This document discusses chemical kinetics and reaction rates. It defines kinetics as the study of reaction rates and mechanisms. Reaction rates describe how quickly reactions occur, while thermodynamics determines whether reactions can occur. The rate of a reaction is the change in concentration of reactants or products over time. Reaction mechanisms involve the molecular steps of reactions. Rate laws are determined experimentally and describe the dependence of reaction rates on reactant concentrations. Integrated rate laws relate concentration to time for reactions and can be used to determine amounts of reactants or products over time. First-order reactions follow integrated rate laws of ln[A] = -kt or [A] = [A]0e-kt, where k is the rate constant and t is time.
Half-reactions indicate the mole ratio of electrons to ions involved in redox reactions. The document discusses how current, time, and charge are related based on 1 mole of electrons equating to 96,500 Coulombs of charge. It provides examples calculating the mass of copper produced from electrolysis and the volume of chlorine gas produced from an industrial electrolysis process based on given values of current and time.
The document discusses reaction kinetics and rate laws. It defines key terms like rate law, order of reaction, and rate constant. The rate law expresses the relationship between the rate of a reaction and the concentrations of reactants raised to powers corresponding to their order. The order of a reaction with respect to a reactant is the exponent on its concentration term in the rate expression. The total order is the sum of all exponents. Examples are provided to demonstrate how to determine orders from rate laws and write rate expressions.
The document summarizes key concepts about rates of reaction from collision theory and kinetic molecular theory. It discusses how the rate of a reaction can be measured by changes in concentration over time and defines instantaneous and average rates. Factors that affect the rate of reaction are concentration, surface area, temperature, and catalysts. Increasing concentration, surface area, or temperature increases the frequency and successful energy of collisions between reactant particles according to collision theory. Catalysts increase the reaction rate by lowering the activation energy needed for reactions.
Enzyme kinetics is the study of chemical reactions catalyzed by enzymes. The reaction rate is measured under varying conditions to investigate the catalytic mechanism of the enzyme. The Michaelis-Menten model describes enzyme kinetics as a two-step process where the substrate forms a reversible enzyme-substrate complex that then decomposes into products. This model derives the Michaelis-Menten equation which relates reaction rate to substrate concentration and defines kinetic parameters like the Michaelis constant Km and maximum reaction rate Vmax.
This document discusses oxidation-reduction (redox) reactions. It defines oxidation as the gain of oxygen, loss of hydrogen, or loss of electrons, while reduction is defined as the loss of oxygen, gain of hydrogen, or gain of electrons. Redox reactions involve both oxidation and reduction occurring simultaneously. Oxidation numbers are assigned to elements to indicate degree of oxidation. Common oxidizing agents are listed that can oxidize Fe2+ to Fe3+, like potassium manganate, chlorine, and nitric acid. Reducing agents like sodium sulfite can reduce Fe3+ back to Fe2+. The document also covers IUPAC nomenclature of inorganic compounds and rules for assigning oxidation numbers.
This document discusses chemical equilibrium, including definitions, characteristics, and factors that affect equilibrium. It defines chemical equilibrium as a state where the forward and reverse reaction rates are equal. Characteristics include the dynamic nature of equilibrium and constant concentrations of reactants and products at equilibrium. Factors that affect equilibrium position include concentration, pressure, temperature, and catalyst additions according to Le Chatelier's principle. The relationship between the equilibrium constant K and standard Gibbs free energy change ΔG° is also described.
- Adsorption occurs when a gas or liquid accumulates on the surface of a solid, forming a film. It differs from absorption which involves diffusion into the bulk.
- The Langmuir adsorption model describes monolayer adsorption on uniform sites but makes assumptions that do not always apply. The BET model extends it to account for multilayer adsorption.
- The Temkin isotherm accounts for indirect interactions between adsorbed molecules which affect heat of adsorption and coverage at high pressures.
This document discusses reaction rates and factors that affect reaction rates. It begins by explaining what reaction rate means and how it can be measured by determining the amount of product formed, reactant used, or time taken for a reaction to complete. It then discusses several factors that affect reaction rates, including temperature, concentration, particle size, and catalysts. Higher temperatures, concentrations, smaller particle sizes, and the presence of catalysts generally increase reaction rates by increasing the frequency and success of particle collisions. The document provides examples and experiments to illustrate these concepts.
Half-reactions indicate the mole ratio of electrons to ions involved in redox reactions. The document discusses how current, time, and charge are related based on 1 mole of electrons equating to 96,500 Coulombs of charge. It provides examples calculating the mass of copper produced from electrolysis and the volume of chlorine gas produced from an industrial electrolysis process based on given values of current and time.
The document discusses reaction kinetics and rate laws. It defines key terms like rate law, order of reaction, and rate constant. The rate law expresses the relationship between the rate of a reaction and the concentrations of reactants raised to powers corresponding to their order. The order of a reaction with respect to a reactant is the exponent on its concentration term in the rate expression. The total order is the sum of all exponents. Examples are provided to demonstrate how to determine orders from rate laws and write rate expressions.
The document summarizes key concepts about rates of reaction from collision theory and kinetic molecular theory. It discusses how the rate of a reaction can be measured by changes in concentration over time and defines instantaneous and average rates. Factors that affect the rate of reaction are concentration, surface area, temperature, and catalysts. Increasing concentration, surface area, or temperature increases the frequency and successful energy of collisions between reactant particles according to collision theory. Catalysts increase the reaction rate by lowering the activation energy needed for reactions.
Enzyme kinetics is the study of chemical reactions catalyzed by enzymes. The reaction rate is measured under varying conditions to investigate the catalytic mechanism of the enzyme. The Michaelis-Menten model describes enzyme kinetics as a two-step process where the substrate forms a reversible enzyme-substrate complex that then decomposes into products. This model derives the Michaelis-Menten equation which relates reaction rate to substrate concentration and defines kinetic parameters like the Michaelis constant Km and maximum reaction rate Vmax.
This document discusses oxidation-reduction (redox) reactions. It defines oxidation as the gain of oxygen, loss of hydrogen, or loss of electrons, while reduction is defined as the loss of oxygen, gain of hydrogen, or gain of electrons. Redox reactions involve both oxidation and reduction occurring simultaneously. Oxidation numbers are assigned to elements to indicate degree of oxidation. Common oxidizing agents are listed that can oxidize Fe2+ to Fe3+, like potassium manganate, chlorine, and nitric acid. Reducing agents like sodium sulfite can reduce Fe3+ back to Fe2+. The document also covers IUPAC nomenclature of inorganic compounds and rules for assigning oxidation numbers.
This document discusses chemical equilibrium, including definitions, characteristics, and factors that affect equilibrium. It defines chemical equilibrium as a state where the forward and reverse reaction rates are equal. Characteristics include the dynamic nature of equilibrium and constant concentrations of reactants and products at equilibrium. Factors that affect equilibrium position include concentration, pressure, temperature, and catalyst additions according to Le Chatelier's principle. The relationship between the equilibrium constant K and standard Gibbs free energy change ΔG° is also described.
- Adsorption occurs when a gas or liquid accumulates on the surface of a solid, forming a film. It differs from absorption which involves diffusion into the bulk.
- The Langmuir adsorption model describes monolayer adsorption on uniform sites but makes assumptions that do not always apply. The BET model extends it to account for multilayer adsorption.
- The Temkin isotherm accounts for indirect interactions between adsorbed molecules which affect heat of adsorption and coverage at high pressures.
This document discusses reaction rates and factors that affect reaction rates. It begins by explaining what reaction rate means and how it can be measured by determining the amount of product formed, reactant used, or time taken for a reaction to complete. It then discusses several factors that affect reaction rates, including temperature, concentration, particle size, and catalysts. Higher temperatures, concentrations, smaller particle sizes, and the presence of catalysts generally increase reaction rates by increasing the frequency and success of particle collisions. The document provides examples and experiments to illustrate these concepts.
1. The document discusses chemical equilibrium, including the concept that at equilibrium the forward and reverse reactions proceed at the same rate, and the amounts of reactants and products remain constant.
2. It introduces the equilibrium constant expression and explains how to write the expression for different chemical equations.
3. Le Châtelier's principle is discussed, that systems at equilibrium will shift in response to changes in conditions to counteract the effect of changes in temperature, pressure, or concentration.
This document summarizes key differences between ideal and real gases. It discusses that real gases interact with each other, taking up physical space and attracting one another, resulting in lower than expected pressures. Real gas properties depart from ideal behavior at low temperatures or high pressures. The partition function approach is used to derive equations of state for both ideal and real gases. For real gases, the virial equation accounts for intermolecular interactions through virial coefficients.
B.tech. ii engineering chemistry unit 5 A electrochemistryRai University
Arrhenius proposed the theory of electrolytic dissociation to explain the properties of electrolytic solutions. The theory states that when an electrolyte dissolves in water, it breaks up into ions - positively charged cations and negatively charged anions. This process is called ionization. Ions are constantly recombining and dissociating, reaching a state of dynamic equilibrium. The extent of ionization depends on an equilibrium constant. Strong electrolytes have a high equilibrium constant and ionize completely, while weak electrolytes have a low constant and only partially ionize.
This document discusses the key concepts of chemical equilibrium. It defines reversible reactions as those that can proceed in both the forward and backward directions simultaneously. At equilibrium, the rates of the forward and reverse reactions are equal and the concentrations of reactants and products remain constant. Several examples of reversible reactions are provided. Characteristics of chemical equilibrium include the constancy of concentrations at equilibrium and the independence of the equilibrium constant from the initial concentrations. Le Chatelier's principle is introduced, which states that if a system at equilibrium experiences a change, it will shift its position to counteract that change. The effects of changing concentration, pressure, temperature, and adding a catalyst are described based on this principle. Industrial processes for maximizing yields of important chemicals
This is a lecture is a series on combustion chemical kinetics for engineers. The course topics are selections from thermodynamics and kinetics especially geared to the interests of engineers involved in combusition
1) Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal. The law of mass action states that the rate of a reaction depends on the concentrations of reactants.
2) The equilibrium constant, K, is defined as the ratio of products over reactants at equilibrium. For heterogeneous reactions involving different phases, the equilibrium constant is expressed in terms of partial pressures.
3) According to Le Chatelier's principle, if a stress is applied to a system at equilibrium, the equilibrium shifts to minimize the effect of the stress. Changes in concentration, pressure, or temperature cause the equilibrium to shift left or right to counteract the applied change.
Ch 8 - Energy, Enthalpy, and Thermochemistry.pdfCharbelRahme2
The document discusses energy, enthalpy, and thermochemistry. It defines key concepts like energy, heat, work, internal energy, enthalpy, and explains how they relate. It also discusses state functions, the first and second laws of thermodynamics, and how to calculate enthalpy changes using standard enthalpies of formation and Hess's law. Experimental methods like calorimetry are also covered to measure energy changes during chemical reactions.
Chem 2 - Chemical Kinetics III - Determining the Rate Law with the Method of ...Lumen Learning
This document discusses determining the rate law for a chemical reaction through initial rate experiments. It explains that the rate of reaction depends on reactant concentrations and describes how to find the orders of each reactant by comparing how the rate changes with concentration changes. The orders are used to define the rate law expression. Experimental data is then used to calculate the numeric rate constant.
Electrochemical Characterization of Electrocatalysts .pptxMabrook Saleh Amer
This document summarizes an electrochemistry workshop presentation on electrocatalyst characterization. It introduces common electrochemical characterization methods like cyclic voltammetry and discusses key figures of merit for evaluating electrocatalyst activity. Examples are provided of electrocatalyst development for important reactions like hydrogen evolution, oxygen evolution, and oxygen reduction. These include developing non-precious metal catalysts and improving catalyst stability and performance through methods like decreasing platinum loading or synthesizing metal phosphides and metal oxides on supports.
This document discusses electrochemistry and provides details about electrochemical cells. It contains the following key points:
1. Electrochemistry is the study of production of electricity from chemical reactions and use of electrical energy to drive non-spontaneous reactions.
2. An electrochemical cell converts chemical energy to electrical energy (galvanic/voltaic cell) or electrical energy to chemical energy (electrolytic cell).
3. A Daniell cell is a voltaic cell that generates a voltage of 1.1V from the redox reaction of zinc and copper. Measurement of electrode potentials and the Nernst equation are also discussed.
These slides are especially made to understand the postulates of quantum mechanics or chemistry better. easily simplified and at one place you will find each of relevant details about the 5 postulates. so go through it & trust me it will help you a lot if you are chemistry or a science student.
well done
This document discusses electronic spectra of metal complexes. It begins by defining quantum numbers related to electron configuration, such as L (total orbital angular momentum) and l (secondary quantum number). It then describes two main types of electronic transitions in coordination compounds: d-d transitions specific to metals, and charge-transfer transitions. The remainder of the document discusses charge-transfer transitions in more detail, defining ligand-to-metal and metal-to-ligand charge transfer, and how solvent polarity affects these transitions.
Thermochemistry is the study of energy changes that occur during chemical reactions and phase changes. Heat (q) is energy that transfers from one object to another due to a temperature difference. Exothermic processes release heat (negative q value), while endothermic processes absorb heat (positive q value). Enthalpy (H) is the heat content of a system at constant pressure, and enthalpy changes (ΔH) quantify the heat absorbed or released during chemical reactions. Thermochemical equations relate the enthalpy change to the heat of reaction. Calorimetry problems use these equations and concepts to calculate energy changes involved in physical and chemical processes.
in this lesson, I mention bond characteristics.in bond characteristics we divide it into three parts the first one is the bond length, then bond energy and the third one is about the bond angle, how and why they are formed this important and basic topic.
Catalysis is the process by which a catalyst increases the rate of a chemical reaction without undergoing any permanent chemical change. A catalyst works by providing an alternative reaction pathway or mechanism that has a lower activation energy. There are two main types of catalysis: homogeneous catalysis where the catalyst is in the same phase as the reactants, and heterogeneous catalysis where the catalyst is in a different phase. Catalysts can be classified as positive or negative depending on whether they increase or decrease the reaction rate.
Ions in solution interact electrostatically, especially at high concentrations over 0.01 M. Cations tend to be surrounded by nearby anions and vice versa. Activity coefficients account for these interactions and allow equilibrium constants for ideal solutions to be applied to non-ideal solutions. Activity coefficients can be computed from the Debye-Hückel equation based on ion characteristics like size and charge.
Thermodynamic and Kinetic aspects of metal complexes.Shivaji Burungale
The document discusses the stability and lability of coordination compounds. It defines labile and inert complexes as those where ligand exchange occurs rapidly or slowly, respectively. Lability is described as not being related to thermodynamic stability. Factors affecting stability and lability are discussed, including the metal ion properties, ligand properties, and thermodynamic considerations. Methods for determining stability constants like mole ratio, continuous variation, and slope-ratio methods are also summarized.
This document discusses key concepts in electrochemistry including:
- Electrochemistry deals with chemical and physical processes involving the production or consumption of electricity.
- Electrode potential is the potential difference that exists between a metal and its ions in solution, arising from their relative tendencies to undergo oxidation or reduction reactions.
- Standard hydrogen electrode is used as a reference electrode to measure standard electrode potentials of other half-cells.
- Standard electrode potential of a half-cell indicates its voltage when connected to the standard hydrogen electrode under standard conditions.
- Electromotive force is the difference in potential between the cathode and anode half-cells of an electrochemical cell.
The document discusses the rate of chemical reactions and factors that affect it. It defines the rate of reaction as the speed at which a chemical reaction occurs. Some reactions are very fast, like wood combustion or nuclear explosions, while others are slow, such as iron rusting. The rate depends on factors like temperature, concentration of reactants, pressure, and surface area. Increasing temperature, concentration, or pressure speeds reactions up by increasing collisions between particles. Thinner solids react faster due to their larger surface area. Catalysts also increase reaction rates without being used up in the reaction. They are important in industry and biology.
1) The chapter discusses tools for studying chemical reactions including equilibrium constants, free energy change, enthalpy, entropy, bond dissociation energy, kinetics and activation energy.
2) It then examines the chlorination of methane as a free-radical chain reaction involving initiation, propagation and termination steps.
3) Key concepts covered include how reaction rate depends on factors like temperature, activation energy and reaction order. Transition state theory and reaction energy diagrams are also explained.
The document discusses kinetics and reaction rates. It defines kinetics as the study of reaction rates, and explains that thermodynamics determines if a reaction will occur but not how fast. It then covers factors that affect reaction rates like temperature, concentration, and catalysts. The document provides examples of calculating rates from concentration data and determining rate laws from experimental results. It also introduces integrated rate laws and how to determine reaction order from different rate law equations.
1. The document discusses chemical equilibrium, including the concept that at equilibrium the forward and reverse reactions proceed at the same rate, and the amounts of reactants and products remain constant.
2. It introduces the equilibrium constant expression and explains how to write the expression for different chemical equations.
3. Le Châtelier's principle is discussed, that systems at equilibrium will shift in response to changes in conditions to counteract the effect of changes in temperature, pressure, or concentration.
This document summarizes key differences between ideal and real gases. It discusses that real gases interact with each other, taking up physical space and attracting one another, resulting in lower than expected pressures. Real gas properties depart from ideal behavior at low temperatures or high pressures. The partition function approach is used to derive equations of state for both ideal and real gases. For real gases, the virial equation accounts for intermolecular interactions through virial coefficients.
B.tech. ii engineering chemistry unit 5 A electrochemistryRai University
Arrhenius proposed the theory of electrolytic dissociation to explain the properties of electrolytic solutions. The theory states that when an electrolyte dissolves in water, it breaks up into ions - positively charged cations and negatively charged anions. This process is called ionization. Ions are constantly recombining and dissociating, reaching a state of dynamic equilibrium. The extent of ionization depends on an equilibrium constant. Strong electrolytes have a high equilibrium constant and ionize completely, while weak electrolytes have a low constant and only partially ionize.
This document discusses the key concepts of chemical equilibrium. It defines reversible reactions as those that can proceed in both the forward and backward directions simultaneously. At equilibrium, the rates of the forward and reverse reactions are equal and the concentrations of reactants and products remain constant. Several examples of reversible reactions are provided. Characteristics of chemical equilibrium include the constancy of concentrations at equilibrium and the independence of the equilibrium constant from the initial concentrations. Le Chatelier's principle is introduced, which states that if a system at equilibrium experiences a change, it will shift its position to counteract that change. The effects of changing concentration, pressure, temperature, and adding a catalyst are described based on this principle. Industrial processes for maximizing yields of important chemicals
This is a lecture is a series on combustion chemical kinetics for engineers. The course topics are selections from thermodynamics and kinetics especially geared to the interests of engineers involved in combusition
1) Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal. The law of mass action states that the rate of a reaction depends on the concentrations of reactants.
2) The equilibrium constant, K, is defined as the ratio of products over reactants at equilibrium. For heterogeneous reactions involving different phases, the equilibrium constant is expressed in terms of partial pressures.
3) According to Le Chatelier's principle, if a stress is applied to a system at equilibrium, the equilibrium shifts to minimize the effect of the stress. Changes in concentration, pressure, or temperature cause the equilibrium to shift left or right to counteract the applied change.
Ch 8 - Energy, Enthalpy, and Thermochemistry.pdfCharbelRahme2
The document discusses energy, enthalpy, and thermochemistry. It defines key concepts like energy, heat, work, internal energy, enthalpy, and explains how they relate. It also discusses state functions, the first and second laws of thermodynamics, and how to calculate enthalpy changes using standard enthalpies of formation and Hess's law. Experimental methods like calorimetry are also covered to measure energy changes during chemical reactions.
Chem 2 - Chemical Kinetics III - Determining the Rate Law with the Method of ...Lumen Learning
This document discusses determining the rate law for a chemical reaction through initial rate experiments. It explains that the rate of reaction depends on reactant concentrations and describes how to find the orders of each reactant by comparing how the rate changes with concentration changes. The orders are used to define the rate law expression. Experimental data is then used to calculate the numeric rate constant.
Electrochemical Characterization of Electrocatalysts .pptxMabrook Saleh Amer
This document summarizes an electrochemistry workshop presentation on electrocatalyst characterization. It introduces common electrochemical characterization methods like cyclic voltammetry and discusses key figures of merit for evaluating electrocatalyst activity. Examples are provided of electrocatalyst development for important reactions like hydrogen evolution, oxygen evolution, and oxygen reduction. These include developing non-precious metal catalysts and improving catalyst stability and performance through methods like decreasing platinum loading or synthesizing metal phosphides and metal oxides on supports.
This document discusses electrochemistry and provides details about electrochemical cells. It contains the following key points:
1. Electrochemistry is the study of production of electricity from chemical reactions and use of electrical energy to drive non-spontaneous reactions.
2. An electrochemical cell converts chemical energy to electrical energy (galvanic/voltaic cell) or electrical energy to chemical energy (electrolytic cell).
3. A Daniell cell is a voltaic cell that generates a voltage of 1.1V from the redox reaction of zinc and copper. Measurement of electrode potentials and the Nernst equation are also discussed.
These slides are especially made to understand the postulates of quantum mechanics or chemistry better. easily simplified and at one place you will find each of relevant details about the 5 postulates. so go through it & trust me it will help you a lot if you are chemistry or a science student.
well done
This document discusses electronic spectra of metal complexes. It begins by defining quantum numbers related to electron configuration, such as L (total orbital angular momentum) and l (secondary quantum number). It then describes two main types of electronic transitions in coordination compounds: d-d transitions specific to metals, and charge-transfer transitions. The remainder of the document discusses charge-transfer transitions in more detail, defining ligand-to-metal and metal-to-ligand charge transfer, and how solvent polarity affects these transitions.
Thermochemistry is the study of energy changes that occur during chemical reactions and phase changes. Heat (q) is energy that transfers from one object to another due to a temperature difference. Exothermic processes release heat (negative q value), while endothermic processes absorb heat (positive q value). Enthalpy (H) is the heat content of a system at constant pressure, and enthalpy changes (ΔH) quantify the heat absorbed or released during chemical reactions. Thermochemical equations relate the enthalpy change to the heat of reaction. Calorimetry problems use these equations and concepts to calculate energy changes involved in physical and chemical processes.
in this lesson, I mention bond characteristics.in bond characteristics we divide it into three parts the first one is the bond length, then bond energy and the third one is about the bond angle, how and why they are formed this important and basic topic.
Catalysis is the process by which a catalyst increases the rate of a chemical reaction without undergoing any permanent chemical change. A catalyst works by providing an alternative reaction pathway or mechanism that has a lower activation energy. There are two main types of catalysis: homogeneous catalysis where the catalyst is in the same phase as the reactants, and heterogeneous catalysis where the catalyst is in a different phase. Catalysts can be classified as positive or negative depending on whether they increase or decrease the reaction rate.
Ions in solution interact electrostatically, especially at high concentrations over 0.01 M. Cations tend to be surrounded by nearby anions and vice versa. Activity coefficients account for these interactions and allow equilibrium constants for ideal solutions to be applied to non-ideal solutions. Activity coefficients can be computed from the Debye-Hückel equation based on ion characteristics like size and charge.
Thermodynamic and Kinetic aspects of metal complexes.Shivaji Burungale
The document discusses the stability and lability of coordination compounds. It defines labile and inert complexes as those where ligand exchange occurs rapidly or slowly, respectively. Lability is described as not being related to thermodynamic stability. Factors affecting stability and lability are discussed, including the metal ion properties, ligand properties, and thermodynamic considerations. Methods for determining stability constants like mole ratio, continuous variation, and slope-ratio methods are also summarized.
This document discusses key concepts in electrochemistry including:
- Electrochemistry deals with chemical and physical processes involving the production or consumption of electricity.
- Electrode potential is the potential difference that exists between a metal and its ions in solution, arising from their relative tendencies to undergo oxidation or reduction reactions.
- Standard hydrogen electrode is used as a reference electrode to measure standard electrode potentials of other half-cells.
- Standard electrode potential of a half-cell indicates its voltage when connected to the standard hydrogen electrode under standard conditions.
- Electromotive force is the difference in potential between the cathode and anode half-cells of an electrochemical cell.
The document discusses the rate of chemical reactions and factors that affect it. It defines the rate of reaction as the speed at which a chemical reaction occurs. Some reactions are very fast, like wood combustion or nuclear explosions, while others are slow, such as iron rusting. The rate depends on factors like temperature, concentration of reactants, pressure, and surface area. Increasing temperature, concentration, or pressure speeds reactions up by increasing collisions between particles. Thinner solids react faster due to their larger surface area. Catalysts also increase reaction rates without being used up in the reaction. They are important in industry and biology.
1) The chapter discusses tools for studying chemical reactions including equilibrium constants, free energy change, enthalpy, entropy, bond dissociation energy, kinetics and activation energy.
2) It then examines the chlorination of methane as a free-radical chain reaction involving initiation, propagation and termination steps.
3) Key concepts covered include how reaction rate depends on factors like temperature, activation energy and reaction order. Transition state theory and reaction energy diagrams are also explained.
The document discusses kinetics and reaction rates. It defines kinetics as the study of reaction rates, and explains that thermodynamics determines if a reaction will occur but not how fast. It then covers factors that affect reaction rates like temperature, concentration, and catalysts. The document provides examples of calculating rates from concentration data and determining rate laws from experimental results. It also introduces integrated rate laws and how to determine reaction order from different rate law equations.
B sc_I_General chemistry U-III(B)Molecular formula and empirical formula Rai University
This document provides information on qualitative analysis of organic compounds. It discusses preliminary tests like physical characteristics, solubility tests, and group tests to deduce the functional groups present in unknown organic samples. It also describes instrumental techniques like chromatography, spectroscopy, and atomic spectroscopy that can be used for further analysis and identification. Quantitative elemental analysis methods are introduced for determining the percentages of carbon, hydrogen, and other elements present in organic compounds.
Kinetics is the study of reaction rates and mechanisms. There are four main factors that affect reaction rates: the nature of reactants, temperature, catalysts, and concentration. The rate of a reaction is measured based on how quickly the concentration of reactants decreases or products increases over time. As reactions proceed, the rate generally decreases as reactant concentrations decrease until the reaction reaches equilibrium or the reactants are used up.
The document discusses chemical kinetics and reaction rates. It defines zero, first, and second-order reactions based on how the reaction rate depends on reactant concentrations. For zero-order reactions, the rate is independent of concentration. For first-order reactions, the rate is directly proportional to one reactant concentration. For second-order reactions, the rate is proportional to the product of two reactant concentrations or the square of one reactant concentration. It also presents the integrated rate laws and methods for determining the order of a reaction from experimental data by plotting concentrations versus time.
This is a lecture is a series on combustion chemical kinetics for engineers. The course topics are selections from thermodynamics and kinetics especially geared to the interests of engineers involved in combusition
The document discusses factors that affect the rate of chemical reactions, including concentration, temperature, surface area, and catalysts. It explains collision theory and activation energy. Exothermic reactions release heat while endothermic reactions absorb heat. Le Chatelier's principle states that chemical equilibriums shift to counteract changes in concentration, temperature, pressure or addition of reactants/products.
The document summarizes key concepts in reaction kinetics and chemical equilibrium. It discusses factors that affect reaction rates, reaction orders, rate laws, and progress curves. It also covers the concepts of chemical equilibrium, equilibrium constants, and factors that can shift equilibrium. Finally, it introduces concepts of energy in chemical reactions including enthalpy, entropy, the first and second laws of thermodynamics, and Gibbs free energy as the driving force for spontaneous reactions.
This document discusses key thermodynamic concepts related to combustion processes, including:
1) Heat of combustion, flame temperature, enthalpy of combustion systems, and equilibrium constants of combustion reactions are the major thermodynamic functions that influence fuel utilization.
2) Heat of combustion represents the potential heat of a fuel and can be used to calculate calorific value. Enthalpy is the heat content of a system at constant pressure.
3) Flame temperature depends on the fuel-oxidant mixture and ranges from theoretical to actual temperatures. The maximum adiabatic flame temperature occurs at slightly excess stoichiometry.
This document provides information about chemical kinetics and reaction rates. It defines chemical kinetics and discusses the two conditions needed for reactions to occur: contact between particles and enough energy for activation. Reaction rate depends on collision frequency and efficiency of reactants and is influenced by nature of reactants, surface area, temperature, concentration, and presence of a catalyst. It also defines catalysts and different types. The relationship between reaction rate and concentration is given by the rate law, and order of a reaction relates to its rate law. Rate laws are written for different reaction examples and mechanisms.
A complete introduction to all things chemical kinetics designed specifically for non-chemists to understand. Fair warning: The presentation is very rigorous in its mathematical treatment, which is makes it a useful reference for looking up equations, but this can unfortunately make it less polished and flowing then a typical presentation. I tried my best to spell everything out clearly, but despite my best efforts it's still pretty dense.
Chem II - Ideal Gas Law (Liquids and Solids)Lumen Learning
The document discusses the ideal gas law, which relates the pressure, volume, temperature, and number of moles of a gas. It describes how the ideal gas law was developed by combining Boyle's law, Charles's law, and Avogadro's law. The ideal gas law equation is PV=nRT, where P is pressure, V is volume, n is moles of gas, R is the gas constant, and T is temperature in Kelvin. An example problem demonstrates how to use the ideal gas law to calculate the moles of gas produced in a chemical reaction.
This document provides an overview of voltammetry and potentiometry techniques. It discusses the history and development of voltammetry, which involves measuring current as a function of applied potential. Common types of voltammetry include linear sweep, cyclic, and stripping voltammetry. The document also describes the basic components of a voltammetry system, including the working, reference, and counter electrodes. Finally, it provides a brief introduction to potentiometry and its applications in titration and measuring concentration, activity, and pH.
This document provides an overview of chemical kinetics and reaction rates. It discusses topics such as reaction rate, rate laws, reaction orders, rate constants, factors that affect reaction rates like temperature, catalysts, and enzyme kinetics. Specific examples are also provided to illustrate concepts like first-order and second-order reactions, reaction mechanisms, and industrial catalytic processes like the Haber process and catalytic converters.
This document summarizes an experiment that investigates the relationship between water pressure and flow rate. Water was flowed from a container through a small hole for timed intervals at varying water heights, and the resulting flow rate was measured. The data showed a proportional relationship between pressure difference and squared flow rate, supporting Bernoulli's equation. The slope of the line of best fit remained constant, as expected for this experimental setup. Limitations included the assumption of zero kinetic energy and challenges closing the hole precisely.
The document discusses reaction rates and kinetics. It defines factors that affect reaction rates such as concentration of reactants, physical state, temperature, and catalysts. It also describes methods for determining reaction rates by measuring changes in concentration over time. Rate laws relate the rate of reaction to concentrations of reactants through rate constants and reaction orders. Integrated rate laws can be used to determine concentrations of reactants over time for reactions of different orders.
It shows the basic facts of catalyst along with its importance in industry along with its long last milestone,its characteristics & application in industry its reaction process and preparation of a solid catalyst.
Chemical kinetics is the study of reaction rates and mechanisms. Key aspects include determining how factors like temperature, pressure, catalysts and light influence reaction rates. Reaction rates are determined by monitoring changes in reactant or product concentration over time. The rate of a reaction depends on the concentrations of reactants and can be modeled using a rate law. Common reaction orders include zero-order, first-order and second-order reactions, which have different relationships between rate and concentration. Catalysts increase reaction rates by providing an alternative reaction pathway with a lower activation energy.
Chemical kinetics: the study of how fast chemical reactions occur.Specifically:
Rates of consumption of reactants and formation of products.
Response of chemical rates to changes in rxn conditions.
Identification of steps through which rxn takes place.
Reasons for study
Prediction of how quickly a rxn approaches equilibrium.
Understanding or elucidation of rxn mechanisms.
Chemical kinetics deals with the rates of chemical reactions and the factors that affect reaction rates. The rate of a reaction depends on variables like concentration, temperature, and presence of a catalyst. A reaction's rate law expresses how its rate depends on the concentrations of reactants. Integrated rate laws relate concentration over time for different reaction orders. Pseudo-first-order reactions occur when one reactant is in excess. The Arrhenius equation describes how temperature affects a reaction's rate constant.
chemical kinetics of chemical reaction (1).pptxFatmaMoustafa6
Chemical kinetics is the branch of chemistry that studies reaction rates and mechanisms. It involves determining how external factors like temperature, pressure, catalysts, and light influence reaction rates. The main goals are to determine reaction rates, study reaction mechanisms, and explore relationships between material structure and reactivity. Reaction rates can be measured by monitoring changes in reactant or product concentration over time. Factors that affect reaction rates include the nature of reactants, concentration of reactants, temperature, catalysts, and the reaction medium. The rate of a reaction is expressed using a rate law, which relates reaction rate to reactant concentrations through rate constants and orders of reaction.
- The document discusses concepts related to chemical kinetics including reaction rates, rate laws, reaction mechanisms, and reaction orders.
- Key concepts covered include determining rate laws through experimental methods, distinguishing between differential and integrated rate laws, and characteristics of reactions that are zero order, first order, or second order.
- Examples are provided to illustrate determining the order of reactions and calculating rate constants from experimental data using integrated rate laws.
This document discusses chemical kinetics and reaction rates. It begins with an introduction to chemical kinetics and defines reaction rate. It then discusses factors that affect reaction rates such as nature of reactants, concentration, temperature, and catalysts. It describes different types of reaction rates and how they are measured. The document also covers rate laws, determining rate orders experimentally, and integrated and differential rate equations for zero, first, and second order reactions. It concludes with an overview of rate theories including the Arrhenius equation.
This document discusses key concepts in chemical kinetics including:
- Rate laws describe how reaction rates depend on reactant concentrations. Rate laws can be determined experimentally from initial rate data.
- Reactions can be zero-order, first-order, or second-order depending on how the rate depends on reactant concentrations.
- Integrated rate laws relate reactant concentrations to time and rate constants. Graphical methods using plots of concentration or transformed concentration vs. time can determine reaction order and rate constants.
1. Chemical kinetics is the branch of chemistry that deals with the rates of chemical reactions and their mechanisms.
2. The rate of a reaction depends on factors like concentration, temperature, pressure, and presence of catalysts.
3. The rate of a reaction can be defined as the change in concentration of a reactant or product per unit time and can be expressed in terms of the rates of appearance or disappearance of reactants and products.
The branch of chemistry, which deals with the study of reaction rates and their mechanisms, called chemical kinetics.
Thermodynamics tells only about the feasibility of a reaction whereas chemical kinetics tells about the rate of a reaction.
For example, thermodynamic data indicate that diamond shall convert to graphite but in reality the conversion rate is so slow that the change is not perceptible at all.
The document discusses factors that affect chemical reaction rates and concepts related to reaction kinetics such as reaction order, rate equations, and rate determining steps. It provides examples of determining the order of reactions from experimental rate data and explains that the slowest step of a reaction, usually the first step, is the rate determining step whose order defines the overall reaction order.
This document is a chapter summary for a chemistry textbook on reaction rates. It defines reaction rates and discusses how rates depend on concentration, temperature, and catalysts. It also covers experimental determination of rates, rate laws, reaction mechanisms, and the effects of temperature. Key equations discussed include the rate law, integrated rate laws for first and second order reactions, the Arrhenius equation relating reaction rate and temperature, and transition state theory to explain the activated complex.
This document provides background information on reaction rates and mechanisms. It discusses how factors like reactant concentrations, temperature, catalysts, and surface area can influence reaction rates. It also defines concepts like the rate law, rate constant, reaction order, energy of activation, and Arrhenius equation. Methods for determining reaction order are described, including by varying reactant concentrations and analyzing integrated rate expressions for zero, first, and second order reactions. The effects of temperature on reaction rates are also addressed through the Arrhenius equation.
This document discusses chemical kinetics and rate of reactions. It defines chemical kinetics as the study of reaction rates and their mechanisms. It then discusses factors that influence reaction rates such as concentration, temperature, pressure, catalysts and more. It defines rate of reaction and discusses how to determine rates. It introduces reaction orders such as zero order, first order and second order reactions. Examples of each type of reaction order are provided along with the appropriate rate equations. Pseudo-first order reactions are also discussed.
kinetics of stability Molecular pharmaceuticsMittalGandhi
This document discusses kinetics of stability and reaction order. It defines key terms like rate, order of reaction, and molecularity. The main types of reaction order discussed are zero order, first order, pseudo first order, and second order. Graphs and equations to determine the rate constant and half-life are provided for each order. Methods for determining the experimental order of a reaction are outlined. Factors that can influence the reaction rate are also summarized. Tables listing the key equations for zero, first, and second order kinetics are included.
Here are the steps to determine the order of the reaction:
1) Plot [X] vs time on a graph. You will get a straight line through the origin, indicating the reaction is first order.
2) Take the log of both sides of the rate law equation:
Rate = k[X]
Log(Rate) = Log(k[X])
3) Plot log(Rate) vs log([X]). You will get a straight line with a slope of 1, confirming the reaction is first order.
Therefore, based on the experimental data and analysis, this reaction is first order with respect to X.
Chemical kinetics is the study of reaction rates and mechanisms. It involves determining:
- Reaction orders and rate laws from initial rates or graphical methods.
- Rate constants and activation energies.
- Elementary reaction steps and overall mechanisms.
The rate of a reaction depends on factors like temperature, concentration, and the presence of catalysts. Reaction rates are quantified by rate laws, which relate the rate to concentrations of reactants raised to their order of reaction. Graphical methods can be used to determine reaction orders from concentration-time data.
This document discusses reaction rates and kinetics concepts including:
- Instantaneous reaction rates can be calculated from the slope of concentration-time graphs at specific points.
- Reaction orders and rate laws can be determined experimentally using methods like the initial rate method or integrated rate law method.
- First-order reactions follow the integrated rate law that the natural log of the concentration is linear with time. Second-order and zero-order reactions also have defining rate laws and kinetics equations.
1) Reaction rates depend on temperature and concentration according to the Arrhenius equation, with the rate constant k increasing exponentially with temperature. This is because higher temperatures cause molecules to move faster and collide more frequently, increasing the probability of reaction.
2) Kinetics can be first order, second order, or other orders depending on how the reaction rate depends on the concentrations of reactants. First order reactions follow exponential decay, while second order reactions have a more complex concentration dependence.
3) Half-lives characterize how long it takes for the concentration of a reactant to reduce to half its initial value, and can be used to determine reaction orders and rate constants from experimental data.
This document discusses chemical kinetics and reaction rates. It begins by defining kinetics as the study of reaction rates and discusses how kinetics provides information about reaction mechanisms. It then describes several factors that affect reaction rates, including physical state of reactants, concentration of reactants, temperature, and presence of catalysts. The document goes on to explain how reaction rates are determined by measuring changes in concentration over time. It also discusses how reaction rates relate to stoichiometry and how reaction rates are affected by changes in concentration. Integrated rate laws for first-order and second-order reactions are presented, along with examples of using these rate laws to determine reaction order from experimental data.
Similar to Tro chapter 13 kinetics spring 2015 (1) (20)
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2. Definitions/Terminology
Kinetics is the study of rates of chemical
reactions and the mechanisms (pathways) by
which they occur.
3
• The rate is the increase in concentration of a
product per unit time or decrease in concentration of
a reactant per unit time.
• Rates have units of +/-[conc/time] (e.g., Molar/sec).
• A mechanism is the series of molecular steps by
which a reaction occurs.
3. Reaction Rates
Thermodynamics determines whether or not a
reaction can occur.
Kinetics describes how quickly a reaction occurs.
Some reactions are thermodynamically favorable (G0
< 0), but are kinetically so slow as to be imperceptible
at room temperature:
4
CH4(g) + O2(g) CO2(g) + H2O(g) G0 = -800 kJ/mol
H2(g) + O2(g) H2O(g) G0 = -230 kJ/mol
• These reactions do not occur at room temperature and
pressure without activation (e.g., a spark or a catalyst).
4. Reaction Rates
5
• Rates define the velocity at which reactants
disappear or products appear
• Rates have units of [concentration/time].
5. Reaction Rates
6
• in many reactions, the coefficients of reactants and products in
the balanced equation are not all the same, for example:
H2 (g) + I2 (g) 2 HI(g)
For every 1 mole of H2 consumed, 1 mole of I2 will also be
consumed, and 2 moles of HI will be produced.
The rate of disappearance of [H2] and [I2] will, therefore, be
½ the rate of appearance of [HI].
• The rate of the overall reaction is, thus, the change in the
concentration of each substance multiplied by
1/[coefficient in the balanced equation]:
Rate = -
D[H2 ]
Dt
= -
D[I2 ]
Dt
= +
1
2
æ
è
ç
ö
ø
÷
D[HI]
Dt
6. Reaction Rates
7
For [H2], the
instantaneous
rate at 50 s is:
For [HI], the
instantaneous
rate at 50 s is:
s
M0.0070Rate
s40
M28.0Rate
s
M0.0070Rate
s40
M56.0
2
1Rate
7. Reaction Rates
Mathematically, the rate of a generic reaction:
aA + bB --> cC + dD can be written as:
For X = A, B, C or D, [X] = [X]t - [X]0, where
[X]t is concentration at time, t after the start of the reaction, and [X]0 is
the initial concentration at the beginning of the reaction, t = 0.
Minus sign indicates that the reactants concentrations decrease with
time.
8
• For example, for a reaction 2A B:
Rate = -(1/2)[A]/t = [B]/t
• Rates are often approximated as instantaneous
rates, and the notation of calculus is used:
Rate = -1/2d[A]/dt = d[B]/dt
The rate equation
8. Reaction Rates
The rate of a simple one-reactant, one-step reaction
is directly proportional to the concentration of the
reacting substance:
Rate has units of conc/unit time (e.g., Ms-1)
[A] is the concentration of A, e.g., in molarity (M).
The proportionality constant, k is called the rate constant.
For this simple expression, k must have units of inverse
time (e.g., s-1).
Rate constants are always positive numerical quantities. 9
A B + C
rate [A] or rate = k[A] the rate law expression
(different from the rate eqn!)
9. Reaction Rates
For the simple expression, rate = k[A]:
If the initial concentration of A is doubled, the
initial rate of the reaction is doubled.
If the initial concentration of A is halved, the
initial rate of reaction is cut in half.
The initial rate is, thus, directly proportional to
the initial [A].
10
10. Reaction Rates
If more than one reactant molecule appears in
the chemical equation for a one-step reaction:
2A B + C
The experimentally determined rate law is:
rate = k[A]2
11
• This relationship means that, If [A] is doubled,
the rate increases by a factor of 4 (= 22).
• If initial [A] is halved, the initial rate decreases
by a factor of 4 (=(1/2)2).
11. Reaction Rates
Rate = k[X] or rate = k[X]2 are examples of a rate
law.
12
Rate laws can only be determined experimentally.
The rate law cannot be determined from the balanced
chemical equation because:
most chemical reactions are not one-step reactions.
12. Reaction Rates
The order of a reaction expresses:
the dependence on the concentrations of each reactant in the rate
law (= the exponent for each concentration).
the sum of the orders for each reactant.
13
• For the reaction:
N2O5(g) 2NO2(g) + 1/2O2(g)
The experimentally determined rate law is:
rate = k[N2O5]
• The reaction is said to be first order in N2O5 and first
order overall.
13. Reaction Rates
First order reactions are dependent on the
concentration of only a single reactant.
First order reactions are common for many
chemical reactions and all simple radioactive
decays.
14
Two examples of first order reactions:
2N2O5(g) 2 N2O4(g) + O2(g)
238U 234Th + 4He
Rate = k[N2O5]
Rate = k[238U]
14. Reaction Rates
15
• For the reaction:
2NO(g) + O2(g) 2NO2(g)
• The experimentally determined rate law is:
rate = k[NO]2[O2]
• This reaction is second order in NO, first
order in O2 and third order overall.
15. Reaction Rates
16
• For the reaction:
(CH3)3CBr(aq) + OH-
(aq) (CH3)3COH(aq) + Br -
(aq)
The experimentally determined rate law is:
rate = k[(CH3)3CBr]
• This reaction is first order in (CH3)3CBr and
zero order in OH- and first order overall.
16. Concentrations of Reactants:
The Rate-Law Expression
Example 1: The following rate data were obtained at 25oC for
the reaction: 2A(g) + B(g) 3C(g)
What are the rate law and rate constant for this reaction?
Experiment
Number
Initial [A]
(M)
Initial [B]
(M)
Initial rate
(M/s)
1 0.10 0.10 2.0 x 10-4
2 0.20 0.10 4.0 x 10-4
3 0.10 0.20 2.0 x 10-4
22
17. Concentrations of Reactants:
The Rate-Law Expression
23
• Compare experiments 1 and 3.
• When [A] is held constant and [B] is doubled, the
rate does not change.
• The reaction is zeroth order in [B].
• The rate law can be written generically as:
Rate = k[A]x[B]y
• The rate law requires experimental
determination of the exponents, x and y.
• Therefore, y = 0 and the rate law reduces to:
Rate = k[A]x
18. Concentrations of Reactants:
The Rate-Law Expression
24
• Next compare experiments 1 and 2.
• If [A] is doubled, the rate increases by a
factor of 2.
• Therefore, (2)x = 2 and x = 1.
• The rate law reduces to:
rate = k[A]
• The reaction is first order in [A] and first
order overall.
20. Concentrations of Reactants:
The Rate-Law Expression
Example 2: The following data were obtained for the
reaction:
2 A(g) + B(g) + 2 C(g) 3 D(g) + 2 E(g)
Experiment
Initial [A]
(M)
Initial [B]
(M)
Initial [C]
(M)
Initial rate
(M/s)
1 0.20 0.10 0.10 2.0 x 10-4
2 0.20 0.30 0.20 6.0 x 10-4
3 0.20 0.10 0.30 2.0 x 10-4
4 0.60 0.30 0.40 1.8 x 10-3
26
From these initial rate data we can determine the rate-
law expression and the rate constant for this reaction.
21. Concentrations of Reactants:
The Rate-Law Expression
27
• Comparing experiments 1 and 3:
• [A] and [B] are held constant, but [C] is
increased by 3-times.
• The rate does not change.
• Therefore, the exponent, z = 0, and the rate
law simplifies to:
Rate = k[A]x[B]y
Rate = k[A]x[B]y[C]z
22. Concentrations of Reactants:
The Rate-Law Expression
28
• Comparing experiments 1 and 2:
• [A] is held constant, but [B] is increased by 3-
times.
• The rate also increases by 3-times.
• Therefore, y = 1, and the rate law further
simplifies to:
Rate = k[A]x[B]
23. Concentrations of Reactants:
The Rate-Law Expression
29
• Comparing experiments 2 and 4:
• [B] is held constant, but [A] is increased by 3-
times.
• The rate also increases by 3-times.
• The reaction is, thus, first order in A, first
order in B and second order overall.
• Therefore, x = 1, and the rate law further
simplifies to:
Rate = k[A][B]
25. Concentrations of Reactants:
The Rate-Law Expression
Example 3: A reaction between compounds A
and B is determined to be first order in A, first
order in B, and second order overall. From the
information given below, fill in the blanks.
Experiment
Initial Rate
(M/s)
Initial [A]
(M)
Initial [B]
(M)
1 4.0 x 10-3 0.20 0.050
2 1.6 x 10-2 ? 0.050
3 3.2 x 10-2 0.40 ? 31
Rate = k[A][B]
29. Concentration vs. Time:
Integrated Rate Laws
The integrated rate equation relates time and
concentration for chemical reactions.
The integrated rate equation can be used to predict the
amount of product that is produced in a given amount of
time.
Provides an alternative to the initial rate method for
obtaining rate constants.
Initially we will look at the integrated rate equation
for first order reactions.
These reactions are 1st order in one reactant and 1st order
overall.
For a generic first order reaction with one reactant:
A products, rate = k[A]
35
or -d[A]/dt = k[A]
30. Concentration vs. Time:
Integrated Rate Laws
where:
t = time elapsed since beginning of reaction.
[A]0= mol/L of A at t = 0 (sometimes called initial conc., [A]i)
[A] = mol/L of A at time t ([A] sometimes written as [A]t)
k = rate constant
Take the integral of this rate law from the beginning (t =
0, [A]0) to [A] at any later reaction time, t.
The integrated rate equation for a 1st order reaction:
36
ln([A]/[A]0) = -kt,
MEMORIZE and be able to
interconvert these equations!
or
[A] = [A]0e-kt
ln([A]0/[A]) = kt,
y = mx + bor ln[A] = -kt + ln[A]0
or ln[A]0 - ln[A] = kt,
31. 2N2O5(g) 2 N2O4(g) + O2(g)
Concentration vs. Time:
Integrated Rate Laws
37
• The reaction is found experimentally to be first order in
N2O5 and first order overall.
• Rate = k[N2O5]
plots of the first order integrated rate equations
for this reaction…
33. Concentration vs. Time:
Integrated Rate Laws
Define the half-time, t1/2, of a reaction as the time
required for half of the reactant to be consumed, i.e., the
time, t, at which [A]=1/2[A]0, then at t1/2, [A]0/[A] = 2.
39
ln 2 = 0.693 = kt1/2
• So, the half-time for any overall first order reaction is
given by:
t1/2 = 0.693/k MEMORIZE (or take ln 2)
• The half-time is independent of [A]0!
[A]0/[A] = 2 = ekt at t = t1/2
34. Concentration vs. Time:
Integrated Rate Laws
Example: Cyclopropane decomposes to propene
according to the following equation:
The reaction is found to be first order in
cyclopropane with k = 9.2 s-1 at 1000 0C. What is
the half time for disappearance of cyclopropane
at 1000 0C?
40
t1/2 = 0.693/9.2 s-1 = 0.075 s
35. Concentration vs. Time:
Integrated Rate Laws
Example: Refer to the previous example.
How much of a 3.0 g sample of cyclopropane
remains after 0.50 seconds?
The integrated rate laws can be use for any unit
that’s proportional to moles or concentration.
This example uses grams rather than mol/L.
41
36. Concentration vs. Time:
Integrated Rate Laws
42
ln [A]0 - ln [A] = kt
[A]0 = 3, t = 0.50 s, k = 9.2 s-1, solve for [A]:
ln 3 - ln [A] = 9.2 s-1(0.5 s)
1.1 - ln [A] = 4.6
ln [A] = -(4.6-1.1) = -3.5
[A] = e-3.5 = 0.03 g ~ 1% remains
Use the integrated first order rate equation:
37. Concentration vs. Time:
Integrated Rate Laws
Example: The half-time for the first order reaction:
CS2(g) CS(g) + S(g)
is 688 hours at 1000 0C. Calculate the rate constant,
k, at 1000 0C and the amount of a 3.0 g sample of
CS2 that remains after 48 hours.
43
k = 0.693/688 hr = 0.00101 hr-1.
39. Concentration vs. Time: The
Integrated Rate Laws
For reactions that are second order with respect to
a single reactant, A, and second order overall, with
rate = k[A]2 = (1/a)(d[A]/dt), the integrated rate law
is:
Where a = stoichiometric coefficient of A in the balanced
overall equation.
If A P, then a = 1.
If 2A P, then a = 2
45
MEMORIZE!
• An alternative form of the 2nd order integrated
rate law:
[A] = [A]0/(1 + akt[A]0).
40. Concentration vs. Time:
Integrated Rate Laws
Half-time, t1/2, for second order reactions:
Use the second order integrated rate-law as a starting
point.
At the half-time, t1/2 is [A] = 1/2[A]0.
46
41. Concentration vs. Time: The
Integrated Rate Equations
If we solve for t1/2:
So the half-time of a second order reaction depends
on [A]0.
Less [A]0, longer half-time
Half-time increases as the reaction proceeds (unlike an
overall first order reaction, where t remains constant).
47
42. Concentration vs. Time:
Integrated Rate Laws
Example: Acetaldehyde, CH3CHO, undergoes gas
phase thermal decomposition to methane and carbon
monoxide.
The rate law was found to be:
rate = k[CH3CHO]2, and k = 2.0 x 10-2 M-1hr-1 at 527 oC.
(a) What is the half-time (t1/2) for disappearance of
CH3CHO if 0.10 mole is injected into a 1.0 L vessel at
527 oC? 48
44. Concentration vs. Time: The
Integrated Rate Laws
(b) How many moles of CH3CHO remain after 200
hours in the 1 L vessel?
50
45. Concentration vs. Time: The
Integrated Rate Laws
(c) What percent of the CH3CHO remains
after 200 hours in the 1 L vessel?
51
46. Concentration vs. Time:
Integrated Rate Laws
52
(a) For the same reaction as the previous example, what is
the half-time (t1/2) for disappearance of CH3CHO if 0.10
mole is injected into a 10.0 L vessel at 527 oC?
● So, if [A]0 decreases by 10-times, t1/2 must increase by 10-
times,
● 10 x 5.0 x 102 hr = 5.0 x 103 hr
● Note that the vessel size is increased by a factor of 10
compared to the previous example, which decreases the
initial concentration by a factor of 10.
● Recall the equation for half time of this 2nd order
reaction (in which the coefficient, a = 1) is:
t1/2 = 1/k[A]0)
47. For an overall zeroth order reaction the rate
expression is:
57
Concentration vs. Time: The
Integrated Rate Laws
Which gives the zeroth order integrated rate
equation:
49. Using Integrated Rate Laws to
Determine Reaction Order
59
y = mx + b
ln [A] = -kt + ln [A]0
Plots of linear form of the integrated rate
equations can determine the reaction order and
rate constant.
1st order:
zeroth order: [A] = -kt + [A]0
2nd order:
50. Using Integrated Rate Laws to
Determine Reaction Order
Example: For the thermal decomposition of ethyl bromide:
C2H5Br(g) C2H4(g) + HBr(g)
60
Time (min) 0 1 2 3 4 5
[C2H5Br] 1.00 0.82 0.67 0.55 0.45 0.37
ln [C2H5Br] 0.00 -0.20 -0.40 -0.60 -0.80 -0.99
1/[C2H5Br] 1.0 1.2 1.5 1.8 2.2 2.7
51. Using Rate Integrated Laws to
Determine Reaction Order
Make an x,y plot for each of the three sets of
values in the table:
1. [C2H5Br] (y-axis) vs. time (x-axis)
If the plot is linear, then the reaction is zeroth order
with respect to [C2H5Br].
61
2. ln [C2H5Br] (y-axis) vs. time (x-axis)
If the plot is linear, then the reaction is first order
with respect to [C2H5Br].
3. 1/ [C2H5Br] (y-axis) vs. time (x-axis)
If the plot is linear, then the reaction is second
order with respect to [C2H5Br].
52. Using Integrated Rate Laws to
Determine Reaction Order
Plot of [C2H5Br] versus time.
62Is it linear?
[C2H5Br]
53. Using Rate Laws to Determine
Reaction Order
Plot of ln [C2H5Br] versus time.
Is it linear?
ln [C2H5Br]
54. Using Integrated Rate Laws
to Determine Reaction Order
Plot of 1/[C2H5Br] versus time.
64Is it linear?
1/[C2H5Br]
55. Using Integrated Rate Laws
to Determine Reaction Order
ln[C2H5Br] vs. time is the only linear plot.
The reaction is, therefore, first order with respect to
[C2H5Br] and first order overall.
From the integrated rate law for a first order reaction:
ln[A] = -kt + ln[A]0, slope = -k.
66
● Slope = y2 - y1/x2 - x1.
• So k = 0.2 min-1
56. Using Rate Equations to
Determine Reaction Order
Example: Concentration-versus-time data for the reaction:
2NO2(g) --> 2NO(g) + O2(g),
are given in the table below. Plot each of these time
functions to determine the rate of the reaction and the value
of the rate constant.
67
Time(min) 0 1 2 3 4 5
[NO2] 1.0 0.53 0.36 0.27 0.22 0.18
ln [NO2] 0.0 -0.63 -1.0 -1.3 -1.5 -1.7
1/[NO2] 1.0 1.9 2.8 3.7 4.6 5.5
57. Using Rate Equations to
Determine Reaction Order
Plot of [NO2] versus time.
69
Is it linear?
[NO2]
58. Using Rate Equations to
Determine Reaction Order
Plot of ln [NO2] versus time.
70
Is it linear?
ln [NO2]
59. Using Rate Equations to
Determine Reaction Order
Plot of 1/[NO2] versus time.
71Is it linear?
1/[NO2]
60. Using Rate Equations to
Determine Reaction Order
1/[NO2] vs. time is the only linear plot.
The reaction is, therefore, second order in
[NO2] and second order overall.
72
● Determine the value of the rate constant from
the slope of the line.
61. Using Rate Equations to
Determine Reaction Order
From the equation for a second order reaction
the slope = a k
In this reaction a = 2.
73
Slope = 0.90 = 2 k
k = 0.45 M-1min-1
62. Collision Theory of
Reaction Rates
Three basic events must occur for a
bimolecular reaction to occur.
The atoms, molecules or ions must:
1. Collide.
2. Collide with enough energy to break and form
bonds.
3. Collide with the proper orientation for a reaction
to occur.
79
63. Collision Theory of
Reaction Rates
For a simple second order reaction between two different
gaseous atoms, A and B:
80
A B
A B
A B
B
A B
A B
A B
A B
4 different possible A-B
collisions
6 different possible A-B
collisions
9 different possible A-B
collisions
A(g) + B(g) products, where rate = k[A][B]
Increasing the number of atoms per unit volume (equivalent
to increasing their concentrations) increases the probability
of A/B collisions per unit time and, thereby increases the
reaction rate.
64. Collision Theory of
Reaction Rates
An example of an effective collision is:
81
X Y
X Y
X Y
X Y
X Y
+
X Y
Colliding molecules (i.e., individual reactants that
contain more than one atom) must be oriented
properly for reaction.
X2(g) + Y2(g) 2 XY(g)
66. Collision Theory of
Reaction Rates
Effective and ineffective molecular collisions for:
NO + N2O --> NO2 + N2
83
Effective
reactants collision
products
Ineffective (wrong orientation)
reactants collision
reactants
(blue spheres = N, red spheres = O)
67. Reaction Mechanisms and the
Rate-Law Expression
Many chemical reactions occur by more than one step.
The sequence of individual chemical steps by which the
reaction occurs is called the reaction mechanism.
Each step consists of either a unimolecular or bimolecular
chemical reaction.
The sum of the individual steps must equal that of the overall
balanced reaction.
Rate laws can be used to propose reaction mechanisms.
84
68. Reaction Mechanisms and the
Rate-Law Expression
The slowest step (the “bottleneck”) in a reaction mechanism
is the rate-determining step (r.d.s.) (sometimes called the
rate-limiting step).
85
Rules for a mechanism to be consistent with the
experimentally determined rate law.
Reaction orders indicate the number of molecules
involved in:
The r.d.s. only
or
The r.d.s. and any steps preceding the r.d.s.
Any step after the r.d.s. cannot be part of the rate law.
Sum of the steps in the mechanism must equal the
balanced equation for the overall reaction.
Intermediates in a mechanism cannot appear in the rate
law.
69. Reaction Mechanisms and the
Rate-Law Expression
Use the experimental rate law to postulate a
mechanism.
Consider the iodide ion-catalyzed
decomposition of hydrogen peroxide to water
and oxygen.
86
2H2O2(aq) ---> 2H2O(l) + O2(g)
I-
70. Reaction Mechanisms and the
Rate-Law Expression
This reaction has been experimentally determined to
be first order in H2O2, first order in I- , and second
order overall, i.e.
Rate = k[H2O2][I-]
87
Step 1, slow: H2O2 + I- --> IO- +H2O
Step 2, fast: IO- + H2O2 --> H2O + O2 + I-
Overall reaction: 2H2O2 --> 2H2O + O2
A mechanism consistent with this rate law is:
71. Reaction Mechanisms and the
Rate-Law Expression
According to this mechanism:
One hydrogen peroxide molecule and one iodide
ion react in the slow step (the r.d.s.).
88
The iodide ion functions as a catalyst: it is
consumed in step 1 and regenerated in step 2.
Hypoiodite (IO-) should not accumulate
appreciably because it is produced in the r.d.s.,
then rapidly consumed in the subsequent fast
reaction with hydrogen peroxide.
(IO- has been detected in very small amounts as a
short-lived reaction intermediate).
72. Reaction Mechanisms and the
Rate-Law Expression
Ozone, O3, reacts very rapidly with nitric oxide, NO,
in a reaction that is first order in each reactant and
second order overall.
89
73. Reaction Mechanisms and the
Rate-Law Expression
One possible mechanism that is consistent with the
rate law is:
90
k1
k2
Rate = k1[O3][NO]
• k2 >> k1, so the slow step is the “bottleneck” which
limits the overall reaction rate
• k1 = k, i.e., the rate constant for the overall reaction
74. Reaction Mechanisms and the
Rate-Law Expression
A mechanism that is inconsistent with the
rate-law is:
91
• Experimentally determined rate laws can, thus,
DISPROVE, but cannot, by themselves, PROVE a
particular mechanism
75. Reaction Mechanisms and the
Rate-Law Expression
92
• Consider the reaction:
2NO(g) + Br2(g) 2NOBr(g)
• The experimentally determined rate law is:
Rate = k[NO]2[Br2]
• One possible mechanism consistent with the rate law is a
simultaneous collision and reaction of two NO molecules
and one Br2 molecule. However, simultaneous three-body
collisions are highly improbable.
• More likely mechanisms would consist of a sequence
of bimolecular steps.
76. Reaction Mechanisms and the
Rate-Law Expression
93
NO + Br2 NOBr2 fast pre-equilibrium
NOBr2 + NO ---> 2NOBr
2NO + Br2 --> 2NOBr overall reaction
• According to this mechanism, the rate law for the slow step is:
Rate = k2[NOBr2][NO]
How to reconcile with the experimental rate law, rate = k[NO]2[Br2]?
• For the fast pre-equilibrium, the rates of the forward and
reverse reactions must be equal: k1[NO][Br2] = k-1[NOBr2].
• Rearrange to [NOBr2] = k1/k-1[NO][Br2], substitute into the rate
law for the r.d.s.:
• Rate = k2(k1/k-1[NO][Br2])[NO] = k[NO]2[Br2], where k = k2k1/k-1
k2
k-1
k1
slow
77. Transition State Theory
Transition state theory postulates that the
reactants form a high energy intermediate,
the transition state, which then rearranges
into the lower energy products.
For a reaction to occur, the reactants must
acquire sufficient energy to form the transition
state.
This energy is called the activation energy, Ea.
94
78. Transition State Theory
95
Energy changes during a chemical reaction.
Ea = activation energy
transition state
ΔE ~ ΔH
2H2(g) + O2(g) 2H2O(g)
80. 97
Transition State Theory
NO2(g) + CO(g) NO(g) + CO2(g)
experimental rate law: Rate = k[NO2]2
• Step 1) is the r.d.s (k1 << k2).
• Step 1) is slower than step 2)
because Ea1 > Ea2.
• The rate law of step 1) is the
same as the rate law of the
overall reaction.
A mechanism consistent with the rate law:
step 1: NO2(g) + NO2(g) NO3(g) + NO(g) Rate1 = k1[NO2]2 slow
step 2: NO3(g) + CO(g) NO2(g) + CO2(g) Rate2 = k2[NO3][CO] fast
81. Transition State Theory
The relationship between the activation energy
for forward and reverse reactions is:
Forward reaction activation energy = Ea
Reverse reaction activation energy = Ea + E
difference = E (~Hrxn, a thermodynamic quantity)
Higher Ea, slower reaction, lower rate constant, k.
In order for a reaction to reach the transition state and
convert to products, the combined kinetic energy of
the colliding molecules must be ≥ Ea.
One way to increase kinetic energy of molecules is to
raise the temperature.
Therefore, the rates of most chemical reaction
increase with increasing temperature. 98
83. Temperature Effects:
Transition State Theory
One method to increase the energy necessary to
break and reform bonds is to heat the molecules.
Recall the reaction of methane with molecular oxygen
is exothermic (H0 < 0) but slow at 25 0C:
102
The reaction can be started with a spark or match.
This provides the initial activation energy necessary to break
the first few bonds.
Afterwards the reaction is self-sustaining, i.e., heat
generated by reaction of the first few molecules provides the
activation energy to break bonds in other molecules.
CH4(g) + O2(g) CO2(g) + H2O(g) H0 = -880 kJ/mol
84. Temperature Effects:
The Arrhenius Equation
103
MEMORIZE and be able to use the Arrhenius equation!
• k is the rate constant
• T is the absolute temperature
• Ea is the activation energy
• R is the gas constant (8.314 J/mol∙K)
• A is the “pre-exponential factor” or “frequency factor”
• The number of approaches to the activation barrier per unit time.
• e(-Ea/RT) is the “exponential factor”
• the fraction of molecules that convert to product after approaching
the activation barrier.
ork = Ae
-Ea
RT
ln k =
-Ea
R
1
T
æ
è
ç
ö
ø
÷ + ln A
85. 104
Temp, K k, M-1∙s-1 Temp, K k, M-1∙s-1
600 3.37 x 103 1300 7.83 x 107
700 4.83 x 104 1400 1.45 x 108
800 3.58 x 105 1500 2.46 x 108
900 1.70 x 106 1600 3.93 x 108
1000 5.90 x 106 1700 5.93 x 108
1100 1.63 x 107 1800 8.55 x 108
1200 3.81 x 107 1900 1.19 x 109
Temperature Effects:
The Arrhenius Equation
Example: Determine the activation energy and frequency factor
for the reaction O3(g) O2(g) + O(g) given the following data:
86. 105
Plot these data as lnk vs I/T:
Temperature Effects:
The Arrhenius Equation
87. 106
Ea = m(-R)
solve for Ea
Ea
= 1.12 x104
K( ) 8.314 J
mol•K( )= 9.31x104 J
mol
Ea
= 93.1 kJ
mol
A = ey-intercept
solve for A 11-11
118.26
sM1036.4
1036.4
A
eA
Temperature Effects:
The Arrhenius Equation
y = mx + b, where m = -(Ea/R) and b = ln(A)
ln k =
-Ea
R
1
T
æ
è
ç
ö
ø
÷ + ln A
88. Temperature Effects:
The Arrhenius Equation
If the Arrhenius equation is written for two
temperatures, T2 and T1 with T2 > T1.
107
2
a
2
1
a
1
-Alnln
and
-Alnln
RT
E
k
RT
E
k
89. Temperature Effects:
The Arrhenius Equation
Subtract one equation from the other:
108
ln k2
- ln k1
= ln A - ln A -
Ea
RT2
- -
Ea
RT1
æ
è
ç
ö
ø
÷
simplify to: ln k2
- ln k1
=
Ea
RT1
-
Ea
RT2
rearrange to: ln
k2
k1
æ
è
ç
ö
ø
÷ =
Ea
R
1
T1
-
1
T2
æ
è
ç
ö
ø
÷
90. Temperature Effects:
The Arrhenius Equation
Consider the rate of a reaction for which Ea = 50
kJ/mol, at 20 oC (293 K) and at 30 oC (303 K).
How much do the two rates differ?
109
91. Temperature Effects:
The Arrhenius Equation
Many reactions have Ea 50 kJ/mol.
For these reactions the rate approximately doubles for
every 10 0C rise in temperature (near room temperature).
This is a good “rule of thumb” for dependence of reaction
rates (strictly speaking, rate constants) on temperature.
110
92. Catalysts
Catalysts increase reaction rates by providing an
alternative reaction pathway with a lower activation
energy.
There is no net consumption of the catalyst during
the reaction it catalyzes.
112
93. Catalysts
Homogeneous catalysts exist in same phase as the
reactants.
Heterogeneous catalysts exist in different phases
than the reactants.
Solid catalysts often catalyze solution reactions.
113
2H2O2 --> O2(g) + 2H2O
MnO2
96. Catalysts
116
Catalytic converters also catalyze oxidation of CO
to CO2
Overall reaction
2 CO(g)+ O2(g) 2CO2(g)
Adsorption
CO(g) CO(surface)
O2(g) O2(surface)
Activation
O2(surface) O(surface)
Reaction
CO(surface) +O(surface) CO2(surface)
Desorption
CO2(surface) CO2(g)
97. Catalysts
The Haber process is used industrially for “fixing” nitrogen
by producing ammonia, NH3, from N2 and H2:
117
N2(g) + 3H2(g) --------> 2NH3(g)
Fe,Fe2O3
• A good illustration of the interplay between kinetics and
thermodynamics.
• This reaction is actually an equilibrium that lies far towards the right
at 25 0C and is also exothermic.
N2(g) + 3H2(g) 2NH3(g) Kc = 108 at 25 0C, H = -92 kJ/mol
• However, the rate of this reaction is negligible at 25 0C.
• The reaction is carried out at 500 0C with catalyst to increase the rate.
• But high temperature favors the reverse reaction (because it’s
exothermic), so high pressure (200-1000 atm) is used to shift the
equilibrium to the right (fewer moles of gaseous product than
reactants).
99. Enzymes
119
Catalase: 2H2O2 O2(g) + 2H2O
Enzymes are biological catalysts (usually proteins)
that:
increase reaction rates
provide specificity, optimized to bind only particular
substrates and convert them to single sets of products.
100. After studying Tro Chapter 13 text and lecture
and working the problems, you should be able
to:
123
1. Determine reaction orders, rate laws, and values of rate constants
given a set of initial rates and initial concentrations.
2. Identify overall orders of reactions and order in each reactant given
a rate law.
3. Know units for zeroth, 1st and 2nd order rate constants.
4. Write the rate laws for simple overall zeroth, 1st and 2nd order
reactions.
5. Write and use the integrated rate equations for zeroth, 1st and 2nd
order rate laws with one reactant.
6. Describe how concentrations change with time or how the rate will
change for a given concentration change for each of the rate laws in
#3.
7. Use these rate laws to determine: half-times, concentrations
remaining after a given reaction time or from a given starting
concentration and rate constants.
8. Identify or draw and describe equations for linear plots of zeroth, 1st
and 2nd order reactions.
9. Determine order of reaction given plots in #8.
101. After studying Tro Chapter 13 text and lecture
and working the problems, you should be able
to:
10. Write reasonable mechanisms given rate laws and vice versa.
11. Define the rate-determining step of a reaction mechanism.
12. Effects of collisions and molecular orientations on reactions.
13. Diffusion-controlled reactions are limited only by collision rates.
14. Define activation energy draw reaction progress diagrams and
label activation energy, transition state and Hrxn.
15. Qualitatively describe temperature dependence of reactions in
terms of molecular kinetic energies.
16. Qualitatively draw and/or interpret Maxwell-Boltzmann distributions
of molecular kinetic energies and label areas equal to and
exceeding activation energy.
17. Memorize and use Arrhenius equation(s) to calculate temperature
dependence of reaction rates and activation energies.
18. Define catalyst and its effect on activation energy and reaction rate.
19. Not responsible for enzyme catalysis or chain reactions.
124
102. Mid-term Exam 4
125
Chapter 18 (Electrochemistry) including:
1. Use the Nernst equation (memorize) to calculate Ecell under non-
standard conditions given half-reaction E0 values, initial concentrations
and the Nernst equation.
2. Calculate values of K and Grxn from E0
cell using the Nernst equation
and G = -nFEcell (memorize).
3. Describe and calculate Eo for a concentration cell.
4. Identify components in an alkaline dry cell given the half reactions and
a diagram of the cell. Show direction of current flow, and explain why
voltage stays fairly constant during discharge.
5. Draw a diagram of a lead acid car battery given the half reactions, and
describe the discharge and recharge cycles.
6. Describe and draw a diagram of a H2/O2 fuel cell.
7. Given E0’ values (pH 7), describe and calculate how energy can be
derived from biological redox reactions.
All of Chapter 13 (kinetics)