Energetics/Thermochemistry

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 • Endothermic • 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 • Exothermic • 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. Endothermic 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. Writing endothermic or exothermic reactions

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) • 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 heatThe 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= 200-20= 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= 200-20=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) 2. Bond Enthalpies: • 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 413 KJ/mol O2(g)O+O 495 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