1. Chapter 7: Energy and Chemical Change Energy is the ability to do work and supply heat Work is motion against an opposing force Kinetic energy is the energy an object has because of its motion For an object of mass m with velocity v

2. The law of conservation of energy states that energy cannot be created or destroyed This can be applied to the collision of two particles with only kinetic energy Potential energy (PE) is the energy of position or internal arrangement Fast moving A transfers energy to B by colliding. The total energy is conserved.

3. KE can be converted into PE and vice versa When the child is at points (a) and (c) they have only PE; at point (b) only KE. Total energy is conserved (PE + KE = constant). When fully compressed or extended only PE; at natural length only KE. Total energy is conserved.

4. Work is required to pull the negatively charged electron away from the positively charged nucleus The gain and loss of PE can be summarized –Pushing or pulling an object against an opposing force requires energy. The objects PE will rise. –When the opposing force is not resisted, the object’s PE falls The SI unit of energy is the joule (J)

5. A 2 kg object moving at 1 meter per second has 1 J of kinetic energy You may also encounter the calorie (cal) The dietary Calorie (note capital), Cal, is actually 1 kilocalorie When a cold and hot object come into contact, they eventually reach thermal equilibrium (the same temperature)

6. The energy that is transferred as heat comes from the object’s internal energy The energy associated with the motion of the object’s molecules is referred to as its molecular kinetic energy The internal energy is often given the symbol E or U We are interested in the change in E:

7. The internal energy change is positive if the system absorbs energy from the surroundings and negative if it releases energy to its surroundings The temperature of an object is related to the average kinetic energy of its atoms and molecules The temperature for curve (1) is lower than for curve (2) because the average kinetic energy is lower.

8. Heat is a transfer of energy due to a temperature difference a)Isolated warm (left) and cold (right) objects b)Thermal contact is made: thermal energy is transferred from left to right c)Thermal equilibrium: the same average KE for molecules in both objects

9. The energy of an object depends only on its current condition The current condition is called the state Internal energy is a state function because it is a measure of energy An important property of state functions is that they are independent from the mechanism or method by which a change occurred

10. The object we are interested in is called the system Everything outside the system is called the surroundings A boundary separates the system from the surroundings The system and surroundings together are called the universe Systems are classified according to what can cross its boundary

11. –Open systems can gain or lose mass and energy across their boundaries –Closed systems can absorb or release energy, but not mass, across their boundaries –Isolated systems cannot exchange energy or matter with their surroundings Consider heat flowing between the system and surroundings The sign of the heat change is used to say whether it was gained or lost

12. –When heat is gained by an object, it is written as a positive number –When heat is lost by an object, it is written as a negative number A spontaneous change is one that continues on its own –Heat flows spontaneously from a warmer to colder object The heat directly gained or lost by an object is directly proportional to the temperature change it undergoes

13. The object’s specific heat (C) relates the heat (q) to the objects temperature change The heat capacity is the amount of heat needed to raise the object’s temperature by one degree Celsius and has the units J/°C C is an extensive property that can be determined from experiment, and is proportional to the sample mass

14. The specific heat capacity (s) is an intensive property, and is unique for each substance For example A large specific heat means the substance releases a large amount of heat as it cools. The heat absorbed or released by an substance is then:

15. Example: The temperature of 251 g of water is changed from 25.0 to 30.0 °C. How much heat was transferred to the water? ANALYSIS: Connect heat to the temperature change. SOLUTION: Note: Heat was absorbed because q is positive

16. Chemical bonds are the net attractive force between nuclei and electrons in compounds Breaking a chemical bond requires energy Making a chemical bond releases energy The potential energy that resides in chemical bonds is called chemical energy The attraction of the electrons for the nuclei in the hydrogen molecule is strong enough to overcome the nucleus-nucleus and electron-electron repulsions.

17. Chemical reactions generally involve both making and breaking chemical bonds The net gain or loss of energy is often in the form of heat Any reaction where heat is a product is called exothermic Reactions that consume energy are called endothermic Reactions can release heat by replacing “weak” bonds with “strong” ones

18. The amount of heat absorbed or released by a chemical reaction is called the heat of reaction A calorimeter can be used to measure the heat of reaction Calorimeters are usually designed to measure heats of reaction under conditions of constant volume or constant pressure Pressure is the amount of force acting on a unit area:

19. Atmospheric pressure is the pressure exerted by the mixture of gases in the atmosphere At sea level the atmospheric pressure is about 14.7 lb/in² Other common pressure units are the atmosphere (atm) and bar: 14.696 lb/in² = 1.0000 atm = 1.0133 bar q v and q p are used to show heats measured at constant volume or pressure, respectfully

20. In reactions where gases are produce or consumed q v and q p can be very different If the volume change is, work (w) is Pressure-volume work (a) A gas confined under pressure. (b) The gas does pressure-volume work on the surroundings when it expands.

21. Note that the work of expansion is negative Work and heat are alternate ways to transfer energy Their sum is the change in internal energy the system undergoes This is a statement of the first law of thermodynamics, which says that energy cannot be created or destroyed

22. Heat and work are not state functions because they depend on the path between the final and initial state Heat and work depend on the path. Both paths give the same value for the internal energy change. In path 1, this appears as heat (q). In path 2 this appears mostly as work (w).

23. The heat produced by a combustion reaction is called the heat of combustion The heats are measured in closed containers because the reactions consume and produce gases The instrument used to measure these heats is called a bomb calorimeter The reaction is run at constant volume so that

24. Heats of reactions in solution are usually run in open containers at constant pressure They may transfer heat and expansion work A bomb calorimeter. The reaction chamber has a constant volume so no work is done. The heat released by the combustion is absorbed by the “bomb” and surrounding water.

25. The heat change measured at constant pressure is the enthalpy, H Enthalpy is also a state function – is negative for an exothermic process – is positive for an endothermic process The the difference in the values of the internal energy and enthalpy change can be large for reactions that consume or release gases

26. The amount of heat that a reaction produces or absorbs depends on the number of moles of reactant that react A set of standard states have been defined for reporting heats of reactions A coffee cup calorimeter can be used to measure heats of reaction at constant pressure. Heat can be released or absorbed, resulting in a change in temperature of the solution.

27. Standard thermodynamic states are: 1 bar pressure for all gases and 1 M concentration for aqueous solutions A temperature of 25 °C (298 K) is often specified as well The standard heat of reaction is the value of the enthalpy change occurring under standard conditions involving the actual number of moles specified the the equation coefficients

28. An enthalpy change for standard conditions is denoted For example, the thermochemical equation for the production of ammonia from it elements at standard conditions is: The physical states are important The law of conservation of energy requires

29. Enthalpy is a state function An enthalpy diagram is a graphical representation of alternate paths between initial and final states Remember to include the physical states of reactants and products in thermochemical equations. Two paths for the formation of carbon dioxide gas. Each give the same enthalpy change.

30. Enthalpy changes for reactions can be calculated by algebraic summation This is called Hess’s Law: The value of the enthalpy change for any reaction that can be written in steps equals the sum of the values of the enthalpy change of each of the individual steps. Enthalpy changes for a huge number of reactions may be calculate using only a few simple rules

31. Rules for Manipulating Thermochemical Equations: 1) When an equation is reversed the sign of the enthalpy change must also be reversed. 2) Formulas canceled from both sides of an equation must be for substances in identical physical states. 3) If all the coefficients of an equation are multiplied or divided by the same factor, the value of the enthalpy change must likewise be multiplied or divided by that factor.

32. An enormous database of thermochemical equations have been compiled: –The standard heat of combustion is the amount of heat released when 1 mol of a fuel completely burns in pure oxygen gas with all products brought to 25 °C and 1 bar Standard heats of combustion are always negative and produce water in liquid form

33. –The standard enthalpy of formation of a substance is the amount of heat absorbed when 1 mole of the substance if formed at 25 °C and 1 bar from its elements in their standard states The standard enthalpy of formation for elements in their standard states are zero These are the values most commonly used to calculated standard enthalpy changes for reactions Standard enthalpies of formation are given in Table 7.2 and Appendix C

34. Hess’s law can be restated in terms of standard enthalpies of formation: Example: Calculate the enthalpy of reaction for 2NO(g)+O 2 (g)  2NO 2 (g) ANALYSIS: Use Hess’s law and Table 7.2 SOLUTION: