CHEMICAL REACTIONS AND ENERGY - 2

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CHEMICAL REACTIONS AND ENERGY - 2

ENTHALPY (H) Accounts for heat flow in chemical reactions happening at constant pressure when work is performed due to P-V change. The enthalpy of a substance is its internal energy plus a term that takes into account the pressure & volume of the substance. H = U + PV

ΔH = ΔU – W qp= ΔU - W In a process carried out at constant pressure:
We cannot calculate the actual value of H. Instead, we can calculate the change in enthalpy. ΔH = ΔU + PΔV (since W = - PΔV) (Equation 1) ΔH = ΔU – W In a process carried out at constant pressure: ΔU = qp + W or (Equation 2) (Equation 1) = (Equation 2) ΔH = qp (at constant pressure) qp= ΔU - W

ONLY at constant Pressure:
ΔH = Qp For this reason, the term “ heat of reaction” and “change in enthalpy” are used interchangeably for rxns studied at constant pressure.

In summary ΔU measures heat lost/gained. ΔH measures heat lost/gained.
At constant volume At constant pressure ΔU measures heat lost/gained. ΔH measures heat lost/gained. The difference between ΔU & ΔH is very small. ΔH is generally satisfactory to use.

Exercise Under standard conditions, 1 mol CO is burnt in a sealed flask w/ constant volume and kJ energy is released. The same amount of CO is burnt under the same conditions in a flask w/ frictionless movable piston, 283kJ of energy is released. What are the values of ΔH, ΔU, and w for both conditions?

Solution: 1st condition--constant volume, w=0 ΔU=qv = -281.75 kJ
ΔH=ΔU-w (w=0)- ΔH= ΔU= kJ, w=0 2nd condition--constant pressure, ΔH= qp= -283kJ ΔU= kJ, ΔH=ΔU-w, w=ΔU-ΔH, w= kJ= kJ

ΔH = Hproducts –Hreactants
For a chemical rxn: ΔH = Hproducts –Hreactants

When Hproducts > Hreactants
ΔH is (+). Heat will be absorbed by the system. Reactants are more stable than the products. reactants have stronger bonds than the products.

When Hproducts < Hreactants
ΔH is (-). Heat will be released by the system. Reactants are less stable than the products. Products have stronger bonds than the reactants.

ENTHALPIES OF RXNS Enthalpy is an extensive property. ΔH is directly proportional to the amounts of reactants consumed in chemical rxns. CH4(g) + 2O2(g) ---> CO2(g) + 2H2O(l) ΔH = -890 kJ 2CH4(g) + 4O2(g) --->2 CO2(g) + 4H2O(l) ΔH = kJ (890x2)

Exercise When 1 mole of methane is burned at constant pressure, 890 kJ of energy is released as heat. Calculate ΔH for a process in which a 5.8 g sample of methane is burned at constant pressure.

Solution: CH4: 16 g/mol When 16g of CH4 burns 890 kJ energy is released 5.8 g of CH ? ________________________________________ ? = kJ of energy

CH4(g) + 2O2(g) ---> CO2(g) + 2H2O(l) ΔH = -890 kJ
2) ΔH for a rxn is equal in magnitude but opposite in sign, to ΔH for the reverse rxn. CH4(g) + 2O2(g) ---> CO2(g) + 2H2O(l) ΔH = -890 kJ CO2(g) + 2H2O(l) ---> CH4(g) + 2O2(g) ΔH = kJ

Which one is greater in amount?
3) ΔH depends on the state (gas, liquid, solid, crystalline structure)of the reactants and products. Therefore, we should write the states of the substances in the rxn equation. I. CH4(g) + 2O2(g) ---> CO2(g) + 2H2O(l) ΔH1 = -890 kJ II. CH4(g) + 2O2(g) ---> CO2(g) + 2H2O(g) ΔH2 = ? kJ Which one is greater in amount?

ΔH1 ( -890 kJ)> ΔH2 ( kJ): H2O(l)  H2O(g) ΔH= + 88 kJ

4) ΔH depends on temperature & pressure of the rxn medium
4) ΔH depends on temperature & pressure of the rxn medium. - We will generally assume that the reactants & products are both at the same temperature, 25 ° C, unless otherwise stated.

5) ΔH is a state function. Therefore, it doesn’t depend on the rxn pathway.

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