3.  I can identify and describe the five factors that affect reaction rates.

4.  There are five general factors that affect the rate of reaction.

5.  The rate of reaction depends on the particular reactants and the complexity of the bonds that have to be broken and formed.  Reactions where the rearrangement of atoms is simple occur faster.  More complex reactions are much slower.

6.  Reactions that involve covalent bonds are usually slower because they are more complex.  The states of a reactant in a chemical reaction can also affect reaction rate.

7.  Reactions that involve gases will be the fastest, followed by those that involve liquids, and reactions that involve solids are the slowest.

8.  Most reactions occur more rapidly with increased temperatures.  The higher the temperature, the faster the molecules move.

9.  The faster molecules move, the more energy they have, and the more often they collide, which helps them react.  Therefore, increasing temperature increases the rate of reaction.

10.  The rate of reaction depends on concentrations of the reactants.  The higher the concentrations, the faster the reaction rate.  The reason for the increase is due to the fact that higher concentrations result in more collisions.

11.  The more surface area a reactant has, the greater the number of particles that are expose for reaction.  A larger surface area increases the frequency at which particle collide.  To increase the surface area of a solid, scientists use powders.

12.  A catalyst is a substance that increases the rate of a reaction without being used up.  Catalysts can be used over and over again.  Catalysts in the human body are called enzymes.

13.  Catalysts are not a reactant or a product.  They can be re-used after a chemical reaction occurs.  Catalysts increase reaction rates because they lower the activation energy for a reaction to occur.

14.  The lower the activation energy, the faster the reaction rate.  It’s not a good idea to speed up all chemical reactions.  Substances that slow a chemical reaction are inhibitors.

16.  I can describe the collision theory.  I can define an effective vs. ineffective collision.  I can define both kinetic and potential energy.  I can interpret an energy diagram.

17.  In order for a chemical reaction to occur, particles must come together.  Combining two substances means forcing their particles to hit or collide.  The collision theory states that molecules must collide to react.

18.  In order for a molecule, atom, or ion to react, it must collide with another molecule, atom, or ion or with the wall of a container.  If there were no collisions among reactants, there would be no reactions.  Not all collisions are effective in leading to the formation of products.

19.  Effective collisions lead to the formation of products.  Ineffective collisions do not lead to the formation of products.  If every collision lead to a reaction, the rates of reaction would be much faster.

20.  For a reaction to occur, particles must be in the correct position that allows the bonds to break and atoms to rearrange.  If the orientation is not right, the molecules bounce off each other.  Colliding particles must also have enough energy so that bonds can be broken and new bonds formed.

21.  Energy is required to break bonds that hold reactants together.  There are two types of energy: Kinetic and potential.  Potential energy is stored and kinetic energy is the energy of motion.

22.  Remember, energy can not be created or destroyed, only transformed.  Particles of matter have both potential and kinetic energy.  The energy to break bonds comes form the energy of the reacting particles.

23.  During a collision, kinetic energy is converted to potential energy as particles are formed, bonds are broken, and atoms rearranged.  Kinetic energy depends on the mass and velocity of a particle.  Particles must have a minimum amount of energy to react.

24.  Activation energy is the energy needed to start a reaction.  Energy diagrams show the changes in energy that occur during a chemical reaction.  The difference between the peak and reactants is the activation energy.

27.  An exothermic reaction releases heat.  Therefore, the products have less energy than the reactants.  An endothermic reaction absorbs heat.  Therefore the products have more energy than the reactants.

28.  The activated complex is a structure that forms in a reaction that is not a reactant or a product.  The activated complex exists where the energy is greatest.  Reactions occur at different rates because of energy.

29.  Particles must collide at the correct orientation to react.  Particles must also have enough energy to overcome the activation energy to react.  Reaction rates are directly related to activation energy because every reaction has a different activation energy.

31.  I can define free energy.  I can describe a spontaneous and nonspontaneous reaction.  I can define both entropy and enthalpy.  I can describe the conditions under which a spontaneous and nonspontaneous reaction occur.

32.  Many chemical and physical processes release energy that can be used to bring about other changes.  Free energy is the energy that is available to do work.

33.  Even though free energy is available, that does not mean it can be used efficiently.  Car engines are only 30% efficient.  That means only 30% of the free energy released from burning gasoline is used to move the car.

34.  The other 70% is lost to friction and heat.  Efforts are constantly underway to increase efficiency.  However, no mechanical process can be made 100% efficient.

35.  Energy can only be obtained from a reaction if the reaction actually occurs.  Not all balanced chemical reactions actually occur in nature.  Spontaneous reactions are reactions that occur naturally and that favor the formation of products at the specified conditions.

36.  All spontaneous reactions release free energy.  Nonspontaneous reactions are reactions that do not favor the formation of products at the specified conditions.  In most reversible reactions, one reaction is spontaneous and one is nonspontaneous.

37.  The terms spontaneous and nonspontaneous do not refer to how fast reactants go to products.  Some spontaneous reactions are extremely slow, and they look like they are nonspontaneous.

38.  Some reactions that are nonspontaneous at one set of conditions may be spontaneous at other conditions.  Changing the temperature or pressure can change a reaction to be spontaneous.

39.  Heat change accompanies most chemical and physical processes.  Both exothermic and endothermic reactions can be spontaneous.  Heat is also called enthalpy.  Enthalpy change is symbolized as H.

40.  To determine whether or not a reaction is spontaneous, we must consider heat and disorder.  The disorder of a system is called entropy.  The more disorder, the more entropy.

41.  The law of disorder states that processes move in the direction of maximum disorder or randomness.  The entropy of gasses is higher than the entropy of liquids.  The entropy of liquids is higher than the entropy of solids.

42.  Entropy increases when solid reactants form liquids.  Entropy also increases when liquid reactants form gases.  Entropy increases when a substance is divided into parts.

43.  Therefore, entropy increases when a solid is dissolved in water because the solute particles are more separated than the solid.  Entropy tends to increase in chemical reactions when the total number of product molecules is greater than the total number of reactant molecules.

44.  Entropy tends to increase when temperature increases.  As the temperature increases, the molecules move faster and faster, which increases disorder.

45.  The size and direction of heat changes and entropy changes determine whether a reaction is spontaneous.  In other words, whether it favors products and releases free energy.

46.  Releasing heat is a favorable condition for a reaction to be spontaneous.  Absorbing heat is an unfavorable condition.  Increasing entropy is a favorable condition for a reaction to be spontaneous.  Decreasing entropy is an unfavorable condition.

47.  Spontaneous reactions occur under the following conditions:  Heat is released and Entropy increases  The entropy increase is larger than the heat absorbed.  The heat released is larger than the decrease in entropy.

48.  Reactions will be nonspontaneous under the following conditions:  The heat absorbed is greater than the increase in entropy.  The heat released is less than the decrease in entropy.  Heat is absorbed and entropy decreases.

50.  Entropy is the measure of disorder of a system.  The symbol for entropy is S.  The units for S is J/K  The standard entropy of a liquid or solid substance at 25 C is designated S o.

51.  The pressure at S o for gases is 101.3 kPa.  Standard entropy changes can be calculated using the following equation: S o = S o (products) – S o (reactants)

52. What is the standard change in entropy for the reaction below? 2NO (g) + O 2(g)  2NO 2(g) S o for NO 2 = 240.5 x S o for NO = 210.6 x S o for O 2 = 205.0 x

54.  In every spontaneous process, some energy becomes available to do work.  This energy is called the Gibbs-free energy change ( G)  It is the maximum amount of energy that can be coupled to another process to do useful work.

55.  The change in Gibbs free energy is related to the change in entropy and the change in enthalpy.  The change in enthalpy is the change in heat.

56.  The equation for Gibbs Free Energy: G = H – T S H= Change in enthalpy S = Change in entropy T = Temperature in Kelvin

57.  All spontaneous reactions release free energy.  G for all spontaneous processes is negative.  All nonspontaneous reactions require work to make them happen.  Therefore, the G for nonspontaneous reactions is positive.

58.  You can calculate G if you know H and S.  However, if H and S are not known, you can still calculate G.

59. Equation: G = G (products) - G (reactants) ofof ofof

60.  Using G to calculate whether a reaction is spontaneous or not only works if the reactants and products are in their standard states.  A reaction that is nonspontaneous under one set of conditions may be spontaneous under another set.

62.  I can explain the term rate.  I can define reaction rates.  I can explain a reaction mechanism.  I can use a rate law to determine what would happen to the reaction rate when concentrations are changed.

63.  The speed at which chemical reactions occur depends on external conditions.  It’s important to know how fast chemical reactions occur and to understand factors that control speed.

64.  Knowing the speed at which food spoils, metal rusts, and dental fillings set is important.  Chemical kinetics is the chemistry concerned with the speed of reactions.  Rates measure change over time.

65.  Chemical changes in a chemical reaction do not occur all at once.  The reaction rate is the rate at which reactants disappear and products appear.  It is the change in concentration of reactants and products in a certain amount of time.

66.  Equation: concentration Avg. Rate = time

67.  The Greek letter delta means change in.  Reaction rates change throughout a chemical reaction so it is important to get an average.

68.  During chemical reactions, the rearrangement of atoms is often complicated.  Chemical reactions do not necessarily occur in one step.  Usually chemical reactions occur in several steps made of simple reactions.

69.  A reaction mechanism is a series of steps that lead from reactants to products.  Detailed reaction mechanisms describe the order in which bonds break and atoms rearrange during a reaction.

70.  The simple steps in a reaction mechanism are called elementary steps.  The product of the first step becomes the reactant of the second step.

71.  An intermediate product is a substance that is produces in one step, but used in a later step.  Determining reaction mechanisms is difficult and time consuming and they are determined by experimentation.

72.  A rate law is an equation that can be used to calculate the reaction rate for any given concentration of reactants.

73.  Equation: rate = k[A] x [B] y  [A] and [B] are molar concentrations.  K is a rate constant.

74.  K has a fixed value for a reaction at a particular temperature and is determined by experiment.  X and y are also determined by experiment.  Not all reactants appear in the rate law.

75.  If changing the concentration of a particular reactant does not change the rate, the reactant does not appear in the law.

76.  Example: NO 2(g) + CO (g)  NO (g) + CO 2(g) Rate = [NO 2 ] 2

77.  CO is not included in the rate law because it does not change the rate.  What would happen to the reaction rate if the concentration increased 5 times?

78.  What would happen if the concentration went from 4.0M to 3.0M? Hint: Take the final concentration and divide by the starting concentration, and then put that number in the rate law.