
Effect: Increasing the surface area of a solid reactant increases the reaction rate.
Explanation: When a solid reactant is ground into smaller pieces (increasing total surface area), more reactant particles are exposed and available for collisions with other reactants. This increases the frequency of collisions.
Note: The energy per individual collision remains the same; the rate increases due to a higher number of potential collision sites.
Concentration of reactants (for solutions or gases):
Effect: Higher concentration of reactants leads to a faster reaction rate.
Explanation: Increasing the concentration of reactants means there are more reactant particles per unit volume. This increases the likelihood and frequency of collisions between reacting particles, leading to more successful collisions.
Catalysts:
Effect: A catalyst significantly increases the rate of a chemical reaction without being consumed in the process.
Explanation: A catalyst provides an alternative reaction pathway or mechanism that has a lower activation energy (Ea) than the uncatalyzed reaction. This allows a greater fraction of reactant particles to possess sufficient energy for effective collisions, thereby speeding up the reaction.
Important: Catalysts participate in the reaction but are regenerated unchanged at the end. They do not change the equilibrium position of a reversible reaction, only how quickly equilibrium is reached.
Pressure (for gaseous reactants):
Effect: Higher pressure for gaseous reactants increases the reaction rate.
Explanation: Increasing the pressure of a gaseous system effectively reduces the volume the gas particles occupy. This leads to a higher concentration of gas particles per unit volume, which in turn increases the frequency of collisions and thus the reaction rate.
Catalysts in detail
Definition: A substance that increases the rate of a specific chemical reaction without being used up in the reaction. It can be recovered chemically unchanged at the end of the reaction.
Mechanism: Catalysts work by providing an alternative reaction pathway that bypasses the original, higher-energy transition state, effectively lowering the activation energy (Ea) required for the reaction to occur. They do not alter the overall enthalpy change (ΔH) or Gibbs free energy change (ΔG) of the reaction.
Behavioral note: Catalysts often function by binding to reactant molecules (forming intermediate species like an enzyme-substrate complex), facilitating bond breaking or forming, and then releasing the products. This temporary interaction allows the reaction to proceed more rapidly and efficiently than it would without the catalyst.
Practical Contexts and Applications
Detoxification methods for natural products (ATSI context): Indigenous Australians developed sophisticated methods to process toxic native plants, such as the cycad fruit, for safe consumption.
Cycad fruit: Highly toxic and carcinogenic due to cycasin and macrozamin content.
Traditional detoxification methods:
Leaching:
Process: Cycad kernels are cut into smaller pieces and often ground to significantly increase their surface area. These prepared kernels are then repeatedly soaked in flowing water for extended periods (days to weeks).
Mechanism: The soluble toxins are gradually leached out (dissolved and washed away) by the water.
Outcome: After leaching, the detoxified kernels are further ground into flour for making bread or cakes.
Fermentation:
Process: Cycad kernels are stored in warm, moist environments, often buried or placed in specific containers, for several months. During this time, naturally occurring microorganisms (bacteria, fungi) ferment the kernels.
Mechanism: Fermentation processes break down or convert the toxic compounds into less harmful or non-toxic substances.
Outcome: The fermented kernels become safe for consumption, representing a clever use of biochemical processes for survival.
Real-world relevance: Understanding reaction types, redox behavior, and reaction rate controls is fundamental in various fields, including:
Predicting product formation: Essential in chemical synthesis and industrial manufacturing.
Optimizing industrial processes: Adjusting temperature, pressure, concentration, and using catalysts to maximize product yield and efficiency (e.g., Haber-Bosch process, catalytic converters).
Assessing environmental and health implications: Understanding corrosion, pollution degradation, metabolic pathways in living organisms, and drug action.
Quick Reference: Key Equations (Cheat Sheet)
Synthesis: 2H2(g) + O2(g) -> 2H2O(l)
Combustion (complete methane): CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l)
Combustion (octane example): C8H18(g) + 12.5O2(g) -> 8CO2(g) + 9H2O(l)
Decomposition (hydrogen peroxide): 2H2O2(l) -> 2H2O(l) + O2(g)
Decomposition (water electrolysis): 2H2O(l) -> 2H2(g) + O2(g)
Decomposition (copper(II) carbonate): CuCO3(s) -> CuO(s) + CO2(g)
Precipitation (lead iodide): 2KI(aq) + Pb(NO3)2(aq) -> PbI2(s) + 2KNO3(aq)
Neutralisation (acid + base): Acid + Base -> Salt + H2O
Acid + carbonate: Acid + Carbonate -> Salt + CO2(g) + H2O(l)
Redox (example half-reactions):
Oxidation: Zn(s) -> Zn2+(aq) + 2e-
Reduction: Cu2+(aq) + 2e- -> Cu(s)
Galvanic cell emf: E°cell = E°cathode - E°anode
Activation energy concept (Ea)