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Thermochemistry Lecture Notes

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Queen's University
CHEM 112
John Carran

Systems: • Open system – open to transfer of heat and matter • Closed – allows transfer of energy, but not matter • Isolated system – closed to transfer of both energy and matter Units: • Energy, U (state function) • Work, w (not state function; half dependent) • Kinetic Energy, eK • Potential Energy, V • Heat, q (not state function) • Calorie, cal – quantity of heat required to change temperature of 1 gram of water by 1 degree C • Joule, J – SI unit for heat (1 cal = 4.184J) ***Pay attention to capital or lowercase symbol/abbreviation*** Thermo Energy • kinetic energy associated with random molecular motion • Generally proportional to temperature • Intensive property Internal energy is transferred between a system and its surroundings as a result of temperature difference Heat flows from hotter to cooler Heat Capacity • quantity of heat required to change temperature of a system by 1 degree C • Specific heat, c (q=mcT)  always positive • Molar heat capacity (q= CT) In the law of conservation of energy, total energy remains constant - qsystem qsurroundings - qsystem -qsurroundingsroundings is absorbing energy) ***look at calculation example in journal*** Bomb Calorimetry – energy transferred only by heat, no work involved The “Coffee-Cup” Calorimeter – assumed to be completely insulated – open to the atmosphere/open to expansion ` *** pay attention to how many moles you are solving for, so that the balanced equation reflects the same values*** Pressure-volume work (involves volume change) • F x d • (m x g) x h • Compression = negative h • Expansion = positive h First Law of Thermodynamics • a system contains only internal energy • a system does not contain heat or work • only occur during a change in the system • delta U = q+ w *** know sign conventions used: when is q and w positive/negative*** • Hard to measure individual state functions, easiest to find change in the state function • In a reversible, you get more work out of the system than a regular system • Constant volume = no work performed *** Know figure 7-15 *** Hess’s Law: • Delta H is an extensive property, therefore how much substance is present. In other words, make sure the coefficients in the reactions properly reflect the molar value • State function – half independent • Enthalpy change for the overall process is the sum of the enthalpy changes for the individual steps (look at “if you’re wondering” box in 7.1) • reference/standard state – most common state that the element occupies at standard temperature • for ionic reactions, you need to make a reference coefficient. (In the power point, pg 53, the hydrogen is assigned an arbitrary value of 0) Enthalpies of formation • can be positive or negative • you must be able to work out a complete reaction, given partial information THERMO 2 – W11L1 Spontaneous process • occur without intervention • directional (entropy, S, provides basis for predicting direction) • equilibrium is established when driving force or tendency of the spontaneous change is expended • Direction of a spontaneous change is often from higher to a lower energy, but there are exceptions Non-spontaneous reactions • require the system to by acted on by an external agent • examples include: ball not rolling up a hill, gases do not collected at one end of a container, heat does not flow from cold body to hot body, water doesn’t freeze at temperature above 0 C. • For chemical systems, the internal energy, U, is equivalent to potential energy Spontaneity: need 2 things to predict: - change in enthalpy, delta H - change in entropy, delta S Thermo 3 – W11L2 Entropy: • states – microscopic energy levels available in a system • microstates, W (omega) – particular way/combinations in which particles are istributed amongst the states. Number of microstates = W • Boltzman constant, k – gas constant per molecule (RNA) • delta S = qrev / T • generally increase entropy: o Pure liquids from solids o Gases formed from solids or liquids o # mo
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