Energy+effects+&+reaction+rates


 * ENERGY EFFECTS **

Chemical changes are usually accompanied by some form of energy change. Some example of energy changes during chemical reactions are the following: i. Heat energy and light energy are released when magnesium burns in air, or when iron or aluminium filings are sprinkled into a gas flame or when wood burns. ii. Heat energy from the combustion of petrol in car engines is converted to mechanical energy. iii. Many reactions in this course release heat energy: the reaction mixture become hotter and heat is transferred to the surroundings. Some other reactions absorb energy: the reaction mixture become colder and takes heat from the surroundings. Many reactions are carried out for the sole purpose of obtaining the energy they release and the chemical products of the reaction may be of little or no value. Energy is neither created nor destroyed; it can only be transformed from one form to another. So if heat energy is released during a chemical reaction, it must have been present in some other form in the reactants. Energy released from reactions when reactants have a higher chemical energy than the products. The chemical energy of a substance is a type of potential energy stored within the substance.
 * CHEMICAL CHANGES IN CHEMICAL REACTIONS **
 * Exothermic and Endothermic Reactions **

** Enthalpy ** is the name given to the stored chemical energy or the heat content of a substance, denoted by the symbol, H. If enthalpy decreases within a reaction a corresponding amount of energy is released to the surroundings. Consider Methane burning in air: CH4(g) + 2O2(g) ® CO (g) + 2H2O (g) + energy i. heat given out corresponds to the decrease in enthalpy. ii. the enthalpy difference between the reactants and the products is equal to the amount of energy released to the surroundings. **__ ENDOTHERMIC REACTION: __** The enthalpy changes that occur in a reaction that absorbs energy, results in the products having a greater enthalpy than the reactants. Carbon disulfide can be manufactured according to the following equation: 4C (s) + 8S (g) ® 4CS2 (g)
 * __ EXOTHERMIC REACTION: __**
 * Exothermic ** - a reaction in which energy is released to the surroundings - the enthalpy is lower for the products than the reactants. The reaction mixture becomes hotter and transfers heat to the surroundings.

i. enthalpy of the products is greater than the reactants. ii. the enthalpy difference between the reactants and the products is equal to the amount of energy absorbed from the surroundings. The enthalpy difference between the reactants and the products is equal to the amount of energy absorbed from the surroundings.

1. For each of the following processes: • State whether the process would be exothermic or endothermic. • Indicate whether the reactants or the products have higher enthalpy. • Identify whether heat would be absorbed or released to the surroundings. a. Burning kerosene _   _    _    b. Mixing solid Ba(OH)2 and solid NH4SCN. The following reaction takes place and the reaction mixture becomes colder. Ba(OH)2 (s) + 2NH4SCN (s) ® Ba2+ (aq) + 2SCN- (aq) + 2NH3 (aq) + 2H2O (l) _   _    _    c. Melting ice at 0 oC. _   _    _    d. Condensing steam at 100 oC. _   _    _    2. In the evaporation of water the enthalpy of the water increases. In view of this fact, explain how perspiration helps the human body to keep cool. _____   _    _____    _    _____    3. In an explosion using dynamite to destroy a city building, describe the following: a. The enthalpy change in the reacting system. _____   _____    _____    b. The energy change in the surroundings. _____   _____    _____   Exothermic or endothermic reactions are associated with a decrease or increase in the enthalpy of the substance involved. The enthalpy of a substance represents the stored chemical energy of the substance involved and depends on the chemical bonding present in the substance.
 * Endothermic ** - a reaction in which energy is absorbed from the surroundings - the enthalpy of the products is greater than that of the reactants. The reaction mixture becomes colder and takes heat from the surroundings.
 * Do Exploring Chemistry Experiment 14 pg. 74 **
 * PROBLEM SET #1 **
 * STORED CHEMICAL ENERGY **

When chemical bonds are broken: i. the breaking of chemical bonds absorbs energy. ii. the formation of chemical bonds releases energy.

The quantity of energy released when a bond forms is equal to the quantity of energy absorbed when such a bond is broken. When a chemical bond is broken, an increase in enthalpy in the system occurs. 2H2 (g) + O2 (g) ® 2H2O (g) i. two H - H bonds and one O = O are broken. ii. four O - H bonds are formed. iii. energy is released from forming the four O - H bonds. iv. the products have a lower enthalpy than the reactants - exothermic reaction. Chemical reactions usually involved the breaking and making of chemical bonds. If the bonds formed are weaker or fewer than those broken, energy will be absorbed. If the bonds formed are stronger or more numerous than those broken, energy will be released.
 * PROBLEM SET #2 **

1. Compare the enthalpies of the following: a. A hydrogen molecule and two separate hydrogen atoms. _____   b. A hydrogen fluoride molecule and separate fluorine and hydrogen atoms. _   2. Identify the bonds which are broken and those which are formed in the following reactions. a. N2 (g) + 3H2 (g) ® 2NH3 (g) b. NH4NO3 (s) ® N2 (g) + 2H2O (g) 3. The reactions between hydrogen gas and chlorine gas to form hydrogen chloride is exothermic. Compare the average energy needed to decompose hydrogen and chlorine molecules with that released when an HCl molecule is formed between hydrogen and chlorine atoms. The symbol D H represents the heat of reaction or change in enthalpy.
 * DESCRIBING ENERGY CHANGES IN CHEMICAL REACTIONS **

D H = H(products) - H(reactants) heat of enthalpy of enthalpy of reaction products of reactants

D H for a //exothermic// reaction has a //negative// value. D H for a //endothermic// reaction has a //positive// values.

// Enthalpy Changes for Exothermic and Endothermic Reactions // • Exothermic reaction • Endothermic reaction • D H is negative • D H is positive • Heat released to surroundings. • Heat absorbed from surroundings.


 * Endothermic ** - the final enthalpy is greater than the initial enthalpy and D H is positive; gain in enthalpy is accompanied by the transfer of energy from the surroundings to the system.
 * Exothermic ** - the final enthalpy is less than the initial enthalpy and D H is negative; loss in enthalpy is accompanied by the release of energy from the system to the surroundings.

i. H2O (g) ® H2 (g) + ½ O2 (g) D H = +242 kJ  242 kJ of heat is absorbed in the decomposition of one mole of steam. ii. S (s) + O2 (g) ® SO2 (g) D H = -242 kJ  242 kJ of heat is released when one mole of S reacts with oxygen to form sulphur dioxide. The heat term is written as a product in exothermic reaction and as a reactant in a endothermic reaction.

a. C (s) + H2O (g) + 131 kJ ® CO (s) + H2 (g) Endothermic reaction; D H = +131 kJ b. 2S (s) + 3O2 (g) ® 2SO3 (g) + 792 kJ Exothermic reaction; D H = -792 kJ


 * PROBLEM SET #3 **

1. For the following reactions: • Draw a diagram showing the enthalpy of the reactants and products. • State whether the reaction is exothermic or endothermic. • Predict whether the temperature of the surroundings would increase or decrease. • Rewrite the equation including the energy term as part of the equation: a. CO (g) + ½ O2 (g) ® CO2 (g) D H = -283 kJ   _ _   _    b. ½H2 (g) + ½ I2 (g) ® HI (g) D H = +26 kJ    _ _   _    2. For the following reactions: • State whether the reaction is exothermic or endothermic. • State the value of D H for the reaction. a. C (s) + H2O (g) + 131 kJ ® CO2 (g) + H2 (g) b. 2S (s) + 3O2 (g) ® 2SO3 (g) + 792 kJ   3. 824 kJ of energy is released when __one__ mole of iron metal is burnt in oxygen gas and converted to iron (III) oxide. Write an equation for this reaction, indicating the heat energy involved which you will need to calculate from the balanced equation. _   _    _    _    _    _    _

Energy changes associated with physical processes are much smaller than those associated with chemical changes. || -44   || When water vapour condenses to liquid, 44 kJ of heat is released for each mole of water condensed. The formation of liquid water from the elements H2 and O2 releases 286 kJ per mole of water formed. In both processes the final product is liquid water: i. in the condensation process weak hydrogen bonds are formed between adjacent molecules - results in a small decrease in the potential energy of the system and a release of heat. ii. the chemical process involves the breaking and making of strong covalent bonds.
 * ENERGY CHANGES AND PHYSICAL PROCESSES **
 * **Process** ||  **Type of Change**  ||  ** D H (kJ)**  ||
 * H2O (g) ® H2O (l) || physical
 * H2 (g) + ½O2 (g) ® H2O (l) || chemical  ||  -286  ||
 * NaOH (s) ® Na+ (aq) + OH- (aq) || physical  ||  -45  ||
 * Na (s) + ½O2 (g) + ½H2 (g) ® NaOH (s) || chemical  ||  -425  ||

H2 (g) + O2 (g) ® H2O2 (l) D H = -188 kJ (Chemical change) H2O2 (g) ® H2O2 (l) D H = -51 kJ (Physical change)

The first equation represents a chemical process while the second represents a physical process. The chemical process involves the making and breaking of strong covalent bonds so that much large amounted of energy are involved and the heat of reaction is considerably larger.

1. Compare the heats of reaction for the following processes: H2 (g) + O2 (g) ® H2O2 (l) D H = -188 kJ   H2O2 (g) ® H2O2 (l) D H = -51 kJ    and explain the difference in terms of the processes involved. __   2. For the following processes: i. State whether the process would be exothermic or endothermic. ii. Indicate whether the reactants or products have the higher enthalpy. iii. Identify whether heat would be absorbed from or released to the surroundings. _____   _    ___    _____    _    ___    _    ___    3. Predict the sign of D H for the following processes: 4. Identify the following reactions as exothermic or endothermic. a. N2 (g) + 3H2 (g) ® 2NH3 (g) D H = -92 kJ _ b. 6CO2 (g) + 6H2O(l) ® C6H12O6 (aq) + 6O2 (g) D H = +2803 kJ _ c. 2C(s) + O2 (g) ® 2CO(g) + 222 kJ _ d. 2H2)(g) + 484 kJ ® 2H2 (g) + O2 (g) _   5. For the reaction: H2 (g) + Cl2 (g) ® 2HCl(g)    6. Instant cold packs used to treat sporting injuries often contain NH4NO3 and a plastic bag of water. When the plastic bag is broken the ammonium nitrate dissolves as follows:    NH4NO3 (s) ® NH4+(aq) + NO3- (aq)    a. Is this reaction exothermic or endothermic? _____    b. Do the products or reactants have a higher enthalpy?    c. Draw a diagram showing the    enthalpy of the reactants and products.   1. Chemicals reactions are usually associated with the absorption or release of ___.__  2. The Law of Conservation of Energy states that  3. Reactions can absorb heat (_ reactions) or can release heat ( reaction).  4. The enthalpy of a substance is described as the  The symbol for enthalpy is .  5. If enthalpy decreases within a chemical reaction, a corresponding amount of energy must be  __to the surroundings. This type of reaction is called an__ _ reaction. 6. If enthalpy increases within a chemical reaction, a corresponding amount of energy must be __from the surroundings. This type of reaction is called an__ _ reaction. 7. If a reaction occurs in a test tube such as the reaction between zinc and HCl, the test tube gets hot. This type of reaction is an _ reaction. 8. What type of reaction does the graph illustrate, endothermic or exothermic? 9. Steam condenses to water: a. Is the process endothermic or exothermic? _ b. Will the reactants or the products have the higher enthalpy? _____ c. Will heat be absorbed or released to the surroundings? _____ 10. a. When chemical bonds are broken, energy is absorbed or released? _ 11. Hydrochloric acid can be produced by the reaction between hydrogen and chlorine gas as shown in the equation: H2 (g) + Cl2 (g) ® 2HCl (g) a. What bonds are broken? __. When broken, energy is__ _ (released or absorbed) b. What bonds are formed? . When formed energy is _____ c. In this case, the energy released from forming HCl bonds is greater than energy to break H-H bonds and Cl-Cl bonds. Is this reaction endothermic or exothermic? 12. Which has more enthalpy, an oxygen molecules or two separate oxygen atoms? ___. Explain why this is so.__ _ 13. Enthalpy is not measured directly but only the changes in it – change in enthalpy symbol. D H = H(_) – H(_). D H for exothermic reactions is negative compared to surroundings, explain why D H for endothermic reactions is positive compared to surroundings, explain why _ __ 14. The heat of reaction is proportional to the number of moles of the substances that react. If twice as many moles react, twice as much heat is released or absorbed. Do the following examples (give heat energy in the form of D H) a. When 1.0 mol of water is formed by neutralisation reaction between hydrogen ions, and hydroxide ions in solution, 57.1 kJ of heat energy is released. How much heat is released when 0.25 mol of water formed? H+ (aq) + OH- (aq) ® H2O (l) + 57.1 kJ 15. Write the following equation so that the change in enthalpy is part of the equation. CaCO3 (s) D CaO (s) + CO2 (g) D H = + 294 kJ mol-1 16. A thermochemical equation expresses a quantitative relationship between amount of substances that react and heat of reaction. H2 (g) + Cl2 (g) ® 2HCl (g) + 185 J 1 mol : 1 mol ® 2 mol with 185 J heat energy released. Write a thermochemical equation for the following: 4HCl(g) + O2 (g) D 2H2O(g) + 2Cl2 (g) D H = + 113 kJ mol-1 17. Heat content is the sum of the and potential energy of the molecules or ions of a specified amount of a substance. If total heat content of reactants is greater than products, energy is ___ and the reaction is__ __. If the total heat content of reactants is less than products, energy is__ and the reaction is  18. In each of the following, determine if the reaction is endothermic or exothermic. a. D H = +44.1 kJ mol-1 _ b. 2H2 (g) + O2 (g) ® 2H2O (g) + 483.6 kJ mol-1 c. D H = -51.8 kJ mol-1 d. ½H2 (g) + ½I2 (g) + 26 kJ ® HI (g) _ 19. Look at the following reactions each which produce liquid water. || Explain why the difference in the change of energy. _   _    _    _    _    _   20. Draw an enthalpy change graph for the following reactions:  a. C (s) + H2O (g) + 131 kJ ® CO (s) + H2 (g) b. 2S (s) + 3O2 (g) ® 2SO3 (g) + 792 kJ
 * PROBLEM SET #4 **
 * 1) Burning magnesium
 * 1) Subliming carbon dioxide
 * 1) Adding concentrated sulfuric acid to water, which results in the temperature of the solution increasing. _____
 * 1) The formation of dew. _____
 * 2) O2 (g) ® 2O(g) _
 * 3) H(g) + I(g) ® HI(g) _____
 * 4) CaCO3 (s) + 178 kJ ® CaO(s) + CO2 (g) _
 * 1) Identify which bonds are broken.
 * 2) Identify which bonds are formed
 * 3) Given that D H = -184 kJ for this reaction, is more energy involved in bond breaking or bond forming?
 * Chemical Reactions Reviews: **
 * 1) When chemical bonds are formed, energy is __?__
 * 2) If bonds formed are weaker or fewer than those broken, energy will be ___?__
 * 3) If bonds formed are stronger or more numerous then those broken, energy will be __?__
 * 1) How much heat is released when 0.4 mol HCl(aq) reacts with 0.4 mol NaOH(aq) to produce sodium chloride and water. 57.1 kJ of energy is released for each mole of water formed?
 * 2) How much heat is released when 100 mL of 0.1 M HNO3 (aq) is mixed with KOH(aq) to form potassium nitrate and water. 57.1 kJ of energy is released for each mole of water formed?
 * H2O (g) ® H2O (l) || D H = -44 kJ physical reaction
 * H2 (g) + ½O2 (g) ® H2O (l) || D H = -286 kJ chemical reaction ||
 * REACTION RATES **

** Rates Of Chemical Reactions ** The rate of chemical reactions is affected by the: i. nature of the reactants ii. concentration of the reactants. iii. state of subdivision of the reactants. iv. temperature v. presence of a catalyst. ** The Concept of Reaction Rate ** A reaction rate can be determined by i. the rate of disappearance of reactants. ii. the rate of appearance of products. Consider the reaction of hydrochloric acid with marble chips: CaCO3 (s) + 2H+ (aq) ® Ca2+ (aq) + CO2 (g) + H2O (l) Changes which take place in this reaction include: i. the mass of CaCO3 (s) decreases. ii. the concentration of H+ (aq) decreases. iii. the concentration of Ca2+ (aq) increases. iv. the volume of CO2 (g) produced increases. The reaction rate of this reaction could be determined by observing the rate of any of these changes. ** The Nature of the Reactants ** By examining many different reactions, chemists have produced some guidelines which are sometimes useful in helping predict reaction rates: i. if a reaction does not involve bonding rearrangement its it likely to be rapid at room temperature. Ag+ (aq) + Cl- (aq) ® AgCl (s) ii. if a reaction involves the breaking of bonds it is likely to be slow at room temperature. CH4 (g) + 2O2 (g) ® 2H2O (g) + CO2 (g)
 * FACTORS THAT AFFECT REACTIONS RATES **

** The Concentrations of the Reactants ** Increasing the concentration increases the rate of reaction. The rate of reaction between HCl (aq) and CaCO3 (aq) can be increased by increasing the concentration of HCl (aq).

** The State of Sub-Division of the Reactants ** By decreasing the size of the particles of the reactants, the rate of reaction increases because it increases the surface area of the reactants. S (l) + O2 (g) ® SO2 (g) A stream of liquid sulphur is sprayed into the combustion chamber through jets causing it to break up into numerous droplets resulting in a much faster rate of reaction. An increase in temperature increases the rate of reaction. A catalyst is a substance that increases the rate of a chemical reaction without being permanently consumed in the reaction.  || MnO2 ||  2H2O2 (aq) O2 (g) + 2H2O (l) MnO2 is a catalysts used in the decomposition of H2O2 into O2 gas and H2O. A catalysts is not required in large amounts to alter the rate of reaction.
 * Temperature **
 * Catalysts **


 * Do Exploring Chemistry Experiments 16, 17 and 18 from page 82. **

1. Which of the following reactions are likely to be rapid at room temperature? a. Ba2+ (aq) + SO42- (aq) ® BaSO4 (s) b. C2H6 (g) + Br2 (g) ® C2H5Br (l) + HBr (g) c. C2H6 (g) + 7/2 O2 (g) ® 2CO2 (g) + 3H2O (g) Answer: _ d. Ce4+ (aq) + Fe2+ (aq) ® Ce3+ (aq) + Fe3+ (aq) 2. The rate of reaction between H2 (g) and I2 (g) to form HI (g) is found to increase when the volume of the system is decreased. Explain. _   3. Why are many metal catalysts used in the form of fine wire mesh? _   4. Explain the following observations: a. Although wheat is usually slow to burn, wheat silos have been known to explode. _   _____    b. Food products are often stored in refrigerators. _____   c. Liquid hydrogen, used as a fuel to launch the space shuttle, is reacted with pure liquid oxygen rather than air. _____   _____    5. In the manufacture of nitric acid, nitric oxide is produced by the oxidation of ammonia according to the following equation: 4NH3 (g) + 5O2 (g) 4NO (g) + 6H2O (g) Under standard conditions of temperature and pressure the rate of this reaction is very slow indeed. Suggest three ways in which the rate of the reaction might be increased.
 * PROBLEM SET #5 **

2HI (g) ® H2 (g) + I2 (g)
 * COLLISION THEORY AND ACTIVATION ENERGY **

For every two molecules of HI (g) which decompose, one molecule of H2 (g) and one molecule of I2 (g) are produced. This equation does not indicate how HI molecules are converted into H2 and I2. A collision occurs between the reactant particles: HI.

i. collide with sufficient energy to disrupt the bonds of the reactant molecules. ii. collide with an orientation that is suitable for the breaking of some bonds and the formation of others.
 * Collision Theory ** - the collision theory assumes that if particles are to react they must first undergo an appropriate collision. The molecules must:
 * Activation energy ** is the minimum energy that is required for a collision to result in a reaction.

It is important that reactions have an orientation requirement since the relative orientation of reactant molecules must be favourable for the breaking of particular bonds and the formation of new bonds in the products. Many collisions between reactant are unsuccessful in producing a reaction.

An energy profile diagram is a way of representing the energy changes that occur during a chemical reaction.
 * Energy Profile Diagrams ** // A Typical Energy Profile Diagram //

Transition State - occurs between the reactant and the product states and is a very high energy state. This high energy state is known as the **transition state** or **activated complex:** i. represents the highest energy state for the reacting system. ii. corresponds to some stage in the reaction at which bond-breaking and bond-formation are taking place. iii. unstable state - temporary. The activation energy of reactants determines as to whether or not a reaction will occur: i. difference in the potential energies of the reactants and the transition state. ii. a barrier which must be overcome for the reaction to take place - activation energy barrier. iii. when reactant molecules collide, a reaction will only occur if the energy of the collision is sufficient to supply the activation energy needed for the reaction to reach the transition state. It should be noted that there is a reaction mechanism that must occur – these are a series of steps for the reaction. Each step that is not reactants or products is called an intermediate. The **rate-determining step** is the slowest intermediate step, reaction can’t go faster than this step. // The Rate of Reaction Depends on the Proportion of Molecular Collisions With Enough Energy to Overcome the Activation Energy Barrier // Indicates that only a relatively small proportions of molecules in this system have energies greater than that of the activation energy barrier. In this system, few collisions would have sufficient energy to be successful in bringing about a reaction - thus reaction rate would be slow.

// Energy Profile Diagrams for Exothermic and Endothermic Reactions // These indicate that an activation energy exists for both exothermic and endothermic reactions. In these particular examples, both reactions are depicted as having the same activation energy. This means that the same amount of energy is needed for a molecular collision to be successful in either direction


 * Reversibility of Reactions **

// Activation Energies for Reverse Reactions // In many reactions there is a significant reaction in the opposite direction in which product molecules react to regenerate the reactants. An example of such a reaction: 2HI (g) D H2 (g) + I2 (g) The activation energy for the reverse reaction is considerably higher for reaction (a) than for reaction (b). Reaction (b) will proceed at a faster rate.

1. What factors determine whether a particular collision between reactant molecules will lead to a reaction? _____   _____    _____    _____    2. A piece of white phosphorous left in the open air at room temperature spontaneously ignites and burn to form P4O10. In contract, a piece of iron rusts very slowly when exposed to the atmosphere. What does this information suggest about the activation energies for these reactions? __   3. a. Draw energy profile diagrams for reaction which have the following activation energies and heats of reaction: b. Assuming all other factors affecting the rates of these reactions are equal, identify the following: • the fastest and slowest forward reaction _____ • the fastest and slowest reverse reaction 3. When nitrogen and hydrogen are heated under conditions of high temperature and pressure ammonia is formed. However, a sample of ammonia when heated produces nitrogen and hydrogen. Write an appropriate equation to present this reversible system. __
 * PROBLEM SET #6 **
 * ||  A   ||   B   ||   C   ||
 * ** Activation Energy (kJ) ** ||  100   ||   50   ||   120   ||
 * ** Heat of Reaction (kJ) ** ||  -50   ||   -80   ||  +50  ||

** Nature of Reactants ** i. slower reactions will have higher activation energies than faster reactions. ii. different activation energy for different reactions are related to the ease with which bond-breaking and reforming processes occur. iii. high activation energy are often associated with reactions in which strong bonds have to be broken. i. increasing the concentration of a reactant in a solution increase the rate at which reactant molecules collide. ii. greater number of collision, faster rate of reaction. iii. same proportion of these collision will be successful as dilute solutions, but the greater the rate of collisions results in a greater number of successful collisions. ** Sub-division ** If the surface area of a solid or liquid is increased, more of the reactant molecules are exposed to collision. i. increase in temperature increases the average kinetic energy of the reacting molecules and change the distribution of molecular kinetic energies. ii. increased velocities of the molecules leads to a greater rate of collision and a greater reaction rate. iii. greater proportion of reactant molecules have sufficient kinetic energy to supply the activation energy needed for reaction - a greater proportion of molecular collisions will be successful.
 * APPLYING COLLISION THEORY **
 * Concentration **
 * Temperature **

** Catalysts ** i. provide an alternative pathway to that available when reactants are used alone - this different pathway has a lower activation energy. ii. less collision energy is required for the reaction to take place - greater proportion of collisions will be successful and the reaction rate will be greater. iii. greater proportion of reactant molecules have sufficient energy to overcome the activation energy barrier.

 || //A Catalyst Provides and Alternative Pathway with a Lower Activation Energy:// ||

// A Catalyst Decreases the Activation Energy of a Reaction - A Greater Proportion of Molecular Collision having Sufficient Energy to Reach the Transition State: // iv. allows reactions to be carried out at much lower temperatures. v. operate by forming intermediate compounds in which the bond-breaking and rearrangement process requires less energy than it the reactants alone were involved.

= Collision Theory and Activation Energy  = Collision theory assumes if particles are to ___ they must undergo appropriate__ Molecules must collide with a. sufficient _ to ___ the bonds of the__ molecules. b. an __that is suitable for the breaking of some bonds and the__ _ of others. ** Energy Profile Diagrams: ** 1. Label the following activation energy diagram with: A. reactants B. transition state C. products D. activation energy (Ea) E. heat of reaction ( D H) F. potential energy G. progress of reaction 3. Explain the following terms: a. activation energy _ b. transition state _ c. Activation complex _ d. reaction mechanism _ e. intermediates f. Rate-determining step: 4. Is a reaction more likely to occur if the activation energy is high or low? 5. Using the energy profile diagram, explain what determines if a reaction will occur. 6. The diagram shows a reaction with a high activation energy and a slow reaction. Draw a similar diagram to show a fast reaction.
 * Summary: Chemical Reaction Rates, Collision Theory & Activation Energy **

7. Draw the following energy profile diagram a. Fast endothermic reaction b. Slow exothermic reaction 8. Some reactions are reversible with the reactants reforming from the products. Each of the following is a reversible reaction. For each of the following state a. if the reaction is exothermic or endothermic. b. if the forward or reverse reaction is more likely to be the faster reaction. _____   _____   ** Rates of Chemical Reactions and applying the Collision Theory: ** Reaction rate determined by rate of disappearance of and rate of appearance of the _. Observing any change within a reaction can determine rate of reaction. The rate of chemical reactions is affected by five factors: 1. __Nature of the reactants__: a. Reactions NOT involving bond re-arrangement will be ___ (rapid or slow at room temp.)__ for example _____ These reactions will have a ___ (higher/lower?) activation energy__ b. Reactions that DO involve bond re-arrangement will be _____ for example These reactions will have a ___ (higher/lower?) activation energy__ c. Strong bonds usually have a ___ (higher/lower?) activation energy__ 2. Concentration of the reactants: a. Increasing concentration of any of the reactants can increase of reaction for example: b. Changing the concentration does NOT change the activation energy. Explain why increasing concentration of reactions increases rate of reaction. 3. State of sub-division of reactants: a. Using powders, sprays, granules – all increase __area of reactants and hence its access to the other reactant/s in the reaction. A simple example is chopping a log of wood into smaller pieces for a fire – increases surface area of wood exposed to oxygen and flame.__ Another example is b. Explain how the collision theory relates to sub-division. 4. Temperature: a. Increase in temperature often results in increase in reaction rate as more particles given enough kinetic energy for successful collisions and formation of product. For example: b. Increasing the temperature increases the energy of the molecules and hence their v___. Increase in velocity leads to a greater rate of__ ___ and a__ greater reaction rate. c. How does greater kinetic energy and velocity relate to activation energy of the reaction? _ _____  _  _____  d. On the following diagram, show what happens to the number of molecules with enough kinetic energy to overcome activation energy barrier when temperature is DECREASED. 5. Catalysts:__ a. Catalysts increase rate of reaction without themselves being used up although they do take part in the reaction. Catalysts are not included in the reaction but is written above the arrow. For example  || chlorophyll ||  photosynthesis: 6CO2 (g) + 6H2O(l) + energy C6H12O6 (aq) + 6O2 (g) here, the catalysts is __. Write the equation for the production of__ oxygen gas with MnO2 as the catalyst: b. A catalysts doesn’t reduce the activation energy pathway but rather provides an _____ pathway for the reaction. This pathway has a activation energy so a greater number of collisions will be _ and reaction rate greater as a greater proportion of molecules have sufficient  __to overcome catalysed reaction pathway.__ c. Reaction can also be carried out at a much __temperature.__ d. Catalysts form an ___ compound in which bond-breaking and re-arrangement processes require__ _ energy than reactants alone. e. Draw (i) a potential energy/progress graph of reaction graph showing with and without a catalyst. (ii) a graph showing the number of molecules with sufficient kinetic energy for a reaction to occur with and without a catalyst. 1. Consider the reaction Mg(s) + 2H+(aq) ® Mg2+(aq) + H2 (g). By observation, describe two ways we could measure the rate of the reaction. 2. Suggest reasons for the difference in rate of the following reactions at room temperature: Fast: Cu2+(aq) + CO32-(aq) ® CuCO3 (s) Slow: C6H12O6 (aq) + 6O2 (g) ® 6CO2 (g) + 6H2O(l) 3. Predict the effect of increasing the pressure of oxygen on the rates of the following reactions. a. 2SO2 (g) + O2 (g) ® 2SO3 (g) b. 2Na (s) + O2 (g) ® Na2)2 (s) _____ c. Ag+(aq) + Cl-(aq) ® AgCl (s) _  d. 2H2O2 (aq) ® O2 (g) + 2H2O(l)  4. When hydrogen gas is prepared in the laboratory, hydrochloric acid is usually added to granulated rather than powdered zinc. Suggest why.  5. When dilute acid is added to sodium thiosulfate solution a precipitate of sulfur is formed.  S2O32-(aq) + 2H+(aq) ® S (s) + SO2 (aq) + H2O (l)  If it takes 30 seconds for a precipitate to appear using 0.05 mol L-1 Na2S2O3 solution, estimate what time it might take if the Na2S2O3 solution was 0.10 mol L-1. Show your reasoning.  6. a. Why must reactant molecules collide with an energy greater than a certain minimum value for  reaction to occur? _  b. Will every collision which has this minimum energy result in the formation of products? Explain.  _  _  7. Use the energy profile diagram to the right to answer the following questions.  a. Is the reaction exothermic or endothermic? _ b. What is the value of D H for the forward reaction? _ c. What is the activation energy for the forward and reverse reactions? _ d. A catalyst for the reaction lowers the activation energy by 10 kJ mol-1. In the presence of the catalyst what is the activation energy of the forward and reverse reactions. _____ e. How does the catalyst affect the rate of the forward and reverse reactions. _ _____  8. For the following reactions draw energy profile diagrams, indicating the potential energies of the reactants and products and the activation energy for the reaction. a. A piece of sodium added to water reacts rapidly to produce hydrogen and sodium hydroxide solution. The reaction is accompanied by the release of considerable quantities of heat. b. A mixture of N2 (g) and O2 (g) when heated very strongly is converted to NO(g). The reaction is endothermic. 9. A common misunderstanding regarding catalysts is that ‘they speed up a reaction but take no part in the reaction’. Explain what is wrong with this statement. ___
 * Questions:**
 * SUMMARY OF ENERGY EFFECTS: **
 * Energy Effects: **

· //Heat of Reaction:// Amount of heat released or absorbed when specified amounts of substances react. · //Exothermic Reactions:// mixture becomes hotter and release heat to surroundings. · //Endothermic Reactions:// mixture becomes colder and absorbs heat from surroundings. · Heat of reaction proportional to number of moles of substance that reacts e.g. if twice as many moles react, twice as much heat released or absorbed. Examples: d. When 1.0 mol of water is formed by neutralisation reaction between hydrogen ions, and hydroxide ions in solution, 57.1 kJ of heat energy released. How much heat is released when 0.25 mol of water formed? · //Thermochemical Equations:// Express a quantitative relationship between amounts of substances that react and heat of reaction e.g. H2 (g) + Cl2 (g) ® 2HCl (g) + 185 J 1 mol : 1 mil ® 2 mol with 185 J heat energy released. · //Internal Energy:// Combined chemical and potential energy of the molecule. Chemical changes cause change in stored chemical energy. · //Heat Content:// Sum of chemical and potential energy of the molecules or ions of a specified amount of a substance. If total heat content of reactants greater than products excess energy released – exothermic. If total heat content of reactants is less than products, energy absorbed – endothermic. · //Enthalpy:// Change in heat content. Symbol H. Enthalpy not measured directly only changes in it - D H. D H = H(products) – H(reactants). ( D - pronounced delta and represents change in magnitude.) D H for exothermic reactions is negative compared to surroundings (why?) D H for endothermic reactions is positive compared to surroundings (why?) · //Origin of heats of reaction:// When bonds broken or formed, heat absorbed or released. When a single reaction takes place, heat of reaction equal to the heat released or absorbed. However, many reactions are a sequence of chemical reactions and the heat of reactions is equal to the algebraic sum of each individual reaction. For example: H2 (g) + Cl2 (g) ® 2HCl (g) D H = 185 J H-H bond broken H = +437 kJ x 1 mol = +437 kJ Cl –Cl bond broken H = +244 kJ x 1 mol = +244 kJ H-Cl bond formed H = -433 kJ x 2 mol = -866 kJ total D H = -185 kJ · In solids, force of attraction between particles is much stronger than gases and therefore magnitude of the heat of reaction is partly dependent on this force. For example heat must be added to solid iodine for sublimation to occur. I2 (s) ® I2 (g) D H = +62 kJ  · When reaction involves different phases, sequence of chemical reactions can become more complex. For example: NH3 (g) + HCl (g) ® NH4Cl (s) For this reason, the subscripts (s), (l), and (g) become very important.
 * 1) How much heat is released when 0.4 mol HCl(aq) reacts with 0.25 mol NaOH(aq)?
 * 2) How much heat is released when 100 mL of 0.1 M HNO3 (aq) mixed with 100 mL of 1.0 M KOH(aq)?
 * 1) heat absorbed when HCl bond broken.
 * 2) heat released when ammonium ion formed
 * 3) heat released when ammonium chloride solid lattice formed.

** Rates of Reaction: ** · //Units:// moles per time unit · //Factors affecting:// Temperature, nature of reactants, concentration, subdivision, catalyst. · //Concept of reaction rate:// determined by (i) rate of disappearance of reactants; (ii) rate of appearance of products. · //Collision Theory:// Collision must occur for reaction proceed. Rate of collisions directly relates to rate of reaction. Appropriate collision orientation on collision. · //Effect of Concentration:// Concentration equals number of particles per unit volume, increase concentration results in increase in collision rate · //Effect of temperature:// For collisions to be effective in causing chemical change they require a minimum energy of collision (//Threshold Energy)//. An increase in temperature results in more pairs of particles possessing this threshold energy. An increase in temperature means an increase in the number of collisions per time unit. But the increase in the number of effective collisions per time unit is a far more important factor. · //Subdivision:// Increase surface are increases rate. · //Catalysts:// provide an alternative pathway for the reaction mechanism to follow (see below). Activation energy of catalysed reaction lower than without catalyst. Not consumed although it does take part in reaction. Does not effect D H. Speeds up forward and reverse reaction. Allows reaction to proceed faster at lower temp. ** Activation Energy: ** · //Activation complex:// transition stage between reaction and produces (top of peak of reaction curve) · //Activation Energy (Ea):// minimum energy required for collision to result in a reaction (energy colliding particles must have to form activated complex) · Fast reactions have a low activation energy, slow reactions have a high activation energy. · //Reaction mechanism//: series of steps for reaction. · //Intermediates//: steps that are not reactants or products. · //Rate//-//determining step//: slowest stope in reaction, speed of reaction cannot go faster than the slowest step. ** Applied Chemistry Reaction Rates ** __ Home Research __ : As part of this course, you need to apply your understanding of Chemistry to the world around you. One such task is to describe and explain examples where rates of reaction have been altered in and around the home. For example, why do you store milk in the fridge and not on the counter, it would certainly be more convenient. Answer the following questions in your own time. It is very likely that one will appear in a test. 1. List the ways in which the rates of chemical reactions can be altered (page 16). 2. Again consider the milk problem. Why don’t you leave milk on the counter? _Explain, using your understanding of reaction rates, why this occurs. _____ 3. If you are adding spices to a cake, you usually use ground spices not whole seeds. Using your understanding of reaction rates, explain why. _ 4. When making a pasta sauce, a cook will always use tomato paste rather than tomato sauce although both are very similar in their ingredients. Using your understanding of reaction rates, explain why. 5. Catalytic converters are an essential part of a car’s exhaust system. Find out what the purpose is, what catalysts are used and how they work. _ _