Download Unit 14 Thermochemistry

Academic Chemistry
Unit 14
Unit 14
Thermochemistry
Name ___________________________
May 5
6
Unit 13
Acids and
Intro to
Thermochemistry
Videos (p.2-3)
Bases Test
9
10
11
12
HW: p. 4-5
13
Thermochemistry
Interpret graphs
Heat of reaction &
Specific Heat
Heat of reaction
(p.6-8)
(p. 11 and 12)
Heat of formation
(p. 15-16)
Lab
(p. 13)
(HW: p. 16)
HW: p. 9-10
HW: Finish p.12*
16
17
Quiz
Heat of fusion &
Phase changes
(p. 18)
(HW: p. 14)
18
19
Review Due
Calorimetry
Cheeto Lab
(p. 20-23)
20
Unit 14
Thermodynamics
Test
(HW: p. 19)
1
Academic Chemistry
Unit 14
Energy and Chemistry: Crash Course Chemistry #17
https://www.youtube.com/watch?v=GqtUWyDR1fg
Everything is made of chemicals except sound, heat, light, etc. All of those things (including chemicals,
sound, heat, light, etc.) is ________________________.
Energy can be found in several forms including mass, __________________, and
__________________ among many other forms.
____________________________________ is energy contained within a system because of its
position.
The First Law of Thermodynamics is also known as the
________________________________________. It states that energy cannot be created and it
cannot be destroyed.
Thermodynamics is the study of heat, energy, and the ability of energy to do work.
Energy is the capacity to ________________ or ___________________________.
Heat is an energy transfer. The symbol for heat is ___.
We can split the universe into 2 parts: the ______________ and the
______________________________.
Positive ΔE means that work is done on the system or heat is transferred ___ the system.
Chemical energy stored in bonds is a kind of __________________________________________.
A reaction in which heat flows out of the system is considered an ______________________ reaction.
A reaction in which heat flows into the system is considered an ______________________ reaction.
2
Academic Chemistry
Unit 14
Enthalpy: Crash Course Chemistry #18
https://www.youtube.com/watch?annotation_id=annotation_785172&feature=iv&src_vid=GqtUWyDR1fg&v=SV7U4
yAXL5I
Any bond between two atoms contains __________________________.
Heat is the energy transferred to the motion of ___________ and molecules.
The amount of heat or work done depends on the ________________ you take.
The change in energy is the same in either case. The change of energy is independent of the pathway. In
chemistry, we call that a __________________________________. The only things that matter for
state functions are the starting state and the ending state.
We are interested in energy being transferred in or out of system because of
____________________ reactions.
In many cases, we are only interested in the _______ part of energy.
Enthalpy is represented by the letter ___.
ΔH = ___ (Change in enthalpy equals heat gained or lost in the reaction)
When a reaction takes place and ___________ changes, that heat is transferring
__________________________ actual chemical bonds.
We measure enthalpy with c_____________________________.
Germain Henri Hess was a chemist who has a law named after him: __________________________.
Hess’s Law says that the total enthalpy _______________ for a reaction doesn’t depend on the
pathway it takes but only depends on its _________________ and _________________ states. So as
long as you start with the same reactants and end with the same products, the enthalpy change is
______________________.
Standard state is just a set of criteria so that chemists can study stuff under the same
___________________________.
Standard state is at T = 278K and P = 1 atm
Standard enthalpy of formation (ΔHf°) is the _____________________________________ lost or
gained when _______ mole of a compound is formed from its constituent elements.
The enthalpy change for a reaction is ________________ to the sum of the enthalpy of formation of
all the products ___________ the sum of the enthalpy of formation of all the reactants.
Σ means “the _____ of”
3
Academic Chemistry
Unit 14
4
Academic Chemistry
Unit 14
5
Academic Chemistry
Unit 14
Thermochemistry
Thermochemistry focuses on the _____________ changes that occur during a __________________ reaction.
• Heat (___) – _________________ that transfers from one object to another because of a ______________
_______________________________ between them. The standard international unit (SI unit) of heat is
the ________________ (______). **Note: Heat always flows from a warmer object to a cooler object.
• Enthalpy (___) – the ______________________________ of a system at a constant __________________.
• Energy – the ___________________________ for doing ____________ or supplying ________________.
Kinetic energy
Potential energy
o due to _________________________
o _______________________________
•
•
o due to _________________________
o stored in the _____________ between atoms
in ___________________________
Law of Conservation of Energy – Energy is neither _______________________ nor
_____________________; but it can be __________________________ from one form to another.
__________ chemical reactions involve a _______________________ or _____________________ of heat.
Exothermic process – __________________________________ to its surroundings (Temperature _____)
Endothermic process – ________________________________ from its surroundings (Temperature_____)
The following graph shows the energy change during the reaction A + B → C.
__________________ energy: the amount of energy which a ______________ has to have in order for a
chemical change to take place.
1) Does the product have more or less energy than the
reactants? ______________,
2) Would the reaction be endothermic or exothermic?
_____________________. Heat is __________________.
6
Academic Chemistry
Unit 14
A thermochemical equation is a ______________________ for a reaction that also includes ______________.
*Thermochemical equations must be balanced.
•
Heat of Reaction (_____) = the ________________of heat for a reaction under constant ______________.
Also known as change in enthalpy.
Direction of heat flow
Sign of ∆H
Reaction Type
Heat flows out of system
(heat on the ________________ side)
Heat flows into the system
(heat on the ________________side)
Practice – Given the following balanced thermochemical reactions, answer the questions.
1: 4Fe (s) + 3O2 (g) → 2Fe2O3 (s) + 1625 kJ
•
Does this reaction release heat or absorb heat?___________________ How much?______________
•
What does kJ mean? _____________________________(measurement of _____________)
•
Endothermic or exothermic?___________________________________
2: C (s) + 2S (s) + 89.3 kJ → CS2 (l)
•
Is heat released or absorbed in this chemical reaction?____________ How much?_______________
•
Endothermic or exothermic?_______________________________
7
Thermochemical equations treat heat change (_____) just like any ________________ or ________________
Chemistry problems involving ∆H are similar to ________________________________ problems; they depend
on the number of _________________ of reactants and products involved, for example:
CaO(s) + H2O(l) → Ca(OH)2 (s) + 65.2 kJ
and
2 CaO(s) + 2 H2O(l) → 2 Ca(OH)2 (s) + 130.4 kJ
You must multiple the heat of reaction by the number of moles.
Practice
H2 (g) + F2 (g) → 2HF(g)
∆H = -536 kJ
Calculate the heat change (in kJ) for the conversion of 2 moles of H2 gas to HF gas at constant pressure.
2Al (s) + Fe2O3 (s)  Al2O3 (s) +
2Fe (s)
∆H = -851 kJ
Calculate the heat change (in kJ) for the thermite reaction of 4 moles of Fe2O3 into Al2O3 at constant pressure
K2O (s)
+ H2O (l)  2KOH (aq)
∆H = 215 kJ
What is the heat change for the above reaction, at constant temperature if you begin with ½ mole of K2O?
8
Academic Chemistry
Unit 14
Homework:
1. What is the law of conservation of energy? ____________________________________________________
_______________________________________________________________________________________
2. Explain how kinetic and potential energy are involved in a chemical reaction. ________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
3. Explain from where the heat in an exothermic reaction comes. (Hint: look at the answers for the last two
questions. _____________________________________________________________________________
_____________________________________________________________________________________
4. Classify these processes as exothermic or endothermic. (Think about whether the “object” is warming
up/accepting heat or cooling down/releasing heat.)
A. Burning alcohol
( exothermic / endothermic )
B. Baking a potato
( exothermic / endothermic )
C. Combustion of gasoline ( exothermic / endothermic )
5. Indicate whether ΔH is positive (+) or negative (-) for the following:
A. N2 + O2 + 43.25 kcal → 2 NO
ΔH =
___
B. 2 C2H6 + 7 O2 → 4 CO2 + 6 H2O + 683.5 kJ
ΔH =
___
C. the reactants contain more enthalpy than the products ΔH =
___
D. the products contain more enthalpy than the reactants ΔH =
___
E. the surroundings lose heat as a reaction occurs
ΔH =
___
F. the temperature increases as a reaction occurs
ΔH =
___
9
Academic Chemistry
Unit 14
Determine if the following reactions are endothermic or exothermic.
6.
N2 (g)
7.
2 C2H6 (g)
8.
C3H8 (g)
+
O2 (g)
+
→
+
43.3 kJ
→
O2 (g)
→
4 CO2 (g)
C3H8 (l)
+
41.8 kJ
2 NO (g)
+
6 H2O (l)
( exothermic / endothermic )
+
683.5 kJ ( exothermic / endothermic )
( exothermic / endothermic )
Each of the following equations (with the ΔH provided) has been rewritten. Find the ΔH for the new equation.
9. Given: CuO (s)
2 Cu (s)
→
Cu (s)
+
½ O2 (g)
ΔH = 37.1 kJ
+
O2 (g)
→
2 CuO (s)
ΔH =
10. Given: C2H2 (g) + 5/2 O2 (g) → 2 CO2 (g) + H2O (l) + 379 kJ
2 C2H2 (g) + 5 O2 (g) → 4 CO2 (g) + 2 H2O (l)
11. Given:
H2O (g)
→
3 H2O (g)
H2O (l)
ΔH =
______
ΔH =
______
ΔH = - 9.72 kJ
→ 3 H2O (l)
ΔH =
______
12. Rewrite the following equations by expressing the energy change as a term in the equation:
a) H2O (g) → H2O (l)
b) 4 Al
+
c) 2 H2SO4
d) H2O (g)
3 O2
→
→
→
2 SO2
H2O (l)
ΔH = -10.76 kJ
2 Al2O3
+
2 H2O
ΔH = -803.8 kJ
+
O2
ΔH = 130.6 kJ
ΔH = - 9.72 kJ
10
Academic Chemistry
Unit 14
11
Academic Chemistry
Unit 14
12
Academic Chemistry
Unit 14
Standard Heat of Formation (_____)
The heat of the reaction can be calculated from standard heats of formation (see table at the end of the
packet) when it is difficult to measure the heat change for a reaction. The standard heat of formation of a
compound is the change in enthalpy that accompanies the formation of one mole of a compound from its
elements with all substances in their standard states at 25°C. All free elements have a ∆Hf0 of _________
Formula:
Guided Practice:
1) What is the standard heat of reaction for the reaction 2CO(g) + O2(g) → 2CO2(g)
∆Hf0 O2(g) =
∆Hf0 CO(g) =
∆Hf0 CO2 (g) =
a. ∆H =
b. This reaction is endothermic/exothermic.
c. Rewrite the equation with the change in enthalpy represented on the correct side of the
thermochemical equation.
2) Calculate the standard heat of reaction for the reaction:
∆Hf0 CaCO3 =
∆Hf0 CaO =
CaCO3(s) → CaO(s) + CO2(g)
∆Hf0 CO2 =
a. ∆H =
b. This reaction is endothermic/exothermic.
c. Rewrite the equation with the change in enthalpy represented on the correct side of the
thermochemical equation.
13
Academic Chemistry
Unit 14
Standard Heats of Formation Homework: ( Use values from the chart on last page of
packet!!!)
1. Use the Standard heats of formation to calculate the standard heats of the reaction for the reaction.
Br2(g) → Br2(l)
∆Hf0 Br2(g) = 30.91 kJ/mol
∆Hf0 Br2(l) = 0.0 kJ/mol
This reaction is endothermic/exothermic.
2. Use the standard heats of formation to calculate the standard heats of reaction for the reaction:
2NO(g) + O2(g) → 2NO2(g)
∆Hf0 NO(g) =
∆Hf0 O2(g) =
∆Hf0 NO2(g) =
3. Calcium carbonate decomposes at high temperature to form carbon dioxide and calcium oxide:
Using the heat of formation table, determine the heat of reaction.
CaCO3  CO2 + CaO
Is this reaction endothermic or exothermic?
4. Determine the heat of reaction when carbon tetrachloride is formed by reacting chlorine with methane.
CH4 + 2 Cl2  CCl4 + 2 H2
Is this process endothermic or exothermic?
5. When potassium chloride reacts with oxygen under the right conditions, potassium chlorate is formed.
Determine the heat of reaction.
2 KCl + 3 O2  2KClO3
Is this process endothermic or exothermic?
6. The following is known as the thermite reaction: 2Al(s) + Fe2O3(s) → Al2O3(s) + 2Fe(s)
a. Find the heat of reaction (ΔH˚) for the thermite reaction.
b. The thermite reaction is highly exothermic. Does your answer support this piece of information?
14
Academic Chemistry
Unit 14
Refresher Specific Heat
The specific heat capacity (____), or simply the _______________________________ of a substance is the amount of heat it
takes to raise the temperature of _________________________ of the substance ________.
Formula:
Cp =
Q = heat energy in Joules, m = mass, ΔT = change in temperature
ΔT = Tfinal – Tinitial
Therefore the units for specific heat are J/(g°C) or cal/(g°C)
To solve for heat energy, rearrange the equation for q:
Q=
Some substances, such as metals, have ______ specific heats. This means it doesn’t take a lot of energy to cause a
____________________________________ change. Other substances, such as water, have high specific heats. It takes more
energy to cause a temperature change.
On a summer day, why does the concrete desk around a swimming pool become hot, while the water stays much cooler?
Refresher Part 2: Phase Changes
Phase changes are changes in the ___________________________________________ of a substance. Phase changes can be
exothermic or endothermic processes.
http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/Heat_Flow.swf
Name of phase change
States of matter involved?
Exothermic or endothermic?
Melting
Solid to liquid
Endothermic
Freezing
Boiling
Condensing
Sublimation*
Deposition**
*Sublimation: solid to gas phase change without passing through the liquid phase (Examples: dry ice, solid air fresheners,
mothballs, “Shrinking” ice cubes)
**Deposition: gas to solid phase change without passing through the liquid phase (Example: frost on a windshield – water
vapor in the air crystallizes on the cold glass)
Vaporization, evaporation, and boiling: What’s the difference?
______________________________ is the process by which a liquid changes to a gas. Evaporation and boiling are two types of
vaporization.
______________________________ is vaporization only at the __________________ of the liquid, at temperatures below the
boiling point. Rate of evaporation depends on temperature, and also on intermolecular forces.
•
A use of evaporation in our bodies is perspiration. How does perspiration help your body cool?
•
How does a fan or a cool breeze help you cool even more?
15
Academic Chemistry
Unit 14
During boiling, vaporization occurs ______________________________________ the liquid. The bubbles you see are bubbles
of _________________ forming from the liquid (it’s not ________). The pressure inside the bubbles equals atmospheric
pressure. The vapor then escapes into the atmosphere.
The boiling point of a liquid at a pressure of 1 atmosphere (sea level) is called the _____________________________
___________________________________. For water, that is _______________.
The boiling point of a liquid changes as the external pressure changes.
• If the external pressure above the liquid is higher than normal, the liquid boils at a __________________ temp.
• If the external pressure above the liquid is lower than normal, the liquid boils at a __________________ temp.
Why don’t foods cook the same at high altitudes? Example: making spaghetti
Phase changes can be represented on a ____________________________________. The heating curve below is for water.
Show where each state of matter exists, label the phase changes, provide values for the temperatures at each phase change
and label the direction arrows as endo- or exothermic. Then write the formula used to show change in heat for each portion of
the graph. Assume standard pressure (1 atm).
Hf = heat of fusion (heat per mass needed to melt a substance) = 334 J/g (for water)
Hv=heat of vaporization (heat per mass needed to vaporize a substance) = 2260 J/g (for water)
Phase changes always occur at ______________________ temperature. For example, the freezing/melting point of water is
0°C. If the temperature is exactly 0°C, there will be a mixture of liquid water and ice present. Because we have both states of
matter (solid and liquid) present at the freezing/melting point, we sat the solid is “in
________________________________________” with the liquid. If you add heat at this point, you can melt all the ice and
then heat the liquid water further if you want. If you take away heat (cool it), you can freeze the rest of the liquid and then
cool the ice further if you want.
16
Academic Chemistry
Unit 14
Specific Heat: Thermal Energy Calculations (use specific heat table in the back of the packet)
Guided Practice:
1. How much thermal energy (J) is needed to raise the temperature of 50.0g H2O from 14°C to 83°C?
2. How much thermal energy (J) must be added to 50.0 kg Al at -5°C to raise its temperature to 125°C?
3. A 500g block of metal absorbs 5016 Joules of thermal energy when its temperature changes from 20°C to 30°C.
Calculate the specific heat of the metal.
Homework:
1. A copper wire has a mass of 165 grams. An electric current runs through the wire for a short time and its temperature
rises from 21°C to 39°C. What quantity of thermal energy has the copper absorbed?
2. How much thermal energy is absorbed by 250 g H2O when it is heated from 10°C to 85°C?
3. A 38kg block of metal is heated from -26°C to 180°C. It absorbs 1,957,000 J of thermal energy during the heating. What
is the specific heat of this metal?
4. A 200 g glass at room temperature, 20°C, is plunged into a hot dishwasher at 80°C. If the temperature of the glass
reaches that of the dishwasher, how much thermal energy does the glass absorb?
5. Five kilograms of ice cubes are moved from the freezer of a refrigerator into a deep freeze. The refrigerator’s freezing
compartment is kept at -4°C. The deep freeze is kept at -17°C. How much thermal energy does the deep freeze’s
cooling system remove from the ice cubes?
17
Academic Chemistry
Unit 14
ENTHALPY OF PHASE CHANGE WS
Use the following information to
solve the phase change problems:
Specific heat of water = 4.18 J/g·°C Specific
heat of ice = 2.03 J/ g·°C
Specific heat of steam = 1.97 J/g·°C
∆Hfusion = 334 J/g
∆Hvaporization = 2259 J/g
Guided Practice:
1. How much heat is required to melt 233 grams of ice into water, from -15°C to room temperature (25°C)?
Ice:
Melting:
Water:
Total:
2. How much heat is required to change 32.5 grams water into steam, from room temperature (25°C) to
115°C?
Water:
Vaporization:
Steam:
Total:
18
Academic Chemistry
Unit 14
Homework:
3. How much heat is needed to melt 1.43 grams of ice into water from -5.34°C to 84.3°C?
Ice:
Melting:
Water:
Total:
4. How much heat is needed to convert 0.232 grams water into steam, from 32.5°C to 112°C?
5. How much heat do you need to add to 3.22 grams H2O to raise the temperature from -23°C to 152°C?
6. How much heat is needed to raise the temperature of 199 grams H2O from -10.3°C to 154°C?
19
Academic Chemistry
Unit 14
Thermochemistry Test Review
Energy (in general)
1) There are two types of energy: _____________________________________________ (energy due to motion) &
________________________________________________________(energy due to position). Thermal energy (heat) is
kinetic/potential (circle one). Chemical energy (energy stored in bonds) is kinetic/potential (circle one).
2) Define the Law of Conservation of Mass:___________________________________________________________________
____________________________________________________________________________________________________
3) A toaster is powered with 1500 J of electric energy. When on, it converts 1000 J to thermal energy, 300 J to light energy,
and the remaining portion to sound energy. How much sound energy is produced?
Enthalpy
4) In your body, blood is at a higher temperature than any other body tissue. So when your hands are cold, how does your
body warm them up? __________________________________________________________________ because heat
always travels from ______ to __________ objects.
5) What is the heat change for the above reaction, at constant pressure if you begin with 282.6 g of K2O? Is this
endothermic or exothermic?
K2O (s)
+ H2O (l)  2KOH (aq)
6) What is the heat of reaction for the following: C3H8 (g)
→
∆H = 215 kJ
C3H8 (l)
+
41.8 kJ
ΔH = ___________________
Energy Diagrams
7) The energy diagram below represents the equation: 2NO(g) + O2(g) → 2NO2(g).
a)
Label the reactants, products, activation energy
b)
Is the reaction endothermic or exothermic?
c)
Explain the energy transfer (Which form of energy existed first, and to which
kind of energy did it change?)
d)
Would ΔH be positive or negative?
20
Academic Chemistry
Unit 14
8) Breaking bonds requires/releases (chose one) energy and would therefore be an exothermic/endothermic (choose one)
process.
9) Use the diagram below to answer the following questions:
a)
What is value of the change in enthalpy? ________________
What letter represents this change? _______
b)
Is this process endothermic or exothermic? _____________
c)
Which letter represents the heat required to break bonds?___
What is this called? _________________________
d)
Which letter represents the activation energy of the reverse
reaction? _______
e)
Which letter represents the enthalpy of the reactants? ____
Which letter represents the enthalpy of the products? ____
Heat of Formation
For questions 10-12: (a) Calculate the standard enthalpy of the reaction for the following reactions using the standard
enthalpies of formation chart in the back of the packet or the other information given in the problem, (b) classify each as either
endothermic or exothermic, and (c) determine which energy diagram best describes the reaction:
Energy Diagram 1:
Energy Diagram 2:
Energy Diagram 3:
10) 4NH3(g) + 5O2(g) → 6H2O(g) + 4NO(g)
a)
b)
21
Academic Chemistry
Unit 14
c) Diagram #____
11) SiO2(g) + 3C (s) + 624.7kJ  SiC (s) + 2CO (g)
a)
b)
c)
Diagram #____
12) Magnesium reacts with hydrochloric acid in a single replacement reaction.
Balanced Equation:
a)
b)
c)
Diagram #____
Specific Heat
13) Define specific heat capacity:
14) The ___________________ that the specific heat capacity, the more resistant the object is to a change in temperature (i.e.
it requires more energy to change the temperature).
15) Determine the initial temperature of an 84.3 gram sample of water after 14,500 J of heat is applied. The final temperature
of the water sample is 100 ºC. The specific heat of water is 4.18 J/gºC.
22
Academic Chemistry
Unit 14
16) A sample of iron is placed in a hot water bath with an initial temperature of 415.0 ºC causing
3.50 x 103 J of heat to be transferred. After several minutes the temperature of the water was recorded at 22.0 ºC.
Calculate the mass of the copper sample.
Metal
Specific Heat
Copper
0.385 J/g x◦C
Aluminum
0.902 J/g x◦C
17) Using the specific heat data from above, will aluminum or copper reach the higher temperature assuming they gain the
same amount of heat?
Phase Changes
18) How much heat is required to change 32.5 grams ice into steam, from -10°C to 115°C?
Specific heat of water = 4.18 J/g·°C
Specific heat of steam = 1.97 J/g·°C
Specific heat of ice = 2.03 J/ g·°C
∆Hfusion = 334 J/g
∆Hvaporization = 2259 J/g
Calorimetry
19) A solution’s temperature increases as the frequency of collisions between reactants increases.
Temperature is a measure of the molecules’ ___________________________ energy.
20) According to the law of conservation of energy, heat lost by the reaction in a calorimeter must
____________ the heat gained/lost (chose one) by the water.
21) Students conduct an experiment where a reaction occurs in a calorimeter. Calculate the heat released
in Joules to the nearest whole number. The specific heat capacity of water is 4.184 J/g⁰C.
Mass
(g)
100.0
Initial
Temperature (ºC)
25.0
Final
Temperature (ºC)
32.0
23
Academic Chemistry
Unit 14
Table of Specific Heats
Substance
H2O (l)
Al (s)
Cu (s)
H2O (s)
Glass (s)
Specific Heat (C) (J/gK)
4.186
0.900
0.386
2.05
0.84
Table of Heats of Formation
24