1. Energetics/Thermochemistry

2. Outline for Energetics
1. endothermic and exothermic reactions
2. units of energy
3. specific heat, and molar heat
4. molar heats of combustion
5. heats or enthalpies of formation ∆Hf
6. Hess’s Law
7. Bond enthalpies (juniors only)

3. Endothermic and exothermic reactions
Reactions which absorb energy or have energy added in order to occur.
Products have more stored energy than reactants
Products are less stable than reactants
Writing the reactions
kJ are written on the reactant side
∆H is positive

4. EXAMPLE OF AN ENDOTHERMIC REACTION

5. Endothermic and exothermic reactions continued
Reactions which release energy
Products have less stored energy than the reactants
Products are more stable than the reactants
Writing the reactions:
The kJ are written on the products side
∆H is negative

6. EXAMPLE OF AN EXOTHERMIC REACTION

8. Writing endothermic or exothermic reactions
Example:
kJ +2H2O2H2 + O2
The kJ are written on the reactant side.
2H2O2H2 + O2 ;∆H=kJ The ∆H is positive.
Exothermic
Example:
2H2 + O2  2H2O + kJ
The kJ are written on the product side.
2H2 + O2  2H2O;
∆H = - kJ The ∆H is negative.

9. Enthalpy or change in heat
Symbol for enthalpy is ∆H.
- ∆H means energy is released such as with an exothermic reaction.
+ ∆H means energy is used or absorbed such as with an endothermic reaction.

10. Units of Energy Joule (J) (the unit used in the math) calorie (c)
1newton.meter2/second2
Is the SI unit for energy
4.18J=1 calorie
1000J = 1 kilojoule (kJ)
calorie (c)
energy required to raise the temperature of 1 gram of water by 1◦C
1000calories = 1kilocalorie (kcal) or 1 food calorie (Calorie)
0.24cal = 1 joule

11. Bomb calorimeter

12. Math with energetics

13. Specific and molar heat The heating or cooling of a substance
1)Heat Capacity- the amount of energy a substance can absorb before its temperature is increased. General equation is:
C = heat absorbed/increase in temp
a) Molar heat capacity: the energy required to raise the temp of 1 mole of a substance by 1◦C. (units =J/mol·C)
b) Specific Heat: the energy required to raise the temperature of 1 gram of a substance by 1◦C
symbol= cp units= J/g°C
examples: 4.18J/gC for H2O, 0.45J/gC for Fe, 0.71J/gC for carbon.

14. Equation for Specific heat
Equation using specific heat
q= cp x m x ∆T where:
cp is specific heat
m=mass in grams
∆T= change in temp.
q= energy in joules

15. Problems with specific heat
Example 1: Find the energy needed to raise the temp of 5.00x102ml of water from 20. C to 100. C. Assume no energy is lost to the surroundings.
Q=cpx m x ∆T Q=?, m= 5.0x102g (1g=1ml for water), ∆T =100-20= 80 C, cp=4.18J/gC
Substitute into the equation:
Q= 4.18J/g C x 5.0x102g x 80. C
Q= 1.7x105J or 170kJ

16. Problems with specific heat
Example 2: A 20.0 g metal sample is heated to 200 C and then dropped into 100.ml of water. Both the metal sample and the water ended up with a final temp of 20 C. Find the metal’s specific heat if the water was 15 C before the metal was placed into it.

17. Example 2 continued In order to find the cp of the metal, there is:
m= 20.0g. ∆T= = 180 C. but Q=?, and cp=?
There are too many variables. However, the metal released its heat into the water so we can find the Q by finding the energy that went to heat the water.
Step 1: Find the energy to heat the water:
Q= ?, m= 100g H2O, cp= 4.18J/gC, ∆T= 20-15C= 5C
Q= 4.18J/gC x100g x 5C Q= 2090 J or 2.1 x 103J
Step 2: Find the specifc heat of the metal:
Q= 2090J, m= 20.0g, cp = ? ∆T= =180C
2090J= cp x 20.0g x 180C, cp= 0.58J/gC

18. Heats or Enthalpies of Formation ∆Hf◦
What is a ∆Hf ◦?
Used to calculate the energy involved in a reaction with out experimenting.
It is the energy content for one mole of a compound.
It is the energy involved in making (forming):
one mole of a compound
From its simplest elements
At 25 C and 1 atm.
These values are used to determine the ∆H for a reaction.
Writing the equation for a ∆Hf
examples; H2O(g): H2 + ½ O2 H2O;
∆Hf =- 242kj/mole This value is from an appendix ◦

19. Bond Enthalpies Bond Enthalpies
Energy is required to break bonds, energy is released when making bonds.
*Exothermic: more energy was released when making bonds in the products than what energy was absorbed to break bonds in the reactants.
*Endothermic: more energy was required to break the bonds in the reactants than what energy was released when making bonds in the products.

20. Bond Enthalpies(continued)
The average energy required to break a covalent bonds.
25 degrees Celsius, 1 ATM, always work w/ gases
-compound is turned into single gaseous atoms, not its simplest, stable form in nature
Ex: C-H(g)C+H KJ/mol
O2(g)O+O KJ/mol
Single bonds require less energy to break than double bonds, < triple bonds.
EXAMPLES: WRITE ON BOARD

21. Molar Heats of Combustion
1.Combustion reactions:
a) requirements for combustion
i) fuel
ii) ignition
iii) O2
b) products of complete combustion
i) energy & light
ii) stable compounds, each w/ oxygen in it
CCO2, HH2O, S SO2

22. Piston in the internal combustion engine

23. continued c) Writing Combustion Reactions
CH 4(g)+2O2  CO2(g)+3H2O(g)+KJ
2. Incomplete Combustion:
a) Why this happens: i) not enough O2
ii) not enough time
iii) not enough surface area
b) products made from incomplete combustion:
CO, ash, soot as well as CO2 and H2O