Download
1 / 13
130 likes | 408 Views
Standard Enthalpies of Formation. Learning Goal: You will be able to write formation equations, find the enthalpies of formation and use them & Hess ’ Law to find the enthalpy of a given reaction. Standard Enthalpy of Formation.
E N D
Standard Enthalpies of Formation Learning Goal: You will be able to write formation equations, find the enthalpies of formation and use them & Hess’ Law to find the enthalpy of a given reaction.
Standard Enthalpy of Formation • the quantity of energy associated with the formation of one mole of a substance from its elements in their standard states (at SATP) • Symbol: ΔH°f Units: kJ/mol • For example: C(s) + ½ O2 (g) + 2 H2 (g) CH3OH (l) Δ H°f = -239.1 kJ/mol 2 Al (s) + 3/2 O2 (g) Al2O3 (s) Δ H°f = -1675.7 kJ/mol • All of the elements on the left sides of the equations are in their standard states (SATP)
Formation Equations • Are equations that show the production of a compound from its elements. How to Write Formation Equations: • Step 1: Write one mole of the product in the state that has been specified. Example: FeSO4 (s) • Step 2: Write the reactant elements in their standard states (HOFBrINClshould be diatomic – all others are monatomic) (reference periodic table. Pg 838). Fe (s) + S (s) + O2 (g)FeSO4 (s) • Step 3: Write coefficients for the reactants to give a balanced equation yielding one mole of product – you may have to use fractions! Fe (s) + S (s) + 2 O2 (g)FeSO4 (s)
Let’s Practice: • Write the formation equation for solid sodium chloride. • Write the formation equation for solid magnesium carbonate
Calculating ΔH • A third way of calculating ΔH is possible using standard enthalpies of formation. • Table C.6 pg 799 gives standard enthalpies of formation for common molecules. • According to Hess’s Law, the enthalpies of known equations may be used to calculate the enthalpy of an unknown reaction. • For example: The formation of hydrogen and carbon monoxide gas from water and graphite: H2O (g) + C (s) H2 (g) + CO (g) ΔH = ? • Use the formation equations for each of the products and reactants to create the target equation.
Calculating ΔH0, continued H2O (g) + C (s) H2 (g) + CO (g) Reactants H2(g) + 1/2 O2(g) --> H2O(g) H0f H2O = - 241.8 kJ/mol C (s) – graphite in standard state H0f = 0 kJ/mol Products C(s) + 1/2 O2(g) --> CO(g) H0f CO = - 110.5 kJ/mol H2 (g)– hydrogen in standard state H0f = 0 kJ/mol H2O(g) --> H2(g) + 1/2 O2(g) H0 = 241.8 kJ/mol C(s) + 1/2 O2(g) --> CO(g) H0 = - 110.5 kJ/mol H2O (g) + C (s) H2 (g) + CO (g) H0 = 131.3 kJ/mol
Using ΔH°f to calculate enthalpy change: where n = the amount in moles of each product or reactant • Example: Calculate the standard change in enthalpy for the thermite reaction: Al(s) + Fe2O3 (s) Al2O3 (s) + Fe(s)
Sample Problem Calculate the heat of combustion of methanol. CH3OH (l) + 3/2 O2 (g) --> CO2 (g) + 2 H2O (g)H0comb = ?
Multistep Energy Calculations Several energy calculations may be required to solve multistep problems • Heat flow q = mcΔT • Enthalpy changes ΔH = n ΔHx • Hess’s Law ΔHtarget = ∑ΔHknown • Enthalpies of Formation
