Chem Notes Ch 8 Flashcards
State function
Property whose value only depends on the initial and final states of the system
Path functions
Properties that depend on the transition of a system from the initial to the final state.
1st law of thermodynamics
Energy cannot be created or destroyed (conservation of energy). It can only be transformed or transferred from one location to another
Universe
Made of two components, the system and the surrounding
Enthalpy (H)
Amount of heat energy contained within a system. Δ H = H products - H reactants
Endothermic
Process in which heat is transferred from surroundings to the system (surrounding will feel cold). Phase change from solid to liquid to gas. Phase change from solid to gas. ΔH>0
Exothermic
Process in which heat is transferred from system to the surroundings (surrounding will feel hot). Phase change from gas to liquid to solid. Phase change from gas to solid. Combustion reaction. ΔH<0
Conduction
The transfer of heat via direct contact
Convection
The transfer of heat due to motion of a liquid or gas
Radiation
The transfer of heat via electromagnetic radiation
Use the property of work to determine whether an expansion or compression is happening
Δ E = q + w
w = -PΔV
ΔE = change in internal energy q = heat w = work P = pressure (always a positive value) ΔV = change in volume
If work is a positive value, then ΔV is a negative value, and the surrounding is doing work on the system to compress it.
If work is a negative value, then ΔV is a positive value, and the system is doing work on the surrounding to expand it.
Under constant pressure systems such as the case is in calorimetry, q and H are the same.
\+q = Endothermic -q = Exothermic \+w = Compression -w = Expansion -ΔE = lost energy to the surrounding
Specific heat (C)
The amount of energy that is required to raise the temperature of 1.0 gram of a substance by 1 ° Celsius. The units for specific heat are J/ g°C. Varies with substance and phase.
q = mCΔT
qsystem + qsurrounding = 0
qsystem = - qsurrounding
q = heat absorbed or released by substance m = mass of the substance (measured in grams) C = specific heat of substance ΔT = change in temperature of substance
qcal = C (heat capacity of calorimeter in kJ/°C) x ΔT qrxn = - qcal
For phase changes, q = (m)(ΔH). ΔH is expressed in kJ/mol. If the phase change involves conversion of solid to liquid, then ΔHfusion will be used. If the phase
change involves conversion of a liquid to a gas, then ΔHvaporization will be used.
Calculating ΔH° using Bond Energies/Bond Enthalpies
Breaking bonds requires energy (endothermic: ΔH > 0) while forming bonds releases energy (exothermic: ΔH < 0). When atoms are bonded together, they are at a lower energy state (more stable) than when they are alone.
ΔH°rxn = ∑ 𝚫𝐇° 𝑩𝒐𝒏𝒅𝒔 𝒃𝒓𝒐𝒌𝒆𝒏 (𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕𝒔) − ∑ 𝚫𝐇° 𝑩𝒐𝒏𝒅𝒔 𝒇𝒐𝒓𝒎𝒆𝒅 (𝒑𝒓𝒐𝒅𝒖𝒄𝒕𝒔)
Calculating ΔH° using enthalpies of formation
The change in enthalpy that happens when one mole of the substance is made in its standard state.
Diatomic molecules in their standard state (all in gas form): Have no fear of ice cold beer. H2 (g), N2 (g), F2 (g), O2 (g), I2 (s), Cl2 (g), Br2 (l).
Standard enthalpy of formation for an element in its standard state is 0
ΔH°rxn = ∑ n𝑯°𝒇 (𝒑𝒓𝒐𝒅𝒖𝒄𝒕𝒔) − ∑ n𝑯°𝒇(𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕𝒔)
Calculating ΔH° using Hess’s law
Sum of enthalpy changes for its intermediate reactions. Reverse reaction = -ΔH°. Multiply equation by co-efficient = ΔH° x (coefficient). Divide equation by co-efficient = ΔH° divide by (coefficient).
