Chemical Kinetics - Rate Of Disappearance

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
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
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
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
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