Notes: Unit 10 Kinetics and Equilibrium - Notes: Unit 10 Kinetics and Equilibrium . Name: KEY IDEAS...

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Regents Chemistry: Mr. Palermo

Notes: Unit 10 Kinetics and Equilibrium

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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: _______________________________________________________________________

Lesson 1: Collision Theory and Factors Affecting Rx Rate

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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

Objective:

Determine what factors affect the rate of reaction


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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


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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


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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.) Pb2+(aq) + S-2(aq) PbS(s)

b.) 2H2(g) + O2(g) 2H2O(l)

c.) N2(g) + 2O2(g) 2NO2(g)

d.) 2KClO3(s) 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

Lesson 2: Energy Changes in Chemical Reactions Review

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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 (110 - 60 =) 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

Objective:

Determine if a reaction is endo or exothermic

Use table I to determine the type of reaction


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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 (100 - 30 =) 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

+571.6kJ (endothermic)


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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

Lesson 3: Potential Energy Diagrams

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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

Objective:

Label potential energy diagrams and determine the type of diagram

represented


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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


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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 ______________________________


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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?

Lesson 4: equilibrium

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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

Types of Equilibrium (all occur in closed systems)

1. Physical (Phase) Equilibrium:

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

Objective:

Determine if a reaction is spontaneous

Determine if entropy increases or decreases in a reaction


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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

Changes in Entropy

ENTROPY AND STATES OF MATTER:


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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?


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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)

Lesson 5: Changing Equilibrium

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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

Objective:

Determine the shift in equilibrium when a stress is placed on a system


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How to determine equilibrium shifts

TRICK FOR EQUILIBIRUM 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:


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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)


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PRACTICE:

2CO(g) + O2(g) 2CO2(g) + 566 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) + 566 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

Notes: Unit 10 Kinetics and Equilibrium - Notes: Unit 10 Kinetics and Equilibrium . Name: KEY IDEAS...

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