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Energy. Energy transfer occurs between a system and its surroundings. System = piece of the universe. Energy. Two types: heat and work. Energy. Heat = energy transfer that is the result of contact between two substances of differing temperatures. Energy.

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Energy

Energy

  • Energy transfer occurs between a system and its surroundings.

  • System = piece of the universe.


Energy1

Energy

  • Two types: heat and work.


Energy2

Energy

  • Heat = energy transfer that is the result of contact between two substances of differing temperatures.


Energy3

Energy

  • Energy can also be transferred between a system and its surroundings.


Energy4

Energy

  • The system may do work on the surroundings, or the surroundings may do work on the system.


Energy5

Energy

  • Work = F*d (a force is used to move an object).

  • Ex. Cu strip being hammered.

  • Ex. Gas engine.


Energy6

Energy

  • The state of a system is defined by its composition, temperature, and pressure.


Energy7

Energy

  • State Property: quality that is dependent only on the state of the system and not how that system reached that state.


Energy8

Energy

  • Example: reactions under constant volume and pressure do not do expansive work - their q values are different!


Conversions

Conversions

  • The joule is the SI unit of energy (1J = 1N*m)


Conversions1

Conversions

  • 1 calorie = 4.184 J

  • Dietary calorie = C = 1000 c = 4184 J

  • 1 kJ = 1000 J


Exo vs endothermic

Exo vs Endothermic

  • Exothermic: reaction releases heat into the surroundings.

  • Endothermic: reaction absorbs heat.


Exo vs endothermic1

Exo vs Endothermic

  • Exothermic: products formed are at lower energy than the reactants. (usually spontaneous reaction)


Exo vs endothermic2

Exo vs Endothermic

  • Endothermic: products formed are at higher energy than the reactants (usually non-spontaneous).


Exo vs endothermic3

Exo vs Endothermic

  • For a reaction at constant P and T: endothermic: q = delta H > 0

  • Exothermic q = delta H < 0


Activation energy

Activation Energy

  • A reaction requires a collision among reactants. Activation energy is the minimal energy required to make this happen.


Activation energy1

Activation Energy

  • The net energy transfer is what determines whether the reaction is exothermic or endothermic.


Calorimetry

Calorimetry

  • The study of heat transfer using a calorimeter (a device that isolates a system and measures its temperature change)


Calorimetry1

Calorimetry

  • Specific heat: the heat needed to raise 1 gram of a substance 1 degree Celsius.

  • Different substances change temperature differently.


Calorimetry2

Calorimetry

  • Specific heat of water = 4.184 J/g*C

  • q = mass*delta T*specific heat


Calorimetry3

Calorimetry

  • Ex. Suppose 652 J of heat is added to 15.0 g of water originally at 20.0o C. What is the final t?

  • t = 30.4o C


Calorimetry4

Calorimetry

  • In a coffee-cup calorimeter, delta H reaction = -q water. Heat given off by the reaction is absorbed by the water in the cup.


Calorimetry5

Calorimetry

  • Ex. Suppose heat is absorbed by 412 g of water, increasing its t from 20.12 to 29.86oC. What is delta H?

  • delta H = -16.8 kJ


Calorimetry6

Calorimetry

  • In bomb calorimeters, some heat is absorbed by the metal as well as the water.


Calorimetry7

Calorimetry

  • q reaction = -q calorimeter = -(C calorimeter) x delta t

  • C calorimeter = total heat capacity of the bomb + water.


Calorimetry8

Calorimetry

  • Ex. Suppose combustion of 1.60 g of methane in a bomb calorimeter raises the t by 5.14o C. What is the q reaction (C = 17.2 kJ/oC)?

  • q reaction = -88.4 kJ


Phase changes

Phase changes

  • Heat lost = heat gained

  • Process: melting to boiling

  • Heat of fusion, Heat of vaporization

  • Q = mass * heat of fusion


Phase changes1

Phase changes

  • Hfof water = 333 J/g

  • Hv of water = 2260 J/g

  • Specific heat of ice = 2.06 J/g*C

  • Steam = 2.02 J/g*C


Phase change

Phase Change

How much energy is required to raise the temperature of 2.50g of water from -3.00 C to 108.00 C?

qtot = 7.58 kJ


Thermochem

Thermochem

  • Rule 1: H is directly proportional to the amount of reactants or products.

  • (stoichiometry)


Thermochem1

Thermochem

  • Rule 2: H for a reaction is equal in magnitude but opposite in sign to H for the reverse reaction or process.

  • A(s) --> A(l) kJ

  • A(l) --> A(s) kJ


Thermochem2

Thermochem

  • Rule 3: Reactions occur in steps. Hess’ Law: If equation 1 + equation 2 = equation 3 (the steps result in the overall net equation), then H3 = H1 +H2 .


Thermochem3

Thermochem

Ex: step 1 - C(s) + 1/2 O2(g) --> CO(g); step 2 - CO(g) + 1/2 O2(g) --> CO2(g) H = -283.0 kJ; Hnet = (-393.5 kJ). What is the net equation and the H of the first step?

H1 = -110.5 kJ, C(s) + O2(g) --> CO2(g)


Heats of formation

Heats of Formation

  • Hof of a compound = H when one mole of the compound is formed from its elements in their stable state.


Thermochem4

Thermochem

  • Ex. 2 Ag(s) + Cl2(g) --> 2 AgCl(s) Ho = -254.0 kJ

  • What is Hof for silver chloride?

  • Hof = -127.0 kJ


Thermochem5

Thermochem

  • Ex. HgO(s) --> Hg(l) + 1/2 O2(g) Ho = +90.8 kJ

  • What is Hof for mercury(II) oxide?

  • Hof = -90.8 kJ


Heats of formation1

Heats of Formation

  • For any thermochemical equation:

  • Ho = f products) - f reactants)

  • Note: the heat of formation of an element in its stable state = 0


Thermochem6

Thermochem

  • Ex. (use Table 8.3)

  • Calculate the standard enthalpy change (o) for the combustion of methane.

  • Ho = -890.3 kJ


Heats of formation2

Heats of Formation

  • Can also apply to ions if you take the standard heat of formation of the hydrogen ion to be zero. (see Table 8.3 bottom)


Thermochem7

Thermochem

  • Ex. Calculate the standard enthalpy change (o) for Zn(s) + 2H+(aq) --> Zn2+ (aq) + H2(g).

  • Ho = -153.9 kJ


Bond energies

Bond Energies

  • Defined as H when one mole of bonds is broken in the gaseous state.

  • Ex. Cl2(g) --> 2Cl(g)

  • H = B.E. Cl-Cl = 243 kJ


Bond energies1

Bond Energies

  • In general, since multiple bonds involve more bonding pairs, they tend to be stronger than single bonds.


Bond energies2

Bond Energies

  • Ex. Use bond energies to to determine if the following is an exothermic or endothermic reaction:

  • 2 CO(g) + O2(g) --> 2CO2(g).

  • exothermic


Thermodynamics

Thermodynamics

  • First law: E = q + w where the change in energy of a system, q = heat flow into the system, and w = work done on the system.


Thermodynamics1

Thermodynamics

  • Esystem = -Esurroundings.

  • At constant volume: w=0, qv=E

  • At constant pressure: w= -PV; q = H = E + ngRT


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