Kinetics & Equilibrium

Unit 11 Section 1 Obj. 1

Speeding up a Reaction

Kinetics is defined as:

Reaction rate is the change in amount of reactant/product per change in time.

Rate = Change in amount of substance

Change in time

Collision theory: Defines 3 conditions that MUST be met for a reaction to occur.

  • Reactants must ______.
  • Collisions must be at the correct ______.
  • Collisions must meet a minimum energy called ______for the reaction to occur.

How does collision theory relate to reaction rates?

There must be an increase in effective collisions in order to increase reaction rate!

(Diagrams from The chemistry that you need to Know K. Deters)

Activated complex (transition state): temporary state of the molecule at the top of the hill

Activation Energy(Ea): the minimum energy needed to get the reaction to occur

To calculate Ea: Subtract the energy of reactants from the energy of the activated complex

Ea = Etransition state - Ereactant

Heat of Reaction (H): Heat of the reaction; can be

  1. Endothermic: (H is positive); products have higher energy than the reactants
  2. Exothermic: (H is negative); products have lower energy than the reactants

To calculate: Subtract the energy of the reactants from the energy of the products

Hrxn = Hproducts - Hreactants

Unit 11 Section 2 Obj. 1 cont.

Factors that Affect Reaction Rate

1.Nature of Reactant: Some substances are just more reactive than others.

2.Surface area: As surface area increases, more effective collisions happen & the reaction rate______.

(Reactant particles must collide. Larger surface means more contact with each other. Greater collision frequency!

3.Concentrations of the Reactants: As reaction concentration increases, more effective collisions happen & the reaction rate ______.

(More reactants mean greater collision frequency !)

1M HCl is less concentrated than 3M HCl

4.Temperature: Molecules at a higher temperature have higher average ______so they move faster with greater collision frequency. ALSO, reactants must have minimal energy (activation energy) to collide. At higher temperatures, the kinetic energy is higher so it contributes more to the needed activation energy.

A rise in 10 C can double the rate of a reaction

Temperature increases NOT ONLY the collision frequency BUT IT ALSO helps the reaction meet its activation energy more effectively.

5.Catalyst: substance that ______the rate of reaction without being used up.

______are catalysts in the body.

Catalysts ______the activation energy by letting it proceed in a different way.

Lower Ea = faster reaction

Why are catalysts important?

Unit 11 Section 3 Obj. 2

What is Equilibrium?

What is happening in the picture?

A reaction at Equilibrium – decomposition (since there is only one reactant, it must decompose)

*note: the reversible reaction shown by the double arrow in the last diagram*

  • ______is a dynamic [ever changing] condition where rates of opposing processes are equal.
  • Can only take place in a closed container/space
  • Types of Equilibrium:
  • Physical Equilibrium (______)
  • Physical Equilibrium (______)
  • Chemical Equilibrium

Phase Equilibrium

  • Rate of one ______is equal to the rate of the opposing phase change.
  • Occurs when two phases exist at the same temperature.
  • Think about those heating/cooling curves and the plateaus where Melting/Freezing and Condensing/Vaporizing took place
  • : Ratevaporization = Ratecondensation

H2O (l)  H2O (g)

In an open container
All liquid molecules will eventually convert to a gas and leave the container
In a closed container
The liquid converts into a vapor but remains just above the remaining liquid (due to the closed lid of the container) creating vapor pressure.

Dynamic Equilibrium

Molecules initially escape from ______.

As the vapor builds up some vapor ______back to a liquid.

Eventually the rate of the forward reaction equals that of the reverse reaction.

Vapor Pressure

The pressure of the ______above a liquid

Factors That Affect Vapor Pressure

Intermolecular Forces

  • The forces between ______
  • The stronger the intermolecular forces the lower the vapor pressure and vice versa
  • Strong Forces such as ______and dipole forces are harder to break creating a ______vapor pressure
  • Weak forces such as London dispersion forces are easier to break creating a ______vapor pressure.

Temperature

  • As temperature increase, molecules have more ______to break their intermolecular forces between liquid molecules

Adding a solute to a solvent

  • Lowers the vapor pressure of a liquid
  • This is due to the solvent not being able to convert to a gas as often

Colligative Properties

  • Adding a ______(making a solution), lowers the vapor pressure which in turn
  • Lowers the Freezing Point of a solution
  • Raises the Boiling Point of a solution
  • The more solute added, the greater the change in freezing point and boiling point

Example

1 M NaCl (aq) vs 1 M C12H22O11(aq) [sugar]

NaCl breaks down into ions in water thus forming 1 M of Na1+ and 1 M Cl1--while sugar does not break down so NaCl and water has a higher boiling point and a lower freezing point

What Factors CANNOT Affect Vapor Pressure?

1)Changing the container size

2)Adding more of the solvent to the container

Solution (Physical) Equilibrium

  • Rate of ______= rate of ______
  • Meaning that the rate of the compound decomposing [separating] equals the rate ions reconnecting into ionic compounds as shown in the diagram to the right.
  • Occurs in saturated solutions

Equilibrium in Chemical Reactions

SO far, we have only talked about chemical reactions that proceed in one direction, from Reactants to Products and then STOP

A + B  C

  • Equilibrium is not reached if one of the products is withdrawn as quickly as it is produced and no new reactants are added.

BUT most chemical reactions are able to proceed in both directions under the appropriate conditions. (closed container) They are REVERSIBLE!

Example:

Fe3O4 (s) + 4 H2 (g)↔ 3 Fe(s) + 4 H2O(g)

WHEN DOES IT STOP? It doesn't.

Reversible Reactions

Represented by a double arrow between reactants and products 

In a closed system, as products are produced they will react in the reverse reaction until the ______of the forward and reverse reactions are ______.

Ratefwd = Raterev

This is called ______.

  • Rate depends on concentration
  • The forward rate of a reaction decreases over time
  • The reverse rate of a reaction increases over time
  • Eventually the 2 rates become equal

The Concept of Equilibrium

Once equilibrium is achieved, the amount of each reactant and product remains constant

CON CON REQUAL: Concentration of reactants and products are ______, not necessarily equal. Rates are Equal!

Depicting Equilibrium

  • A ______ identifies that a reaction is in equilibrium.

Unit 11 Section 4 Obj. 2 cont.

Writing Equilibrium Equations

The equilibrium expression is ratio of ______over ______.

Only use ______or ______ substances in the expression; never use pure solids or liquids

Writing an Equilibrium Expression

/ The equilibrium expression for this reaction would be… /
The lower case letter = coefficient
The upper case = chemical formula / K is the equilibrium constant

Practice: Write Equilibrium expressions for the following

  1. SnO2 (s) + 2CO (g)  Sn (s) + 2CO2 (g)
  1. CaCO3 (s)  CaO (s) + CO2 (g)
  2. Zn (s) + Cu2+ (aq)  Cu (s) + Zn2+ (aq)

What Does the Value of K Mean?

  • If K > 1, the reaction is product-favored; product predominates at equilibrium.
  • If K < 1, the reaction is reactant-favored; reactant predominates at equilibrium

Unit 11 Section 5 Obj. 3

Le Chatelier's Principle

  • Whenever a______is applied to a reaction at equilibrium, the reaction will ______its point of equilibrium to offset the stress.
  • Stresses include:

•Change Temperature

•Change Pressure

•Changes concentrations [product or reactant]

Change in Concentrations[of a Reactant or Product]

  • Adding a reactant or product shifts the equilibrium away from the ______.
  • Increase product – makes more reactant
  • Increase reactant – makes more product
  • Removing a reactant or product shifts the equilibrium towards the ______.
  • Decrease product – makes more product
  • Decrease reactant – makes more reactant

Example: The Haber process [Change in Concentration]

N2 (g) + 3H2 (g)  2Nh3 (g)

If H2 is added while the system is at equilibrium, the system must respond to counteract the added H2

•That is, the system must consume the H2 and produce products until a new equilibrium is established.

•Therefore, [H2] and [N2] will decrease and [NH3] increases.

Effects of Volume & Pressure on Gaseous Equilibrium

Volume & pressure have a inverse relationship

Increasing pressure (decrease in ______) favors the direction that has fewer moles of gas.

Decreasing ______(increase in volume) favors the direction that has greater moles of gas.

In a reaction with the same number of product and reactant moles of gas, pressure has no effect.

Example: Increase Volume

N2O2 (g)  2NO2 (g)

If the volume of the container increases while the system is at equilibrium(pressure of the gas decreases), the system must respond to counteract the decreased pressure

•That is, the system will shift to the right ;more moles of product will form until a new equilibrium is established.- this raises the pressure back to where it originally was

•Therefore, [NO2] will increase and [N2O4] decreases.

Example: Increase Pressure

N2O2 (g)  2NO2 (g)

An increase in pressure (by decreasing the volume) favors the formation of colorless N2O4.

•The instant the pressure increases, the system is not at equilibrium and the concentration of both gases has increased.

•The system moves to reduce the number moles of gas: shifts left

•A new equilibrium is established!

Effects of Temperature Change

The equilibrium constant is temperature dependent.

  • For an endothermic reaction, H > or H is positive (+); heat can be considered as a reactant.
  • For an exothermic reaction, H < 0 or H is negative (-); heat can be considered as a product.

Removing heat (i.e. cooling the vessel), favors towards the decrease:

  • if endothermic, cooling favors the reverse reaction,
  • If exothermic, cooling favors the forward reaction.

Adding heat (i.e. heating the vessel) favors away from the increase:

  • If endothermic, adding heat favors the forward reaction,
  • If exothermic, adding heat favors the reverse reaction.

Example: Various Changes

ΔH =+50KJ; If the temperature is lowered, the reaction will shift to the left and the solution will change to pink making more Co(H2O)62+

Example: The Haber process

Stress Added / Shift / Stress Added / Shift
[a] Increase [N2] / [e] Increase pressure
[b] Decrease [H2] / [f] Increase volume
[c] Increase [NH3] / [g] Increase temperature
[d] Decrease [NH3] / [h] Decrease temperature

OVERVIEW: Equilibrium shifts due to stresses:

Concentration increase shifts ______ from increase

Concentration decrease shift toward decrease

Increase [] in pressure shifts in direction of ______.

Decrease [] in pressure shifts in direction of more gas molecules

Increase [] in temperature favors endothermic reaction; shift ______from heat

Decrease [] in temperature favors exothermic reaction shift towards heat

Effect of Catalyst:

Addition of ______increases the rate of both the forward and reverse reactions.

  • There is no change in concentrations but equilibrium is reached more rapidly.

Does not changethe value of K [aka no shift to reaffirm equilibrium]

The Haber Process: Industrial Production of Ammonia

Application of LeChatelier’s Principle: N2 (g) + 3 H2 (g) 2 NH3 (g) + 92 kJ

Increase pressureShift Right

Decrease Temp Shift Right

Remove NH3 add N2 and H2Shift Right

Maximum yields of NH3 occurs under high pressures, low temperatures and by constantly removing NH3 and adding N2 & H2