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Enthalpy. Most chemical and physical changes occur under essentially constant pressure (reactors open to the Earth’s atmosphere) very small amounts of work are performed as system expands or contracts

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Enthalpy

Enthalpy

  • Most chemical and physical changes occur under essentially constant pressure (reactors open to the Earth’s atmosphere)

    • very small amounts of work are performed as system expands or contracts

    • the change in internal energy occurs primarily, or exclusively, as heat that is gained or lost.


Enthalpy1

Enthalpy

  • If a process occurs at constant pressure and the only work done is PV work, the heat flow is described by the enthalpy of the system.

  • Enthalpy (H):

    • a state function defined by the equation:

      H = E + PV

      (Question: Are P and V state functions?)


Enthalpy2

Enthalpy

  • Although the enthalpy of a system cannot be measured, the change in enthalpy (D H) can.

    • heat gained or lost by a system when a process occurs at constant pressure

      D H = Hfinal - Hinitial = qP

      where qP = heat gained/lost at constant pressure


Enthalpies of reaction

Enthalpies of Reaction

  • The change in enthalpy can be found using:

    D H = Hfinal - Hinitial

  • For a chemical reaction:

    • Hfinal = H products

    • Hinitial = H reactants

  • The enthalpy change for a chemical reaction is:

    D H = Hproducts - Hreactants


Enthalpies of reaction1

Enthalpies of Reaction

  • The enthalpy change that accompanies a chemical reaction is called the enthalpy of reaction

    • D Hrxn

    • Also called heat of reaction


Enthalpy3

Enthalpy

  • For a chemical reaction,

    D H = Hproducts - Hreactants

  • D H < 0 (negative)

    • heat is lost by system

      • exothermic

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (l)


Enthalpy4

Enthalpy

  • For a chemical reaction,

    DH = Hproducts - Hreactants

  • D H > 0 (positive)

    • heat gained/absorbed

      by the system

      • endothermic

CO2 (g) + 2 H2O (l) CH4 (g) + 2 O2 (g)


Enthalpies of reaction2

Enthalpies of Reaction

P

  • If D Hrxn = positive

    • endothermic

      • heat must be added

  • If D Hrxn = negative

    • exothermic

      • heat is given off

H

R

time

R

H

P

time


Enthalpies of reaction3

Enthalpies of Reaction

  • DHrxn is associated with a specific chemical reaction.

    • extensive property

      • Depends on the amount of material

  • Thermochemical equationsare balanced chemical equations that show the associated enthalpy change

    • balanced equation

    • enthalpy change (DHrxn)


Enthalpies of reaction4

Enthalpies of Reaction

  • An example of a thermochemical equation:

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

  • The coefficients in the balanced equation show the # moles of reactants and products that produced the associated DH.

    • If the number of moles of reactant used or product formed changes, then the DH will change as well.


Enthalpies of reaction5

Enthalpies of Reaction

  • For the following reaction:

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

    -890. kJ-890. kJ

    1 mol CH42 mol O2

    -890. kJ-890. kJ

    1 mol CO22 mol H2O


Enthalpies of reaction6

Enthalpies of Reaction

  • Guidelines for Using Thermochemical Equations:

    • Enthalpy is an extensive property

      • The magnitude of DH is directly proportional to the amount of reactant consumed or product produced


Enthalpies of reaction7

Enthalpies of Reaction

  • The thermochemical equation for burning 1 mole of CH4 (g):

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

    • When 1 mole of CH4 is burned, 890. kJ of heat are released.

    • When 2 moles of CH4 are burned, 1780. kJ of heat are released.


Enthalpies of reaction8

Enthalpies of Reaction

Example: How much heat is gained or lost when 10.0 g of CH4 (g) are burned at constant pressure?

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

Given:10.0 g CH4 (g)

-890. kJ

1 mol CH4

Find: heat


Enthalpies of reaction9

Enthalpies of Reaction

Heat = 10.0 g CH4 x1 mole CH4 x-890. kJ

16.0 g CH41 mol CH4

= -556 kJ

Is this an endothermic or exothermic process?


Enthalpies of reaction10

Enthalpies of Reaction

Example: How much heat is gained or lost when 10.0 g of water are formed at constant pressure in the following reaction?

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

Given:10.0 g H2O (g)

-890. kJ

2 mol H2O

Find: heat


Enthalpies of reaction11

Enthalpies of Reaction

Heat = 10.0 g H2O x1 mole H2O x-890. kJ

18.0 g H2O2 mol H2O

= -247 kJ

The negative sign indicates that heat was released to the surroundings.


Enthalpies of reaction12

Enthalpies of Reaction

  • Guidelines for Using Thermochemical Equations (cont).

    • The enthalpy change for a reaction is equal in magnitude but opposite in sign to the DH for the reverse reaction.

      2 H2O2 (l) 2 H2O (l) + O2(g) DH = -196 kJ

      2 H2O (l) + O2(g) 2 H2O2 (l) DH = +196 kJ


Enthalpies of reaction13

Enthalpies of Reaction

  • Guidelines for Using Thermochemical Equations (cont):

    • The enthalpy change for a reaction depends on the physical state of the reactants and products.

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

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


Enthalpies of reaction14

Enthalpies of Reaction

  • Why does the DHrxn depend on the physical state of the reactants and products?

    • Energy is either absorbed or released when chemicals change from one physical state to another.

H2O (l) H2O (g) DH = + 44 kJ


Calorimetry

Calorimetry

  • The enthalpy change associated with a chemical reaction or process can be determined experimentally.

    • measure the heat gained or lost during a reaction (or process) at constant P

      • Measure the change in temperature


Calorimetry1

Calorimetry

  • Calorimetry:

    • the experimental measurement of heat gained or lost during a chemical reaction or process

  • Calorimeter

    • an instrument used to measure the heat gained or lost during a chemical reaction or process.


Calorimetry2

Calorimetry

  • Calorimetry is used to experimentally determine:

    • Heat capacity or specific heat

    • DHrxn

      • Enthaply change for a reaction

    • DHfusion

      • Enthalpy change when a substance goes from the liquid to the solid state

    • DHvaporization

      • Enthalpy change when a substance goes from the liquid to the gas state


Calorimetry3

Calorimetry

If you leave your keys and your chemistry book sitting in the sun on a hot summer day, which one is hotter?

Why is there a difference in temperature between the two objects?


Calorimetry4

Calorimetry

  • The temperature increase observed when an object absorbs a certain quantity of energy is determined by its heat capacity (C).

  • Amount of heat required to raise the temperature of an object 1oC (or 1 K)

  • As heat capacity increases, more heat must be added to produce a specific temperature increase.


Calorimetry5

Calorimetry

  • For pure substances, heat capacity is usually given for a specified amount of substance:

    • Molar heat capacity:

      • amount of heat required to raise the temperature of 1 mole of a substance by 1oC

    • Specific heat:

      • amount of heat required to raise the temperature of 1 g of a substance by 1oC


Calorimetry6

Calorimetry

  • Specific Heat = quantity of heat transferred

    mass x temp change

    = q

    mass x DT

  • Molar Heat = quantity of heat transferred

    Capacitymoles x temp change

    = q

    mol x DT


Calorimetry7

Specific heat = q

m x DT

Calorimetry

Example:If 418 J is required to increase the temperature of 50.0 g of water by 2.00 K, what is the specific heat of water?

Given:q = 418 J

m = 50.0 g

DT = 2.00K

Find: specific heat


Calorimetry8

  • Common units for specific heat are:

    • Jcal

    • g.Kg.oC

Calorimetry

Specific heat = 418 J =4.18 J

50.0 g x 2.00 Kg.K


Calorimetry9

Molar heat capacity = q

mol x DT

Calorimetry

Example:What is the molar heat capacity of aluminum if it takes 9.00 J to raise the temperature of 5.00 g of aluminum from 298.0 K to 300.0 K?

Given:q = 9.00 J

m = 5.00 g

DT = 2.0 K

Find: molar heat

capacity


Calorimetry10

  • Common units for molar heat capacity are:

    • Jcal

    • mol.Kmol.oC

Calorimetry

Molar heat capacity =

9.00 J x 27.0 g Al =24 J

5.00 g x 2.0 Kmol Almol.K


Calorimetry11

Calorimetry

Example:If the specific heat of Al (s) is 0.90 J/g.K, how much heat is required to raise the temperature of 10.0 kg of Al from 25.0oCto 30.0oC?

Given:C = 0.90 J/g.K

m = 10.0 kg

DT = 5.0 K

Find: heat

C = q

m x DT


Calorimetry12

Calorimetry

q = C x m x DT

q = 0.90 J x10.0 kg x1000 g x 5.0 K

g.K1 kg

= 45,000 J = 45 kJ

C = q

m x DT

This equation is one that you will use OFTEN in calorimetry.


Calorimetry13

Specific heat = q

m x DT

Calorimetry

Example:If the specific heat of Fe(s) is 0.45 J/g.K, what change in temperature would be observed when 1.0 kJ of heat is added to 45 g of Fe(s)?

Given:C = 0.45 J/g.K

m = 45 g

q = 1.0 kJ

Find:DT


Calorimetry14

Calorimetry

Specific heat = q

m x DT

DT =1.0 kJ x 1000 J x g.K

45 g1 kJ 0.45 J

=49 K

DT = q

m x C


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