Notes: Unit 10 Kinetics and Equilibrium

Transcription

1 Name: Regents Chemistry: Mr. Palermo Notes: Unit 10 Kinetics and Equilibrium

2 Name: KEY IDEAS Collision theory states that a reaction is most likely to occur if reactant particles collide with the proper energy and orientation. (3.4d) The rate of a chemical reaction depends on several factors: temperature, concentration, nature of reactants, surface area, and the presence of a catalyst. (3.4f) Some chemical and physical changes can reach equilibrium. (3.4h) At equilibrium the rate of the forward reaction equals the rate of the reverse reaction. The measurable quantities of reactants and products remain constant at equilibrium. (3.4i) LeChatelier s principle can be used to predict the effect of stress (change in pressure, volume, concentration, and temperature) on a system at equilibrium. (3.4j) Energy released or absorbed by a chemical reaction can be represented by a potential energy diagram. (4.1c) Energy released or absorbed during a chemical reaction (heat of reaction) is equal to the difference between the potential energy of the products and the potential energy of the reactants. (4.1d) A catalyst provides an alternate reaction pathway, which has a lower activation energy than an uncatalyzed reaction. (3.4g) VOCABULARY For each word, provide a short but specific definition from YOUR OWN BRAIN! No boring textbook definitions. Write something to help you remember the word. Explain the word as if you were explaining it to an elementary school student. Give an example if you can. Don t use the words given in your definition! Reaction Rate: Entropy: Potential Energy: Catalyst: Activation Energy: Activated Complex: Spontaneous Reaction: LeChatelier s Principle:

3 Lesson 1: Collision Theory and Factors Affecting Rx Rate Objective: Determine what factors affect the rate of reaction KINETICS: Study of the RATE or SPEED at which REACTIONS occur A Reaction is the BREAKING and REFORMING of BONDS to make entirely new compounds as products. EFFECTIVE COLLISIONS: In order for a reaction to occur, reactant PARTICLES MUST COLLIDE (effectively) with the following: 1.) 2.) Example: H2 + I2 2HI 1

4 Lesson 1: Collision Theory and Factors Affecting Rx Rate Factors Affecting Reaction Rate SIX FACTORS that affect the rate of reaction by changing the number of effective collisions that take place between particles. The MORE EFFECTIVE COLLISIONS, THE FASTER THE REACTION! 1. TYPE OF SUBSTANCE: substances react Easily break into IONS when you dissolve them. Example: AgNO3 (s) Ag + + NO3 - substances react Requires more energy/time to break bonds Example: H2 (g)+i2 (g) 2 HI (g) 2. CONCENTRATION: Concentration reaction rate (speed) a. More particles increases chance of effective collisions 2

5 Lesson 1: Collision Theory and Factors Affecting Rx Rate 3. TEMPERATURE: temperature Reaction Rate: a. Increases # of effective collisions Reactants have more energy when colliding 4. PRESSURE (GASES ONLY) pressure, reaction rate (affects GASES ONLY!) a. Due to an increase in concentration 5. SURFACE AREA: in surface area the reaction rate. a. Due to more exposed particles that can react (more effective collisions) 6. CATALYST: Substance that rxn rate without being consumed in the rxn 3

6 Lesson 1: Collision Theory and Factors Affecting Rx Rate SUMMARY: Ionic solutions have faster reactions than molecule compounds. (bonding) Temp. Rate conc. rate surface area rate Pressure rate, P rate Catalysts speed up reactions. PRACTICE: At room temperature which reaction would be expected to have the fastest reaction rate? a.) Pb 2+ (aq) + S -2 (aq) b.) 2H2(g) + O2(g) c.) N2(g) + 2O2(g) d.) 2KClO3(s) PbS(s) 2H2O(l) 2NO2(g) 2KCl(s) + 3O2(g) PRACTICE: Under what conditions will the rate of a chemical reaction always decrease? a.) The concentration of the reactants decreases and the temp decreases b.) The concentration of the reactants decreases, and the temp increases c.) The concentration of the reactants increases and the temp increases d.) The concentration of the reactants increases, and the temp increases CHECK YOUR UNDERSTANDIND: Given the reaction: Zn(s) + 2HCl(aq) ZnCl2 + H2(g) The reaction occurs more slowly when a single piece of zinc is used than when the same mass of powdered zinc is used. Why does this occur? a.) The powdered zinc is more concentrated b.) The powdered zinc has a greater surface area c.) The powdered zinc requires less activation energy d.) The powdered zinc generates more heat energy 4

7 Lesson 2: Energy Changes in Chemical Reactions Review Objective: Determine if a reaction is endo or exothermic Use table I to determine the type of reaction HEAT OF REACTION ΔH: The amount of HEAT ENERGY LOST or GAINED throughout a REACTION ΔHheat of reaction = Hproducts - Hreactants (ΔH = enthalpy) TYPES OF CHEMICAL REACTIONS: ENDOTHERMIC REACTIONS: Heat is by Energy stored in chemical bonds of products ΔH is (+) A + B + ENERGY ---> C + D EXAMPLE: Reaction A + B C If HA =40kJ and HB =20kJ, then reactants have a total of 60kJ If HC =110kJ, then ( =) 50kJ of heat must have been absorbed by the reactants. Rewritten: A + B + 50kJ C Total energy on both sides are equal (law of conservation of energy) EXAMPLE from TABLE I 1

8 Lesson 2: Energy Changes in Chemical Reactions Review EXOTHERMIC REACTIONS: Heat is as a ΔH is (-) More stable reaction Spontaneous A + B ---> C + D + ENERGY Example: Reaction A + B C If HA =60kJ and HB =40kJ,then reactants have a total of 100kJ If HC =30kJ, then ( =) 70kJ of heat must have been released as a product. Rewritten: A + B C + 70kJ Total energy on both sides are equal (law of conservation of energy) EXAMPLE from TABLE I Example: Reverse reactions on Table I What is the ΔH of the following reaction? Is this exothermic or endothermic? 2H2O(l) 2H2(g) + O2(g) ***For reverse reactions switch signs of ΔH kJ (endothermic) 2

9 Lesson 2: Energy Changes in Chemical Reactions Review PRACTICE: What is the heat of reaction (ΔH) of the following? Is this exothermic or endothermic? 2H2(g) + O2(g) 2H2O(l) CHECK YOUR UNDERSTANDING: Fill in the table using table I 3

10 Lesson 3: Potential Energy Diagrams Objective: Label potential energy diagrams and determine the type of diagram represented POTENTIAL ENERGY DIAGRAMS: Shows the change in potential (stored) energy during a chemical reaction Types of Potential Energy Diagrams: EXOTHERMIC POTENTIAL ENERGY DIAGRAM: than reactant side WHY? Energy is released as a product, so the net amount of potential energy decreases. (-) ΔH ENDOTHERMIC POTENTIAL ENERGY DIAGRAM than reactant side WHY? Energy is absorbed by the reactants, so the net amount of potential energy increases. (+) ΔH 1

11 Lesson 3: Potential Energy Diagrams Parts of the Potential Energy Diagram You must be able to label these so label the diagrams in your notes as we go through this!!! ACTIVATION ENERGY: Minimum energy required for a reaction to occur (energy needed to get over the hill) ACTIVATED COMPLEX: Highest energy point of reaction Temporary Where bonds are broken and reformed 2

12 Lesson 3: Potential Energy Diagrams HEAT OF REACTANTS/PRODUCTS: Amount of potential energy possessed by the reactants and products HEAT OF REACTION ΔH Amount of energy lost/gained in a reaction SUBTRACT heat of products minus reactants ΔH = H(products) H(reactants) EFFECTS OF ADDING A CATALYST the activation energy (energy needed to start the reaction) Reaction occurs 3

13 Lesson 3: Potential Energy Diagrams PRACTICE: Is this an endothermic or exothermic reaction? How do you know? PRACTICE: Which arrow represents the activation energy of the forward reactants? CHECK YOUR UNDERSTANDING: Is the ΔH positive or negative for this reaction? 4

14 Lesson 4: equilibrium Objective: Determine if a reaction is spontaneous Determine if entropy increases or decreases in a reaction EQUILIBRIUM: The of the forward reaction is to the rate of the reverse reaction Equilibrium is Dynamic (in constant motion) Equilibrium is represented by double arrow 1. Physical (Phase) Equilibrium: Types of Equilibrium (all occur in closed systems) Rate of forward phase change equals rate of reverse phase change Ratevaporizing = Ratecondensing Ex. Water is vaporizing at the same rate it is condensing 2. Chemical Equilibrium : The rate of the forward reaction is equal to the rate of the reverse reaction Ex. The Haber Process 1

15 Lesson 4: equilibrium 3. Solution Equilibrium: The rate of dissolving equals the rate of precipitating. Ex. Will a reaction happen on its own once it s started? SPONTANEOUS REACTIONS: A reaction that happens on its own once initiated (lower activation energy) Increase in Recall. Entropy The randomness (disorder) of the system. The More substances the more entropy The higher the temperature the more entropy ENTROPY AND STATES OF MATTER: Changes in Entropy 2

16 Lesson 4: equilibrium EXAMPLE: (determine if there is an increase or decrease in entropy) 1. KClO3(s) 2KCl(s) + 3O2(g) 2. KCl(l) KCl(s) 3. CO2(s) CO2(g) PRACTICE: (determine if there is an increase or decrease in entropy) 1. H + (aq) + C2H3O2 - (aq) HC2H3O3(l) 2. H + (aq) + OH (aq) H2O(l) PRACTICE: What point on the heating curve has the most entropy? PRACTICE: Which reaction will occur spontaneously? 3

17 Lesson 4: equilibrium PRACTICE: According to table I which reaction will occur spontaneously? a. N2(g) + 2O2(g) 2NO2(g) b. 2H2(g) + O2(g) 2H2O(g) CHECK YOUR UNDERSTANDING: (determine if there is an increase or decrease in entropy) 1. H2(g) + Cl2(g) 2HCl(g) 2. H2O(g) H2O(s) 4

18 Lesson 5: Changing Equilibrium Objective: Determine the shift in equilibrium when a stress is placed on a system LE CHATELIER s PRINCIPLE: If a system at equilibrium is subjected to a STRESS, the equilibrium will shift in the direction that relieves that stress Causes a change in concentration of both the reactants and products until the equilibrium is re-established. TYPES OF STRESS: Concentration temperature pressure for gases 1

19 Lesson 5: Changing Equilibrium TRICK FOR EQUILIBIRUM SHIFTS: How to determine equilibrium shifts When you a stress the equilibrium shifts from the stress When you (take away) a stress the equilibrium will shift back the decrease to replace it. CONCENTRATION: TEMPERATURE: 2

20 Lesson 5: Changing Equilibrium PRESSURE (GASES ONLY): CATALYST: on equilibrium because both the forward and reverse reactions will be affected equally (both will speed up). EXAMPLE: The Haber Process N2 (g) + 3 H2 (g) 2 NH3 (g) + heat a) [N2] shift towards products (right) b) [H2] shift towards reactants (left) c) [NH3] shift towards reactants (left) d) [NH3] shift towards products (right) e) pressure shift towards products (right) f) pressure shift towards reactants (left) g) temperature shift towards reactants (left) h) temperature shift towards products (right) 3

21 Lesson 5: Changing Equilibrium PRACTICE: 2CO(g) + O2(g) 2CO2(g) kj 1. If concentration of CO is increased what direction will the equilibrium shift? 2. What happens to the concentration of CO2? 3. What happens to the concentration of O2? PRACTICE: 2CO(g) + O2(g) 2CO2(g) kj 1. If O2 is removed, what direction does the equilibrium shift? 2. What happens to the concentration of CO? 3. What happens to the concentration of CO2? CHECK YOUR UNDERSTANDING: Given the equation representing a reaction at equilibrium: N2(g) + 3H2(g) <==>2NH3(g) What occurs when the concentration of H2(g) is increased? (1) The equilibrium shifts to the left, and the concentration of N2(g) decreases. (2) The equilibrium shifts to the left, and the concentration of N2(g) increases. (3) The equilibrium shifts to the right,and the concentration of N2(g) decreases. (4) The equilibrium shifts to the right, and the concentration of N2(g) increases 4