Thermodynamics Heating/Cooling Curves

1. ThermodynamicsHeating/Cooling Curves 1 SAVE PAPER AND INK!!! When you print out the notes on PowerPoint, print "Handouts" instead of "Slides" in the print setup. Also, turn off the backgrounds (Tools>Options>Print>UNcheck "Background Printing")!

2. Credits 2 • Some information comes from Zumdahl, Steven, and Susan Zumdahl. Chemistry. Boston: Houghton Mifflin, 2003.

3. What is energy? 3 • One interesting definition of energy is that which is needed to oppose natural attractions (for example, gravity and electrostatic attractions) • We will define energy as the capacity to do work or to produce heat. • The law of conservation of energy states that energy can be converted from one form to another but can neither be created nor destroyed.

4. More Definitions 4 • Potential energy is energy due to position or composition. • The kinetic energy of an object is energy due to the motion of the object and depends on the mass of the object, m, and its velocity, v: KE = 1/2mv2 • Heat involves the transfer of energy between two objects due to a temperature difference.

5. Yup, Even MORE Definitions 5 • The system is the part of the universe on which we wish to focus attention. • The surroundings include everything else in the universe. • When a reaction results in the evolution of heat, it is said to be exothermic; that is, energy flows outof the system. • Reactions that absorb energy from the surroundings are said to be endothermic.

6. It seems easy, but… 6 • All of this is counterintuitive! • If a reaction is exothermic, a beaker will feel hot to you (the surroundings) but the temperature is actually dropping (in the system)! • If a reaction is endothermic, a beaker will feel cold to you (the surroundings) but the temperature is actually rising (in the system)! • The variable for heat is q, and when q = -x, the rxn is exothermic. When q = +x, the rxn is endothermic.

7. What is enthalpy? 7 • A less familiar property of a system is its enthalpy, H, which is defined as H = E +PV where E is the internal energy of the system, P is the pressure of the system, and V is the volume of the system. • BUUUUUUUT, a change in enthalpy is equal to heat (q) at constant pressures. • And when do we change the pressure in the lab? Never! So for us, enthalpy and heat are the same thing. (And therefore, +H means endothermic, and –H means exothermic)

8. What is calorimetry? 8 • The device used experimentally to determine the heat associated with a chemical reaction is called a calorimeter. • Calorimetry, the science of measuring heat, is based on observing the temperature change when a body absorbs or discharges energy as heat.

9. Heating/Cooling Curves 9 http://library.thinkquest.org/C006669/media/Chem/img/Graphs/HeatCool.gif

10. Uhhhh, OK…? 10 • Notice that there are two features to the graph on Slide 9, pieces with flat slopes and pieces with positive, linear slopes • Flat slopes occur during phase changes. • Upward slopes occur during temperature changes. • YES, this means that temperature does not change during a phase change!

11. Why Not? 11 • It’s easiest through an example: water boiling • For water to boil, the particles must go from a state where the particles have strong IMFs between them which holds them somewhat together into a state where the particles are totally free of any IMFs and act independently. • To break all of those attractions for every single molecule takes a lot of extra energy. ALL the molecules have to change state before the temperature will rise again!

12. 12 Enthalpy of Vaporization • The enthalpy of vaporization is the energy required to help a substance change from a liquid to a gas (+Hvap) or from a gas to a liquid (-Hvap). • It is usually given in kJ/mol, and therefore, the heat required to vaporize a substance can be calculated by q = nHvap, where n is moles and Hvap is the enthalpy of vaporization

13. Enthalpy of Fusion 13 • The enthalpy of fusion is the energy required to help a substance change from a liquid to a solid (-Hfus) or from a solid to a liquid (+Hfus). • It is usually given in kJ/mol, and therefore, the heat required to melt or freeze a substance can be calculated by q = nHfus, where n is moles and Hfusis the enthalpy of fusion

14. So what about when temp. IS changing? 14 • Different substances have different abilities for heating up quickly (or not) • The heat capacity, C, of a substance, which is a measure of this property, is defined as C = heat absorbed/increase in temperature • If the heat capacity is given per gram of substance, it is called the specific heat capacity, and its units are in J/g°C • If the heat capacity is given per mole of the substance, it is called the molar heat capacity, and it has the units J/mol °C

15. And how do I use this “specific heat”? 15 • During a temperature change ( NOT a phase change), the amount of heat being transferred is calculated as q = mCΔT, where m is the mass of the substance, C is the specific heat capacity, and ΔT is the change in temperature. • NOTE: The same substance has different C values for the different states of matter! Cice ≠ Cwater ≠ Csteam

16. Homework 16 • REMEMBER: There are two equations! One if for phase changes, and one is for temperature changes. Some questions will require a combination of both, and/or the use of the law of conservation of energy. • USE the internet to find values for the specific heat capacity and enthalpy of fusion or vaporization for each substance.

17. Homework Questions 17 • 1) How many joules are required to heat 250 grams of liquid water from 0C to 100C? • 2) How many joules are required to melt 100 grams of copper? • 3) How many joules are given off when 120 grams of water are cooled from 14C to -45C?