Download Johnny Xie Period 5 Chapter 6 Thermochemistry 6.1 The Nature of

Johnny Xie
Period 5
Chapter 6 Thermochemistry
6.1 The Nature of Energy: the capacity to do work or produce heat
1.
Law of Conservation of Energy: energy can neither be created nor destroyed; it
converts/transfers from one to another. = Energy of the universe is constant.
2.
Energy classification: Potential Energy (PE)—energy due to position or composition
Kinetic Energy (KE)—E due to motion, depends on mass and velocity.
3.
Temperature is a property that reflects motions of particles in the object. Heat involves the
transfer of energy due to difference in temperature.
4.
Work: force over a distance. There are two ways to transfer energy: through work or heat.
5.
Pathway: the path where energy is transferred. Energy is independent of pathway.
6.
State function/property: property of the system that depends only on its present state, not
pathway. Energy is a state function, but work and heat are NOT state functions.
7.
Chemical energy: In Thermochem, chemists divide the universe into:
A. System: the part being focused on (reactants, products)
B. Surrounding: everything else besides the system in the universe (container, room etc.)
8.
Exothermic: heat out of the system; stored potential energy is released as thermal energy like
heat. In these chemical reactions, products are more stable.
9.
Endothermic: heat into the system In these chemical reactions, products are less stable.
10. Nature tends to favor states that have lower energy rather than those that have higher energy.
11. First Law of Thermodynamics: Energy of the universe is constant.
12. Internal energy: sum of kinetic and potential energies of all the particles in the system. The
internal energy of a system can be changed by a flow of work, heat, or both.
13. Internal Energy: ΔE= q + w
- ΔE change in internal energy, q: heat, w: work
- Thermo quantities always have 2 parts: number (magnitude), unit. (For energy, joules, J)
- Sign reflects the system’s point of view of the energy flow.
- When q is +, system’s energy will increase; (Endothermic)
When q is -, system’s energy will decrease. (Exothermic)
- ΔE<0, energy out of system into surrounding; ΔE>0, energy into system from surrounding
- w +: surrounding does work on system; w- : system does work on surroundings.
14.
For gases, w= P ΔV, P: pressure, ΔV: change in volume = final Volume - initial Volume
Expanding gas, ΔV+: volume increasing, w-, w= - PΔV
Compressing gas, ΔV: -, volume decreasing, w+ (work flows into the system) w= +PΔV
6.2 Enthalpy and Calorimetry (the study of heat)
1.
Enthalpy: H= E+ PV. A state function, E: internal energy. P: pressure. V: volume,
2.
When the only work is P V, ΔH = energy flow as heat = heat of the reaction (ΔH=heat at constant
pressure)
3.
ΔH= H products – H reactants;
ΔH +: endothermic; ΔH - : exothermic
4.
Calorimeter: device used experimentally to measure heat
5.
Definition of Heat capacity= heat absorbed / increase in temperature
6.
Specific heat capacity: energy conserved to raise 1g of a substance by 1 degree Celsius
Molar heat capacity: energy conserved to raise 1 mol of a substance by 1 degree Celsius
7.
Constant-pressure Calorimetry: used to determine changes in enthalpy for reactions in solutions
No pressure involved.
Ex. A coffee-cup calorimeter made of two Styrofoam cups.
8.
Enthalpy of the rxn: ΔH= s m ΔT (s, specific heat, m, mass of solution, T: temperature)
9.
Extensive property: depend on the amount of substance ex. Heat of reaction.
10. Constant –volume Calorimetry: ΔV=0 (PΔV=0), no work is done
ΔE= q+w, w=0. So
ΔE=q.
Ex. Bomb calorimeter.
6.3 Hess’s Law
1.
Hess’s Law—A principle stating that change in enthalpy is the same whether the reaction takes
one step or multiple steps.
2.
Characters of Enthalpy changes (used to Hess’s Law)

If a reaction is reversed, the sign of ΔH is reversed.

Magnitude of ΔH is directly proportional to the quantities of reactants and products in rxn.
If a balanced eq. is multiplied by an integer, the value of ΔH is multiplied by the same integer.
3.
Techniques: work backwards from the required reaction, multiple the rxn by a number or reverse
any rxn as needed, use the final solution as guidance, trial and error.
6.4 Standard Enthalpies of Formation
1.
ΔHf。
: Change in enthalpy that accompanies the formation of one mole of a compound from its
elements with all substances in standard states.
Definitions of Standard States:
Compound:
A)
GAS: 1atm
B) LIQUID/ SOLID: pure liquid or pure solid
C) SOLUTION: concentration of 1M
Element:
D) For am element, exist under 1 atm at 25℃
2.
Calculating ΔHf
。(using the formation table)
A) When the rxn is reversed, also reverse the sign of ΔH
B) When the balanced eq. is multiplied by an integer, multiply the value of ΔH by the same
integer too.
C) The formula: ΔHf。reaction =ΔHf。(products) -ΔHf。(reactants)
D)
ΔHf。for an element is 0. ex. ΔHf。of O2 is 0.
6.5-6.6 Present Sources of Energy; New Energy Sources
- Fossil fuels: burning plants or decayed bodies to fossil fuels and claim energy from them
- Petroleum: thick, dark liquid composed of hydrocarbon; Natural gas: methane, butane etc.
- Coal: remains of plants that were buried and subjected to high pressure and heat for a long time
- The Greenhouse effect caused by CO2 and other harmful gases to the earth.
- New sources include: coal conversion, hydrogen, oil shale, ethanol, methanol, seed oils etc.
Multiple choice questions:
1.
PbS(s) +3/2 O2(g) PbO (s)+ SO2 (g). ΔH= - 414 kJ/mol. What is the enthalpy change for the
burning of 1.00kg of solid lead (II) sulfide?
(A) -1730 kJ
2.
(B) -414 kJ
(D) +414 kJ
For the reaction shown, which is closest to the value of ΔH? [Cr3+(aq) -143 kJ/mol; Ni2+(aq) -54
2Cr3+ (aq) + 3Ni(s)  2Cr(s) + 3Ni2+ (aq)
kJ/mol]
(A) 124 kJ
3.
(C) – 895 kJ
(B) 89 kJ
(C) -89 kJ
(D) -124 kJ
Burning of the solid carbon: C(s) + O2 (g) CO2(g)
ΔH= -239.5 kJ/mol.
What is the relationship between the enthalpy change and the heat of reaction for the burning of
solid carbon at atmospheric pressure?
(A). The enthalpy change is less than the heat of reaction at atmospheric pressure.
(B). The enthalpy change is equal to the heat of reaction at atmospheric pressure.
(C ). The enthalpy change is greater than the heat of reaction at atmospheric pressure.
(D). The enthalpy change is equal to the internal energy change, but not to the heat of reaction.
4.
Referring to the data below, calculate standard enthalpy change,
, for the reaction at 25 ℃
(2000 AP)
(A) + 200kJ
(B) +86 kJ
(C ) -200kJ
(D) -86kJ
5. Given the following thermochemical data:
N2O4 (g)  2NO2 (g)
= 57. 93 kJ
2NO(g) + O2 (g)  2NO2 (g)
= -113. 14 kJ
Determine the heat of reaction
(A) 432.87 kJ
2NO (g) + O2  N2O4 (g)
(B) -55.21 kJ
(C ) -171.07 kJ
(D) -85.54 kJ
6. An insulated cup contains 75.0 g of water at 24.00 ℃. A 26.00 g sample of a metal at 85.25 ℃ is
added. The final temperature of the water and metal is 28.34 ℃ What is the specific heat of the metal?
(A) 0.876 J/ g℃
(B) 0.920 J/g℃
(C ) 1.23 J/g℃
(D) 0. 812 J/g℃
7. (Reminder: Review entropy, ΔS, chapter 16) The evaporation of liquid is expected to have:
(A). a positive ΔH and a negative ΔS
(B). a negative ΔH and a negative ΔS
(C ).a positive ΔH and a positive ΔS
(D). a positive ΔH and a negative ΔS
8. PbO (S) + C (s)  Pb (s) + CO (g).
ΔH= + 106.8 kJ/mol.
This is___reaction, which ___energy.
(A). an endothermic, releases
(B). an exothermic; absorbs
(C ). an exothermic; releases
(D), an endothermic; absorbs
9. Calculate the work involved in expanding a gas from an initial state of 1.00L and 10.0 atm of
pressure to 10.0L and 1.0 atm pressure. The work is____.
(A) -11.0 L atm (B) -10.0 L atm
(C ) -9.00 L atm
(D) -8.00 L atm
10. A calorimeter has a heat capacity of 1265 J/℃. A reaction causes the temperature of the calorimeter
to change from 22.34℃ to 25.12℃. How many joules of heat were released in this process?
(A)-3518J
(B) 3518J
(C ) 3522 J
(D) 3517J