Chemical Kinetics - Rate of Disappearance
- Introduction
- Chemical kinetics: study of the speed at which chemical reactions occur
- Rate of disappearance: rate at which reactants are consumed
- Definition of rate of disappearance
- Rate of disappearance: negative rate of change of concentration of a reactant over time
- Measured in units of concentration/time (mol/Ls)
- Rate of disappearance equation
- For a general reaction: aA + bB → cC + dD
- Rate of disappearance of reactant A: -1/a * d[A]/dt
- Determining rate of disappearance experimentally
- Measure change in concentration of reactant over time
- Plot concentration vs. time
- Determine gradient (slope) of the resulting graph to find the rate of disappearance
- Example 1: Rate of disappearance of hydrogen peroxide
- Reaction: 2H2O2 → 2H2O + O2
- Rate of disappearance of H2O2: -1/2 * d[H2O2]/dt
- Example 2: Rate of disappearance of nitrogen dioxide
- Reaction: 2NO2 → 2NO + O2
- Rate of disappearance of NO2: -1/2 * d[NO2]/dt
- Example 3: Rate of disappearance of carbon monoxide
- Reaction: 2CO + O2 → 2CO2
- Rate of disappearance of CO: -1/2 * d[CO]/dt
- Factors affecting rate of disappearance
- Concentration of reactants: higher concentration leads to faster rate of disappearance
- Temperature: higher temperature increases the rate of disappearance
- Catalysts: presence of catalysts can increase the rate of disappearance
- Rate laws and rate constants
- Rate law: mathematical expression relating rate of disappearance to reactant concentrations
- Rate constant: proportionality constant in the rate law equation
- Conclusion and summary
Rate Laws
- Definition of rate law
- Rate law: equation that relates the rate of a chemical reaction to the concentrations of the reactants
- General form: rate = k[A]^m[B]^n
- Order of reaction: sum of the exponents (m + n)
- Determining rate laws experimentally
- Method of initial rates
- Measure initial rates with different initial concentrations of reactants
- Determine how rate changes as concentrations change
- Integrated rate laws
- Use data from concentration vs. time experiments to derive rate laws
- Method of initial rates
- First-order reactions
- Rate law for a first-order reaction: rate = k[A]
- Integrated rate law: ln[A] = -kt + ln[A]0
- Half-life for a first-order reaction: t1/2 = 0.693/k
- Second-order reactions
- Rate law for a second-order reaction: rate = k[A]^2
- Integrated rate law: 1/[A] = kt + 1/[A]0
- Half-life for a second-order reaction: t1/2 = 1/(k[A]0)
- Zero-order reactions
- Rate law for a zero-order reaction: rate = k[A]^0 = k
- Integrated rate law: [A] = -kt + [A]0
- Half-life for a zero-order reaction: t1/2 = [A]0/2k
- Determining the order of a reaction
- Method of initial rates can determine the orders of reactants
- Compare the change in rate with different initial concentrations of reactants
- Determining the rate constant (k)
- Use experimental data and the rate law to solve for the rate constant
Reaction Mechanisms
- Definition of reaction mechanism
- Reaction mechanism: step-by-step sequence of elementary reactions that make up an overall chemical reaction
- Elementary reactions and molecularity
- Elementary reaction: a single step in a reaction mechanism
- Molecularity: the number of reactant particles involved in an elementary reaction
- Unimolecular: one reactant particle
- Bimolecular: two reactant particles
- Termolecular: three reactant particles (rare)
- Rate-determining step
- Rate-determining step: slowest step in a reaction mechanism that determines the overall rate of the reaction
- Determines the rate law for the overall reaction
- Reversible reactions and equilibrium
- Reversible reactions: reactions that can proceed in both the forward and reverse directions
- Equilibrium: when the rates of the forward and reverse reactions are equal
- Reaction intermediates
- Reaction intermediates: species that are formed in one step of a reaction mechanism and consumed in a subsequent step
- Catalysts
- Catalyst: substance that increases the rate of a chemical reaction without being consumed in the reaction
- Provides an alternate reaction pathway with lower activation energy
- Summary of reaction mechanisms
- Reaction mechanisms explain the steps and intermediates involved in a chemical reaction
- The rate-determining step determines the overall rate of the reaction
- Catalysts can increase the rate of a reaction by providing an alternate pathway
Chemical Kinetics - Rate Laws
- Definition of rate laws
- Rate law: equation that relates the rate of a chemical reaction to the concentrations of the reactants
- General form: rate = k[A]^m[B]^n
- Order of reaction: sum of the exponents (m + n)
- Determining rate laws experimentally
- Method of initial rates
- Measure initial rates with different initial concentrations of reactants
- Determine how rate changes as concentrations change
- Integrated rate laws
- Use data from concentration vs. time experiments to derive rate laws
- Method of initial rates
- Example: First-order reaction
- Rate law: rate = k[A]
- Integrated rate law: ln[A] = -kt + ln[A]0
- Half-life: t1/2 = 0.693/k
- Example: Second-order reaction
- Rate law: rate = k[A]^2
- Integrated rate law: 1/[A] = kt + 1/[A]0
- Half-life: t1/2 = 1/(k[A]0)
- Example: Zero-order reaction
- Rate law: rate = k[A]^0 = k
- Integrated rate law: [A] = -kt + [A]0
- Half-life: t1/2 = [A]0/2k
- Determining the order of a reaction
- Method of initial rates can determine the orders of reactants
- Compare the change in rate with different initial concentrations of reactants
- Determining the rate constant (k)
- Use experimental data and the rate law to solve for the rate constant
Chemical Kinetics - Reaction Mechanisms
- Definition of reaction mechanism
- Reaction mechanism: step-by-step sequence of elementary reactions that make up an overall chemical reaction
- Elementary reactions and molecularity
- Elementary reaction: a single step in a reaction mechanism
- Molecularity: the number of reactant particles involved in an elementary reaction
- Unimolecular: one reactant particle
- Bimolecular: two reactant particles
- Termolecular: three reactant particles (rare)
- Rate-determining step
- Rate-determining step: slowest step in a reaction mechanism that determines the overall rate of the reaction
- Determines the rate law for the overall reaction
- Reversible reactions and equilibrium
- Reversible reactions: reactions that can proceed in both the forward and reverse directions
- Equilibrium: when the rates of the forward and reverse reactions are equal
- Reaction intermediates
- Reaction intermediates: species that are formed in one step of a reaction mechanism and consumed in a subsequent step
- Catalysts
- Catalyst: substance that increases the rate of a chemical reaction without being consumed in the reaction
- Provides an alternate reaction pathway with lower activation energy
- Summary of reaction mechanisms
- Reaction mechanisms explain the steps and intermediates involved in a chemical reaction
- The rate-determining step determines the overall rate of the reaction
- Catalysts can increase the rate of a reaction by providing an alternate pathway
Chemical Equilibrium
- Introduction to chemical equilibrium
- Chemical equilibrium: state in which the rates of the forward and reverse reactions are equal
- Equilibrium constant (K): ratio of the product concentrations to the reactant concentrations at equilibrium
- Writing the equilibrium constant expression
- Equilibrium constant expression: K = [C]^c[D]^d / [A]^a[B]^b
- [C], [D]: molar concentrations of products
- [A], [B]: molar concentrations of reactants
- c, d, a, b: coefficients of balanced equation
- Equilibrium constant expression: K = [C]^c[D]^d / [A]^a[B]^b
- Manipulating the equilibrium constant expression
- Multiplying a reaction by a factor: raise K to the power of that factor
- Reversing a reaction: take the reciprocal of K
- Combining reactions: multiply the K values
- Equilibrium concentrations and ICE tables
- Initial change equilibrium (ICE) tables: determine unknown concentrations at equilibrium
- Use stoichiometry and equilibrium concentrations to fill in the table
- Le Chatelier’s principle
- Le Chatelier’s principle: if a system at equilibrium is subjected to a stress, it will adjust to relieve that stress
- Changes in concentration, pressure, or temperature can shift the equilibrium position
- Factors affecting equilibrium
- Concentration changes: increase in reactants or decrease in products shifts equilibrium towards the products
- Pressure changes: increase in pressure favors the side with fewer moles of gas
- Temperature changes: exothermic reactions favor the products, endothermic reactions favor the reactants
- Solubility equilibrium
- Equilibrium involving the dissolution of a solid solute in a solvent
- Solubility product constant (Ksp): equilibrium constant for a solubility equilibrium
- Common ion effect
- Common ion effect: solubility of a slightly soluble salt decreases when a common ion is present in the solution
Acid-Base Equilibria
- Arrhenius definition of acids and bases
- Arrhenius acids: substances that increase the concentration of H+ ions in solution
- Arrhenius bases: substances that increase the concentration of OH- ions in solution
- Brønsted-Lowry definition of acids and bases
- Brønsted-Lowry acids: substances that donate protons (H+ ions)
- Brønsted-Lowry bases: substances that accept protons
- Conjugate acid-base pairs
- Conjugate acid: formed when a base accepts a proton
- Conjugate base: formed when an acid donates a proton
- Conjugate acid-base pairs differ by one proton
- Acid dissociation constant (Ka)
- Ka: measures the extent of acid dissociation in water
- Strong acids have large Ka values, weak acids have small Ka values
- pH and pOH
- pH: measure of the hydrogen ion concentration in a solution
- pOH: measure of the hydroxide ion concentration in a solution
- pH + pOH = 14 for water at 25°C
- Acid-base titrations
- Titrations: used to determine the concentration of an acid or base in a solution
- Use the stoichiometry of the reaction to calculate the unknown concentration
- Buffer solutions
- Buffer: solution that resists changes in pH when small amounts of acid or base are added
- Made from a weak acid and its conjugate base, or a weak base and its conjugate acid
- Buffer capacity: ability of a buffer to resist changes in pH
Electrochemistry
- Oxidation-reduction reactions (redox reactions)
- Oxidation: loss of electrons
- Reduction: gain of electrons
- Reducing agent: species that donates electrons (undergoes oxidation)
- Oxidizing agent: species that accepts electrons (undergoes reduction)
- Balancing redox equations
- Half-reaction method: separate the overall reaction into two half-reactions, balance each half-reaction, and combine them
- Electrochemical cells
- Electrochemical cell: device that uses redox reactions to convert chemical energy into electrical energy
- Consists of two half-cells: anode (site of oxidation) and cathode (site of reduction)
- Galvanic cells
- Galvanic (voltaic) cells: spontaneous redox reactions that produce electrical energy
- Electrons flow from anode to cathode through an external circuit
- Cell notation
- Cell notation: shorthand representation of an electrochemical cell
- Anode | Anode Solution || Cathode Solution | Cathode
- Cell notation: shorthand representation of an electrochemical cell
- Standard electrode potential (E°)
- Standard electrode potential: measure of the tendency of a half-reaction to occur as a reduction
- E° values are compared to the standard hydrogen electrode (SHE) at 25°C
- Nernst equation
- Nernst equation: relates the cell potential to the concentrations of reactants and products
- E = E° - (0.0592/n) * logQ, where Q is the reaction quotient
Nuclear Chemistry
- Introduction to nuclear chemistry
- Nuclear reactions involve changes in the nucleus of an atom
- Examples: radioactivity, nuclear decay, nuclear fission, nuclear fusion
- Types of nuclear decay
- Alpha decay: emission of an alpha particle (helium nucleus)
- Beta decay: emission of a beta particle (electron or positron)
- Gamma decay: emission of a gamma ray (high-energy electromagnetic radiation)
- Nuclear reactions and equation balancing
- Nuclear reactions are balanced based on mass number and atomic number
- Sum of mass numbers and atomic numbers must be equal on both sides of the equation
- Half-life
- Half-life: time it takes for half of a radioactive substance to decay
- Used to determine the radioactive decay rate and the age of radioactive materials
- Radioactive dating
- Radioactive isotopes can be used to determine the age of ancient artifacts and fossils
- Carbon-14 dating: used to date organic materials up to 50,000 years old
- Nuclear power and nuclear reactors
- Nuclear fission: splitting of heavy atomic nuclei to release energy
- Nuclear reactors use controlled fission reactions to generate heat and electricity
- Safety measures and waste disposal are important considerations
Organic Chemistry
- Introduction to organic chemistry
- Organic chemistry: study of compounds containing carbon
- Carbon can form long chains and different structural arrangements, leading to diverse organic compounds
- Functional groups
- Functional groups: specific arrangements of atoms within organic molecules that give them characteristic chemical behaviors
- Examples: alcohols, aldehydes, ketones, carboxylic acids, esters, amines
- Nomenclature of organic compounds
- IUPAC nomenclature: systematic way of naming organic compounds based on their structure and functional groups
- Prefixes and suffixes are used to indicate the various components of the compound
- Isomerism
- Isomers: compounds with the same molecular formula but different structural arrangements or spatial orientations
- Structural isomers: differ in the arrangement of atoms
- Stereoisomers: differ in the spatial arrangement of atoms
- Reactions of organic compounds
- Organic compounds undergo a wide range of reactions, including substitution, addition, elimination, and oxidation-reduction reactions
- Polymerization
- Polymerization: process of forming large molecules (polymers) by combining smaller units (monomers)
- Addition polymerization: monomers add together without the loss of any atoms
- Condensation polymerization: monomers join together and a small molecule (often water) is eliminated
- Organic synthesis
- Organic synthesis: process of designing and creating new organic compounds
- Requires knowledge of reaction mechanisms, reactivity of functional groups, and manipulation of reaction conditions
Carbohydrates and Lipids
- Introduction to carbohydrates
- Carbohydrates: organic compounds composed of carbon, hydrogen, and oxygen in a 1:2:1 ratio
- Functions: energy storage, structural support, cell recognition
- Monosaccharides
- Monosaccharides: simple sugars that cannot be chemically hydrolyzed to yield smaller carbohydrates
- Examples: glucose, fructose, ribose
- Disaccharides
- Disaccharides: formed by the linking of two monosaccharide units via a glycosidic bond
- Examples: sucrose, lactose, maltose
- Polysaccharides
- Polysaccharides: long chains of monosaccharide units
- Functions: energy storage (starch, glycogen), structural support (cellulose, chitin)
- Introduction to lipids
- Lipids: hydrophobic organic compounds that are insoluble in water
- Functions: energy storage, insulation, structural component of cell membranes, hormone production
- Fatty acids and triglycerides
- Fatty acids: long hydrocarbon chains with a carboxylic acid group
- Triglycerides: formed by the esterification of three fatty acids to a glycerol molecule
- Functions: energy storage, thermal insulation
- Phospholipids and cell membranes
- Phospholipids: composed of a glycerol molecule, two fatty acid chains, and a phosphate group
- Form the main structural component of cell membranes
Proteins and Enzymes
- Introduction to proteins
- Proteins: large complex molecules composed of amino acid building blocks
- Functions: structural support, enzyme catalysis, transport, defense, regulation
- Amino acids
- Amino acids: contain an amino group, carboxyl group, and a side chain (R group)
- Differ in their R group, which determines their properties and functions
- Peptide bonds and polypeptides
- Peptide bond
Chemical Kinetics - Rate of Disappearance Introduction Chemical kinetics: study of the speed at which chemical reactions occur Rate of disappearance: rate at which reactants are consumed Definition of rate of disappearance Rate of disappearance: negative rate of change of concentration of a reactant over time Measured in units of concentration/time (mol/Ls) Rate of disappearance equation For a general reaction: aA + bB → cC + dD Rate of disappearance of reactant A: -1/a * d[A]/dt Determining rate of disappearance experimentally Measure change in concentration of reactant over time Plot concentration vs. time Determine gradient (slope) of the resulting graph to find the rate of disappearance Example 1: Rate of disappearance of hydrogen peroxide Reaction: 2H2O2 → 2H2O + O2 Rate of disappearance of H2O2: -1/2 * d[H2O2]/dt Example 2: Rate of disappearance of nitrogen dioxide Reaction: 2NO2 → 2NO + O2 Rate of disappearance of NO2: -1/2 * d[NO2]/dt Example 3: Rate of disappearance of carbon monoxide Reaction: 2CO + O2 → 2CO2 Rate of disappearance of CO: -1/2 * d[CO]/dt Factors affecting rate of disappearance Concentration of reactants: higher concentration leads to faster rate of disappearance Temperature: higher temperature increases the rate of disappearance Catalysts: presence of catalysts can increase the rate of disappearance Rate laws and rate constants Rate law: mathematical expression relating rate of disappearance to reactant concentrations Rate constant: proportionality constant in the rate law equation Conclusion and summary