1. Visualizing Molecular Motion Figure 6-14 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 6Slide 1 of 41 6-7 Kinetic Molecular Theory of Gases Particles are point masses in constant, random, straight line motion. Particles are separated by great distances. Collisions are rapid and elastic. No force between particles. Total energy remains constant.

2. Gas Nonideal Behaviour Class discussion of evidence for. Intermolecular forces. Finite molecular size. Condensed phases – they exist! (Cogito ergo sum? H 2 O(g) liquefies therefore I am?) V vs T graph for both ideal and real gases.

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4. Synthetic Chemistry Chemical reactions are used to transform elements and simple compounds into more complex and important substances such as tetracycline (antibiotic), ammonium nitrate (fertilizer) and epitaxial silicon (computer chips). The synthetic chemistry of carbon is extremely varied. Nature produces complex organic molecules, often hydrocarbons, which can also be used as important energy sources.

5. Thermochemistry Chemical reactions can be employed to supply energy in various forms (usually heat, light and electricity). Eg. Ethanol – can be used as a solvent, as a synthetic building block or to provide energy through direct combustion, use in fuel cells, etc. C 2 H 5 OH(l) + 3O 2 (g) → 3CO 2 (g) + 3H 2 O(l)

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7. How much heat is released depends on the amount of ethanol burned and can be reported in kJ∙mol -1, kJ∙ kg -1, etc. The quantity of heat produced will also vary with the temperature of the reactants and products as well as their phases. When one mole of C 2 H 5 OH is burned under constant pressure conditions about 726 kJ of energy are transferred (lost) to the surroundings. Write ΔH o Combustion = -726 kJ∙mol -1

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9. In thermodynamics the system is a particular set of substances or objects being studied. The surroundings (of our system) is made up of the rest of the universe. There are three types of system. Open system – can exchange matter and energy with the surroundings. Closed system –can exchange energy but not matter with the surroundings. Isolated system – no exchange of matter or energy with the surroundings is possible (Hot coffee in a thermos bottle).

10. Slide 10 of 57 6-1 Getting Started: Some Terminology Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 7 FIGURE 7-1 Systems and their surroundings

11. Energy Units! Energy units – usually joules, J, kJ, MJ. 1 J = 1 kg∙m 2 ∙s -2 and 1 calorie = 4.184 J (the dietary calorie is 1000 times larger!) Energy changes are associated with both chemical and physical processes. For example, to melt a solid (fusion) or to vaporize a liquid energy must be supplied to our system. H 2 O(s) → H 2 O(l) ΔH o Fusion = 6.01 kJ.mol -1 H 2 O(l) → H 2 O(g) ΔH o Vaporization = 44.0 kJ∙mol -1 Which process in important at “Tim’s” ?

12. Heating and Cooling Curves Heating and cooling curves are simple to construct unless there are phase changes to consider. In class we will construct two heating curves for constant pressure processes. Curve 1: Temperature vs heat supplied (q) for the heating of Al(s) from 15 O C to 111 O C. Curve 2: Temperature vs heat supplied for the transformation of ice at -45 O C to steam at a temperature of 165 O C.

13. Specific Heats In drawing the two heating curves we considered specific heats (specific heat capacities) – the amount of heat needed to raise the temperature of 1 g of a substance in a particular phase by 1 degree. We see q = mc∆t or q = ms∆t …….. When employing these equations m, s and ∆t must refer to a particular substance and a particular phase.

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