Book: Ch. 6 Flashcards
Define: thermodynamics
Study of energy and its transformations.
Define: thermochemistry
Branch of thermodynamics that deals with heat in chemical and physical change.
Define: system vs. surroundings [in thermodynamics]
System: the part of the universe we are focusing on.
Surroundings: everything else.
Define: internal energy (E) of a system
The sum of all the energies of all the components of a system.
What is the equation for change in the internal energy of a system?
∆E = E_final - E_initial = E_products - E_reactants
Define: heat (q)
Thermal energy. The energy transferred as a result of a difference in temperature between the system and the surroundings.
What is the equation for the total change in a system’s internal energy in terms of q and w?
∆E = q + w
Define: law of conservation of energy (first law of thermodynamics)
∆E_universe = ∆E_sys + ∆E_surroundings = 0
Define: joule (J)
SI unit of energy. 1 J = 1 kg m²/s²
Define: calorie (cal)
1 cal = 4.184 J
Define: state function
A property dependent only on the current state of the system, not on the path the system takes to reach that state. Examples: internal energy, composition, pressure, volume, temperature.
Define: pressure-volume work (PV work)
A type of mechanical work done when the volume of the system changes in the presence of an external pressure: w = -P∆V
Define: enthalpy (H)
For reactions at constant pressure, enthalpy eliminates the need to deal with PV work. H = E + PV
Define: change in enthalpy (∆H)
Change in the system’s internal energy plus the product of the pressure and the change in volume: ∆H = ∆E + P∆V
How is ∆H related to heat?
∆H = q_p, or enthalpy equals the heat absorbed or released at constant pressure.
Define: exothermic process
Releases heat as a product. Decrease in the enthalpy of the system: ∆H < 0
Define: endothermic process
Absorbs heat as a reactant. Increase in the enthalpy of the system: ∆H > 0
Define: heat capacity (C)
The quantity of heat required to change the temperature of an object by 1 K: q/∆T = constant = heat capacity (C) [J/K]
Define: specific heat capacity (c)
Quantity of heat required to change the temperature of 1 gram of a substance or material by 1 K: specific heat capacity (c) = q/m∆T [J/gK]
Given the specific heat capacity of an object being heated, the mass or temperature change can be found along with the heat absorbed by what equation?
q = cm∆T
Define: molar heat capacity (C_m)
Quantity of heat required to change the temperature of 1 mole of a substance by 1 K: C_m = q/n∆T [J/mol K]
Define: calorimeter
A device used to measure the heat released or absorbed by a physical or chemical process.
What is a coffee-cup calorimeter?
For processes that take place at a constant pressure, heat transferred may be measured with one of these: you use the equation -q_sys = q_water [in the calorimeter]
How is the heat released during the combustion of a sample measured with a bomb calorimeter?
-q_rxn = q_calorimeter, where q_calorimeter = C_calorimeter ∆T_calorimeter
Define: thermochemical equation
Balanced equation that includes the enthalpy change of the reaction (∆H).
Define: Hess’s law
∆H_overall = ∑ ∆H_i
Define: standard states
Established set of specific conditions used to reduce the number of variables when studying a reaction.
Define: standard enthalpy of a reaction
∆Hº_rxn is the enthalpy change of a reaction when it is measured at the standard state. It is also called the standard heat of reaction.
Define: formation equation
1 mol of a compound forms from its elements. Like C + 2H₂ → CH₄
Define: standard enthalpy of formation
∆Hº_f is the enthalpy change for a formation equation when all the substances are in their standard states.
What is the equation for the standard enthalpy of a reaction?
∆Hº_rxn = ∑m∆Hº_f (products) - ∑n∆Hº_f (reactants), where m and n are the amounts of products and reactants given by the coefficients in the balanced equation.
