The document provides an outline and notes on the topics of rate of reaction, collision theory, factors affecting reaction rates, activation energy, endothermic and exothermic reactions, equilibrium, and Le Chatelier's principle. Key points include:
- Reaction rates can be determined by measuring how quickly reactants are used up or products are formed. Temperature, catalysts, surface area, and concentration can impact reaction rates.
- Collision theory states that reactants must collide with sufficient energy and correct orientation for a reaction to occur, forming an intermediate activated complex.
- Equilibrium is reached when the rates of the forward and reverse reactions are equal, though concentrations of reactants and products may not be equal.
-
The document discusses chemical equilibrium and reversible reactions. It defines chemical equilibrium as a state where the forward and reverse reactions are proceeding at the same rate, such that the concentrations of reactants and products remain constant. It describes characteristics of equilibrium such as it being dynamic, having equal forward and reverse reaction rates, and requiring a closed system. It also introduces Le Châtelier's principle, which states that disturbances to a system at equilibrium cause the equilibrium to shift in a direction that counteracts the applied stress.
The document summarizes key concepts relating to chemical kinetics and chemical equilibrium. It discusses how the rates of chemical reactions are determined by measuring changes in concentration over time. It also explains how reaction rates are affected by molecular collisions, activation energy, nature of reactants, concentration, temperature, and presence of catalysts. The document introduces chemical equilibrium as a dynamic steady state and defines equilibrium constants. It describes Le Chatelier's principle, explaining how changing concentrations, temperature, or pressure shifts equilibrium.
Lecture 2 By MUHAMMAD FAHAD ANSARI 12 IEEM 14fahadansari131
This document summarizes the key points about types of chemical reactions and factors that affect reaction rates. It discusses the 5 main types of chemical reactions: synthesis, decomposition, single-replacement, double-replacement, and oxidation-reduction. It also outlines characteristics of each type and provides examples. Additionally, the document covers concepts like activation energy, catalysts, temperature effects, and concentration effects on reaction rates. Finally, it briefly discusses chemical equilibrium and kinetic molecular theory of gases.
The document discusses equilibrium constants (Kc) and how to calculate them using concentrations of reactants and products at equilibrium. It provides examples of calculating Kc values for reactions, including determining initial and change in concentrations. It also discusses using Kc to predict the direction a reaction will proceed based on comparing the reaction quotient (Q) to Kc.
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.
The document discusses chemical kinetics and factors that influence reaction rates. It covers topics like rate laws, reaction order, rate constants, and factors affecting reaction rates such as temperature, concentration of reactants, and presence of catalysts. The rate of a reaction is important to study as it provides information about manipulating reaction conditions and determining reaction mechanisms and products.
It is fully based on the notes provided by the K V Sangathan. For the revision to students they are short but enough to clear the concept of equillibrium.hope you like them.give your reviews.
THANK YOU
The document discusses thermochemistry and concepts related to heat, temperature, enthalpy, and thermochemical equations. It provides definitions and examples of:
- Specific heat and how to calculate it using the formula q=m*c*ΔT
- Enthalpy change (ΔH), enthalpy of reaction, enthalpy of formation (ΔHf), and enthalpy of combustion (ΔHc)
- How to use Hess's law to calculate enthalpy changes from individual reaction steps
The document discusses chemical equilibrium and reversible reactions. It defines chemical equilibrium as a state where the forward and reverse reactions are proceeding at the same rate, such that the concentrations of reactants and products remain constant. It describes characteristics of equilibrium such as it being dynamic, having equal forward and reverse reaction rates, and requiring a closed system. It also introduces Le Châtelier's principle, which states that disturbances to a system at equilibrium cause the equilibrium to shift in a direction that counteracts the applied stress.
The document summarizes key concepts relating to chemical kinetics and chemical equilibrium. It discusses how the rates of chemical reactions are determined by measuring changes in concentration over time. It also explains how reaction rates are affected by molecular collisions, activation energy, nature of reactants, concentration, temperature, and presence of catalysts. The document introduces chemical equilibrium as a dynamic steady state and defines equilibrium constants. It describes Le Chatelier's principle, explaining how changing concentrations, temperature, or pressure shifts equilibrium.
Lecture 2 By MUHAMMAD FAHAD ANSARI 12 IEEM 14fahadansari131
This document summarizes the key points about types of chemical reactions and factors that affect reaction rates. It discusses the 5 main types of chemical reactions: synthesis, decomposition, single-replacement, double-replacement, and oxidation-reduction. It also outlines characteristics of each type and provides examples. Additionally, the document covers concepts like activation energy, catalysts, temperature effects, and concentration effects on reaction rates. Finally, it briefly discusses chemical equilibrium and kinetic molecular theory of gases.
The document discusses equilibrium constants (Kc) and how to calculate them using concentrations of reactants and products at equilibrium. It provides examples of calculating Kc values for reactions, including determining initial and change in concentrations. It also discusses using Kc to predict the direction a reaction will proceed based on comparing the reaction quotient (Q) to Kc.
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.
The document discusses chemical kinetics and factors that influence reaction rates. It covers topics like rate laws, reaction order, rate constants, and factors affecting reaction rates such as temperature, concentration of reactants, and presence of catalysts. The rate of a reaction is important to study as it provides information about manipulating reaction conditions and determining reaction mechanisms and products.
It is fully based on the notes provided by the K V Sangathan. For the revision to students they are short but enough to clear the concept of equillibrium.hope you like them.give your reviews.
THANK YOU
The document discusses thermochemistry and concepts related to heat, temperature, enthalpy, and thermochemical equations. It provides definitions and examples of:
- Specific heat and how to calculate it using the formula q=m*c*ΔT
- Enthalpy change (ΔH), enthalpy of reaction, enthalpy of formation (ΔHf), and enthalpy of combustion (ΔHc)
- How to use Hess's law to calculate enthalpy changes from individual reaction steps
This document discusses various topics in thermochemistry including:
- Enthalpy changes in chemical reactions and how they are measured using calorimetry. Exothermic and endothermic reactions are explained.
- Hess's law, which states that the enthalpy change of a reaction is independent of the reaction pathway. It can be used to calculate enthalpy changes.
- Standard enthalpies of formation and how they allow calculation of enthalpy changes using Hess's law and bond dissociation enthalpies.
- Measuring enthalpy changes using bomb calorimetry and coffee cup calorimetry. Limitations of each method are discussed.
This document discusses thermochemistry and energy changes that occur during chemical reactions. It defines exothermic and endothermic reactions, and how to construct energy level diagrams to represent them. Specific heats of reaction like combustion, precipitation, displacement, and neutralization are also explained. Experiments to determine various heats of reaction are described. The relationships between the heat of reaction and type of reactants, as well as the number of carbons in alcohols are also summarized.
I Hope You all like it very much. I wish it is beneficial for all of you and you can get enough knowledge from it. Clear and appropriate objectives, in terms of what the audience ought to feel, think, and do as a result of seeing the presentation. Objectives are realistic – and may be intermediate parts of a wider plan.
There are two types of equilibrium: physical and chemical. Physical equilibrium occurs when a system's physical state does not change over time, such as when a solid melts at a fixed temperature. Chemical equilibrium is reached when the rates of the forward and reverse reactions are equal and the concentrations of reactants and products remain constant. The equilibrium constant (K) is a ratio of product to reactant concentrations raised to their stoichiometric coefficients that remains constant at equilibrium. K can be used to determine whether the equilibrium favors reactants or products.
This document provides an overview of the states of matter and phase changes. It discusses gases, liquids, and solids. For gases, it covers the gas laws, kinetic molecular theory, and intermolecular forces. For liquids, it discusses surface tension, vapor pressure, and factors that affect boiling point such as intermolecular forces and molecular shape. It also describes the different types of solids including molecular, ionic, metallic, polymeric, and network solids.
The document discusses chemical equilibrium. It defines equilibrium as a state where the rates of the forward and reverse reactions of a chemical process are equal, resulting in no net change in the concentrations or properties of the system. It provides examples of physical and chemical equilibrium processes. It describes key characteristics of equilibrium like dynamic nature, constant concentrations and temperatures, and the relationship between reaction rates and equilibrium constants.
The document discusses chemical equilibrium, including:
- Chemical equilibrium is a state where reactants and products are present at constant concentrations.
- Physical and chemical equilibrium can involve physical or chemical processes respectively.
- Reversible reactions can proceed in both directions at equilibrium, while irreversible reactions only proceed in one direction.
- Equilibrium can be homogeneous, with all substances in one phase, or heterogeneous, with substances in multiple phases.
- The law of mass action and equilibrium constants relate reaction rates and concentrations at equilibrium.
- Le Chatelier's principle states that if stress is applied to a system at equilibrium, it will respond to reduce the effect of the stress.
1. A liter of gasoline contains 8000 calories of energy. A person uses an average of 2000 calories per day. Excess calories are stored as fat.
2. Calorimetry is used to determine the energy content of substances by measuring heat changes. Specific heat and heat capacity allow calculation of heat from temperature changes.
3. Enthalpy (H) quantifies heat flow during chemical reactions. Standard enthalpies of formation provide a reference scale for enthalpy values.
Chapter 18.1 : The Nature of Chemical EquilibriumChris Foltz
This document provides information about chemical equilibrium, including definitions, concepts, and examples. It defines chemical equilibrium as a state where the rates of the forward and reverse reactions are equal and the concentrations of reactants and products remain constant. The equilibrium constant, K, is introduced as a ratio of product concentrations over reactant concentrations raised to their stoichiometric coefficients. Examples are provided to demonstrate how to write equilibrium expressions and calculate K values or concentrations at equilibrium.
1. This document summarizes key concepts from a chapter on chemical reactions including formula weight, moles, balancing equations, stoichiometry, percent yield, and oxidation-reduction reactions.
2. Key topics covered include calculating molar mass, determining the limiting reactant and theoretical yield of a reaction, and identifying oxidation and reduction in redox reactions.
3. The chapter also discusses solubility rules, writing net ionic equations, and the heat absorbed or released by exothermic and endothermic reactions.
The document discusses chemical equilibrium, which occurs when the forward and reverse reactions of a chemical reaction proceed at the same rate. At equilibrium, the concentrations of reactants and products remain constant. The equilibrium constant, K, is a ratio of products over reactants that characterizes the position of equilibrium. A large K value indicates the reaction favors products, while a small K value indicates the reaction favors reactants.
This document discusses chemical equilibrium, including:
- Reactions reach equilibrium when concentrations of reactants and products remain constant over time.
- The equilibrium constant, K, quantifies the position of equilibrium and can be used to calculate concentrations at equilibrium.
- Equilibrium expressions can involve gas concentrations or pressures, and heterogeneous equilibria only include gases and dissolved substances in expressions.
- Knowing K allows prediction of whether a reaction will occur and the direction a system will shift to reach equilibrium.
Thermochemistry is the study of heat changes in chemical reactions. There are several key concepts:
1) Exothermic reactions release heat to the surroundings and have a negative ∆H value, while endothermic reactions absorb heat from the surroundings and have a positive ∆H value.
2) The first law of thermodynamics states that energy is conserved and can be converted between different forms but not created or destroyed. ∆Esystem = q + w, where q is heat and w is work.
3) Enthalpy (H) is a state function that measures the heat absorbed or released during physical and chemical changes at constant pressure. The enthalpy change, ∆H
The document summarizes key concepts about chemical equilibria including:
1) The equilibrium constant K describes the position of chemical equilibrium and can be written in terms of concentrations or pressures.
2) K expressions are written based on reaction stoichiometry and do not include solids/liquids.
3) Le Châtelier's principle states how changing conditions affects the equilibrium position.
This document discusses several key concepts in thermochemistry and thermodynamics:
- The first law of thermodynamics relates heat, work, and changes in internal energy. Enthalpy is a useful state function for chemical reactions.
- Calorimetry can be used to measure heat flows and determine enthalpy changes during chemical and physical processes. Bomb calorimetry and coffee cup calorimetry are described.
- Enthalpy changes accompany chemical reactions and phase changes. The sign and magnitude of enthalpy changes can indicate whether a process is exothermic or endothermic.
- Hess's law allows calculation of enthalpy changes from thermochemical data using stepwise processes
The fundamentals of chemical equilibrium including Le Chatier's Principle and solved problems for heterogeneous and homogeneous equilibrium.
**More good stuff available at:
www.wsautter.com
and
http://www.youtube.com/results?search_query=wnsautter&aq=f
This document discusses the topic of chemical equilibrium. It provides an introduction to key concepts such as reversible reactions occurring at the same rate and the equilibrium constant. It also discusses factors that affect chemical equilibrium such as concentration, pressure, temperature and catalysts. Le Chatelier's principle is explained, which states that applying stress to a system at equilibrium causes it to counteract the change. Specific examples are provided to illustrate how changing concentration, pressure and temperature impacts the position of equilibrium.
This document provides an overview of states of matter, gas laws, intermolecular forces, liquids, solids, and phase changes. It discusses the key properties and behaviors of gases, liquids, and solids, including gas laws, vapor pressure, boiling point, factors that affect boiling point, types of solids, and phase changes. The document also provides examples and practice problems to illustrate these concepts.
Thermochemistry deals with the heat involved in chemical and physical changes. It is a branch of thermodynamics that studies energy and its transformations. Enthalpy (H) is a measure of the total energy of a system at constant pressure and can be used to determine the heat of a reaction. Calorimetry experiments allow measurement of heat changes through determination of temperature changes of a system and surroundings using equations such as q = cmΔT. Bomb calorimetry and coffee cup calorimetry are two common techniques used to directly measure the heat of chemical reactions.
This document discusses reaction rates and chemical equilibrium. It begins by defining reaction rates and factors that influence reaction rates such as temperature, concentration, surface area, and catalysts. It then explains collision theory and the role of activation energy in reactions. The document also covers Le Chatelier's principle, how stresses such as concentration, temperature, and pressure affect chemical equilibrium. It defines equilibrium constants and discusses solubility equilibrium, including solubility product constants and the common ion effect. Finally, it introduces entropy, the role of entropy in spontaneous reactions, and free energy.
The document discusses kinetics and reaction rates. It defines kinetics as the branch of chemistry that studies the speed or rate of chemical reactions. It explains that reaction rates can be measured by changes in concentration, temperature, or pressure over time. The rate depends on factors like the nature of reactants, concentration, temperature, catalysts, surface area, and pressure. Reactions may occur in multiple steps through reaction intermediates rather than a single step. The collision theory and concept of activation energy are introduced to explain why certain collisions result in reactions. Reaction coordinate diagrams are used to illustrate the energy changes in reactions.
This document discusses various topics in thermochemistry including:
- Enthalpy changes in chemical reactions and how they are measured using calorimetry. Exothermic and endothermic reactions are explained.
- Hess's law, which states that the enthalpy change of a reaction is independent of the reaction pathway. It can be used to calculate enthalpy changes.
- Standard enthalpies of formation and how they allow calculation of enthalpy changes using Hess's law and bond dissociation enthalpies.
- Measuring enthalpy changes using bomb calorimetry and coffee cup calorimetry. Limitations of each method are discussed.
This document discusses thermochemistry and energy changes that occur during chemical reactions. It defines exothermic and endothermic reactions, and how to construct energy level diagrams to represent them. Specific heats of reaction like combustion, precipitation, displacement, and neutralization are also explained. Experiments to determine various heats of reaction are described. The relationships between the heat of reaction and type of reactants, as well as the number of carbons in alcohols are also summarized.
I Hope You all like it very much. I wish it is beneficial for all of you and you can get enough knowledge from it. Clear and appropriate objectives, in terms of what the audience ought to feel, think, and do as a result of seeing the presentation. Objectives are realistic – and may be intermediate parts of a wider plan.
There are two types of equilibrium: physical and chemical. Physical equilibrium occurs when a system's physical state does not change over time, such as when a solid melts at a fixed temperature. Chemical equilibrium is reached when the rates of the forward and reverse reactions are equal and the concentrations of reactants and products remain constant. The equilibrium constant (K) is a ratio of product to reactant concentrations raised to their stoichiometric coefficients that remains constant at equilibrium. K can be used to determine whether the equilibrium favors reactants or products.
This document provides an overview of the states of matter and phase changes. It discusses gases, liquids, and solids. For gases, it covers the gas laws, kinetic molecular theory, and intermolecular forces. For liquids, it discusses surface tension, vapor pressure, and factors that affect boiling point such as intermolecular forces and molecular shape. It also describes the different types of solids including molecular, ionic, metallic, polymeric, and network solids.
The document discusses chemical equilibrium. It defines equilibrium as a state where the rates of the forward and reverse reactions of a chemical process are equal, resulting in no net change in the concentrations or properties of the system. It provides examples of physical and chemical equilibrium processes. It describes key characteristics of equilibrium like dynamic nature, constant concentrations and temperatures, and the relationship between reaction rates and equilibrium constants.
The document discusses chemical equilibrium, including:
- Chemical equilibrium is a state where reactants and products are present at constant concentrations.
- Physical and chemical equilibrium can involve physical or chemical processes respectively.
- Reversible reactions can proceed in both directions at equilibrium, while irreversible reactions only proceed in one direction.
- Equilibrium can be homogeneous, with all substances in one phase, or heterogeneous, with substances in multiple phases.
- The law of mass action and equilibrium constants relate reaction rates and concentrations at equilibrium.
- Le Chatelier's principle states that if stress is applied to a system at equilibrium, it will respond to reduce the effect of the stress.
1. A liter of gasoline contains 8000 calories of energy. A person uses an average of 2000 calories per day. Excess calories are stored as fat.
2. Calorimetry is used to determine the energy content of substances by measuring heat changes. Specific heat and heat capacity allow calculation of heat from temperature changes.
3. Enthalpy (H) quantifies heat flow during chemical reactions. Standard enthalpies of formation provide a reference scale for enthalpy values.
Chapter 18.1 : The Nature of Chemical EquilibriumChris Foltz
This document provides information about chemical equilibrium, including definitions, concepts, and examples. It defines chemical equilibrium as a state where the rates of the forward and reverse reactions are equal and the concentrations of reactants and products remain constant. The equilibrium constant, K, is introduced as a ratio of product concentrations over reactant concentrations raised to their stoichiometric coefficients. Examples are provided to demonstrate how to write equilibrium expressions and calculate K values or concentrations at equilibrium.
1. This document summarizes key concepts from a chapter on chemical reactions including formula weight, moles, balancing equations, stoichiometry, percent yield, and oxidation-reduction reactions.
2. Key topics covered include calculating molar mass, determining the limiting reactant and theoretical yield of a reaction, and identifying oxidation and reduction in redox reactions.
3. The chapter also discusses solubility rules, writing net ionic equations, and the heat absorbed or released by exothermic and endothermic reactions.
The document discusses chemical equilibrium, which occurs when the forward and reverse reactions of a chemical reaction proceed at the same rate. At equilibrium, the concentrations of reactants and products remain constant. The equilibrium constant, K, is a ratio of products over reactants that characterizes the position of equilibrium. A large K value indicates the reaction favors products, while a small K value indicates the reaction favors reactants.
This document discusses chemical equilibrium, including:
- Reactions reach equilibrium when concentrations of reactants and products remain constant over time.
- The equilibrium constant, K, quantifies the position of equilibrium and can be used to calculate concentrations at equilibrium.
- Equilibrium expressions can involve gas concentrations or pressures, and heterogeneous equilibria only include gases and dissolved substances in expressions.
- Knowing K allows prediction of whether a reaction will occur and the direction a system will shift to reach equilibrium.
Thermochemistry is the study of heat changes in chemical reactions. There are several key concepts:
1) Exothermic reactions release heat to the surroundings and have a negative ∆H value, while endothermic reactions absorb heat from the surroundings and have a positive ∆H value.
2) The first law of thermodynamics states that energy is conserved and can be converted between different forms but not created or destroyed. ∆Esystem = q + w, where q is heat and w is work.
3) Enthalpy (H) is a state function that measures the heat absorbed or released during physical and chemical changes at constant pressure. The enthalpy change, ∆H
The document summarizes key concepts about chemical equilibria including:
1) The equilibrium constant K describes the position of chemical equilibrium and can be written in terms of concentrations or pressures.
2) K expressions are written based on reaction stoichiometry and do not include solids/liquids.
3) Le Châtelier's principle states how changing conditions affects the equilibrium position.
This document discusses several key concepts in thermochemistry and thermodynamics:
- The first law of thermodynamics relates heat, work, and changes in internal energy. Enthalpy is a useful state function for chemical reactions.
- Calorimetry can be used to measure heat flows and determine enthalpy changes during chemical and physical processes. Bomb calorimetry and coffee cup calorimetry are described.
- Enthalpy changes accompany chemical reactions and phase changes. The sign and magnitude of enthalpy changes can indicate whether a process is exothermic or endothermic.
- Hess's law allows calculation of enthalpy changes from thermochemical data using stepwise processes
The fundamentals of chemical equilibrium including Le Chatier's Principle and solved problems for heterogeneous and homogeneous equilibrium.
**More good stuff available at:
www.wsautter.com
and
http://www.youtube.com/results?search_query=wnsautter&aq=f
This document discusses the topic of chemical equilibrium. It provides an introduction to key concepts such as reversible reactions occurring at the same rate and the equilibrium constant. It also discusses factors that affect chemical equilibrium such as concentration, pressure, temperature and catalysts. Le Chatelier's principle is explained, which states that applying stress to a system at equilibrium causes it to counteract the change. Specific examples are provided to illustrate how changing concentration, pressure and temperature impacts the position of equilibrium.
This document provides an overview of states of matter, gas laws, intermolecular forces, liquids, solids, and phase changes. It discusses the key properties and behaviors of gases, liquids, and solids, including gas laws, vapor pressure, boiling point, factors that affect boiling point, types of solids, and phase changes. The document also provides examples and practice problems to illustrate these concepts.
Thermochemistry deals with the heat involved in chemical and physical changes. It is a branch of thermodynamics that studies energy and its transformations. Enthalpy (H) is a measure of the total energy of a system at constant pressure and can be used to determine the heat of a reaction. Calorimetry experiments allow measurement of heat changes through determination of temperature changes of a system and surroundings using equations such as q = cmΔT. Bomb calorimetry and coffee cup calorimetry are two common techniques used to directly measure the heat of chemical reactions.
This document discusses reaction rates and chemical equilibrium. It begins by defining reaction rates and factors that influence reaction rates such as temperature, concentration, surface area, and catalysts. It then explains collision theory and the role of activation energy in reactions. The document also covers Le Chatelier's principle, how stresses such as concentration, temperature, and pressure affect chemical equilibrium. It defines equilibrium constants and discusses solubility equilibrium, including solubility product constants and the common ion effect. Finally, it introduces entropy, the role of entropy in spontaneous reactions, and free energy.
The document discusses kinetics and reaction rates. It defines kinetics as the branch of chemistry that studies the speed or rate of chemical reactions. It explains that reaction rates can be measured by changes in concentration, temperature, or pressure over time. The rate depends on factors like the nature of reactants, concentration, temperature, catalysts, surface area, and pressure. Reactions may occur in multiple steps through reaction intermediates rather than a single step. The collision theory and concept of activation energy are introduced to explain why certain collisions result in reactions. Reaction coordinate diagrams are used to illustrate the energy changes in reactions.
UNIT 8 CHEMICAL KINETICS.pptxUNIT 8 CHEMICAL KINETICS.pptxfatema220366
This document provides an overview of chemical kinetics and key concepts related to reaction rates. It discusses representation of reaction rates, factors that affect reaction rates like concentration, temperature, and catalysts. Activation energy and the activated complex are explained. Different order reactions and molecularity are defined. Rate laws, rate constants, and half-life are also covered. Examples of zero order, first order and second order reactions are given to illustrate concepts.
Chemical energy is energy stored in the bonds of chemical compounds. This energy is released when a chemical reaction takes place, transforming the original substances into new ones. Exothermic reactions release thermal energy (heat) into their surroundings, causing the temperature to increase. Examples include combustion and neutralization reactions. Endothermic reactions absorb energy from their surroundings.
This document provides an overview of chemical reactions and energetics for a 10th grade IGCSE course. It discusses exothermic and endothermic reactions in relation to energy changes and temperature. It also covers the factors that affect reaction rates, including concentration, particle size, catalysis, and temperature. The document defines oxidation and reduction in terms of electron transfer and identifies redox reactions. It provides examples of exothermic and endothermic reactions and discusses how catalysts can lower the activation energy and increase reaction rates. Interactive links are included to illustrate and reinforce the concepts.
CHEMICAL REACTION
CHEMICAL EQUATION
CHEMICAL FORMULA
BALANCING
TYPES OF CHEMICAL REACTION
COLLISION THEORY
FACTORS AFFECTING THE RATE OF CHEMICAL REACTION
This document provides information about stoichiometry concepts including mole ratios, limiting reagents, theoretical and actual yields, and percent yields. It gives examples of mole-mole, gram-gram, and gram-liter conversions using balanced chemical equations. Heat of reaction is discussed, including signs for exothermic and endothermic reactions. Sample problems are provided for limiting reagents, percent yields, and heat of reaction calculations. Worksheets are assigned to practice these stoichiometry skills in preparation for a test involving mole-mole, mass, volume, heat, yield, and limiting reagent conversions and calculations.
Core & Extension - Chemical Rxns - Reversible Rxns I.pptxMathandScienced
1. The document discusses reversible chemical reactions and chemical equilibrium.
2. It explains that some reactions can go in both the forward and reverse directions under certain conditions, and that at equilibrium the rates of the forward and reverse reactions are equal.
3. The document describes how changing conditions like temperature, pressure, or concentration can shift the equilibrium position by favoring the endothermic or exothermic direction.
Unit 1 Chemical reactions in our surroundingsalekey08
This document discusses chemical reactions and stoichiometry. It begins by defining physical and chemical changes, and chemical equations which represent chemical reactions symbolically. It then discusses various topics related to chemical reactions including reactants and products, balancing equations, reaction rates and factors that affect them. It also classifies common reaction types and discusses energy changes. Finally, it introduces stoichiometry, which uses mole ratios determined from balanced chemical equations to calculate amounts of reactants and products.
This document discusses chemical reactions and how to write and balance chemical equations. It defines chemical reactions as when new substances are formed with new properties, whereas physical changes only alter physical properties. To identify if a reaction occurred, one looks for evidence like new colors, gas formation, or heat/light production. Chemical equations represent reactions, with reactants on the left and products on the right. Formulas replace word names and physical states are indicated. Equations are balanced by adding coefficients to ensure equal numbers of each type of atom/ion on both sides of the equation. Several examples demonstrate how to write word equations, convert to formulas, and balance chemical equations.
This document defines and describes chemical reactions. It explains that in a chemical reaction, atoms rearrange to form new compounds, with old bonds breaking and new bonds forming. Chemical reactions can be represented by chemical equations that show the reactants and products. Reactions can be exothermic, releasing energy, or endothermic, absorbing energy. They can also vary in reaction rate. The document also covers the law of conservation of matter, balancing chemical equations, and the main types of chemical reactions: synthesis, decomposition, single replacement, double replacement, and combustion reactions.
1) A chemical reaction is a process where one or more substances are destroyed and one or more new substances are created.
2) There are five main types of chemical reactions: single replacement, double replacement, synthesis, decomposition, and combustion.
3) Evidence for a chemical reaction includes evolution of light or heat, temperature changes, formation of gases or precipitates, and color changes.
The document provides an introduction to chemical reactions, including definitions, parts of reactions, types of reactions, and evidence of reactions. It explains that a chemical reaction is a process where reactants are destroyed and products are formed, with the following key points:
- Reactants undergo chemical change to form products, as indicated by the yield arrow (→).
- Chemical equations must obey the law of conservation of mass, with the same number and type of atoms on both sides of the reaction.
- The five main types of chemical reactions are synthesis, decomposition, single replacement, double replacement, and combustion.
- Several tests can indicate if a chemical reaction has occurred, such as gas evolution, temperature change, color change,
1) A chemical reaction is a process where one or more substances are destroyed and one or more new substances are created.
2) There are five main types of chemical reactions: single replacement, double replacement, synthesis, decomposition, and combustion.
3) Evidence for a chemical reaction includes evolution of light or heat, temperature changes, formation of gases or precipitates, and color changes.
1) A chemical reaction is a process where one or more substances are destroyed and one or more new substances are created.
2) There are five main types of chemical reactions: single replacement, double replacement, synthesis, decomposition, and combustion.
3) Evidence for a chemical reaction includes evolution of light or heat, temperature changes, formation of gases or precipitates, and color changes.
This chapter discusses reaction kinetics and mechanisms. It explains that reaction mechanisms show the specific steps by which reactants are converted to products. Collision theory is used to interpret reactions, where molecules must collide with sufficient energy and proper orientation for a reaction to occur. The activation energy is the minimum energy needed to form an activated complex or transition state. Reaction mechanisms must satisfy the overall balanced chemical equation and agree with the experimentally determined rate law. Catalysts are discussed as substances that increase the rate of reactions without being consumed.
This document discusses chemical equations and balancing chemical reactions. It explains that a chemical equation describes a chemical change and is made up of reactants on the left and products on the right, with coefficients showing quantities. Balancing equations involves adjusting these coefficients to satisfy the law of conservation of mass, ensuring the same number and type of atoms enter and leave the reaction. Several examples of balancing equations are provided.
This document discusses chemical kinetics and factors that affect reaction rates. It defines chemical kinetics as the branch of chemistry concerned with the rate of chemical reactions and the mechanisms by which they occur. Some key points made include:
- Thermodynamics determines if a reaction can occur but not how fast, while kinetics describes the reaction rate.
- A reaction may be thermodynamically favored but proceed too slowly to be observed without a catalyst.
- Factors that increase reaction rates include higher concentrations of reactants, higher temperatures, and the presence of catalysts.
This document discusses balancing chemical equations. It explains that a chemical equation describes a chemical reaction by showing the reactants and products. It also notes that balancing a chemical equation establishes the quantitative relationship between reactants and products by using coefficients. There are three main steps to balancing an equation: 1) writing the unbalanced equation, 2) balancing the equation by applying the law of conservation of mass so each element has the same number of atoms on both sides, and 3) indicating the states of matter of the reactants and products using abbreviations like (g) or (s). An example problem of balancing the equation for the reaction of tin oxide with hydrogen gas to form tin and water vapor is provided.
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2. Outline:
•Rate
•Reactions & collision theory
•Graphs w/ activation energy of endo & exo
•4 factors that affect rates of reaction
•Read chapter 17
•Students should take notes & copy
diagrams (might have space after vocab)
3. Sample notes
• What is rate?
• What factors affect rate?
• What is collision theory?
• Samples of exothermic and endothermic
graphs of reaction rate
4. Lab/demos
Zinc in HCl, 2 concentrations, Zn,
different surface areas, matches
WHEN is it done?
Catalyst demo (shows activity complex*)
Heat 300 ml ofA to 80C then add 100 mlH202 observe then add 30 ml of catalyst CoCl2
Surface area bottle
5. What is rate?
• Describe these rates:
– When we light a match
– We put zinc in HCl (2 different concentrations)
• How do we know when the reaction is
done?
6. Rate of reaction
• can be determined by
measuring the rate at which
reactants are used up or the rate
at which products are formed.
7. • A would be increasing rate: likeexplosive (gun powder)
• B is a constant rate (like how soaps work)
• C is slowing down, like what refrigerators would do
• KEY: Chemists work on controlling rates.
Discuss rates at
various points.
8. Collision theory
• Reactants have to strike each other with
sufficient energy and in the correct
configuration in order for the reaction to
occur.
• Kicker will share an example of students at a dance.
• Reactants must hit with:
– 1. sufficient energy
– 2 correct config. Or orientation
• In order for a reaction to progress.
9. Consider a simple reaction involving a collision
between two molecules - ethene, CH2=CH2, and
hydrogen chloride, HCl.
These react to give
chloroethane.
As a result of the
collision between
the two molecules,
the double bond
between the two
carbons is
converted into a
single bond. A
hydrogen atom
gets attached to
one of the carbons
and a chlorine
atom to the other.
10. During a reaction
• An intermediate particle is formed. It is
neither reactants nor products. This
happens when there is enough energy &
proper orientation of the reactants. This is
called an activated complex. This is an
intermediate that does not stick around
long.
See catalyst demo & color change!
11. Discuss activation
energy
See page 597 graph.
What causes paper to
burn?
Kicker will burn some paper
and discuss the graph it
would make.
Wave it in the air, rub it like
two sticks, then use a
heat source
13. The activation barrier must be crossed before reactants
are converted to products. The activated complex is a
temporary arrangement of particles that has sufficient
energy to become either reactant or products.
16. This is an exothermic reaction. The products have
less energy than the reactants. But it took a lot of
activation energy to get it going. Most of these
reactions are self sustaining as they have enough
energy to keep going.
17. This is an endothermic reaction. The products have more energy than the
reactants. These reactions are not self-sustaining because not enough
energy is released to keep the reaction going. Endothermic reactions need a
constant supply of energy to keep going.
Some examples are ice melting, evaporation, photosynthesis
19. SAMPLE QUIZ ?s
What do you have at A,
B, & C. What is D? Is
this exo or endo?
20. I get this, but how
can we control the
rate of the reaction?
21. 4 methods of increasing the rate
of reaction.
• Temperature
• Catalyst
• Surface area
• Concentration.
22. Temperature
• If temp. increases, kinetic energy increases and
so do the number of collisions.
– If temperature decreases, the opposite is true
• A 10 C increase in temp will approximately
double the rate of reaction
• A 30 C increase will cause an increase in rate of 8
times faster. (note that it is exponential)
23. Catalyst• This is a substance
which changes the
rate, without
undergoing any
permanent chemical
change itself. It does
this by lowering the
activation energy by
allowing a a different
path for the reaction
to occur.
24. Write this reaction
• The rate of the decomposition of hydrogen
peroxide (dihydrogen dioxide) into water
and oxygen gas can be increased by adding
the catalyst manganese dioxide. (Put the catalyst
above the arrow)
25. Video of how this catalyst works.
H2O2 MnO2 u ruuuuuu H2O + O2
29. Some reactions are simply faster
because of the nature of the
substances.
• Ions in solution can react faster
• AgNO3 + Cl- --> AgCl (s)
• Covalent with covalent are usually slow because
you have to break tough bonds before other bonds
can form.
• N2 + H2 -->NH3
• Br2(l) + C6H11 (aq) --> C6H11Br + HBr
30. Assignment
Do pg 3 NOW
Assign pg 4 & 5 in packet
Read pg 595-600 & 626
Read Iodine Clock Lab in packet pg 6
(get ONLY 80 ml each A & B)
WE will DEMO temp.
31.
32.
33. Assignment
• You need to read equilibrium pages in our
book
• Pg 601-604 & 612-619*ignore math for now, we’ll do that later
• DO pg 628 # 4, 5, 6, 8, 9 & study for a quiz
35. Use the worksheet as notes
Equilibrium Ch 17
Review rate… How is it measured?
2Zn + 2HCl --> H2 + 2ZnCl
Collision theory
Factors affecting rate
36. 1. gas given off -
Mg + 2HCl → McCl2 + H2(g)
2. insoluble product is formed -
AgNO3(aq) + KCl(aq) → KNO3(aq) + AgCl (s)
3. molecular product formed -
HCl(aq) + NaOH(aq) → NaCl(aq) + HOH (L)
DO YOU REMEMBER double displacement when there
was no reaction.
Most reactions are irreversible, that is, they go
to completion.
37. Equilibrium is the
• State at which the concentration of all
reactants and products remain constant.
– See pg 602
38. These equilibrium rxns. Are
reversible
A + B C + D
A + B C + D is forward rxn
C + D A + B is reverse rxn
If A + B is forming C + D as fast as C + D is
forming A + B, the reaction is at equilibrium.
The rate forward and backward is equal.
• NOTHING SAYS THAT THE
CONCENTRATIONS ARE EQUAL THOUGH!
39. Demo… shake clear--->blue
– WHAT AM I DOING TO MAKE IT GO?
– As we agitate it it becomes oxygenated. This shows reversibility
40. At equilibrium,
• it appears as if nothing is happening, but
reactants are still forming products and
products are forming reactants at the same
rate.
• NOTE: The concentration of P & R are not
necessarily equal. Long arrow points to
where you have more.
42. See worksheet graphs
– Graphically. See pg.
– Look carefully at the axis.
– Why do some come down? (green & purple)?
43. Test Tube & Straw Device
• When are we at equilibrium?
• PG 602: EQUILIBRIUM: a dynamic state
where the concentrations of the R and P
remain constant over time, as long as
conditions are not changed. (such as
pressure, temp, # particles)
44. Which do we have more of?
• H20 + HC2H3O2 H3O +
+ C2H3O2
-1
• H20 + HClO4 H3O +
+ ClO4
-1
45. • Discuss INDUSTRY. This is only 78.2%
complete, What if I want to make more
• H2+ I2 2HI
46. Homo Vs Hetero
• Homogenous all the same state of matter
• Heterogeneous P & R are in different states
of matter.
• Just and FYI in case you hear about these in
your reading.
47. Le Chateleir’s principle.
IF STRESS IS APPLIED TO A SYSTEM IN
DYNAMIC EQUILIBRIUM, THE SYSTEM
CHANGES TO RELIEVE THE STRESS.
“Things move in the direction that relieves the stress”
or
“the system rolls with the punches”
48.
49. Shifts (see worksheet)
• Concentration (soln. &gases only)
• More, concentrated will shift to make more.
• See Pg 612-613
• Removing a product always pulls the
reaction toward the product.
• “Removing something, pulls it that way!”
50. Temperature (soln & gases)
A + B + heat C
endothermic in the forward direction
A + B C + heat
exothermic in the forward direction
Adding heat favors the endothermic reaction. WHY?
51. DEMO
• 2NO2 (g) N2O4 (g) + heat
• N2O4 is clear
• The reverse reaction is
endothermic.
• LET’S TRY COOLING IT, HEATING IT
52. Pressure and volume (for gases only)
2A(g) + B(g) 2C(g)
• A change in pressure affects only a system
that has an unequal number of moles of
gaseous reactants and products.
• If you have more pressure the system will
adjust to take up less space. (pg 617)
58. Keq lecture pg 605-610, 620-
• Keq is a value the tells us which direction a
eversible reaction is favored. Toward
products or toward reactants.
• USE HANDOUT first in packet, then see ppt.)
– Book probs.pg 628… 18-21, 38-42, 65, 68, 59
– Know what Keq values mean!
– Be able to write expressions & do calculations
for Keq. (look ahead at Ksp)
63. • PRACTICE: 1. Calculate the Keq for
this reaction given the concentrations at a
constant temperature are as follows:
[NO2] = 8.8 M [NO] = 1.2 M [O2] = 1.6 M
2NO(g)+O2 (g) 2NO⇌ 2 (g)
64. • PRACTICE: 2. write the equilibrium
equation for this reaction:
2Na(s)+CuCl2 (aq) Cu⇌ (s)+ 2NaCl(aq)
65. • PRACTICE: 3. Calculate Keq for the
above reaction (practice 2) if the
concentrations are as follows:
[CuCl2] = 0.050M & [NaCl] = 0.50M
2Na(s)+CuCl2 (aq) Cu⇌ (s)+ 2NaCl(aq)
66. SO WHAT DO Keq VALUES MEAN?
• Recall that the equations for Keq is a fraction. So, if the value of Keq >
1 (greater than 1), then that means the numerator is larger than the
denominator. So which direction would be favored in the reaction?
___________ (product or reactant?)
• Keq >1 favors products of forward reaction
• Keq<1 favors reverse reaction (reactants)
74. You may see this. . . Ksp
(pg 622)
• This the the constant of solubility we write
these the same way.
75. Assignment
Worksheets & BOOK QUESTIONS on your
own paper to turn in tomorrow!
Practices in packet
Pg 629..: 18-21, 29, 31,
32, 33, 34, 38-42, 65, 68,
(59 challenge)
76. Work on problems
• Remainder of Quick reviews by tomorrow
• Rate /equilibrium quiz
• Still: Nuclear information