🧪AP Chemistry Unit 4 – Chemical Reactions

Key Concepts

• Chemical reactions involve the breaking and forming of chemical bonds between atoms or molecules
• Reactants are the starting materials that undergo change in a chemical reaction, while products are the substances formed as a result
• Conservation of mass states that the total mass of the reactants must equal the total mass of the products in a chemical reaction
  • Atoms are neither created nor destroyed during a chemical reaction, they are simply rearranged
• Stoichiometry is the quantitative study of the amounts of reactants and products involved in a chemical reaction
• Activation energy is the minimum energy required for a chemical reaction to occur and is often provided by heat, light, or a catalyst
• Catalysts are substances that speed up chemical reactions without being consumed in the process by lowering the activation energy
• Reaction rates describe how quickly a chemical reaction proceeds and can be influenced by factors such as temperature, concentration, and surface area

Types of Chemical Reactions

• Synthesis reactions involve the combination of two or more reactants to form a single product ($A + B \rightarrow AB$)
• Decomposition reactions involve the breakdown of a single reactant into two or more products ($AB \rightarrow A + B$)
• Single displacement reactions occur when one element replaces another in a compound ($A + BC \rightarrow AC + B$)
• Double displacement reactions involve the exchange of ions between two compounds ($AB + CD \rightarrow AD + CB$)
• Combustion reactions are rapid reactions that produce heat and light, often involving oxygen as a reactant ($CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$)
• Acid-base reactions involve the transfer of protons ($H^+$) from an acid to a base, forming a salt and water ($HCl + NaOH \rightarrow NaCl + H_2O$)
• Redox reactions involve the transfer of electrons between species, resulting in changes in oxidation states ($2Mg + O_2 \rightarrow 2MgO$)

Balancing Chemical Equations

• Chemical equations represent the reactants, products, and their stoichiometric coefficients in a reaction
• Balanced equations have equal numbers of each type of atom on both sides of the arrow, following the law of conservation of mass
• Coefficients are used to balance equations and indicate the relative amounts of reactants and products
  • Coefficients cannot be fractions and must be the lowest whole number ratio possible
• Subscripts in chemical formulas represent the number of atoms of each element in a molecule and should not be changed when balancing equations
• Balancing equations is done by adjusting coefficients, not subscripts, to ensure equal numbers of each atom type on both sides
• Steps for balancing equations:
  1. Write the unbalanced equation
  2. Count the number of each atom type on both sides
  3. Balance the equation by adjusting coefficients, starting with the most complex molecule
  4. Recount atoms and make final adjustments to ensure the equation is balanced

Reaction Rates and Kinetics

• Reaction rate is the speed at which a chemical reaction proceeds, typically measured in concentration change over time ($\text{rate} = \frac{\Delta[\text{product}]}{\Delta t}$)
• Factors affecting reaction rates include temperature, concentration, pressure (for gases), surface area, and the presence of catalysts
  • Increasing temperature, concentration, pressure, or surface area generally increases reaction rates
• Collision theory states that reactions occur when reactant particles collide with sufficient energy (activation energy) and proper orientation
• The rate law expresses the relationship between the reaction rate and the concentrations of reactants, often in the form: $\text{rate} = k[A]^m[B]^n$
  • $k$ is the rate constant, $[A]$ and $[B]$ are reactant concentrations, and $m$ and $n$ are the reaction orders
• Reaction order is the power to which the concentration of a reactant is raised in the rate law and can be determined experimentally
• Zero-order reactions have rates independent of reactant concentrations, while first-order and second-order reactions have rates proportional to the first and second powers of reactant concentrations, respectively

Chemical Equilibrium

• Chemical equilibrium is a dynamic state in which the forward and reverse reaction rates are equal, resulting in no net change in reactant or product concentrations
• The equilibrium constant ($K$) is the ratio of the product of the product concentrations to the product of the reactant concentrations, each raised to their stoichiometric coefficients
  • For the general reaction $aA + bB \rightleftharpoons cC + dD$, the equilibrium constant is $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$
• Le Chatelier's principle states that when a system at equilibrium is disturbed, it will shift in the direction that counteracts the disturbance to re-establish equilibrium
  • Changing concentration, pressure, volume, or temperature can shift the equilibrium position
• Equilibrium shifts can be predicted using the reaction quotient ($Q$) and comparing it to the equilibrium constant ($K$)
  • If $Q < K$, the reaction will shift towards the products; if $Q > K$, it will shift towards the reactants
• Catalysts do not affect the equilibrium constant or the equilibrium concentrations of reactants and products, but they do speed up the rate at which equilibrium is reached

Thermodynamics in Reactions

• Thermodynamics is the study of heat and its relationship to chemical reactions and other forms of work
• Enthalpy ($H$) is a measure of the total heat content of a system, and the change in enthalpy ($\Delta H$) during a reaction indicates whether heat is absorbed (endothermic, $\Delta H > 0$) or released (exothermic, $\Delta H < 0$)
• Entropy ($S$) is a measure of the disorder or randomness of a system, and the change in entropy ($\Delta S$) during a reaction indicates whether the system becomes more ordered ($\Delta S < 0$) or more disordered ($\Delta S > 0$)
• Gibbs free energy ($G$) is a measure of the usable energy in a system, and the change in Gibbs free energy ($\Delta G$) determines the spontaneity of a reaction
  • If $\Delta G < 0$, the reaction is spontaneous; if $\Delta G > 0$, the reaction is non-spontaneous; if $\Delta G = 0$, the system is at equilibrium
• The relationship between $\Delta G$, $\Delta H$, and $\Delta S$ is given by the equation: $\Delta G = \Delta H - T\Delta S$, where $T$ is the absolute temperature in Kelvin
• Hess's law states that the overall enthalpy change of a reaction is the sum of the enthalpy changes of the individual steps, regardless of the pathway taken
• Calorimetry is the measurement of heat transfer during chemical reactions or physical processes, often using a calorimeter to determine the enthalpy change

Practical Applications

• Chemical reactions are the basis for many important processes in industry, medicine, and everyday life
• Combustion reactions are used to generate energy in power plants, vehicles, and heating systems
  • Fossil fuels (coal, oil, natural gas) undergo combustion to release heat, which is then converted to electrical or mechanical energy
• Synthesis reactions are used to produce a wide range of products, from pharmaceuticals to plastics and fertilizers
  • The Haber-Bosch process, which synthesizes ammonia ($N_2 + 3H_2 \rightleftharpoons 2NH_3$), is crucial for the production of fertilizers and explosives
• Electrochemical reactions, a type of redox reaction, are the basis for batteries and fuel cells, which convert chemical energy into electrical energy
  • In a lead-acid battery, the reaction between lead, lead dioxide, and sulfuric acid generates an electric current
• Catalysts are used in many industrial processes to increase reaction rates and efficiency, reducing energy consumption and waste
  • Catalytic converters in vehicles use platinum, palladium, and rhodium to convert harmful exhaust gases into less harmful substances
• Equilibrium principles are applied in the production of chemicals, such as the synthesis of ammonia, sulfuric acid, and methanol
  • The Contact process for producing sulfuric acid ($2SO_2 + O_2 \rightleftharpoons 2SO_3$) relies on optimizing equilibrium conditions to maximize yield

Common Challenges and Tips

• Balancing chemical equations can be challenging, especially for complex reactions with multiple reactants and products
  • Start by balancing the most complex molecule, then move on to simpler ones, adjusting coefficients as needed
• Determining reaction order and writing rate laws requires careful analysis of experimental data
  • Plot concentration vs. time data on a graph and analyze the shape of the curve to determine the order with respect to each reactant
• Predicting the direction of equilibrium shifts can be tricky when multiple factors change simultaneously
  • Consider each factor individually, then combine their effects to determine the overall shift
• Remembering the signs and meanings of $\Delta H$, $\Delta S$, and $\Delta G$ is essential for understanding thermodynamics
  • Associate "H" with "Heat," "S" with "Scatter" (disorder), and "G" with "Go" (spontaneity) to help remember their meanings
• When solving equilibrium problems, be sure to use the correct equilibrium constant expression based on the balanced equation
  • Pay attention to the stoichiometric coefficients and the phases of the reactants and products (solids and liquids are not included in the expression)
• Practice, practice, practice! Work through many different types of problems to reinforce your understanding of the concepts and improve your problem-solving skills
  • Focus on the key concepts and principles, and try to connect them to real-world examples to make the material more relatable and easier to remember