Thermochem Flashcards by Jennie Rose

kinetic energy associated with

random motion of particles

thermal energy

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potential energy related to the

arrangement of particles

chemical energy

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an object, or collection of objects, being studied

system

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everything outside the system that can exchange

energy and/or matter with the system

surrounding

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heat is * and * property

energy transferred due to temperature difference

extensive property

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temperature is * and * property

measure of average ke or intensity of heat?

intensive property

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exothermic
endothermic
qsystem:

exo: q sys < 0
endo: q sys > 0

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energy transferred as heat that is required to raise

the temperature of 1 g of a substance by 1 K

specific heat capacity

J/g-K

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energy transferred as heat that is required to raise

the temperature of 1 mole of a substance by 1 K

molar heat capacity

J/mol-K

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if two objects have the same mass, the object having the larger specific heat capacity will undergo the * temperature change for a given amount of energy transferred

if two objects have different masses, but the same specific heat capacity, the * object will undergo a smaller temperature change for a given amount of energy transferred

if two objects have different masses and different specific heat capacities, the temperature change must be calculated from the *

smaller

larger

conservation of energy

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larger specific heat capacity = * temperature change

smaller

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q = mc(delta T)

if heat of vap, sub or fus; q = ?

q = heat of (v,f,s) * mass

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work associated with a change in volume (ΔV)
that occurs against a resisting external pressure (P)
formula and unit

pressure-volume work
P - m^3
(1 Pa = 1 kg/m-s^2) 
delta V = m^3
J = kg-m^2/s^2

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If work was done by
the system:
If work was done on
the system:

If work was done by
the system: w < 0 (expansion)
If work was done on
the system: w > 0 (compression)

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First law of thermodynamics:
Energy can neither be created nor destroyed.
It can only be transformed into another form of
energy through the interaction of *

heat (q),

work (w), and internal energy (U).

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First law of thermodynamics:
the energy change for a system (ΔU) is
the sum of the energy transferred as * and the energy transferred as * between the system and its surroundings

heat (q), work (w)

ΔU = q + w

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The * in a chemical system is the sum
of the potential and kinetic energies inside the system,
that is, the energies of the atoms, molecules, or ions in
the system.

internal energy (U)

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q>0 (+)
q<0 (-)
w>0 (+)
w<0 (-)

endothermic, U increases
exothermic, U decreases
work done on the system, U increases
work done by the system, U decreases

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The total energy of the universe is *

formula

constant
ΔUuniverse = 0
ΔUuniverse = ΔUsystem + ΔUsurroundings
ΔUsystem = - ΔUsurroundings

thermodynamic property wherein changes in these
quantities depend only on the initial and final states
examples:

STATE FUNCTIONS

ΔT, ΔU, ΔV, ΔH, ΔP

quantities that depend on the pathway taken to
get from the initial condition to the final condition
examples:

PATH FUNCTIONS

q, w

measure of energy transferred as heat in constant pressure

formula:

enthalpy

ΔH = qp= ΔU + PΔV

SIGN CONVENTIONS
- ΔH =
+ΔH =

ΔH = exothermic process

+ΔH = endothermic process

If a system is at constant volume:

ΔUsystem =

ΔUsystem = q(system) + w(system) = qv

If the system is NOT at constant volume: | ΔUsystem =

ΔUsystem = qp + w(system)

``` qv = qp + w Since w = -PΔV and qv = ΔU, then ΔU = qp – PΔV qp = ? ΔH = ? ```

``` qp = ΔU + PΔV ΔH = ΔU + PΔV ```

ØAt constant volume: qv = ? ØAt constant pressure: qp = ?

ØAt constant volume: qv = ΔU ØAt constant pressure: qp = ΔH

ΔH = qp= ΔU + PΔV sign conventions *ΔH = exothermic process *ΔH = endothermic process

negative ΔH = exothermic process | positive ΔH = endothermic process

pure, unmixed reactants in their standard states* have formed pure, unmixed products in their standard states unit:

STANDARD REACTION ENTHALPY, ΔrH° | units: kJ/mol (or kJ/mol-rxn)

the most stable form of the substance in the physical state that exists at a pressure of 1 bar and at a specified temperature (usually 25 °C)

STANDARD STATE

enthalpy changes are * to each reaction § * equation + states are important!! § enthalpy change depends on the number of * of reaction (number of times it is carried out) § for chemical reactions that are the reverse of each other, ΔrH° values are numerically the *, but *in sign

specific balanced moles same, opposite

heat is shown as * of the chemical reaction | heat transferred is an * property

part | extensive

measurement of energy evolved or absorbed as heat in a chemical or physical process apparatus: * à types of calorimetry: § constant pressure: § constant volume:

calorimetry calorimeter § constant pressure coffee cup calorimeter § constant volume bomb calorimeter

measure the amount of energy transferred as heat under constant-pressure conditions = enthalpy change, ΔH formula

àcoffee cup calorimeter | qrxn + qsoln = 0

evaluate the energy released by the combustion of fuels and the caloric value of food constant volume = no P-V work done on/by the system = energy is transferred as heat only = internal energy, ΔU formula

bomb calorimeter qrxn + qbomb + qwater = 0

``` qrxn + qbomb + qwater = 0 qrxn = unknown qwater = * qbomb = * Cbomb = * of the bomb in * ```

``` qwater = mwater Cwater ΔT qbomb = Cbomb ΔT ``` heat capacity in J/K

How to calculate for ΔrH°

1. Using bond energies 2. Using Hess’s Law 3. Using heats of formation

Bond energy During a chemical reaction, reactant bonds are broken and product bonds are made. Breaking bonds requires energy Making bonds releases energy

(endothermic, BE > 0) | exothermic, BE < 0

Using bond energy theorem, how to compute for | ΔrH°?

ΔrH° = Σ𝐵𝐸 − Σ𝐵𝐸

Breaking bonds requires energy : endothermic, * Making bonds releases energy : exothermic, * If bonds are stronger in products than reactants, ΔrH° is * If bonds are stronger in reactants than products, ΔrH° is *

BE > 0 BE < 0 negative positive

if a reaction is the sum of two or more other reactions, ΔrH° for the overall process is the sum of the ΔrH° values of those reactions

Hess's law

regardless of the multiple stages or steps of a reaction, the total enthalpy change for the reaction is the sum of all changes a manifestation that enthalpy is a *

skl HAHAHAHA state function!

àenthalpy change for the formation of 1 mol of a compound directly from its component elements in their standard states

STANDARD MOLAR ENTHALPY OF FORMATION, 𝚫fH0

ΔfH0 for an element in its * state is zero

standard

ΔfH0 can often be used to compare | * of related compounds

stabilities

enthalpy change for a reaction, ΔrH0, under standard conditions may be calculated from ΔfH0 values equation: the equation is a consequence of the definition of ΔfH0 and Hess’s Law

ΔrH° = ΣnΔfH° − ΣnΔfH°

is the reaction product- or reactant-favored at equilibrium? LOOK AT THE ΔrH0 VALUE! product-favored = * reactant-favored = *

``` product-favored = negative ΔrH0 (exothermic) reactant-favored = positive ΔrH0 (endothermic) ```