1. Energy and Chemical Reactions
Q4U2

2. What is Energy?
Law of Conservation of Energy: energy can not be created or destroyed, only changed from one form to another.
Two Main Forms of Energy
Kinetic energy is motion energy. 
Defined as the energy of a moving object.
A thrown football, a speeding automobile, a waterfall, or a rock falling from a cliff are examples of objects that have kinetic energy. 
Potential energy is energy stored in matter. 
Potential energy appears in many different forms
Defined as the energy in matter due to its position or the arrangement of its parts.

3. Chemical Potential Energy
Rubber bands or Gasoline both have potential energy
The energy stored in molecules is called chemical potential energy.
This energy is stored in bonds
The chemical makeup (arrangement of molecules) of gasoline makes it a good fuel source (source of potential energy).
The energy stored in gasoline is released by burning it (combustion).
The airplane motor uses this released energy to turn a propeller.
During combustion, chemical bonds are broken and reformed creating new products and releasing energy.

4. Chemical Potential Energy
Is Related to heat lost or gained in chemical reactions
Heat is released, given off, in an exothermic reaction
Heat is absorbed, taken in, in an endothermic reaction
The reaction between H2 and O2 is highly exothermic
The energy from it powers cars

5. Combustion

6. Enthalpy
The heat content of a chemical system is called the enthalpy (symbol: H)
The enthalpy change (ΔH) is the amount of heat released or absorbed when a chemical reaction occurs at constant pressure.
ΔH is specified per mole of substance
Standard units are:
kJ mol-1 (kJ/mol) or
kcal mol-1 (kcal/mol)
A calorie is the amount of heat needed to raise the temp of 1g of water by 1 ˚C
1 calorie (1 cal) = joules (4.184 J)
1kilocalorie = 1 Calorie =1000 calories
The Bond Enthalpy is the energy required to break a chemical bond. It is usually expressed in units of kJ mol-1, measured at 298 K.
The Standard unit for energy is the joule
Energy changes are measured under standard laboratory conditions (STP)

7. Bond Enthalpy
Bond Enthalpy is the energy required to break a chemical bond. It is usually expressed in units of kJ/mol
ΔH = ∑ ΔH(bonds broken) - ∑ ΔH(bonds formed)
Add up all the energies of the broken bonds; add up all the energies of the bonds that are reformed and subtract one from the other.

8. EXAMPLE #1
Find ΔH for the following reaction given the following bond energies:
ΔH = ∑ ΔH(bonds broken) - ∑ ΔH(bonds formed)
Bond   Bond Energy(kJ/mol)
H-H
O=O
O-H
Figure out which bonds are broken and which bonds are formed.
2 H-H bonds are broken.
1 O=O bond is broken
2 O-H bonds are formed per water molecule, and there are 2 water molecules formed, therefore 4 O-H bonds are formed
Substitute the values given into the equation:

9. Example #2 Calculate the enthalpy of the combustion of Methane
CH4 + O2  CO2 + H2O
Balance the equation
CH4 + 2O2  CO2 + 2H2O
Write the Lewis Dot Structure H
H-C-H, O=O, O=O  O=C=O + H-O-H, H-O-H
Count bonds broken and bonds Formed
Bonds broken (reactants) Bonds Formed(Products)
4 C-H 413 kJ/mol each C=O 799 kJ/mol each
2 O=O 495 kJ/mol each O-H kJ/mol each
Calculate the sum of the bond enthalpies of products and reactants
∑ ΔH(bonds broken) = (4 X 413) + (2 X 495) = 2642
∑ ΔH(bonds formed) = (2 X 799)+ ( 4 X 463) = 3450
Calculate the enthalpy of reaction
ΔH = ∑ ΔH(bonds broken) - ∑ ΔH(bonds formed)
2642 – 3450 = -808 kJ/mole

10. Complete Problems 1-4 HO Solutions H-C-C-C-H + O=O  O=C=O + H-O-H
Broken bonds(ΔH each) Formed bonds(ΔH each)
8 C-H kJ/mol C=O 799 kJ/mol
2 C-C kJ/mol O-H 463 kJ/mol
5 O=O kJ/mol
∑ΔH(bonds broken)=(8 X413)+(2X348)+(5X495)=6475
∑ ΔH(bonds formed) = (6 X 799)+ ( 8 X 463) = 8498
ΔH = ∑ ΔH(bonds broken) - ∑ ΔH(bonds formed)
6475 – 8498 = kJ/mole

11. Solution Problem 2 HO Solutions H-C-C-O-H + O=O  O=C=O + H-O-H
Broken bonds(ΔH each) Formed bonds(ΔH each)
5 C-H kJ/mol C=O 799 kJ/mol
1 C-C kJ/mol O-H 463 kJ/mol
1 C-O kJ/mol
1 O-H kJ/mol
3 O=O 495 kJ/mol
∑ΔH(bonds broken)=(5 X413) (3X495)=4719 kJ/mol
∑ ΔH(bonds formed) = (4 X 799)+ ( 6 X 463) = 5974 kJ/mol
ΔH = ∑ ΔH(bonds broken) - ∑ ΔH(bonds formed)
4719 – 5974 = kJ/mole

12. Solution Problem 3 HO Solutions H-C=C-H + O=O  O=C=O + H-O-H
Broken bonds(ΔH each) Formed bonds(ΔH each)
4 C-H kJ/mol C=O 799 kJ/mol
2 C=C kJ/mol O-H 463 kJ/mol
5 O=O 495 kJ/mol
∑ΔH(bonds broken)=(4 X413)+(2X 839)+(5X495)= 5805kJ/mol
∑ ΔH(bonds formed) = (8 X 799)+ ( 4 X 463) = 8744kJ/mol
ΔH = ∑ ΔH(bonds broken) - ∑ ΔH(bonds formed)
5805 – 8744 = kJ/mole

13. Solution Problem 4 HO N2 + 3H2  2NH3
Broken bonds(ΔH each) Formed bonds(ΔH each)
1 N=N kJ/mol N-H 391 kJ/mol
3 H-H kJ/mol
∑ΔH(bonds broken)=941+(3X 436)= 2249kJ/mol
∑ ΔH(bonds formed) = (6 X 391) = 2346kJ/mol
ΔH = ∑ ΔH(bonds broken) - ∑ ΔH(bonds formed)
2249– 2346 = -97 kJ/mole

14. ΔH =ƸΔH(products) - ƸΔH(reactants)
Enthalpy Change
Exothermic Reaction
Endothermic Reaction
Energy is released.
Energy is absorbed.
Energy is a product of the reaction.
Energy is a reactant of the reaction.
Energy of the reactants is greater than
the energy of the products
ΔH(reactants) >ΔH(products)
Energy of the reactants is
less than the energy of the products ΔH(reactants) < ΔH(products)
ΔH =ƸΔH(products) - ƸΔH(reactants)
 ΔH = Negative number
   ΔH = Positive number

15. Heat of Reaction
Heat of Reaction: The difference in energy between the Energy of the Products and the Energy of the Reactants.
Δ H = H Prod - H Rxts

16. Exothermic Reaction
This is an exothermic reaction, the energy of the products is less than the energy of the reactants
The energy was lost to the environment as heat!

17. Endothermic Reaction
The Products have a higher energy level than reactants
Reactants are located on the flat portion to the left of the peak.
Products are located on the flat portion to the right of the peak.
The activated complex is located at the peak of the reaction diagram.

18. An Activated Complex
An activated complex is an intermediate state that is formed during the conversion of reactants into products.
It is the structure that results at the maximum energy point along the reaction path.
The activation energy of a chemical reaction is the difference between the energy of the activated complex and the energy of the reactants.
The lower the level of energy in a system the more stable it is(less likely to react).

19. Activated Complex

20. Activation Energy (Ea)
The amount of energy required to start the reaction (Ea)
The difference between the energy level of the reactants and the maximum energy in the reaction progress equals the activation energy

21. Catalysts and Activation Energy
Catalysts lower the activation energy of a reaction allowing it to start more easily
They do not effect the amount or type of product formed!

22. Entropy Example of ice melting
Entropy is a measure of how much of the energy of a system is potentially available to do work and how much of it is potentially manifest as heat.(heat is often lost to the environment)
Manufacturing equipment: heats the room because of heat loss to surroundings
Example of ice melting
the difference in temp. between a warm room (the surroundings) and a cold glass of ice water (the system) is equalized as heat from the warm surroundings are transferred to the cooler ice and water mixture.
In an isolated system such as the room and ice water together, the dispersal of energy from warmer to cooler regions always results in a net increase in entropy
The more disorder the higher the entropy!

23. Entropy and Energy Complex (lower entropy) (higher energy) endothermic
exothermic
(higher entropy) Chaos
(lower energy)

24. Answer the following questions based on the potential energy diagram shown here:
Does the graph represent an endothermic or exothermic reaction?
Label the position of the reactants, products, and activated complex.
Determine the heat of reaction, ΔH, (enthalpy change) for this reaction.
Determine the activation energy, Ea for this reaction.
How much energy is released or absorbed during the reaction?
How much energy is required for this reaction to occur?

25. Answers A) The reaction is endothermic
B) Reactants are located on the flat portion to the left of the peak. Products are located on the flat portion to the right of the peak. The activated complex is located at the peak of the reaction diagram.
C) The enthalpy change is +50 kJ
D) The activation energy is 200 kJ
E) 50 kJ is absorbed in this reaction
F) An input of 200 kJ is required for this reaction to occur

26. ΔH Calculations Practice
Consider a general reversible reaction such as: A + B ↔ C + D 
Given the following potential energy diagram for this reaction, determine ΔH and Ea for both the forward and reverse directions.
Is the forward reaction endothermic or exothermic?    

27. Solution
Activation Energy (Ea) is the difference in energy between the activated complex and the reactants
Large energy barriers separates the reactants and the products so that only very energetic molecules can pass over this barrier.
The forward reaction is exothermic.

28. Mass Energy Relationships
When 2 moles of H2 react with 1 mole of O2 to produce 2 moles of H2O (g), heat is released.
2 H2 (g) + O2 (g)  2 H2O(g) + 484kj
The heat needed to form a compound from its elements is called the energy of formation (∆Hf)
This information is found in tables, for 1 mole of different compounds

29. Mass Energy Relationships
If the ∆Hf of each reactant and each product is known, a balanced equation can be used to calculate the energy released or absorbed during a reaction.
The difference between the sum of the energies of the products and the sum of the energies of the reactants is the energy absorbed(+) or released(-) by the reaction
∆Hf(reaction) = ∑∆Hf(products) - ∑∆Hf(reactants)
∑ stands for the sum

30. Mass-Energy Relationships
How much energy is produced by the reaction of 16.0g of Fe2O3 with excess aluminum metal according to the equation:
Fe2O3(s) + 2Al(s)  Al2O3(s) + 2Fe kj
Steps to Solve
Begin with balanced equation above
Convert grams of Fe2O3(s) to moles
16.0g Fe2O3 1 mol Fe2O3
160g Fe2O3
From the equation, determine the number of moles and the energy involved. Calculate the answer
16.0g Fe2O3 1 mol Fe2O kj = 85.2 kj produced
160g Fe2O mol Fe2O3

31. Practice Problems (1-2)
A rocket fuel is prepared by reacting hydrazine and dinitrogen tetroxide according to the equation:
2N2H4 (l) + N2O4 (l)  3N2(g) +4H2O(g) kj
calculate the heat released when 3200 grams of hydrazine are consumed in the rocket engine.
The dissociation of ammonia into its elements is an endothermic reaction, absorbing 92.2kj of energy according to the equation:
2NH3(g) kj  3H2(g) + N2(g)
How much energy will be required to decompose 85.0g of ammonia?

32. Solutions to Practice Problems(1-2)
3200.0g N2H4 1 mol N2H kj = 120,000kj produced
32.0 g N2H4 1 mol N2H4
85.0gNH mol NH kj = 231 kj produced
17.0 g NH mol NH3

33. Show Video Clip (29min)
Chemistry Connections: Specific Heat Capacity and Calorimetry
nly=1&guidAssetId=928367f4-9cf5-4ed6-bae5- eded91b7dac9

34. Measuring Energy Changes
The heat required to change the temperature of a substance depends upon the amount and nature of the substance as well as the extent of the temperature change
For example:
1g of water requires 4.18joules of energy to change the temp1 ̊C.
It takes only 0.987j to raise the temp of AlF3 1 ̊C
Energy is transferred between a reaction and its surroundings.
The amount transferred can be calculated:
heat gained/lost= (mass)(change in temp)(specific heat)
q=(m)(∆T)(Cp) q is heat gained/lost
m is mass(g)
∆T is change in temp
Cp is the specific heat

35. Specific Heat A physical property of matter.
The temperature of matter is a direct measure of the motion of the molecules
The greater the motion the higher the temperature:

36. Specific Heat
Motion requires energy: The more energy matter has the higher temperature it will also have.
Typically this energy is supplied by heat.
Heat loss or gain by matter is equivalent to energy loss or gain
How much will the temperature of an object increase or decrease by the gain or loss of heat energy?
The answer is given by the specific heat (S) of the object.

37. Chemical and Physical change
Heat is gained or lost during chemical and physical changes
Calorimetry is a technique used to measure those changes in heat.
Important definitions:
Specific heat of water (Shw )=4.184 J / ( g oC )
This is the amount of heat required to raise the temp of 1 gram of water, 1 degree Celsius
Heat of fusion of ice (Hfi )=334 J / g
This is the amount of heat required to melt1 gram of ice at 0.00 degree Celsius (solid to liquid)
Heat of vaporization of water (Hvw)=2256 J / g
This is the amount of heat required to VAPORIZE 1 gram of water at degrees Celsius (liquid to gas)

38. Let’s Practice Using these Concepts!
How much heat is needed to melt a 60.0g ice cube at 0.0°C?
Heat of fusion of ice (Hfi )=334 J / g
Remember this is 1 gram of ice, we have 60.0g!
334 J g = J g

39. More Practice
How much heat is needed to melt a 45.0g ice cube at 0.0°C?
334 J g = J g
How much heat is needed to vaporize 16.3g of water at degrees Celsius?
(Hvw)=2256 J / g
16.3 g J = J g

40. Specific Heat
Take an object of mass m, put in x amount of heat and carefully note the temperature rise, then S is given by
In this definition mass is usually in either grams or kilograms and temperature is either in kelvin or degrees Celsius.
Note that the specific heat is "per unit mass".
This means the volume does not matter

41. Example problem finding specific heat
11 grams of substance is heated from 20.0˚C to 30.0˚C. The substance absorbed 4253 J of energy. What is the specific heat of the substance?
S= X/(mass)(∆T)
S = J
(11.0g)(10˚C)
=38.6 J/g ˚C

42. Specific Heat Example Problems
How much heat is required to raise the temperature of 68.0g of AlF3 from 25.0˚C to 80.0˚C ?
Steps to Solve
Find specific heat of AlF3 from table
Cp of AlF3 = J/g ֹ C˚
q=(m)(∆T)(Cp)
q=(68.0g) (80.0 – 25.0˚C) (0.8948J/g ˚C)
q= 3350 J

43. Calorimitry
A Calorimeter containing water is used to measure the heat absorbed or released in a chemical reaction.
The temperature change in the water is used to measure the amount of heat absorbed or released by the reaction
Law of conservation of energy says energy is neither created or destroyed.
In an insulated system, any heat lost by one quantity of matter must be gained by another.
Energy flows from the warmer material to the cooler material until they are both the same temperature.
Heat lost = Heat gained
m(∆T) Cp= m(∆T) Cp
The heat evolved from the reaction changes the temperature of a working substance (usually water) with a known heat capacity.
A measurement of the temperature rise in the surroundings (calorimeter body) allows a determination of the heat crossing the boundary between the system (where the reaction takes place) and the surroundings (where the temperature change is measured).

44. Calorimeter
A Calorimeter is an instrument for measuring the heat of a reaction during a well defined process
A much simpler, but less accurate calorimeter, which, by its construction, is necessarily constant pressure
A constant Volume calorimeter

45. Example of a Calorimetry Problem
Suppose a piece of lead with a mass of 14.9g at a temp of 92.5˚C is dropped into an insulated container of water. The mass of water is 165g and its temperature before adding the lead is 20.0˚C. What is the final temperature of the system? Cp lead = J/g ֹ ˚C
Remember Heat lost = Heat gained
m(∆T) Cp= m(∆T) Cp (The lead will lose energy)
The heat lost by lead is q= m(∆T) Cp = 14.9 g X (92.5 ˚C – Tf ) X J/gֹ ˚C
The heat gained by water is q= m(∆T) Cp = 165g X (Tf ˚C) X 4.18 J/gֹ ˚C
The heat gained must equal the heat lost:
165g X(Tf ˚C)X 4.18 J/gֹ ˚C=14.9 g X(92.5 ˚C– Tf )X J/gֹ ˚C
(Tf ˚C) X 690. J/˚C = (92.5 ˚C – Tf ) X 1.90J/˚C
(690. J/˚C )(Tf ) – J = 176 J – (1.90J/˚C )(Tf )
(690. J/˚C )(Tf ) + (1.90J/˚C )(Tf ) = 176 J J
(690. J/˚C J/˚C )(Tf ) = J
Tf = 20.2 ˚C

46. Practice Problems
How much heat is absorbed by a reaction that lowers the temperature of 500.0g of water in a calorimeter by 1.10˚C ?
Aluminum reacts with iron(III) oxide to yield aluminum oxide and iron. Calculate the heat given off in the reaction if the temperature of kg of water in the calorimeter increases by 3.0˚C.
Burning 1.00g of a fuel in a calorimeter raises the temperature of 1.0 kg of water from 20.0˚C to 28.05˚C. Calculate the heat given off in this reaction. How much heat would one mole of fuel give off, assuming the molar mass of 65.8g/mol?

47. Solutions qw = m(∆T) (Cw ) = 500.g(1.10˚C)(4.184J/g˚C) = 2300 J
= 2.30 kJ
= (1.0 x 103g)(3.00˚C)(4.184J/g˚C)
= J
= 12.6 Kj
∆T = ˚C – 20.0 ˚C = 8.05 ˚C
for 1g: qw = m(∆T) (Cw )
= (1. 0 x 103g)(8.05˚C)(4.184J/g˚C)
= J
= 33.7 kJ for 1.0g
for 1 mol: qw = (33.7 kj/1g)(65.8g/1mol)
= 2220kJ/mol

48. Writing thermochemical equations
An equation which shows both mass and heat relationships between products and reactants is called a thermochemical equation.
Example
2 H2(g) + O2(g) ----> 2 H2O(l) ΔH = kJ
The magnitude of ΔH is directly proportional to the amount of reactants or products.
A + 2 B ----> C ΔH = -100 kJ
1/2 A + B ----> 1/2 C ΔH = -50 kJ
ΔH for a reaction is equal in magnitude but opposite in sign for the reverse reaction.
C ----> A + 2 B ΔH = +100 kJ

49. Hess' Law
Hess' Law states that the heat evolved or absorbed in a chemical process is the same whether the process takes place in one or in several steps. This is also known as the law of constant heat summation.
Hess's law can be applied to calculate enthalpies of reactions that are difficult to measure.

50. Hess Law: Calculating enthalpy
For a chemical equation that can be written as the sum of two or more steps, the enthalpy change for the overall equation equals the sum of the enthalpy changes for the individual steps.
No matter how you go from given reactants to products (whether in one step or several), the enthalpy change for the overall chemical change is the same.
The heat transferred in a given change is the same whether the change takes place in a single step or in several steps.

51. Hess Law: Calculating enthalpy
In order to use Hess's Law:
1) If a rxn is reversed, the sign of ΔH is reversed.
    S(s) + O2 (g) ---> SO2 (g)  ΔH = -296 kJ
    SO2 (g) ---> S(s)+ O2 (g)  ΔH = +296 kJ
2) If the coefficients in a balanced equation are multiplied by some number, the value of ΔH must be multiplied by that same number.
S(s) + O2 (g) ---> SO2 (g) ΔH = -296 kJ
2 S(s) + 2O2 (g)--->2 SO2 (g) ΔH = (2)(-296 kJ) = -592 kJ

52. Hess Law: Calculating enthalpy
Example
S(s) + O2 (g) ---> SO2 (g)   ΔH = -296 kJ
SO2 (g) + 1/2 O2 (g) ---> SO3 (g)   ΔH = kJ
____________________________________________
S(s) + 1 1/2 O2 (g) ---> SO3 (g)    ΔH = kJ