Kinetics and Equilibrium Kinetics Kinetics is the part

Kinetics and Equilibrium

Kinetics • Kinetics is the part of chemistry that examines the rates of chemical reactions. • Collision theory is the concept of kinetics that says in order for a reaction to take place reactant particles must collide.

Factors Affecting Rate of Reaction • Nature of the Reactants – reactions involve the breaking and formation of bonds. Covalently bonded substances react slower than ionic substances. Covalently bonded substances have more bonds and need greater energy to be broken.

Factors Affecting Rate of Reaction • Concentration – most chemical reactions will occur at a faster rate if the concentration of one or more reactants is increased. KMT would predict this.

Factors Affecting Rate of Reaction • Surface area – more surface area of a substance exposed will create more chances for reactant particles to collide.

Factors Affecting Rate of Reaction • Pressure – pressure has little to no effect on solids and liquids. Increased pressure on a gas will cause gas particles to become more concentrated. Increased concentration increased rate of reaction.

Factors Affecting Rate of Reaction • The presence of a Catalyst – catalyst are substances that increase the rate of a reaction by providing a different and easier pathway for a reaction (lowering the activation energy). Catalysts take part in a reaction, but they are unchanged when the reaction is complete.

Factors Affecting Rate of Reaction • Temperature – increased temp. will cause the molecules of a substance to move faster. Particles moving faster plus increased collisions will increase the possibility of the rate of reaction to be faster.

Collision Theory “You can’t react if you don’t collide!” Bumper cars – need to get bumped out of your seat to have a reaction!

Collision Theory “You can’t react if you don’t collide!” 1. Proper orientation Side collisions just move you. You need to hit head on for the bonds to be broken and new bonds be formed A 2 + B 2 2 AB

Collision Theory “You can’t react if you don’t collide!” 2. Speed Particles need to have enough speed when they collide. They need that MINIMUM amount of energy Activation Energy Ea

Activation Energy Ea Amount of energy needed to “kick -start” a reaction Example: A match will not light on its own, the match must have an addition of mechanical energy (striking the box) so it will light

Potential Energy Diagram Activation Energy

Potential Energy Diagram • Shows the potential energy change that occurs during a chemical reaction. • A reaction will only occur if reactants have enough energy to collide effectively. • Reactants must collide “properly” in order for an activated complex to form.

Endothermic Potential Energy Diagram • Product has more potential energy than the reactants. • Gain in energy • ∆H is positive

Exothermic Potential Energy Diagram • Product has lower potential energy than reactants. • There was a release of energy • ∆H is negative

Equilibrium • Equilibrium – the forward and reverse reactions occur at the same rate • Equilibrium equation uses a double arrow • State of balance between two opposite processes taking place at the same rate. • Can only occur in a system in which neither reactants nor products can leave the system.

• The quantities of the reactants and products may not always be equal; they are usually not. • Equilibrium is involved in physical processes such as phase changes and dissolving.

Phase Equilibrium • Can exist between the solid and liquid phases of a substance. • At the melting point of a solid or the freezing point of a liquid phase • Melting can occur at same time as freezing • Also, in a closed system, rate of evaporation can equal rate of condensation.

Solution Equilibrium • Solids in liquids exist in equilibrium in a saturated solution • When the rate of dissolving and recrystallization are equal, equilibrium exists and the solution is saturated. • Temperature increasing or decreasing will change the rate of the forward or reverse reaction until a new equilibrium is reached.

Chemical Equilibrium • When you mix reactants at first and there are no products present, only the forward reaction can occur. • As time progresses, the reactants decrease and the forward reaction slows, while the concentration of the products increase. • This causes the rate of the reverse reaction to increase.

Chemical Equilibrium • The forward reaction will continue to slow as the reverse reaction proceeds until equilibrium is met. • That is chemical equilibrium.

Le Chatelier’s Principle • If there is any change in temperature, concentration or pressure on an equilibrium system it is called a stress. • If there is a change in concentration where an addition to one or more reactants occurs, the rate of the forward reaction would increase. • Addition of concentration on the reactant side would cause an increase of product.

Concentration Changes • Since the forward reaction is favored, the equilibrium is said to “shift to the right. ” • As more products form, the rate of the reverse reaction increases until equilibrium is reached (when the concentration of products equals the concentration of reactants) • Equilibrium is when the rate of the forward reaction is equal to the rate of the reverse reaction.

Concentration Changes • CH 4 + H 20 3 H 2 + CO • If CH 4 is increased, the rate of the “forward” reaction will increase. • If CH 4 is reduced, the rate of the forward reaction will decrease. ***When the concentration of a substance decreases, the reaction that produces that substance is favored.

Temperature Changes • Le Chatelier’s principle says that a system will undergo changes to reduce a stress. • When temperature is raised both the forward and reverse reaction is favored, but not equally. • When heat is added to an endothermic reaction, the reactants will be favored.

Temperature Changes • Since heat is a product of the endothermic reaction, the reverse reaction absorbs some of the heat. • When the reverse exothermic reaction absorbs the heat, the reverse reaction becomes favored and more reactants are formed. • 4 NH 3(g) + 5 O 2(g) <==> 4 NO(g) + 6 H 20(g) + heat

Temperature Changes • The opposite when taking heat away. • Taking heat away from the reaction will cause a stress on the product side. The stress will favor the forward reaction causing more product to form. • 4 NH 3(g) = 5 O 2(g) <==> 4 NO(g) + 6 H 20(g) + heat

Pressure Changes

Pressure Changes • Pressure changes have NO EFFECT on rate of reaction when solids and liquids are involved. • HOWEVER, increasing pressure on gas increases the concentration of the gas.

Pressure Changes • Increasing pressure favors the side with fewer gas molecules and vice versa. • Pressure change will not affect a system that has no gas molecules, or that has equal # of gas molecules on both sides. • Catalysts have no effect on system at equilibrium.

Important Concepts to Remember: • The rates of reactions can be changed by changes in concentration, temperature, and a catalyst. • At equlibrium: – Rate of forward reaction = Rate of reverse reaction – Concentration of reactants and products are constant, but the amounts are not equal.

Entropy and Enthalpy • There are two fundamental tendencies in nature that will tell us whether or not chemical or physical changes will occur. • These two fundamental tendencies are entropy and enthalpy.

Enthalpy • Enthalpy says that there is a tendency for nature to change to a state of lower energy. • Exothermic reactions move toward lower energy state • The reactants release energy and the products have less potential energy than the reactants.

Enthalpy • At any given temperature, the particles in a system are more likely to collide with enough energy in the exothermic reaction than in the endothermic reaction. • With this energy change we can expect reactions to go in the exothermic reaction.

Enthalpy Diagrams Exothermic reactions energy course of reaction

Energy Level Diagrams Exothermic reactions energy reactants course of reaction

Energy Level Diagrams Exothermic reactions energy reactants products course of reaction

Energy Level Diagrams Exothermic reactions energy reactants energy given out ∆H is negative products course of reaction

Energy Level Diagrams Endothermic reactions energy course of reaction

Energy Level Diagrams Endothermic reactions energy reactants course of reaction

Energy Level Diagrams Endothermic reactions energy products reactants course of reaction

Energy Level Diagrams Endothermic reactions energy products reactants energy taken in ∆H is positive course of reaction

Summary Table Exothermic reactions Endothermic reactions

Summary Table Exothermic reactions Endothermic reactions Energy is taken in Energy is given out to from the surroundings

Summary Table Exothermic reactions Endothermic reactions Energy is taken in Energy is given out to from the surroundings ∆H is negative ∆H is positive

Summary Table Exothermic reactions Endothermic reactions Energy is given out to the surroundings Energy is taken in from the surroundings ∆H is negative ∆H is positive Products have less Products have more energy than reactants

∆H How much energy is given out or taken in? • Energy is needed to break chemical bonds • Energy is given out when bonds are made ∆H is the difference between the energy needed to break the bonds in the reactants, and the energy given out when new bonds are made in the products

Entropy • Entropy is the measure of the disorder or randomness of a system. • Nature has a tendency to change to a state of disorder. • The greater the disorder the higher the entropy. • From “neat” to “messy”

Entropy • We know that particles need to collide in a “special” way to form a regular arrangement of molecules. • This regular arrangement is highly organized. • But particles can collide to produce more disorder. • It is expected for reactions to go from lower entropy (great order) to high entropy (greater disorder).

Entropy • We expect reactions to go in the direction of high entropy or great disorder. • When we look at physical state, we expect solid species which are highly organized to move to the liquid phase which is less organized or has a higher entropy and then to a gas.

Entropy • For chemical changes, compounds represent great order or low entropy than that of free elements from which they are composed • Which room has higher entropy?

The Equilibrium Expression • Equilibrium Expression is a mathematical expression that shows the relationship of reactants and products in a system at equilibrium. • The concentration of products over concentration of reactants expressed in moles per liter. • Each concentration is raised to the power of its coefficient in the balanced equation.

Equilibrium Expression • The equilibrium constant is represented by (Keq) which remains the same for a particular reaction at a specified temperature. • There are four steps to writing equilibrium expressions

Equilibrium Constant • 1. Write the balanced equation for the system. • 2. Place the products as factors in the numerator of a fraction and the reactants as factors in the denominator. • 3. Place a square bracket around each formula. The square bracket means molar concentration. • 4. Write the coefficient of each substance as the power of its concentration. The resulting expression is the equilibrium expression, which should be set to equal to the Keq for that reaction.

Sample Calculation • Write the equilibrium expression for the equilibrium system of nitrogen, hydrogen, and ammonia. • Reactants: N 2, H 2 Products: NH 3 • N 2(g) + 3 H 2(g) 2 NH 3(g) + heat • Keq = [NH 3]2 [H 2]3[N 2]

Equilibrium Expression • Changes in concentration will not cause a change in the equilibrium constant, nor will the addition of a catalyst cause any change. • Only a change in temperature will affect the value of Keq • If the value of Keq is large, the products are present in a larger concentration than reactants. The products are favored. • If the value is small, the opposite is true, reactants are favored.

Equilibrium Expression Remember: only temperature changes will change the equilibrium constant! For exothermic reactions: an increase in temp will decrease K and vice versa. For endothermic reactions, an increase in temp will increase K and vice versa.