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Energy and Chemical ReactionsEnergy and Chemical Reactions

Chapter 6Chapter 6

James P. Joule 1818-1889. Discoveredmechanical equivalent of heat, which led to the First Law of Thermodynamics.

Germain Henri Hess 1802-1850.Hess’s Law.

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The Nature of EnergyThe Nature of Energy

Kinetic energy is the energy of motion.

Potential energy is energy which is available, such as an object that can fall a distance, potentially releasing kinetic energy when it hits the ground.

Kinetic and Potential EnergyKinetic and Potential Energy

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The Nature of EnergyThe Nature of Energy

We sometimes use the calorie instead of the joule:1 cal = 4.184 J (exactly)

A nutritional Calorie:1 Cal = 1000 cal = 1 kcal

221 mvEk (in J)

Energy UnitsEnergy Units

SI Unit for energy is the joule, J:

For example, energy of motion (kinetic energy):

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The Nature of EnergyThe Nature of Energy

In chemistry, we’re interested in the internal energies of reactants and products and how it changes (ΔE) in going from: reactants products

We usually can only measure ΔE, not E

All objects and substances have an internal energy, E (this is a form of potential energy, and in this course iscommonly called “chemical energy”).

“Δ” means change in

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Energy cannot be created or destroyed.

First Law of ThermodynamicsFirst Law of ThermodynamicsTotal energy is conservedTotal energy is conserved

In chemistry, we are concerned with heat andchemical energy.

The first law of thermodynamics tells us that thesum of heat and chemical energy must beconstant (at constant pressure).

Heat at constant pressure is called enthalpy.

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Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.Can only measure the change in enthalpy:

H = Hfinal - Hinitial = qP

EnthalpyEnthalpy

We will use enthalpy (H) and energy (E) interchangeably

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Endothermic and Exothermic ProcessesEndothermic and Exothermic Processes

Endothermic: system absorbs heat from surroundings. (ΔH = +)

Exothermic: system transfers heat to the surroundings. (ΔH = -)

An endothermic reaction feels cold.An exothermic reaction feels hot.

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For a reaction

Hrxn = H(products) - H (reactants)

Enthalpy is an extensive property (magnitude H is

directly proportional to amount):

CH4(g) + 2O2(g) CO2(g) + 2H2O(g) H = -802 kJ

2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) H = -1604 kJ

When we reverse a reaction, we change the sign of H:

CO2(g) + 2H2O(g) CH4(g) + 2O2(g) H = +802 kJ

Change in enthalpy depends on state (gas, liquid, solid):

H2O(g) H2O(l) H = -44 kJ

Enthalpies of Reaction Enthalpies of Reaction

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Heat Capacity and Specific HeatHeat Capacity and Specific HeatCalorimetry = measurement of heat flow.Calorimeter = apparatus that measures heat flow.Heat capacity = the amount of energy required to raise the temperature of a substance (by 1o).Molar heat capacity = heat capacity of 1 mol of a substance (energy required to raise temp. of one moleof a substance (by 1o)Specific heat = specific heat capacity = heat capacity of 1 g (energy required to raise temp of one gram by 1o.Heat absorbed by system, q, is given by:

q = (specific heat) (grams of substance) T (Be careful of the sign of q)

Calorimetry Calorimetry

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Calorimetry Calorimetry Some specific heats (in J/g-K)…

N2(g): 1.04

Al(s): 0.90

Fe(s): 0.45

CH4(g): 2.20

CaCO3(s): 0.82

H2O(l): 4.18

NOTE the very high specificheat for liquid water.

This has a moderating influenceon the weather and the temperature of living organisms.

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Constant-Pressure Constant-Pressure CalorimetryCalorimetryAtmospheric pressure is constant!

H = qP

Heat emitted by reaction (qrxn) is absorbed by solution (qsoln)

So, qrxn = -qsoln

Calorimetry Calorimetry

qrxn = -qsoln = -(specific heat of soln) (grams of solution) T.

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Calorimetry Calorimetry

9.55g sample of solid NaOH dissolves in 100.0 g water in coffee-cup calorimeter. Temp rises from 23.6 to 47.4oC. Calculate ΔH (kJ/mol) for the soln process: NaOH(s) Na+ (aq) + OH- (aq)Assume specific heat of soln is same as that of pure water (which is 4.18 J/g-K).

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qrxn = -Ccal ΔT

Bomb Calorimetry (Constant-Volume Calorimetry)Bomb Calorimetry (Constant-Volume Calorimetry)Calorimetry Calorimetry

Used to study combustion processes

where Ccal is the heat capacity of the calorimeter (to be determined experimentally)

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• Hess’ lawHess’ law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step.

Hess’ Law Hess’ Law

Note that: H1 = H2 + H3

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• A consequence of Hess’ Law is that reactions can be added just like algebraic equations.

• For example:

Hess’ Law Hess’ Law

CH4(g) + 2O2(g) CO2(g) + 2H2O(g) H = -802 kJ 2H2O(g) 2H2O(l) H = -88 kJ

CH4(g) + 2O2(g) CO2(g) + 2H2O(l)H = -890 kJ

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Enthalpies of FormationEnthalpies of Formation

• If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Ho

f .

This is also called the heat of formation.

Standard conditions (standard state): 1 atm and 25 oC (298 K).

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Enthalpies of FormationEnthalpies of Formation• Standard enthalpy of formation of the most stable form of an element

is zero. For example, it is zero for C(s), H2(g), Pb(s), O2(g), etc

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Hrxn =H1 + H2 +H3

Enthalpies of FormationEnthalpies of FormationUsing Enthalpies of Formation to Calculate Using Enthalpies of Formation to Calculate Enthalpies of ReactionEnthalpies of Reaction.We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formationExample:Combustion of 1 mol of propane gas, C3H8 (g)

Can construct alternate path of three steps shown as 1, 2 and 3

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reactants products rxnf fH m H n H

Enthalpies of FormationEnthalpies of Formation

Using Enthalpies of Formation to Calculate Using Enthalpies of Formation to Calculate Enthalpies of ReactionEnthalpies of ReactionFor a reaction:

(a) 2 SO2(g) + O2(g) 2 SO3 (g) -- what is ΔH?

(c) 4 FeO (s) + O2 (g) 2 Fe2O3 (s) -- what is ΔH?

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Foods and FuelsFoods and FuelsFoodsFoods

1 Nutritional Calorie = kcal

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Foods and FuelsFoods and FuelsFuelsFuels U.S.: 1.0 x 106 kJ of fuel per day.Most from petroleum and natural gas.Remainder from coal, nuclear, and hydroelectric.Fossil fuels are not renewable.