Thermochemistry . DO NOW. The specific heat of ethanol is 2.44 J/ g°C . How many kilojoules of energy are required to heat 50.0 g of ethanol from -20.0°C to 68°C?. Objective . Describe how calorimeters are used to measure heat flow . Construct thermochemical equations .
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H = qq = H = m x C x T
Note: We cannot calculate the actual value of enthalpy, only the change in enthalpy
CaO(s) + H2O(l) Ca(OH)2(s) + 65.2 kJ
H = Hproducts – Hreactants
2CO(g) + O2(g) 2CO2(g) H = -566.8 kJ
2CO(g) + O2(g) 2CO2(g) + 566.8 kJ
Negative sign means energy is released
Change is down
ΔH is <0
= Exothermic (heat is given off)
2CO + O2
Products2CO(g) + O2(g) → 2CO2(g) + 566.8 kJ
566.8kJ given off
2CO2(g) 2CO(g) + O2(g) H = +566.8 kJ
2CO2(g) + 566.8 kJ 2CO(g) + O2(g)
Positive sign means energy is absorbed
Change is up
ΔH is > 0
= Endothermic (heat is absorbed)
2CO(g) + O2(g)
CaCO3→ CaO + CO2
2CO2(g) + 566.8 kJ →2CO(g) + O2(g)
Rewrite chemical equation as a thermochemical equation.
Exothermic or endothermic reaction?
If 3 moles of O2 react with excess CH4 how much heat will be produced?
Convert moles to desired unit
Convert to moles
Start with known value
1 mol CH4
10. 3 g CH4
16.05 g CH4
1 mol CH4
= 514 kJ
Ratio from balanced equation
ΔH = -514 kJ, which means the heat is released for the reaction of 10.3 grams CH4
CaO(s) + H2O(l) Ca(OH)2(s) ΔH = -65.2 kJ
CaO + H2O
ΔH = -65.2 kJ