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Internal Energy אנרגיה פנימית

Internal Energy (U) of a system is the total energy contained within the system, partly as kineticenergy and partly as potential energy

Kinetic energy involves three types of molecular motion:

Collectively, these are sometimes called thermal energy

Translationהעתקה

Rotationסיבוב

Vibrationויבראציה

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Potential energy involves intramolecular interactions (i.e., bonds):

and intermolecularinteractions:

Internal Energy אנרגיה פנימית

Internal Energy (U) of a system is the total energy contained within the system, partly as kineticenergy and partly as potential energy

Intramolecular forcesמולקולריים-כוחות תוך

Intermolecular forcesמולקולריים-כוחות בין

Collectively, these are sometimes called chemical energy

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ThermochemistryThermochemistry is the study of energy changes that occur during chemical reactions

Universe Focus is on heat and matter transfer between thesystem מערכת

System

Surroundings

Surr

ound

ings

Surroundings and the surroundings סביבה

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ThermochemistryThere are three types of systems in thermochemistry:

OPENמערכת פתוחה

MatterMatter

EnergyEnergy

CLOSEDמערכת סגורה

Matter

EnergyEnergy

Matter

ISOLATEDמערכת מבודדת

MatterMatter

EnergyEnergy

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Energy Transfer Mechanisms

• Energy can be transferred between the system and its surroundings as:

– Heat: the reaction in the system changes the temperature of the surroundings

– Work: the reaction in the system causes work to be done (i.e., force is moving through a distance)

• Different types of work: expansion/compression, electrical, etc.

• There is no such thing as “negative energy” nor “positive energy”; the sign of work (or heat) signifies the direction of energy flow.

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Heat (q) חוםHeat is energy transfer resulting from thermal differences between the system and surroundings

Heat “flows”spontaneously from higher T → lower T

Heat “flow” stops at thermal equilibrium

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Heat Transfer Illustrated

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Heat Transfer Mechanism Illustrated

Inelastic molecular collisions are responsible for heat transfer:

1. More energetic molecules …

2. … transfer energy to less energetic molecules.

Heat transfer demo

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Work (w) עבודהWork is an energy transfer between a system and its surroundings

expansion

compression

Compression: work done ON the system⇒ System gains energy (+w)

Expansion: work done BY the system⇒ System loses energy (-w)

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For now we will consider only pressure-volume work.w = –PΔV

Pressure-Volume Work

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• The state of a system: its exact condition at a fixed instant.– State is determined by the kinds and

amounts of matter present, the structure of this matter at the molecular level, and the prevailing pressure and temperature.

• A state function is a function whose value depend only the present state of a system, and does not depend on how the state was reached (i.e., does not depend on the history of the system).

State Functions פונקציות מצב

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• Law of Conservation of Energy – in a physical or chemical change, energy can be exchanged between a system and its surroundings, but no energy can be created or destroyed. חוק שי מור ה אנרגיה

⇒ The internal energy change of a system is simply the difference between its final and initial states:

ΔU = Ufinal – Uinitial

⇒ If energy change occurs only as heat (q) and/or work (w), then:

ΔU = q + w

First Law of Thermodynamics

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Think from the point of view of the system:• Energy entering a system carries a positive sign:

– heat absorbed by the system (q > 0), or– work done on the system (w > 0)

• Energy leaving a system carries a negative sign– heat released by the system (q < 0)– work done by the system (w < 0)

First Law: Sign Convention

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• qrxn is the quantity of heat exchanged between a reaction system and its surroundings.

• An exothermic reaction releases heat– In isolated systems, system T ↑.– The system goes from higher to lower energy; qrxn<0.

• An endothermic reaction absorbs heat– In isolated systems, system T ↓.– The system goes from lower to higher energy; qrxn>0.

Heats of Reaction (qrxn) חום תגובה

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Conceptualizing an Exothermic ReactionSurroundings are at 25 °C

1. Hypothetical situation: all heat is instantly released to the

surroundings. Heat = qrxn

2. Typical situation: some heat is released to the surroundings,

some heat is absorbed by the solution.

3. In an isolated system, all heat is absorbed by the solution.

Maximum temperature rise.

25 °C

32.2 °C35.4 °C

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Internal Energy Changeat Constant Pressure

• For a system where the reaction is carried out at constant pressure, ΔU = qP – PΔV or ΔU + PΔV = qP

• Most of the thermal energy is released as heat.

• Some work is done to expand the system against the surroundings (push back the atmosphere).

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• Enthalpy is defined as:H = U + PV

• Enthalpy change is thus:ΔH = ΔU + Δ(PV)

• For a process carried out under a constant pressure:ΔH = ΔU + PΔV⇒ ΔH = qp

• For a process carried out under a constant pressure and in which the volume does not change (⇒ no work is done):ΔH = ΔU

Enthalpy and Enthalpy Change

Most reactions occur at constant pressure, so for most reactions, the heat

evolved equals the enthalpy change.

The evolved H2 pushes back the atmosphere;

work is done at constant pressure.

Mg + 2 HCl → MgCl2 + H2

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• Enthalpy is an extensive property.– It depends on how much of the

substance is present.

• Since U, P, and V are all state functions, enthalpy H must be astate function also.

• Enthalpy changes are unique for each reaction.

Properties of Enthalpy

Enthalpy change depends only on the initial and

final states. In a chemical reaction we call the initial

state the ____ and the final state the ____.

Two logs on a fire give off twice as much heat as does one log.

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• Values of ΔH are measured experimentally.• ΔH < 0 ⇒ exothermic reactions.• ΔH > 0 ⇒ endothermic reactions.

Enthalpy Diagrams

A decrease in enthalpy during the reaction;

ΔH < 0.

An increase in enthalpy during the reaction;

ΔH > 0.

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• ΔH changes sign when a process is reversed. • Therefore, a cyclic process has the value ΔH = 0.

Reversing a Reaction

Same magnitude; different signs.

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Given the equation

(a) H2(g) + I2(s) → 2 HI(g) ΔH = +52.96 kJ

calculate ΔH for the reaction

(b) HI(g) → ½ H2(g) + ½ I2(s).

The complete combustion of liquid octane, C8H18, to produce gaseous carbon dioxide and liquid water at 25 °C and at a constant pressure gives off 47.9 kJ of heat per gram of octane. Write a chemical equation to represent this information.

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• For problem-solving, heat evolved (exothermic reaction) can be thought of as a product. Heat absorbed (endothermic reaction) can be thought of as a reactant.

• We can generate conversion factors involving ΔH.• For example, the reaction:

ΔH in Stoichiometric Calculations

H2(g) + Cl2(g) → 2 HCl(g) ΔH = –184.6 kJ

can be used to write:

–184.6 kJ

1 mol H2 1 mol Cl2

–184.6 kJ

2 mol HCl

–184.6 kJ

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What is the enthalpy change associated with the formation of 5.67 mol HCl(g) in this reaction?

H2(g) + Cl2(g) → 2 HCl(g) ΔH = –184.6 kJ

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Hess’s Law ofConstant Heat Summation

• The heat of a reaction is constant, regardless of the number of steps in the process

ΔHoverall = sum (ΔH’s of individual reactions)

• When it is necessary to reverse a chemical equation, change the sign of ΔH for that reaction

• When multiplying equation coefficients, multiply values of ΔH for that reaction

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Hess’s Law: An Enthalpy Diagram

We can find ΔH(a) by subtracting ΔH(b) from ΔH(c)

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Calculate the enthalpy change for reaction (a) given the data in equations (b), (c), and (d).

(a) 2 C(graphite) + 2 H2(g) → C2H4(g) ΔH = ?(b) C(graphite) + O2(g) → CO2(g) ΔH = –393.5 kJ(c) C2H4(g) + 3 O2 → 2 CO2(g) + 2 H2O(l) ΔH = –1410.9 kJ(d) H2(g) + ½ O2 → H2O(l) ΔH = –285.8 kJ

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Calculating Enthalpy Changes• We would like to calculate the change in enthalpy in a reaction

like that:ΔH = Hproducts – Hreactants

• Problem: we can only measure enthalpy changes (i.e., we can’t obtain absolute values of enthalpy of compounds).

ΔH = Hproducts – Hreactants

= Hproducts – Href. – Hreactants + Href.

= (Hproducts – Href.) – (Hreactants – Href.)

1. ΔH of a certain reaction that forms the products from reference molecules

2. ΔH of a certain reaction that forms the reactants from the same reference molecules

What could be a common reference?

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Standard Enthalpies of Formation• The common reference: the elements from which both reactants

and products are made!

• We define the standard state of a substance as the state of the pure substance at 1 atm pressure and the temperature of interest(usually 25 °C).

• The standard enthalpy change (ΔH°) for a reaction is the enthalpy change in which reactants and products are in their standard states.

• The standard enthalpy of formation (ΔHf°) of a substance is the enthalpy change of forming 1 mol of a substance from its component elements in their standard states. אנתלפי ית היצ י רה

ΔHf° of a pure element = 0

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Standard Enthalpiesof Formation at 25 oC

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Using standard enthalpies of formation to calculate ΔH0 of a reaction

2. ΔHf0(products)

is known

reactants

products

elements ⇒ ΔHf0 = 0

3. ΔH0 can be calculated!

• Example: an exothermic reaction:

enth

alpy

1. ΔHf0(reactants)

is known

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Calculations Based onStandard Enthalpies of Formation

• In other words …1. Add all of the values for ΔHf° of the products

(multiplied by the corresponding stoichiometric coefficients).2. Add all of the values for ΔHf° of the reactants

(multiplied by the corresponding stoichiometric coefficients).3. Subtract #2 from #1

( )[ ] ( )[ ]∑∑ Δ×−Δ×=Δreactants

0r

products

0p

0rxn reactantsproducts ff HcHcH

Summation over all products

Stoichiometric coefficients of products

Summation over all reactants

Stoichiometric coefficients of reactants

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Synthesis gas is a mixture of carbon monoxide and hydrogen that is used to synthesize a variety of organic compounds. One reaction for producing synthesis gas is

3 CH4(g) + 2 H2O(l) + CO2(g) → 4 CO(g) + 8 H2(g) ΔH° = ?Use standard enthalpies of formation to calculate the standard enthalpy change for this reaction.

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The combustion of isopropyl alcohol, common rubbing alcohol, is represented by the equation

2 (CH3)2CHOH(l) + 9 O2(g) → 6 CO2(g) + 8 H2O(l) ΔH° = –4011 kJUse this equation and data from Table 6.2 to establish the standard enthalpy of formation for isopropyl alcohol.

Without performing a calculation, determine which of these two substances should yield the greater quantity of heat per mole upon complete combustion: ethane, C2H6(g), or ethanol, CH3CH2OH(l).

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Important Exhothermic Reactions• Combustion of fossil fuels:

– Coal: C(s) + O2(g) → CO2(g)

– Natural gas: CH4 + 2 O2(g) → CO2(g) + 2 H2O(l)

– Petroleum: C8H18(l) + 25/2 O2(g) → 8 CO2(g) + 9 H2O(l)

• Food: Fuels for the body– Carbohydrates (starches and sugars), fats, and proteins– During digestion, carbohydrates are converted into the simple sugar

glucose (C6H12O6)C6H12O6(s) + 6 O2(g) → 6 CO2(g) + 6 H2O (l)

High-grade coal yields 30 kJ/gram of coal

ΔH0 = –890.3 kJ/mol ≡ –55.5 kJ/g

ΔH0 = –5450 kJ/mol ≡ –52.4 kJ/g

ΔH0 = –2803 kJ/mol ≡ –15.6 kJ/g

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• Thermochemistry concerns energy changes in physical processes or chemical reactions

• Thermochemistry includes the notion of a system and its surroundings; the concepts of kinetic energy, potential energy, and internal energy; and the distinction between two types of energy exchanges: heat (q) and work (w)

• Internal energy (U) is a function of state

• Enthalpy (H) is a function based on internal energy, but modified for use with constant-pressure processes

Summary of Concepts

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Summary of Concepts

• The first law of thermodynamics relates the heat and work exchanged between a system and its surroundings to changes in the internal energy of a system

• The concepts of standard state, a standard enthalpy change, and a standard enthalpy of formation are important in thermochemical calculations

• Some practical applications of thermochemistry deal with the heats of combustion of fossil fuels and the energy content of foods