chm 12901 general chemistry with a biological focus unit 4 study guides

Key Concepts

• Chemical reactions involve the rearrangement of atoms to form new substances with different properties than the original reactants
• Reactants are the starting materials that undergo a chemical change to form products, the resulting substances formed from the reaction
• Conservation of mass states that the total mass of the reactants must equal the total mass of the products in a chemical reaction
• Stoichiometry is the quantitative study of the relative 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 a reaction by lowering the activation energy without being consumed in the reaction
• Reaction rates describe how quickly a chemical reaction proceeds and can be influenced by factors such as temperature, concentration, and surface area
• Chemical equilibrium is a dynamic state where the forward and reverse reactions occur at equal rates, resulting in no net change in the concentrations of reactants and products

Types of Chemical Reactions

• Synthesis reactions involve the combination of two or more reactants to form a single product (2H2 + O2 -> 2H2O)
• Decomposition reactions involve the breakdown of a single reactant into two or more products (2H2O2 -> 2H2O + O2)
• Single displacement reactions occur when one element replaces another element in a compound (Zn + 2HCl -> ZnCl2 + H2)
• Double displacement reactions involve the exchange of ions between two compounds to form two new compounds (NaCl + AgNO3 -> AgCl + NaNO3)
• Combustion reactions are rapid oxidation reactions that produce heat and light, often involving hydrocarbons reacting with oxygen (CH4 + 2O2 -> CO2 + 2H2O)
• Acid-base reactions involve the transfer of protons (H+) from an acid to a base, forming a conjugate acid-base pair (HCl + NaOH -> NaCl + H2O)
• Redox reactions involve the transfer of electrons between species, resulting in changes in oxidation states (2Na + Cl2 -> 2NaCl)

Balancing Chemical Equations

• Chemical equations represent the reactants, products, and their stoichiometric relationships in a reaction
• Balanced equations have equal numbers of each type of atom on both the reactant and product sides
• Coefficients are used to balance equations and indicate the relative amounts of each species involved in the reaction
• The law of conservation of mass is maintained in balanced chemical equations
• Subscripts in chemical formulas cannot be changed when balancing equations, as they represent the actual composition of the compound
• Balancing equations involves adjusting coefficients through trial and error or using algebraic methods until the number of each type of atom is equal on both sides
• Balanced equations provide information about the mole ratios of reactants and products, which is essential for stoichiometric calculations

Reaction Rates and Kinetics

• Reaction rate is the change in concentration of a reactant or product per unit time and is influenced by factors such as temperature, concentration, pressure, and catalyst presence
• Collision theory states that reactions occur when reactant particles collide with sufficient energy (activation energy) and proper orientation
• Increasing temperature increases the average kinetic energy of particles, leading to more collisions with sufficient activation energy and faster reaction rates
• Increasing reactant concentration increases the frequency of collisions, resulting in faster reaction rates
• Increasing surface area of solid reactants exposes more particles to collisions, leading to faster reaction rates
• Catalysts lower the activation energy of a reaction, allowing more collisions to result in successful reactions and increasing reaction rates
• Reaction mechanisms describe the step-by-step sequence of elementary reactions that occur to convert reactants into products
• Rate laws express the relationship between the reaction rate and the concentrations of reactants, often determined experimentally

Equilibrium in Chemical Reactions

• Chemical equilibrium is a dynamic state where the forward and reverse reactions occur at equal rates, resulting in no net change in the concentrations of reactants and products
• The equilibrium constant (K) is the ratio of the product of the concentrations of the products raised to their stoichiometric coefficients divided by the product of the concentrations of the reactants raised to their stoichiometric coefficients
• The value of K indicates the position of equilibrium: K > 1 means products are favored, K < 1 means reactants are favored, and K = 1 means reactants and products are equally favored
• Le Chatelier's principle states that when a system at equilibrium is disturbed, the system will shift to counteract the disturbance and re-establish equilibrium
• Increasing the concentration of a reactant or decreasing the concentration of a product will shift the equilibrium towards the products
• Increasing the concentration of a product or decreasing the concentration of a reactant will shift the equilibrium towards the reactants
• Changing the temperature of a system at equilibrium will shift the equilibrium in the direction that opposes the temperature change (endothermic reactions are favored by increased temperature, while exothermic reactions are favored by decreased temperature)
• Changes in pressure or volume affect gaseous equilibria: increasing pressure or decreasing volume will shift the equilibrium towards the side with fewer moles of gas

Thermodynamics of Reactions

• Thermodynamics is the study of energy changes in chemical reactions and physical processes
• Enthalpy (H) is the total heat content of a system and is a state function, meaning its value depends only on the current state of the system, not the path taken to reach that state
• Exothermic reactions release energy to the surroundings and have a negative change in enthalpy (ΔH < 0)
• Endothermic reactions absorb energy from the surroundings and have a positive change in enthalpy (ΔH > 0)
• Entropy (S) is a measure of the disorder or randomness of a system and is also a state function
• The second law of thermodynamics states that the entropy of the universe always increases in a spontaneous process
• Gibbs free energy (G) is a measure of the usable energy in a system and is defined as G = H - TS, where T is the absolute temperature
• A negative change in Gibbs free energy (ΔG < 0) indicates a spontaneous process, while a positive change (ΔG > 0) indicates a non-spontaneous process
• At equilibrium, ΔG = 0, and the system has reached a state of maximum entropy and minimum free energy

Biological Applications

• Enzymes are biological catalysts that lower the activation energy of biochemical reactions, allowing them to occur at physiological temperatures and pressures
• Enzyme-substrate complexes form when the active site of an enzyme binds to a specific substrate, facilitating the reaction
• Michaelis-Menten kinetics describes the relationship between enzyme concentration, substrate concentration, and reaction rate in enzyme-catalyzed reactions
• Cellular respiration is a series of enzyme-catalyzed reactions that break down glucose to release energy in the form of ATP, involving glycolysis, the citric acid cycle, and oxidative phosphorylation
• Photosynthesis is a series of redox reactions that convert light energy into chemical energy in the form of glucose, involving light-dependent reactions and the Calvin cycle
• Metabolism refers to the sum of all chemical reactions in a living organism, including anabolic (building) and catabolic (breaking down) processes
• pH plays a crucial role in biological systems, affecting enzyme activity, protein structure, and cellular functions
• Buffers are solutions that resist changes in pH when small amounts of acid or base are added, helping maintain homeostasis in biological systems

Lab Techniques and Safety

• Always wear appropriate personal protective equipment (PPE) such as lab coats, gloves, and safety goggles when working with chemicals
• Use a fume hood when working with volatile or hazardous substances to minimize exposure to fumes and vapors
• Properly label all chemicals and solutions with the name, concentration, and any relevant safety information
• Use caution when handling glassware, as it can break or crack, causing cuts or spills
• Never pipette by mouth; always use a pipette aid or bulb to avoid accidental ingestion of chemicals
• Dispose of chemical waste according to the guidelines provided by your instructor or lab supervisor, using appropriate waste containers
• Know the location and proper use of safety equipment such as eye wash stations, safety showers, and fire extinguishers
• Familiarize yourself with the safety data sheets (SDS) for the chemicals you will be using, which provide information on hazards, handling, and emergency procedures
• Report any accidents, spills, or injuries to your instructor or lab supervisor immediately, no matter how minor they may seem