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Academic Chemistry Unit 14

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Unit 14 Thermochemistry

Name ___________________________

May 5 6

Unit 13

Acids and Bases Test

Intro to Thermochemistry

Videos (p.2-3) HW: p. 4-5

9 10 11 12 13 Thermochemistry

(p.6-8) HW: p. 9-10

Interpret graphs

(p. 11 and 12) HW: Finish p.12*

Heat of reaction & Heat of formation

(p. 13) (HW: p. 14)

Specific Heat

(p. 15-16) (HW: p. 16)

Heat of reaction

Lab

16 17 18 19 20

Quiz

Heat of fusion & Phase changes

(p. 18) (HW: p. 19)

Calorimetry Cheeto Lab

Review Due (p. 20-23)

Unit 14 Thermodynamics

Test


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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.


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Enthalpy: Crash Course Chemistry #18

https://www.youtube.com/watch?annotation_id=annotation_785172&feature=iv&src_vid=GqtUWyDR1fg&v=SV7U4yAXL5I

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”


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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 __________________.


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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?_______________________________

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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?


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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 = ___


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Determine if the following reactions are endothermic or exothermic.

6. N2 (g) + O2 (g) + 43.3 kJ → 2 NO (g) ( exothermic / endothermic )

7. 2 C2H6 (g) + O2 (g) → 4 CO2 (g) + 6 H2O (l) + 683.5 kJ ( exothermic / endothermic )

8. C3H8 (g) → C3H8 (l) + 41.8 kJ ( 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) → Cu (s) + ½ O2 (g) ΔH = 37.1 kJ

2 Cu (s) + O2 (g) → 2 CuO (s) ΔH =

10. Given: C2H2 (g) + 5/2 O2 (g) → 2 CO2 (g) + H2O (l) + 379 kJ ΔH = ______

2 C2H2 (g) + 5 O2 (g) → 4 CO2 (g) + 2 H2O (l) ΔH = ______

11. Given: H2O (g) → H2O (l) ΔH = - 9.72 kJ

3 H2O (g) → 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) ΔH = -10.76 kJ

b) 4 Al + 3 O2 → 2 Al2O3 ΔH = -803.8 kJ

c) 2 H2SO4 → 2 SO2 + 2 H2O + O2 ΔH = 130.6 kJ

d) H2O (g) → H2O (l) ΔH = - 9.72 kJ


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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: CaCO3(s) → CaO(s) + CO2(g)

∆Hf0 CaCO3 = ∆Hf0 CaO = ∆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.


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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 ∆Hf

0 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) = ∆Hf

0 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?


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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?

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/Heat_Flow.swf


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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.


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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?


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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:


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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?


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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) ∆H = 215 kJ

6) What is the heat of reaction for the following: C3H8 (g) → 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?


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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)


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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.


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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.

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 ice = 2.03 J/ g·°C

Specific heat of steam = 1.97 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)

Initial Temperature (ºC)

Final Temperature (ºC)

100.0 25.0 32.0

Metal Specific Heat

Copper 0.385 J/g x◦C

Aluminum 0.902 J/g x◦C


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Table of Specific Heats Substance Specific Heat (C) (J/gK) H2O (l) 4.186 Al (s) 0.900 Cu (s) 0.386 H2O (s) 2.05 Glass (s) 0.84

Table of Heats of Formation

namestaff.katyisd.org/sites/khschem/publishingimages/pages... · · 2016-05-09academic chemistry...

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