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Chapter 5. Principles of Chemical Reactivity. Basic Principles. Thermodynamics: The science of heat and work Energy: the capacity to do work -chemical, mechanical, thermal, electrical, radiant, sound, nuclear -affects matter by raising its temperature, eventually causing a state change

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Chapter 5

Chapter 5

Principles of Chemical Reactivity


Basic principles
Basic Principles

Thermodynamics:The science of heat and work

Energy: the capacity to do work

-chemical, mechanical, thermal, electrical, radiant, sound, nuclear

-affects matter by raising its temperature, eventually causing a state change

-All physical changes and chemical changes involve energy

Potential Energy:

energy that results from an object’s position

-gravitational, chemical, electrostatic

Kinetic Energy:

energy of motion


Basic principles1
Basic Principles

Law of Energy Conservation:

Energy can neither be created nor destroyed

-a.k.a. The first law of thermodynamics

-The total energy of the universe is constant

Temperature vs. Heat:

  • Temperature is the measure of an object’s heat energy

  • Heat ≠ temperature


The measurement of heat
The Measurement of Heat

Thermal Energy

depends on temperature and the amount (mass or volume) of the object

-More thermal energy a substances has the greater the motion its atoms/molecules have

-Total thermal energy of an object is the sum of the individual energies of all atoms/molecules/ions that make up that object

SI unit: Joule (J)

1 calorie = 4.184 J

English unit = BTU


Converting calories to joules
Converting Calories to Joules

Convert 60.1 cal to joules


Basic Principles

System: object or collection of objects being studied

  • In lab, the system is the chemicals inside the beaker

    Surroundings: everything outside of the system that can exchange energy with the system

  • The surroundings are outside the beaker

    Universe: system plus surroundings

    Exothermic:heat transferred from the system to the surroundings

    Endothermic: heat transferred from the surroundings to the system


Specific heat capacity c

J

g.K

Specific Heat Capacity (C)

amount of heat required to raise the temperature of 1gram of a substance by 1 kelvin

SI Units: Specific heat capacity = J /g.K

Specific heat of water = 4.184


Heat transfer
Heat Transfer

Heat transfer equation

used to calculate amounts of heat (q) in a substance

K

J

J

g.K

g

q1 + q2 + q3 … = 0 or qsystem + qsurroundings = 0


Heat transfer1
Heat Transfer

Calculate the amount of heat to raise the temperature of 200 g of water from 10.0 oC to 55.0 oC


Heat transfer2
Heat Transfer

Calculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C


Heat transfer3
Heat Transfer

A 1.6 g sample of metal that appears to be gold requires 5.8 J to raise the temperature from 23°C to 41°C. Is the metal pure gold?

J

g.K

Specific heat of gold is 0.13

Therefore the metal cannot be pure gold.


Changes of state
Changes of State

occurs when enough energy is put into a substance to over come molecular interactions

Solid-liquid:

molecules in a solid when heated move about vigorously enough to break solid-solid molecular interactions to become a liquid

Liquid-gas:

molecules in a liquid when heated move about more vigorously enough to break liquid-liquid molecular interactions to become a gas

Note: This happens in reverse by removing heat energy


Energy and changes of state
Energy and Changes of State

Heat of fusion:

heat needed to convert a substance from a solid to a liquid (at its melting/freezing point)

333 J/g for water

Heat of vaporization:

heat needed to convert a substance from a liquid to a gas (at its boiling/condensation point)

2256 J/g for water

Example: Calculate the amount of heat involved to convert 100.0 g of ice at -50.0°C to steam at 200.0°C.


The first law of thermodynamics
The First Law of Thermodynamics

This law can be stated as, “The combined amount of energy in the universe is constant”

Also called-The Law of Conservation of Energy:

  • Energy is neither created nor destroyed in chemical reactions and physical changes.


The first law of thermodynamics1
The First Law of Thermodynamics

There are two basic ideas for thermodynamic systems:

  • Chemical systems tend toward a state of minimum potential energy

    Some examples of this include:

    -H2O flows downhill

    -Objects fall when dropped

    Epotential = mg(h)


State function
State Function

The value of a state function is independent of pathway

-An analogy to a state function is the energy required to climb a mountain taking two different paths:

E1 = energy at the bottom of the mountain

E1 = mgh1

E2 = energy at the top of the mountain

E2 = mgh2

E = E2-E1 = mgh2 – mgh1 = mg(h)

Examples of state functions:

Temperature, Pressure, Volume, Energy, Entropy, and enthalpy

Examples of non-state functions:

Number of moles, heat, work


The first law of thermodynamics2
The First Law of Thermodynamics

  • Chemical systems tend toward a state of maximum disorder

    Common examples of this are:

    - A mirror shatters when dropped and does not reform

    - It is easy to scramble an egg and difficult to unscramble it

    - Food dye when dropped into water disperses


The first law of thermodynamics3
The First Law of Thermodynamics

Thermodynamic questions:

  • Will these substances react when they are mixed under specified conditions?

  • If they do react, what energy changes and transfers are associated with their reaction?

  • If a reaction occurs, to what extent does it occur?


Changes in internal energy e or d u
Changes in Internal Energy (E or DU)

all of the energy contained within a substance

  • all forms of energy such as kinetic, potential, gravitational, electromagnetic, etc.

    First Law of Thermodynamics states that the change in internal energy, E, is determined by the heat flow (q) and the work (w)

    E = q + w

    book: DU = q + w


Changes in internal energy e
Changes in Internal Energy (E)

DE = Eproducts – Ereactants

DE = q + w at constant pressure:

w = -P x DV

q > 0 if heat is absorbed by the system

q < 0 if heat is absorbed by the surroundings

w > 0 if surroundings do work on the system

w < 0 if system does work on the surroundings


Changes in internal energy e1
Changes in Internal Energy (E)

If 1200 joules of heat are added to a system, and the system does 800 joules of work on the surroundings, what is the:

  • energy change for the system, Esys?

  • energy change of the surroundings, Esurr?


Changes in internal energy e2
Changes in Internal Energy (E)

E is negative when energy is released by a system

-Energy can be written as a product of the process


Changes in internal energy e3
Changes in Internal Energy (E)

E is positivewhen energy is absorbed by a system undergoing a chemical or physical change

  • Energy can be written as a reactant of the process


Enthalpy changes for chemical reactions

Exothermic reactions:

release energy in the form of heat to the surroundings (DH < 0)

-heat is transferred from a system to the surroundings

Endothermic reactions:

gain energy in the form of heat from the surroundings (DH > 0)

-heat is transferred from the surroundings to the system

For example, the combustion of propane:

Combustion of butane:

Enthalpy Changes for Chemical Reactions


Enthalpy change d h
Enthalpy Change (DH)

Heat content of a substance at constant pressure

- Chemistry is commonly done in open beakers on a desk top at atmospheric pressure

- Therefore enthalpy change (H) is the change in heat content: H = qp at constant pressure

DE = qv at constant volume

If DE and DH < 0:

energy is transferred to the surroundings

If DE and DH > 0:

energy is transferred to the system

- Enthalpy and energy differ by the amount of work

DE = DH + w and w = -PDV


Enthalpy changes for chemical reactions1
Enthalpy Changes for Chemical Reactions

Exothermic reactions generate specific amounts of heat

  • Because the potential energies of the products are lower than the potential energies of the reactants

    Endothermic reactions consume specific amounts of heat

  • Potential energies of

    the reactants are lower

    than the products

    DH for the reverse reaction

    is equal, but has the

    opposite sign to the forward

    reaction


Thermochemical equations

balanced chemical reaction with the H value for the reaction

H < 0 designates an exothermic reaction:

heat is a product, the container feels hot

H > 0 designates an endothermic reaction:

heat is a reactant, the container feels cold

Thermochemical Equations


Hess s law
Hess’s Law

If a reaction is the sum of two or more other reactions, DH for the overall process is the sum of the DH for the component reactions

  • Hess’s Law is true because H is a state function

    If we know the following H values:

?

Target:

1.

2.


Hess s law1
Hess’s Law

We can calculate the H for the reaction by properly adding (or subtracting) the H for reactions 1 and 2

  • Notice that the target reaction has FeO and O2 as reactants and Fe2O3 as a product

  • Arrange reactions 1 and 2 so that they also have FeO and O2 as reactants and Fe2O3 as a product

    • Each reaction can be doubled, tripled, or multiplied by a half, etc.

    • Hvalues are then doubled, tripled, etc.

    • If a reaction is reversed the Hvalueis changed to the opposite sign

4 FeO 4 Fe 2 O2

1.

+1088 kJ

+

2.

1


Hess s law2
Hess’s Law

Given the following equations and H values

calculate H for the reaction below:


Hess s law3
Hess’s Law

-The + sign of the H value tells us that the reaction is endothermic.

-The reverse reaction is exothermic, i.e.


Standard enthalpy of formation
Standard Enthalpy of Formation

Thermochemical standard state conditions

The thermochemical standard T = 298.15 K

The thermochemical standard P = 1.0000 atm

- Be careful not to confuse these values with STP

Thermochemical standard states of matter

For pure substances in their liquid or solid phase the standard state is the pure liquid or solid

- For aqueous solutions the standard state is 1.00 M concentration

For gases the standard state is the gas at 1.00 atm of pressure

- For gaseous mixtures the partial pressure must be 1.00 atm


Standard molar enthalpies of formation h f o
Standard Molar Enthalpies of Formation, Hfo

The enthalpy change for the formation of one mole of a substance formed directly from its constituent elements in their standard states

  • The symbol for standard molar enthalpy of formation is Hfo (KJ/mol)

    The standard molar enthalpy of formation for MgCl2 is:


Standard molar enthalpies of formation h f o1
Standard Molar Enthalpies of Formation (Hfo)

Standard molar enthalpies of formation have been determined for many substances and are tabulated in Appendix L

Standard molar enthalpies of elements in their most stable forms at 298.15 K and 1.000 atm are zero.

Example: The standard molar enthalpy of formation for phosphoric acid is -1279 kJ/mol. Write the equation for the reaction for which Hof = -1279 kJ

Note: P in standard state is P4 (s)

Phosphoric acid in standard state is H3PO4(s)


Standard molar enthalpies of formation
Standard Molar Enthalpies of Formation

Calculate the enthalpy change for the reaction of one mole of H2(g) with one mole of F2(g) to form two moles of HF(g) at 25oC and one atmosphere.


Standard molar enthalpies of formation1
Standard Molar Enthalpies of Formation

Calculate the enthalpy change for the reaction in which 15.0 g of aluminum reacts with oxygen to form Al2O3 at 25oC and one atmosphere.

You do it!


Hess s law4
Hess’s Law

Version II

For chemical reaction at standard conditions:

  • the standard enthalpy change is the sum of the standard molar enthalpies of formation of the products minus the sum for the reactants

    - each enthalpy of formation is multiplied by its coefficient in the balanced chemical equation

Final - Initial


Hess s law5
Hess’s Law



C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)

Calculate the DH°f for the following reaction from

the data in appendix L:


Hess s law6
Hess’s Law

Given the following information, calculate Hfo for H2S(g)

You do it!


Calorimetry
Calorimetry

An experimental technique that measures the heat transfer during a chemical or physical process

Constant pressure calorimetry:

A styrofoam coffee-cup calorimeter is

used to measure the amount of heat

produced (or absorbed) in a reaction

  • This is one method to measure

    qP (called DH) for reactions in solution

    qreaction + qsolution = 0

    Note: Assuming no heat transfer to the surroundings


Calorimetry1

If an exothermic reaction is performed in a calorimeter, the heat evolved by the reaction is determined from the temperature rise of the solution

This requires a two part calculation

When we add 25.00 mL of 0.500 M NaOH at 23.000oC to 25.00 mL of 0.600 M CH3COOH already in the calorimeter at the same temperature, the resulting temperature is observed to be 25.947oC. Determine heat of reaction and then calculate the change in enthalpy (as KJ/mol) for the production of NaCH3COO.

CH3COOH(aq) + NaOH(aq) NaCH3COO(aq) + H2O(l)

Calorimetry


Calorimetry2
Calorimetry heat evolved by the reaction is determined from the temperature rise of the solution

Constant volume calorimetry:

Or Bomb calorimetry measures measure the amount of heat produced (or absorbed) in a chemical reaction

-this method is used for measuring qv (DE)

qreaction + qbomb + qwater = 0


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