Enthalpy 3

THREE types of systems

1) open system 2) closed system 3) isolated system

The enthalpy change for a reaction :

>>depends on the STATE of the REACTANTS and PRODUCTS (i.e. solid, liquid, or gas) .....2H2O(g) ΔH = − 802.4 kJ/mol .... 2H2O(l) ΔH = − 890.4 kJ/mol

By convention q & w POSITIVE SIGN WHEN:

>q :the system gains heat(HEAT GOING IN ) >w:when work is done ON the system (WORK GOING IN).

State functions 3

Changes in Energy, Pressure, Volume, and Temperature

constant-volume bomb or bomb calorimeter.

Constant volume calorimetry is carried out in a device ( isolated system) >VOLUME IS CONSTANT but pressure is not. >typically used to determine HEATS OF COMBUSTION EX. The sample to be combusted is in a small cup : - (the bomb). -cup is surrounded by a known amount of water ***When combustion is ignited, the heat generated is ABSORBED by the water, causing a measurable temperature change in the water.

EXOTHERMIC

Hproducts(heat) < Hreactants ΔH = −890.4 kJ/mol >heat GIVEN OFF BY the surroundings

ENDOTHERMIC

Hproducts> Hreactants(heat) ΔH = +6.01 kJ/mol >heat ABSORBED BY THE SYSTEM from the surroundings

Hess's Law

If a reaction is carried out in a series of steps, DH for the reaction will be equal to the sum of the enthalpy changes for each of the chemical steps. >Recall that DH is a state function, so only the initial and final states matter, not the path taken to get there. DH = Hfinal - Hinitial So if we want to figure out DH for H3-H1 We can do that by summing steps in an alternative pathway: = (H3-H2) + (H2-H1)

Units of energy

SI unit is the joule (J). 1J = 1kg *m2 /s2

Universe

System + Surroundings

"coffee-cup" calorimeter 2

The REACTANTS & PRODUCTS are the SYSTEM being studied. The WATER they are dissolved in, and the calorimeter, are part of the SURROUNDINGS. qsys = -CsmDT = -qsurr I.E. qrxn = -CsmDT = -qsoln >Since we are measuring the heat gained or lost by the solution (surroundings), we must change the sign to get the correct answer for qrxn .

ΔHf°

The most stable elemental form at standard conditions is used for the formation reaction. **For these forms, ΔHf° =0

"coffee-cup" calorimeter

This type of calorimeter is not tightly sealed, so >pressure is assumed to remain at ATMOSPHERIC PRESSURE

system

a part of the universe that is of specific interest.

closed system

allows the transfer of energy but not mass.

average bond enthalpies

are used for polyatomic molecules.

widely used non-SI unit

calorie (cal) 1 cal = 4.184 J 1000 cal = 1 kcal A nutritional "Calorie" = 1 kcal

open system

can exchange mass and energy with the surroundings

enthalpy (H)

defined by the equation: H = U + PV A note about SI units: PRESSURE: pascal; 1Pa = 1 kg/(m . s2) VOLUME: cubic meters; m3 PV: 1kg/(m . s2) x m3 = 1(kg . m2)/s2 = 1 J ENTHALPY: joules U, P, V, AND H = STATE FUNCTIONS

Enthalpy is an EXTENSIVE property:

dependent on the amount of matter involved. > ΔH as a value refers to the ENTIRE reaction, as written. H2O(l) → H2O(g) ΔH = +44kJ/mol (DOUBLES AMOUNT OF MATTER) 2H2O(l) → 2H2O(g) ΔH = +88 kJ/mol (DOUBLE ENTHALPY)

isolated system

does not exchange either mass or energy with its surroundings.

Pressure-volume (PV) work

done when there is a volume change under constant pressure. w = −PΔV P is the external opposing pressure. ΔV is the change in the volume of the container. DV = Vfinal - Vinitial, so DV is positive when it expands. **Work is done BY the system, so w is NEGATIVE

first law of thermodynamics

energy can be converted from one form to another, but cannot be created or destroyed in any physical or chemical change.

bond enthalpy

enthalpy change associated with breaking a bond in 1 mole of gaseous molecule. H2(g) → H(g) + H(g) ΔH° = 436.4 kJ/mol

standard enthalpy of reaction (ΔH °rxn)

enthalpy of a reaction carried out under standard conditions. ΔH °rxn = ΣnΔH f°(products) - ΣmΔH f°(reactants) *where n and m are the stoichiometric coefficients for the reactants and products.

endothermic process

heat is transferred from the surroundings to the system

exothermic process

heat is transferred from the system to the surroundings

ΔH:

is EQUAL is magnitude but OPPOSITE in SIGN for the REVERSE reaction (must be to obey the Law of Conservation of Energy!)

enthalpy of bond formation

is always a NEGATIVE number *But all the average bond enthalpies are written as POSITIVE numbers. >Calculating as reactants minus products gives the correct sign for DHo.

Any energy gained or lost by the system

is lost or gained from the surroundings. ΔUsys = -ΔUsurr >U= interal energy >"sys" & "surr" denote SYSTEM & SURROUNDINGS >ΔU = Uf - Ui; the difference in the energies of the initial and final states; ΔU is a state function.

State functions 2

magnitude of change depends only on the initial and final states of the system, not on the path the system took to get there. *Whether you take the stairs or use the elevator, Delevation is the same.

Calorimetry

measurement of heat changes >using a CALORIMETER

State functions

properties that are determined by the state of the system, regardless of how that condition was achieved

qrxn = - (Ccal)(DT) = -qcal

qcal = −qrxn Since the bomb is metal, it too can absorb heat, so the SPECIFIC HEAT of the CALORIMETER itself must be known. This is measured in a separate experiment to determine Ccal.

surroundings

rest of the universe outside the system

thermochemical equation

shows the enthalpy change EX. ΔH = +6.01 kJ/mol

Thermodynamics

study of the interconversion of heat and other kinds of energy. *THREE types of systems:

specific heat (Cs) of a substance

the amount of heat required to raise the temperature of 1 g of the SUBSTANCE by 1°C (= 1K). [J/(g • °C) or J/(g • K) ] C=SPECIFIC HEAT heat (q) :) associated with a temperature change M=MASS

heat capacity (C )

the amount of heat required to raise the temperature of an OBJECT (whatever its mass is) by 1°C >The "OBJECT" may be a given quantity of a particular substance [ J/°C or J/K] C=HEAT CAPACITY heat (q) :) associated with a temperature change M=MASS

enthalpy of reaction (ΔHrxn) (or heat of reaction

the difference between the enthalpies of the products and the enthalpies of the reactants: ΔHrxn = H(products) - H(reactants) *Assumes reactions in the lab occur at constant pressure. ΔH > 0 (positive), heat is absorbed by the system: endothermic process ΔH < 0 (negative), heat is released by the system: exothermic process

standard enthalpy of formation (ΔH f°)

the heat change that results when 1 mole of a compound is formed from its constituent elements in their standard states. e.g. N2 (g) + H2 (g) > NH3(g) *superscripted degree sign denotes standard conditions. >Atmospheric pressure >298 K (25oC) >>"f" stands for formation. >>ΔH f° for many substances are tabulated in Appendix 2 of the textbook.

Thermochemistry

the study of heat in chemical reactions.

Heat

transfer of thermal energy *Heat can be either absorbed or released during a process.

ΔH = qp

we can measure the enthalpy of a reaction by measuring heat gain/loss at constant pressure. >USING COFFEE CUP CALORIMETER

For any process, the CHANGE in ENTHALPY is:

ΔH = ΔU + Δ(PV) If pressure is CONSTANT: ΔH = ΔU + PΔV\ [qp = ΔH] = for a constant-pressure process

Using average bond enthalpies, the enthalpy for a gas phase reaction can be estimated using

ΔH° = ΣBE(reactants) - ΣBE(products) >where BE = average bond enthalpy ΔH° = total energy input - total energy released

Overall change in the SYSTEM'S INTERNAL ENERGY:

ΔU = q + w q= heat added or released by the system w = work done on or by the system.

Under conditions of constant PRESSURE

ΔU = q + w ΔU = q − PΔV [qP = ΔU + PΔV] qp= HEAT AT CONSTANT PRESSURE

If a reaction is carried out at constant VOLUME:

ΔV = 0 and NO WORK is done w = −PΔV = 0 SO, ΔU = q + 0 AND, [ qV = ΔU ] qv= HEAT AT CONSTANT VOLUME