Chapter 5 Thermochemisty

​Chapter 5  Thermochemistry

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• Introduction

• 5.1 Energy Basics

• 5.2 Calorimetry

• 5.3 Enthalpy

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LEARNING OUTCOMES

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• Define energy, distinguish types of energy, and describe the nature of energy changes that accompany chemical and physical changes

• Distinguish the related properties of heat, thermal energy, and temperature

• Define and distinguish specific heat and heat capacity, and describe the physical implications of both

• Perform calculations involving heat, specific heat, and temperature change

​Thermochemistry is a branch of thermodynamics.

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The key difference between thermochemistry and thermodynamics is that thermochemistry is the quantitative study of the relation between heat and chemical reactions whereas thermodynamics is the study of laws associated with the relation between heat and chemical reactions.

Multiple Choice

Which will heat up faster?

• copper

• granite

• iron

• basalt

Multiple Choice

Which will heat up slower?

• water

• lead

• granite

• iron

Multiple Choice

Which of the following best explains why the sand at the beach is hotter than the water?

• Sand has a higher specific heat than water.

• Sand has a lower specific heat than water.

• There is more water than sand at the beach.

• There is more sand than water at the beach.

Multiple Choice

If water at $$100\degree C$$   is mixed with the SAME amount of water at $$50\degree C$$  the resulting mixture will have a temperature of

• $$150\degree C$$  

• $$100\degree C$$  

• $$75\degree C$$  

• $$50\degree C$$  

Explanation Slide...

ΔQ = m x c x ΔT

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ΔQ = 0.03 kg x 4200 J/kg°C x 12 °C

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ΔQ = 1512 J

​Calorimetry

​calorimeter: a device used to measure the amount of heat absorbed or released in a chemical or physical process.

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calorimetry: the process of measuring the amount of heat involved in a chemical or physical process.

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​Enthalpy

​enthalpy (H) sum of a system’s internal energy and the mathematical product of its pressure and volume

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enthalpy change (ΔH)heat released or absorbed by a system under constant pressure during a chemical or physical process

​Formation reactions are chemical reactions that form one mole of a substance from its constituent elements in their standard states.

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By standard states, we mean a diatomic molecule, if that is how the element exists and the proper phase at normal temperatures (typically room temperature). The product is one mole of substance, which may require that coefficients on the reactant side be fractional (a change from our normal insistence that all coefficients be whole numbers). For example, the formation reaction for methane (CH4) is:

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C(s) + 2H2(g) → CH4(g)

​Which of the following are proper formation reactions?

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H2(g) + Cl2(g) → 2HCl(g)

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Si(s) + 2F2(g) → SiF4(g)

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CaO(s) + CO2 → CaCO3(s)

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Write formation reactions for each of the following.

1. FeO(s)

2. C2H₆(g)

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​Hess Law

Some text here about the topic of discussion

​Hess’s law states that when chemical equations are combined algebraically, their enthalpies can be combined in exactly the same way.

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Two corollaries immediately present themselves: If a chemical reaction is reversed, the sign on ΔH is changed. If a multiple of a chemical reaction is taken, the same multiple of the ΔH is taken as well.

Multiple Choice

Using the equations below:

C_(s) + O_2(g) → CO_2(g) ∆H = –390 kJ

Mn_(s) + O_2(g) → MnO_2(s) ∆H = –520 kJ

what is ∆H (in kJ) for the following reaction?

MnO_2(s) + C_(s) → Mn_(s) + CO_2(g)

• 910

• 130

• -130

• -910

Multiple Choice

The following equations show the oxidation of carbon and carbon monoxide to carbon dioxide.

C_(s) +O_2(g) → CO_2(g) ΔH = –x kJ mol^–1

CO_(g) + O_2(g) → CO_2(g) ΔH = –y kJ mol^–1

What is the enthalpy change, in kJ mol^–1, for the oxidation of carbon to carbon monoxide?

C_(s) + O_2(g) → CO_(g)

• x + y

• -x - y

• y - x

• x - y

Multiple Choice

Hess' Law makes use of which principle to calculate the enthalpy change of a reaction?

• The law of conservation of energy

• The law of conservation of matter

• The law that you will always find a lost item in the last place you look for it

• Murphy's law

Multiple Select

Which of the following statements are true for the reaction:

SO_2(g) + ¹/₂O_2(g) ↔ SO_3(g) ΔH = –92 kJ mol^-1

Where ↔ indicates that the reaction can proceed in the forward and the reverse direction.

• The forward and reverse reaction both produce 92 kJ of energy.

• Oxidising 2 moles of SO₂ would produce twice as much energy.

• The reverse reaction has an enthalpy of +92 kJ mol^-1.

• Collecting the SO3 produced in the liquid state would not change the measured enthalpy.

Multiple Choice

In order to find the enthalpy of combustion of C₃H₈ how must the enthalpy changes be arranged?

• ΔH₃ = ΔH₁ + ΔH₂

• ΔH₂ = ΔH₃ - ΔH₁

• ΔH₁ = ΔH₂ - ΔH₃

• 0 = ΔH₁ + ΔH₂ + ΔH₃

Multiple Choice

The enthalpies of combustion of C(s), H2(g) and C4H9OH(l) (in kJmol^-1) are as follows  
C(s) + O₂(g)   ->  CO₂(g)                                               ∆H=a
H₂(g) + ½O₂(g)   ->   H₂O(l)                                        ∆H=b
C₄H₉OH(l) + 6O₂(g)   ->   4CO₂(g) + 5H₂O(l)  ∆H=c
What is the enthalpy change for the reaction shown below?
  4C(g) + 5H₂(l) + ½O₂(g)   ->   C₄H₉OH(l)

• c – 4a – 5b
• 2a + 10b - c
• 4a + 5b - c
• 2a + 5b + c