2. Energy Chapter 15

3. Mr. Smith’s favorite definition for chemistry How energy interacts with matter

4. What is Energy? Energy is the capacity to do work. Unit- Joule (J) or Calories (cal) 1cal=4.184J Physical and chemical changes always involve energy changes. Energy can be absorbed or released during these changes.

5. Absorbing and Releasing energy Any change in matter in which energy is released is exothermic. Examples-freezing water, heat packs Na (s) + Cl 2(g)  NaCl (s) + ENERGY Any change in matter in which energy is absorbed is endothermic. Examples-melting ice, boiling water sunlight + 6CO 2 (g) + H 2 O(l)  C 6 H 12 O 6 (aq) + 6O 2 (g)

6. Conservation of energy Energy is never lost, only converted to another form of energy Not entirely true, the exception is... So the real rule is –Conservation of mass and energy

7. Forms of energy

8. Heat – form of energy Kinetic energy – energy of movement (temperature) Heat is energy that is in the process of flowing from a warmer object to a cooler object. Measured in Joules (J) or calories (cal) Link

9. Heat The flow of thermal energy from one object to another. Heat always flows from warmer to cooler objects. Ice gets warmer while hand gets cooler Cup gets cooler while hand gets warmer

10. Heat Capacity Some things heat up or cool down faster than others. Sand heats up and cools down faster than water

11. Specific Heat Capacity (Cp) is the amount of heat required to raise the temperature of 1 g of a material by one degree (C or K). 1) Cp water = 4.184 J / g°C 2) Cp sand = 0.664 J / g ° C This is why sand heats up quickly during the day and cools quickly at night while water takes longer.

12. Why does water have such a high specific heat? Water molecules have strong bonds with each other; therefore it can absorb more heat. Metals have weak bonds and can not absorb as much heat. water metal

13. How to calculate changes in thermal energy Q = m C p ΔT Q = change in thermal energy/heat (J) m = mass of substance (g) C p = heat capacity of material (J/g°C) Δ T = change in temperature (T f – T i )

14. General Rules General Rule #1 - The greater the specific heat value, the less the temperature will rise when a given heat energy is absorbed Not only does the specific heat value describe how much heat may be absorbed by a substance before its temperature rises, it also describes the ability of a substance to deliver heat to a cooler object. General Rule #2 - As the specific heat value decreases, the ability to deliver heat to cooler object increases. (So lower C feels hotter) For example, imagine holding two hot pieces of metal - X (C=2 J/g°C) and Y (C=3 J/g°C). If the hot piece of metal X was held in one hand and the hot piece of metal Y in the other hand, the hand holding the metal X would get hotter. Because metal X has a specific heat less than metal Y, the metal X sample transfers heat to a cooler object (your hand) more readily.

15. General Rules Why do different materials possess different specific heat values? One reason for the variation is that each substance is made up of atoms that have different masses. The mass of each copper atom is larger than the mass of each aluminum atom, for example. Therefore a given mass (such as 58 g of copper) has fewer atoms than the same mass of aluminum. When heat is added to 58 g of copper, fewer atoms need to be put in motion (remember temperature is related to kinetic energy). Thus, less heat is needed to increase the kinetic energy of the atoms in the sample, and raise the temperature by 1°C. As a result, the specific heat value for copper is lower than the specific heat of aluminum. Copper and zinc have identical specific heat values. This is due to the similar mass of the atoms. General Rule #3 - The larger the metal atom, the lower is specific heat value.

16. How to measure energy? Cannot measure energy directly But you can measure temperature –Indirectly gives you heat (Q) via Q=mCΔT Calorimetry

17. Heating Curve of Water Animation

18. Cooling Curve of Water

19. Heat Curve – shows physical changes caused by heat

20. During Temperature Change: Kinetic energy changes –speed of molecules Only 1 state of matter present (solid, liquid, OR gas)

21. During Phase Change: Mixture of 2 states of matter NO temperature change (constant) ex. Boiling point, melting point  Heat curve levels out  Molecules are not moving faster  Energy is being used to loosen/break IMF’s

22. How to measure heat energy Calorimetry - Chapter 15.2

23. System and Surroundings Chemists use the terms system and surroundings to keep track of energy changes. System - all of the components that are being studied at a given time. Surroundings - everything outside the system.

24. How does energy changes take place? Energy can be changed into different forms. Example: Potential energy changed to kinetic energy. Energy can be transferred from one thing to another. Example: The heat given off by a hot cup of coffee (system) is transferred to the surrounding air (surroundings).

25. A calorimeter is used to help measure the specific heat of a substance. First, mass and temperature of water are measured Then heated sample is put inside and heat flows into water Δ T is measured for water to help get its heat gain This gives the heat lost by the substance Knowing its Q value, its mass, and its Δ T, its Cp Cp can be calculated

26. A 150 g sample of water (initially at 45.0°C) is mixed with a 200.g sample of water (initially at 84.0°C). Specific Heat of water is 4.18J/g°C. Find the final temperature of the system. (67.3°C)

27. A 82.12 g sample of an unknown element (initially at 757.00 o C) is dropped into 354 g of water (initially at 25.11 o C). The final temperature of the system is 50.66 o C. The specific heat of water is 4.18J/gC. Find the specific heat of the unknown element.