Chemical Kinetics and Reaction Mechanisms
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Questions and Answers

What is the definition of an elementary step in a chemical reaction?

  • A step in a reaction that involves the formation of a stable intermediate.
  • A step in a reaction that is always the slowest step in the overall mechanism.
  • A step in a reaction that requires either 1, 2, or 3 particle collisions. (correct)
  • A step in a reaction that occurs at a constant rate regardless of the concentration of reactants.

Which factor does NOT directly affect the rate of a chemical reaction?

  • The chemical nature of the reactants
  • The volume of the reaction vessel (correct)
  • The presence of a catalyst
  • The concentration of the reactants

Which of the following is NOT an example of how surface area can affect the rate of a chemical reaction?

  • A reaction involving a gas will proceed faster in a closed container. (correct)
  • Powdered sugar dissolves faster than sugar cubes.
  • Burning a log in small pieces will produce more heat than burning one large log.
  • A finely ground metal reacts faster than a large piece of metal.

How does the activation energy relate to the rate of a chemical reaction?

<p>The lower the activation energy, the slower the reaction rate. (B)</p> Signup and view all the answers

What is the relationship between the exponents in the rate equation and the coefficients in the rate-determining step of a reaction mechanism?

<p>The exponents in the rate equation correspond to the coefficients in the rate-determining step. (A)</p> Signup and view all the answers

What is an activated complex?

<p>An unstable species with partially broken and formed bonds at the peak of potential energy. (B)</p> Signup and view all the answers

Which of these factors is NOT a primary factor affecting the rate of a heterogeneous reaction?

<p>The concentration of reactants. (A)</p> Signup and view all the answers

What was the significance of the dark background with discrete lines observed when light from elements was passed through a spectrometer?

<p>It provided evidence for the existence of quantized energy levels in atoms, suggesting that electrons could only exist at specific energy levels. (D)</p> Signup and view all the answers

According to Bohr's model, what happens when an atom gains energy?

<p>The atom's electrons jump to higher energy levels. (B)</p> Signup and view all the answers

What was the major problem with Bohr's model that prevented it from explaining the spectra of atoms with more than one electron in their outer shell?

<p>It did not consider the interactions between electrons in multi-electron atoms. (D)</p> Signup and view all the answers

What is the main reason why electrons in atoms don't constantly emit light according to classical physics?

<p>The electrons only emit light when they transition between energy levels, not when they are in a stable orbit. (B)</p> Signup and view all the answers

Why was the Bohr model considered a significant advancement in understanding atomic structure?

<p>It was the first theoretical model to explain why electrons are confined to specific energy levels. (A)</p> Signup and view all the answers

What is the minimum kinetic energy required for a reaction to occur called?

<p>Threshold energy (D)</p> Signup and view all the answers

What does Hess’s Law state about reaction enthalpies?

<p>The enthalpy change for a reaction equals the sum of the enthalpy changes for its steps. (B)</p> Signup and view all the answers

What must be considered when writing a standard enthalpy of formation equation?

<p>The coefficients must yield one mole of product from the elements. (D)</p> Signup and view all the answers

What is the standard enthalpy of formation for an element in its standard state?

<p>It is zero. (A)</p> Signup and view all the answers

If the coefficient of a chemical equation is doubled, how is ΔH affected?

<p>ΔH is doubled. (A)</p> Signup and view all the answers

What is the formula for calculating ΔH based on standard enthalpies of formation?

<p>ΔH° = sum of enthalpies of products - sum of enthalpies of reactants. (B)</p> Signup and view all the answers

Which of these defines the average rate of reaction?

<p>The change in concentration of products over time. (C)</p> Signup and view all the answers

What is the primary focus of chemical kinetics?

<p>Investigating ways to alter the speed of reactions. (D)</p> Signup and view all the answers

In a thermochemical equation, what is important about the ΔH value?

<p>It must have the correct sign indicating whether the reaction is exothermic or endothermic. (B)</p> Signup and view all the answers

What does the rate constant 'k' represent in the rate law equation?

<p>The proportionality constant specific to a reaction (C)</p> Signup and view all the answers

If the rate of a reaction depends on [reactant] raised to the power of 2, what is the effect of doubling the initial concentration on the rate?

<p>The rate quadruples (C)</p> Signup and view all the answers

What is the overall order of reaction if the order with respect to NO is 1 and with respect to F2 is also 1?

<p>2 (C)</p> Signup and view all the answers

Which of the following describes the half-life of a reaction?

<p>The time required for the concentration of the reactant to halve (D)</p> Signup and view all the answers

How does increasing the concentration of reactants affect the reaction rate?

<p>It increases the reaction rate (A)</p> Signup and view all the answers

For first-order reactions, what is the relationship between the natural logarithm of concentrations at different times?

<p>ln[𝐴]𝑡 = -kt + ln[𝐴]𝑜 (A)</p> Signup and view all the answers

What effect does an increase in temperature generally have on reaction rates?

<p>It speeds up the reaction (C)</p> Signup and view all the answers

In a second-order reaction, how is the relationship between rate and concentration expressed?

<p>1/[A]t = kt + 1/[A]𝑜 (A)</p> Signup and view all the answers

What did Louis de Broglie suggest about electrons?

<p>Electrons can behave like both waves and particles. (A)</p> Signup and view all the answers

According to the uncertainty principle proposed by Werner Heisenberg, what is impossible to determine simultaneously?

<p>The exact position and speed of an electron. (B)</p> Signup and view all the answers

What does the principal quantum number (n) represent?

<p>The energy level of the electron. (C)</p> Signup and view all the answers

Which orbital has a dumb-bell shape and can have a maximum of three orientations?

<p>P orbital (A)</p> Signup and view all the answers

What is a consequence of the spin quantum number in relation to electrons?

<p>No two electrons can have the same set of four quantum numbers. (B)</p> Signup and view all the answers

Flashcards

Enthalpy Change (ΔH)

The change in enthalpy for a reaction, expressed in kJ, indicating whether the reaction releases or absorbs heat.

Thermochemical Equation

A chemical equation showing the enthalpy change (ΔH) alongside the balanced equation, with a positive sign for endothermic reactions and a negative sign for exothermic reactions.

Standard Enthalpy of Formation (ΔH°f)

The enthalpy change (ΔH) associated with the formation of one mole of a substance from its elements in their most stable forms under standard conditions.

Calculating ΔH using Standard Enthalpies of Formation

The enthalpy change for a reaction calculated by adding the standard enthalpies of formation of products and subtracting the standard enthalpies of formation of reactants.

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Potential Energy Diagram

A diagram depicting the energy changes that occur during a chemical reaction, showing the relative energy levels of reactants, products, and transition states.

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Hess's Law

The principle that the enthalpy change for a reaction is the same whether the reaction occurs in one step or in a series of steps.

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Rate of Reaction

The rate of a reaction, which is a measure of how quickly reactants are consumed or products are formed.

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Chemical Kinetics

The study of the rates and mechanisms of chemical reactions.

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Order of reaction

The exponent in the rate law equation that represents the effect of the initial concentration of a specific reactant on the reaction rate.

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Rate constant (k)

The rate constant is a proportionality constant in the rate law equation. It represents the specific speed of a reaction at a particular temperature.

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Overall order of reaction

The sum of the exponents in the rate law equation. It indicates the overall sensitivity of the reaction rate to changes in reactant concentrations.

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Half-life of a reaction

The time it takes for the concentration of a reactant to decrease to half its initial concentration.

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Integrated rate law equation

An equation that relates the concentration of a reactant to time. It allows you to predict the concentration of a reactant at any time during the reaction.

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First order integrated rate law

The integrated rate law equation for a first-order reaction. It describes the exponential decay of a reactant's concentration over time.

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Second order integrated rate law

The integrated rate law equation for a second-order reaction. It shows how the reciprocal of a reactant's concentration changes linearly with time.

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Half-life of a first order reaction

The time it takes for the concentration of a reactant to decrease to half its initial concentration in a first-order reaction. For a first-order reaction, the half-life is constant and independent of the initial concentration.

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Activation Energy

The minimum amount of energy reactants need to overcome to start a reaction.

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Elementary Step

A single step in a chemical reaction that involves collisions between 1, 2, or 3 molecules.

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Reaction Mechanism

A series of elementary steps that make up the overall reaction.

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Reaction Intermediates

Molecules formed during a reaction mechanism that are short-lived and react further to form products.

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Activated Complex

An unstable, high-energy state where bonds are partially broken and formed during a reaction.

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Rate-Determining Step

The slowest step in a multi-step reaction that determines the overall rate of the reaction.

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Catalyst

A substance that speeds up a chemical reaction without being consumed in the process.

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Surface Area

The amount of surface area exposed in a heterogeneous reaction, affecting how quickly reactants interact.

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Threshold Energy

The minimum kinetic energy required for molecules to collide and form an activated complex, essentially converting kinetic energy into activation energy.

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Maxwell-Boltzmann Distribution

The distribution of kinetic energies of molecules in a sample at a given temperature. Represented by a graph where the vertical axis shows the number of molecules and the horizontal axis shows the kinetic energy.

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Homogeneous Catalyst

A catalyst that exists in the same physical state as the reactants. It speeds up a reaction without being consumed itself.

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Heterogeneous Catalyst

A catalyst that exists in a different physical state than the reactants. It increases the reaction rate but doesn't appear in the chemical equation.

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Rate Law

A mathematical equation that describes the relationship between the rate of a reaction and the concentrations of the reactants, including the rate constant (k) and the order of the reaction (n and m).

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Bohr Model

This model suggested that electrons could only exist at certain energy levels, explaining atomic spectra and the periodic table.

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Atomic Spectrum

The unique pattern of light emitted by an element when its atoms are excited.

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Excitation

The process of an atom gaining energy and moving to a higher energy level.

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De-excitation

The process of an atom losing energy and returning to a lower energy level.

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Energy Levels

The specific energy levels that electrons can occupy within an atom.

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Electron Waves

Electrons can behave like waves, with their properties described by wave functions.

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Wave Function

A mathematical function that describes the probability of finding an electron in a specific region of space.

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Quantum Numbers

A set of three numbers that describe the energy level, shape, and orientation of an electron's orbital.

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Orbitals

Regions of space around the nucleus where electrons are most likely to be found, determined by the shape of the wave function.

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Pauli Exclusion Principle

The principle stating that no two electrons in an atom can have the same four quantum numbers, meaning each orbital can hold a maximum of two electrons.

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Study Notes

Exam Notes

  • Grade 12 Chemistry Exam, Wednesday, January 29, 2025
  • Exam duration: 2 hours (120 minutes)
  • Exam worth: 30%
  • Exam is broken down into parts covering different types of questions: multiple choice, problem-solving, and short answer.

Mark Breakdown

  • Part A: Multiple Choice (K) - 30 marks, 35 minutes.
  • Part B: Problem Solving & Calculations (T) - 40 marks, 50 minutes.
  • Part D: Short Answer (C) - 17 marks, 20 minutes.
  • Part E: Short Answer (A) - 13 marks, 15 minutes.

Unit 1: Energy Changes and Rates of Reaction

  • Unit covers topics on energy changes and reaction rates.
  • Covers chapters 5 and 6.
  • Includes definitions of kinetic, potential, and chemical energy.
  • Discusses different types of systems (open, closed, isolated).
  • Includes important terms relating to energy changes within chemical systems.

Unit 5: Energy Changes and Rates of Reaction Continued.

  • Covers molar enthalpies, enthalpy change, and standard enthalpy change.
  • Includes specific heat capacity (c) as a calculation, related to heat (q), mass (m), and temperature change (ΔT).
  • Includes enthalpy changes for physical, chemical, and nuclear changes.
  • Covers representing enthalpy changes using thermochemical equations and ΔH values.
  • Detailed calculation/formula for enthalpy change (ΔH).
  • Covers Hess's Law of Additivity of Reaction Enthalpies,
  • Covers Hess's Law for enthalpy changes of steps/reactions.
  • Describes enthalpy changes for reactions using the terms endothermic or exothermic.
  • Presents methods to calculate/measure enthalpy change.

Unit 6: Rate of Reaction

  • Describes chemical kinetics.
  • Provides definitions for rate of reaction and average rate of reaction.
  • Defines theories related to reaction rates and measuring reaction rates.
  • Covers rate law, order of reaction, and rate constants.
  • Identifies factors affecting reaction rates.
  • Explains different methods of obtaining rate law values.
  • Includes a discussion on factors affecting rate of reactions (temperature, surface area, concentration, and presence of a catalyst).

Unit 7: Hess's Law and Enthalpy of Formation Calculation

  • Explains Hess's Law of Heat Summation.
  • Discusses standard enthalpy of formation, with formulas
  • Provides definitions of enthalpy and related concepts
  • Presents the use of enthalpy in chemical calculations.
  • Provides detailed descriptions of the calculations and how to apply them to various problems.

Unit 8: Rate Laws and Order of Reactions Continued

  • Covers theoretical explanations of Factors Affecting Reaction Rates.
  • Discusses concepts of activation energy and the rate-determining step.
  • Defines concepts of rate laws, reaction mechanisms, and reaction intermediates.
  • Explains concepts of collision theory, including activation energy, effective collisions and reaction mechanisms.
  • Explains order of reaction with respect to reactants, concentration, etc.
  • Explains half-life of a reaction.
  • Discusses integrated rate law equations for first and second-order reactions.
  • Discusses different methods of calculating rate laws.

Unit 9: Chemical Equilibrium Systems

  • Discusses dynamic equilibrium −Describes reversible reactions −Defines and explains reaction quotient (Q) and equilibrium constant (K).
  • Provides calculations relating to the equilibrium constant
  • Covers systems involved with solids, liquids and gasses.
  • Describes qualitative and quantitative changes in equilibrium systems using Le Chatelier's Principle.
  • Discusses concepts of various changes in equilibrium.

Unit 10: Acid-Base Equilibria

  • Discusses acid-base equilibria, including definitions of acids, bases, and self-ionization of water (the autodissociation of water).
  • Covers Arrhenius definitions, along with Bronsted-Lowry definitions, for acids and bases.
  • Includes definitions of conjugate acid-base pairs.
  • Covers concepts of ionization constant (Ka) and base constant (Kb) for weak acids and bases.
  • Covers neutralization reactions, concepts of titration and equivalence points.
  • Includes a discussion of percentage ionization calculations.
  • Covers acid-base properties of salt solutions and includes definitions of hydrolysis reactions and how to determine if a salt solution will be acidic, basic, or neutral.

Unit 11: Quantitative Changes in Equilibrium Systems

  • Describes calculations for equilibrium constant (K) and reaction quotient (Q).
  • Provides discussion and calculations on how to determine whether a chemical system is at equilibrium based on the relative magnitudes of Q and K.
  • Includes discussion of calculations using the quadratic formula.
  • Covers Le Chatelier's principle (and factors influencing equilibrium reactions).

Unit 12: Electrochemistry

  • Covers the concepts of oxidation-reduction reactions.
  • Provides definitions of oxidation and reduction, and concepts of oxidizing and reducing agents.
  • Covers electron transfer theory and relates it to redox reactions, including balancing oxidation-reduction (redox) reactions using half-reaction method.
  • Introduces oxidation numbers and provides examples and calculations of oxidation numbers for various atoms for ionic and covalent compounds.

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Grade 12 Chemistry Exam Notes PDF

Description

Test your knowledge on chemical kinetics and the various factors influencing reaction rates. Explore concepts like elementary steps, activation energy, and Bohr's model. This quiz will help reinforce your understanding of reaction mechanisms and their significance in chemistry.

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