Presentation on theme: "Kinetics and Equilibrium Chapter 15. I: Definitions Activation Energy: the minimum amount of energy needed to produce an activated complex Heat of Reaction:"— Presentation transcript:
I: Definitions Activation Energy: the minimum amount of energy needed to produce an activated complex Heat of Reaction: the amount of heat released or absorbed in a reaction Exothermic Reaction: a reaction that releases heat energy
Endothermic Reaction: a reaction that absorbs heat energy Activated Complex: the temporary, unstable, intermediate union of reactants Equilibrium: a dynamic chemical condition in which opposing reactions are proceeding at equal rates, producing an apparent constant condition.
II: Kinetics Chemical Kinetics deals with: 1) The rates of chemical reactions 2) The pathway by which the reaction occurs (how)
A Reaction Rate depends on several factors: in order for particles to react, they must collide with each other. Each collision must have enough energy and hit in the right direction- they must be effective.
The nature of the reactants The nature of the reactants: solutions of ionic solids react more quickly than non-ionic solids because ions are created. Polar reacts more quickly than non-polar.
The concentration of the reactants The concentration of the reactants: ↑ the concentration ↑ the rate of the reaction. For gases, an increase in pressure brings the particles closer together and acts like an increase in concentration.
Surface Area Surface Area: ↑ surface area ↑the rate of the reaction. A powdered substance will always react faster than a solid of the same amount.
The temperature of the system The temperature of the system: ↑ temperature ↑ the rate of a reaction. Why? Because more energy speeds up particles and increases the number of collisions. The more collisions, the more effective collisions, the faster the reaction.
Presence of a Catalyst Presence of a Catalyst: a substance that ↑ the rate of a reaction without changing itself. **Will not begin the reaction, it just lowers the required activation energy by providing a new pathway for reacting particles.
Roles of Energy in Reactions – Activation Energy: the minimum energy needed to cause a reaction to begin – Heat of Reaction: the difference between the potential energy of products and of the reactants Formula: ∆H = H products – H reactants
Exothermic Reaction Reactions that release energy. The potential energy of the products is lower than the potential energy of the reactants. The sign for the Heat of Reaction (∆H) is negative. Potential Energy Diagram:
Endothermic Reaction Reactions that absorb energy. The potential energy of the products is greater than the potential energy of the reactants. The sign for the Heat of Reaction (∆H) is positive. Potential Energy Diagram:
Catalyst Potential Energy Diagram with a Catalyst
Equilibrium Dynamic Equilibrium: When the forward and the reverse rates of a reaction are equal. Phase Equilibrium: In a closed container, when the rate of melting is equal to the rate of freezing; or when the rate of evaporation is equal to the rate of condensation. H 2 O (s) ↔ H 2 O (l) or H 2 O (l) ↔ H 2 O (g)
Solution Equilibrium: In a saturated solution (in a closed container), when all that can be dissolved in liquid has been reached, the rate of dissolving and the rate of recrystallizing are equal; C 12 H 22 O 11 (s) ↔ C 12 H 22 O 11 (aq) or CO 2 (g) ↔ CO 2 (aq) Chemical Equilibrium: when the forward and reverse reactions occur at equal rates. ↔ The products can become the reactants again. ** The concentration of the reactants and the products are kept constant but do not have to be equal!!!
Le Chatelier’s Principle Changes to concentration, pressure or temperature are known as applied stresses. This principle describes what happens to a system under such stresses.
Concentration Changes If there is an increase in concentration, the system will want to use up what has been increased, therefore the reaction will shift away from the increase. If there is a decrease in concentration the system will want to create more of what was removed, therefore the reaction will shift towards the decrease. At equilibrium, if you add more reactants, equilibrium is upset and must return to the value so equilibrium will shift to the product side of the reaction *put in more = shift to the other side *take out = remain on the same side
Pressure Changes **Effects GASES** a) Increase the pressure, the system will move to the side with the fewest # of molecules (count the coefficients). b) Decrease pressure, the system will move to the side with the most # of molecules (count the coefficients).
Temperature Changes If the system is exothermic: -an increase in temperature will shift to reactant side -a decrease in temperature will shift to product side If the system is endothermic: -an increase in temperature will shift to product side -a decrease in temperature will shift to reactant side
Effect of a Catalyst Since catalysts will increase both the forward and the reverse reaction, there is no change in equilibrium
Spontaneous Reactions One that occurs in nature under a given set of conditions. Gibbs Free Energy Change: ∆G tendency of a reaction to proceed to a minimum energy and a maximum entropy. Least energy, greatest disorder. Whether or not reactions proceed seems to depend on the balance of two basic principles:
The drive toward greater stability (reduced potential energy) -∆H The drive toward less organization (increased entropy) +∆S – Entropy: the amount of randomness or disorder in a system; the symbol for the change in entropy is ∆S ∆S solid > ∆S liquid > ∆S gas
Gibbs Free Energy Change Reaction: ∆G = ∆H - T∆S -∆G is Spontaneous For a reaction to occur spontaneously – more likely exothermic reactions with an increase in entropy.
-∆G = a spontaneous reaction +∆G = a non spontaneous reaction ∆G is 0 = a reaction at equilibrium.