Study Notes

Study Notes

CHEM 1101

Study Notes 1901 Thermochemistry Work 𝑤 = 𝐹 × 𝑑 (𝑓𝑜𝑟𝑐𝑒 × 𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒) Conversions 1𝑗𝑜𝑢𝑙𝑒 = 1𝑘𝑔𝑚2 𝑚−2 1 𝑐𝑎𝑙𝑜𝑟𝑖𝑒 = 4.184 𝑗𝑜𝑢𝑙𝑒𝑠 First Law of Energy is never created nor destroyed, it is conserved. Thermodynamics I.e. εE is constant. If system loses energy surroundings gain proportionate amount. ∆𝑈 = 𝑈𝑓𝑖𝑛𝑎𝑙 − 𝑈𝑖𝑛𝑖𝑡𝑖𝑎𝑙 Summary of +ve q = heat gained -ve q = heat lost Energy Exchange +ve w = work done on system -ve w = work done by system +ve ∆𝑈 = energy gain -ve ∆𝑈 = energy loss State functions Absolute value of internal energy of system usually indeterminate. Internal energy is a state function. This means that it depends only on the state of the system not on the path at which the system arrived at that state. Q and W are not state functions i.e. Whether a battery shorts out or is discharged the change in energy is the same, but Q and w are different. Enthalpy Is equal to the internal energy of the system plus the product of pressure and volume ( H = U + PV). At constant pressure is equal to the amount of heat gained or lost. Enthalpy is an extensive property (i.e. it depends on the amount of reactants/products). Calorimetry The measurement of heat flow. This is measured by the magnitude of the temperature change resultant from heat flow. At a constant volume the ∆𝑈 and not the ∆𝐻 is always measured Heat Capacity The heat capacity is the amount of energy required to raise the temperature of a substance by one and the Specific Kelvin / one degree (Celsius). The specific heat capacity is the amount of energy required to raise the Heat Capacity temperature of 1g of substance by 1K.

𝑐=

𝑞

𝑚∆𝑇

∆𝑓 𝐻 is the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms. ∆𝑓 𝐻 Is the enthalpy of formation under standard conditions. Different states mean different enthalpies of formation. E.g. 𝐶3 𝐻8 + 5𝑂2 → 3𝐶𝑂2 + 4𝐻2 𝑂 = 𝐶3 𝐻8 → 3𝐶 + 4𝐻2 = −∆𝑓 𝐻 𝑜𝑓 𝐶3 𝐻8 = 3𝐶 + 3𝑂2 → 3𝐶𝑂2 = 3(∆𝑓 𝐻 𝑜𝑓 𝐶𝑂2 ) = 4𝐻2 + 2𝑂2 → 4𝐻2 𝑂 = 4(∆𝑓 𝐻 𝑜𝑓 𝐻2 𝑂)

Enthalpy of Formation

Spontaneous Processes Real Processes 2nd Law of Thermodynamics Boltzmann’s Equation

1

Total ∆𝑓 𝐻 = ∑ 𝑛∆𝑓 𝐻𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 − ∑ 𝑛∆𝑓 𝐻𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 ∴ = [3(−393.5) + 4(−285.8)] − [(−10385) + 5(0)] = 2220𝑘𝐽 Can be changed by different temperatures and pressures. For example ice → melts spontaneously when T>0 All real processes are irreversible. Any irreversible process results in an overall increase in entropy, whereas a reversible process results in no change in entropy. Total entropy of the universe increases in any spontaneous reaction. Molecules are in constant motion. If we consider one mole of a particular gas then, at any instant, the particular set of positions and energy’s of the individual gas molecules is called a microstate of the thermodynamic system. A microstate is a single possible arrangement of the positions and kinetic energies of the gas molecules when the gas is in a specific thermodynamic state. 𝑆 = 𝑘𝑙𝑛𝑊 (𝑆 = 𝑒𝑛𝑡𝑟𝑜𝑝𝑦 𝑊 = 𝑛𝑜. 𝑜𝑓 𝑚𝑖𝑐𝑟𝑜𝑠𝑡𝑎𝑡𝑒𝑠 𝑘 = 1.38 × 10−23 𝐽𝐾 − ) Entropy is a measure of the number of microstates that are associated with a particular macroscopic state. Entropy increases for processes where gases are formed from either solids or liquids, liquids of solutions are formed from solids or the number of gas molecules increases during a chemical reaction. Entropies increase with mass and with the number of atoms in a molecules as well as vibrational degrees of freedom.