chapter 4: aqueous reactions and solution...

Unit 5 Student Notes Packet 1

Introduction to Thermodynamics

5.1 – The Nature of Energy

The study of energy and its transformations is called __________________________.

Thermochemistry looks at relationships between chemical reactions and energy changes involving ________

Units of Energy: The SI unit of energy is the _______________ (J).

Because a joule is not a large amount of energy, we will often use the unit ___________________ (kJ).

Another unit used to describe energy is the ___________________ (cal). This is not the same amount as a food calorie that you see on nutrition labels. (A food Calorie is 1000 calories, or 1 kcal.)

Types of Energy: We are familiar with type types of energy: kinetic and potential. Describe each type below:

o Kinetic energy –

o Potential energy –

Transferring Energy: Work and Heat

There are two common ways we experience energy changes in our everyday lives: Work and Heat

o Work (w) –

Force (F) –


Work is a product of the force, F, and the _________________ d, an object is moved.

o Heat –

Energy: _______________________________________________________________

Important formulas and relationships

1 kJ = 1000 J 1 cal = 4.184 J 1 “Cal” = 1 kcal = 1000 cal*Cal is the food calorie

w = F x d

F = m x g (Force = mass x gravity; g = 9.8 m/s2)

5.2 – First Law of Thermodynamics

First Law of Thermodynamics:

_______________________________________________________________

Any energy lost by the system must be gained by the surroundings or vice versa.

Internal energy – a sum of all the ____________________ and _________________ energy of all the components in the system

The change in internal energy (ΔE) is the difference in energy before and after work is done or heat is transferred:

ΔE =

Relating ΔE to Heat and Work

Because energy is a combination of heat and work, the change in energy is given by the heat added or released by the system, q, and the work done on or by the system, w.

ΔE =


5.3 Enthalpy

Enthalpy (H) is a function of thermodynamics.

Accounts for _________ __________ in chemical changes occurring at:-Constant pressure-No forms of work are performed other than pressure-volume (P-V) work

ΔH represents the _______________ in enthalpy

ΔH =

Endothermic – _____________________________________________________________

Exothermic – ______________________________________________________________


Heat and Work

Identify each energy exchange as heat or work and determine whether the sign of heat or work (relative to the system) is positive or negative:

a) An ice cube melts and cools the surrounding beverage. (The ice cube is the system.)

b) A metal cylinder is rolled up a ramp. (The metal cylinder is the system.)

c) Steam condenses on skin, causing a burn. (The condensing steam is the system.)

Internal Energy Problems (ΔE)

1. Calculate ∆E for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1.4 kJ of work is done on the system.

2. A system absorbs 196 kJ of heat and the surroundings do 117 kJ of work on the system. What is the change in internal energy of the system?

3. A system releases 622 kJ of heat and does 105 kJ of work on the surroundings. What is the change in internal energy of the system?

4. The gas in a piston (defined as the system) warms and absorbs 655 J of heat. The expansion performs 344 J of work on the surroundings. What is the change in internal energy for the system?

5. The air in an inflated balloon (defined as the system) warms over a toaster and absorbs 115 J of heat. As it expands, it does 77 kJ of work. What is the change in internal energy for the system?


PRESSURE-VOLUME WORK

When work is done involving a gas, work is a function of pressure. Work done by the expansion of a gas Pressure is defined as force per unit of area, so when the volume of a gas is changed work was either done

on the gas or by the gas. work = −P∆V

101.3 J = 1 L·atm

Examples:

1. To inflate a balloon you must do pressure-volume work on the surroundings. If you inflate a balloon from a volume of 0.100 L to 1.85 L against an external pressure of 1.00 atm, how much work is done (in Joules)?

2. A cylinder equipped with a piston expands against an external pressure of 1.58 atm. If the initial volume is 0.485 L and the final volume is 1.245 L, how much work (in J) is done?

3. When fuel is burned in a cylinder equipped with a piston, the volume expands from 0.255 L to 1.45 L against an external pressure of 1.02 atm. In addition, 875 J is emitted as heat. What is ΔE for the burning of the fuel?

4. The average human lunch expands by 0.50 L during each breath. If this expansion occurs against an

external pressure of 1.0 atm, how much work (in J) is done during the expansion?


QUANTIFYING HEAT (CALCULATING WITH HEAT CAPACITY)

Temperature is a _______________________ of the thermal energy within a sample of matter.

Heat is the _______________ of thermal energy

Thermal energy always flows from__________________ to _________________ temperatures.

Heat transfer stops when the system and surroundings reach the same temperature, a condition called

___________________ ________________________.

When a system absorbs heat (q), its temperature changes by ΔT:

Heat absorbed by the system is directly proportional to its corresponding temperature change:

q ΔT

Heat capacity (C) – a measuring of the system’s ability to absorb thermal energy without undergoing a large change in temperature

o Constant used in the formulao The higher the heat capacity, the smaller the change in temperature when heat is absorbed

C= q∆T

=¿

Specific heat capacity (Cs)– the amount of heat required to raise the

temperature of _____________ of the substance by ________.o Each substance has its own specific heat capacity (or

“specific heat”)

o Water’s specific heat = 4.184 J/°C

Molar heat capacity – the amount of heat required to raise the temperature of _____________ of a substance by __________.

ΔT

HEAT (q) SYSTEM

SubstanceSpecific Heat

Capacity (J/g·°C)

Elements Lead 0.128 Gold 0.128 Silver 0.235 Copper 0.385 Iron 0.449 Aluminum 0.903Compounds Ethanol 2.42 Water 4.18Materials Glass (Pyrex) 0.75 Granite 0.79 Sand 0.84


TEMPERATURE CHANGES AND HEAT CAPACITY

q=m×C s×∆T

Examples:

1. Suppose you find a penny (minted before 1982, when pennies were almost entirely copper) in the snow. How much heat is absorbed by the penny as it warms from the temperature of the snow, which is -8.0 °C, to the temperature of your body, 37.0 °C? Assume the penny is pure copper and has a mass of 3.10 g.

2. How much heat is required to warm 150 g of sand from 25.0°C to 100.0°C? (The specific heat of sand is 0.84 J/°C.)

3. How much heat is required to warm 1.50 L of water from 25.0°C to 100.0°C? (Assume the density of 1.0 g/mL for water.)

4. Suppose that 25 grams of each substance is initially at 27.0°C. What is the final temperature of each

substance upon absorbing 2.35 kJ of heat?

a. Gold c. Aluminum

b. Silver d. Water


THERMAL ENERGY TRANSFER

Heat transfers from the hotter object to the colder object. If we assume that the two objects are thermally isolated from everything else, then the heat lost by one

substance _____________________________ the heat gained by the other. This is according to the law of conservation of _____________________. If we define one substance as the system, and one as the surroundings, we can say:

Suppose a block of metal initially at 55°C is submerged into water initially at 25°C. Thermal energy transfers as heat from the _____________ to the _____________.

The metal will get __________________ and the water will get

__________________ until the two substances reach the same

temperature (a.k.a. _________________________________________).

The exact temperature change that occurs depends on the ____________________ of the metal and the

water, and on their ____________________________________________.

Since q=mC s∆T , it follows that…

Examples:

1. A 32.5 g cube of aluminum initially at 45.8°C is submerged into 1-5.3 g of water at 15.4°C. What is the final temperature of both substances at thermal equilibrium? (Assume that the aluminum and the water are thermally isolated from everything else.)

qmetal = - qwater


2. A block of copper of unknown mass has an initial temperature of 65.4°C. The copper is immersed in a beaker containing 95.7 g of water at 22.7°C. When the two substances reach thermal equilibrium, the final temperature is 24.2°C. What is the mass of the copper block?

3. A 31.1 g wafer of pure gold, initially at 69.3°C, is submerged into an unknown mass of water at 64.2 g of

water at 27.8°C in an insulated container. What is the final temperature of both substances at thermal equilibrium?

Q: Substances A and B, initially at different temperatures, come in contact with each other and reach thermal equilibrium. The mass of substance A is twice the mass of substance B. The specific heat capacity of substance B is twice the specific heat capacity of substance A. Which statement is true about the final temperature of the two substances once thermal equilibrium is reached?

a) The final temperature will be closer to the initial temperature of substance A than substance B.b) The final temperature will be closer to the initial temperature of substance B than substance A.c) The final temperature will be exactly midway between the initial temperature of substances A and B.


CALORIMETRY: Measuring ∆ H rxn°

Bomb calorimetry – occurs at constant _______________________ and measures __________ for a reaction

Performed in a bomb calorimeter, a piece of equipment designed to measure _________ for _____________________________________________.

Reaction occurs in a sealed container. Burn a sample inside the bomb and measure the temperature. ΔT is rlated to the heat absorbed by the

calorimeter (qcal) by the equation:

If no heat escapes from the calorimeter, then the amount of heat gained by the calorimeter exactly equals the heat released by the ____________________:

Examples:1. When 1.010 g of sucrose (C12H22O11) undergoes combustion in a bomb calorimeter, the temperature rises

from 24.92°C to 28.33°C. Find ΔErxn for the combustion of sucrose in kJ/mol sucrose. The heat capacity of the bomb calorimeter, determined in a separate experiment, is 4.90 kJ/°C.

2. When 1.550 g of liquid hexane (C6H14) undergoes combustion in a bomb calorimeter, the temperature rises from 25.87°C to 38.13°C. Find ΔErxn for the reaction in kJ/mol hexane. The heat capacity of the bomb calorimeter, determined in a separate experiment, is 5.73 kJ/°C.


THERMOCHEMICAL EQUATIONS

Balanced chemical equation that also includes _________, or heat of reaction

Also known as “enthalpy of reaction” Extensive property – means that it depends on the ___________________ of material undergoing the

reaction. The ∆ H rxn

° is for the ________________________________ amounts of reactants and products for the reaction AS WRITTEN.

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g) ∆H rxn° = -2044 kJ

The equation tells us that when ___ mol of C3H8 reacts with ____ mol of O2 to form ____ mol of CO2 and ____ mol of H2O, 2044 kJ of heat is _______________.

We can use the coefficients with the ∆H rxn° as a ratio of quantities:

To determine the amount of released when a certain mass of C3H8 is combusted, use the following plan:

Examples:1. An LP gas tank in a home barbeque contains 13.2 kg of propane, C3H8. Calculate the heat (in kJ) associated

with the complete combustion of all the propane in the tank.

2. Ammonia reacts with oxygen according to the equation:4NH3 (g) + 5O2 (g) 4NO (g) + 6H2O (g) ∆ H rxn

° = -906 kJCalculate the heat (in kJ) associated with the complete reaction of 155 g of NH3.

3. What mass of butane in grams is necessary to produce 1.5 x 103 kJ of heat? What mass of CO2 is produced?


C4H10 (g) + 132 O2 (g) 4 CO2 (g) + 5 H2O (g) ∆H rxn

° = -906 kJ

4. Consider the thermochemical equation for the combustion of acetone (C3H6O), the main ingredient in nail polish remover.

C3H6O (l) + 4 O2 (g) 3 CO2 (g) + 3 H2O (g) ∆H rxn° = -1790 kJ

If a bottle of nail polish remover contains 177 mL of acetone, how much heat is released by its complete combustion? The density of acetone is 0.788 g/mL.

5. Charcoal is primarily carbon. Determine the mass of CO2 produced by burning enough carbon (in the form of charcoal) to produce 5.00 x 102 kJ of heat.

C (s) + O2 (g) CO2 (g) ∆H rxn° = -393.5 kJ

6. What mass of natural gas (CH4) must burn to emit 267 kJ of heat?CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (g) ∆ H rxn

° = -802.3 kJ


CALORIMETRY: Measuring ∆ H rxn°

Coffee-cup calorimetry – occurs at constant ______________________

and measures __________ for a reaction

Used for aqueous reactions, or reactions _______________

_________________________________________________.

Coffee cup calorimeter – equipment that insulates the reaction

from the external environment

The reaction occurs in a

_________________________________ of solution within the

calorimeter.

Thermometer measures heat change of the solution. We assume the heat gained by the solution equals

the heat lost by the reaction (or vice versa).

Since the reaction happens under constant pressure (open to the atmosphere), qrxn = ΔHrxn. To find the ΔHrxn per mole, then divide by the number of moles actually reacted.

Often times, you’ll make the following assumptions: o The total volume of combined solutions is

the sum of the two solutions. Use density to determine mass

o The specific heat of the aqueous solution is the specific heat of water.

Examples:

1. When a student mixes 50 mL of 1.0 M HCl and 50 mL of 1.0 M NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from 21.0°C to 27.5°C. Calculate the enthalpy change for the reaction in kJ/mol HCl. Assume the following: the total volume of the solution is 100.0 mL, the density is 1.0 g/mL, and the specific heat of the solution is that of water.


2. When 50.0 mL of 0.100 M AgNO3 and 50.0 mL of 0.100 M HCl are mixed in a constant-pressure calorimeter, the temperature of the mixture increases from 22.3°C to 23.1°C. The temperature increase is caused by the following reaction:

AgNO3(aq) + HCl(aq) AgCl(s) + HNO3(aq) ΔH = ?

Calculate ΔH for this reaction in kJ/mol AgNO3, assuming that the combined solution has a mass of 100.0 g and a specific heat of 4.18 J/g·°C.

3. When a 6.50-g sample of solid sodium hydroxide dissolves in 100.0 g of water in a coffee cup calorimeter, the temperature rises from 21.6°C to 37.8°C. Calculate ΔH (in kJ/mol NaOH) for the solution process:

NaOH(s) Na+(aq) + OH-(aq) ΔH = ?

Assume the specific heat of the solution is the same as pure water. NOTE: Mass of solution includes solute AND solvent.

4. When a 4.25-g sample of solid ammonium nitrate dissolves in 60.0 g of water in a coffee-cup calorimeter, the temperature drops from 22.0°C to 16.9°C. Calculate the ΔH (in kJ/mol NH4NO3) for the solution process:

NH4NO3 (s) NH4+ (aq) + NO3

-(aq) ΔH = ?Assume that the specific heat of the solution is the same as that of pure water.


POTENTIAL ENERGY DIAGRAMS (ENTHALPY OF A REACTION)

Diagrams relative energies of reactants and products in a reaction. The difference between the levels represents the change in enthalpy (ΔH).

Exothermic

- energy is released

- products have a _____________ energy

- ΔH arrow points _________

Endothermic

- energy is absorbed

- products have a _____________ energy

- ΔH arrow points _________


RELATIONSHIPS INVOLVING ΔHrxn

I. If a chemical equation is multiplied by some factor, then ΔHrxn is also multiplied by the

______________________________________.

A + 2B C ΔH1

2A + 4B 2C ΔH2 = ______________

II. If a chemical equation is reversed, then ΔHrxn _______________________________.

Because ΔH is a “state function”, the value only depends on the initial and final states of the system:

ΔH = Hfinal – Hinitial

When a reaction is reversed, the final state becomes the initial state and vice versa. Thus, the sign of ΔH changes:

A + 2B C ΔH1

C A + 2B ΔH2 = ________

Examples:

1. Use the thermochemical equation to determine the ΔH for the following:

2 H2 (g) + O2 (g) 2H2O (g) ΔH = -483.6 kJ

a. 6H2 + 3O2 6H2O ΔH = ___________________

b. 2H2O 2H2 + O2 ΔH = ___________________

c. H2O H2 + ½ O2 ΔH = ___________________

d. H2 + ½ O2 H2O ΔH = ___________________

2. Use the thermochemical equation to determine the ΔH for the following:

C3H8 (g) + 5 O2 (g) 3 CO2(g) + 4 H2O(g) ΔH = -2043 kJ

a. 2 C3H8 (g) + 10 O2 (g) 6 CO2(g) + 8 H2O(g) ΔH = ___________________

b. 6 CO2(g) + 8 H2O(g) 2 C3H8 (g) + 10 O2 (g) ΔH = ___________________

c. 3 C3H8 (g) + 15 O2 (g) 9 CO2(g) + 12 H2O(g) ΔH = ___________________


III. If a chemical equation can be expressed as the sum of a series of steps, then ΔHrxn for the overall equation is the _________ of the heats of reactions for each step.

Hess’s Law – the change in enthalpy for a stepwise process is the sum of the enthalpy changes of the steps.

Summing the steps in a chemical reaction to determine ΔHrxn:

A + 2B C ΔH1

C 2D ΔH2

A + 2B 2D ΔH3 = _____________

3. Find ΔHrxn for the reaction: C (graphite) C (diamond) ΔH = ?

Use these reactions with known ΔH‘s:

C (graphite) + O2 CO2 (g) ΔH = -393.5 kJCO2 (g) C (diamond) + O2 (g) ΔH = +395.4 kJ

4. Find ΔHrxn for the reaction: CO (g) + 2 H2 (g) CH3OH (g) ΔH = ?

Use these reactions with known ΔH‘s:

2 CO (g) 2 C (s) + O2 (g) ΔH = + 221.0 kJ 2 C (s) + O2 (g) + 4 H2 (g) 2 CH3OH (g) ΔH = - 402.4 kJ


Hess’s Law Calculations:

C (s) + H2O(g) CO2 (g) + H2(g) ΔHrxn = ?

Find ΔHrxn for the reaction between C(s) and H2O(g) using the listed reactions with known ΔHrxn:

A. C (s) + O2 (g) CO2 (g) ΔH = – 393.5 kJB. 2 CO (g) + O2 (g) 2 CO2 (g) ΔH = – 566.0 kJC. 2 H2 (g) + O2 (g) H2O (g) ΔH = – 483.6 kJ

Think: How can we manipulate the reactions with known Δ H in a way to accomplish the following? Get the reactants of interest on the left side Get the products of interest on the right Cancel out all other species (by having them

once on the left and once on the right)

Options:1. Flip the reactions so they go the other direction (and

change the sign of ΔH)2. Multiply the coefficients to change the number of

each species as needed (and multiply ΔH by the same)3. Flip the equation AND multiply the coefficient (and

change the sign AND multiply ΔH by the factor)


Hess’s Law Problems:

1. Calculate ΔH for the reaction 4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g), from the following data.A. N2 (g) + O2 (g) 2 NO (g) ΔH = -180.5 kJB. N2 (g) + 3 H2 (g) 2 NH3 (g) ΔH = -91.8 kJC. 2 H2 (g) + O2 (g) 2 H2O (g) ΔH = -483.6 kJ

2. Calculate ΔH for the reaction CH4 (g) + NH3 (g) HCN (g) + 3 H2 (g), from the reactions.A. N2 (g) + 3 H2 (g) 2 NH3 (g) ΔH = -91.8 kJB. C (s, graphite) + 2 H2 (g) CH4 (g) ΔH = -74.9 kJ/moleC. H2 (g) + 2 C (s, graphite) + N2 (g) 2 HCN (g) ΔH = +270.3 kJ

3. Calculate ΔH for the reaction 2 Al (s) + 3 Cl2 (g) 2 AlCl3 (s) from the following data.A. 2 Al (s) + 6 HCl (aq) 2 AlCl3 (aq) + 3 H2 (g) ΔH = -1049. kJB. HCl (g) HCl (aq) ΔH = -74.8 kJ/moleC. H2 (g) + Cl2 (g) 2 HCl (g) ΔH = -1845. kJD. AlCl3 (s) AlCl3 (aq) ΔH = -323. kJ/mole

4. Calculate ΔH for the reaction C2H4 (g) + H2 (g) C2H6 (g), from the following data.A. C2H4 (g) + 3 O2 (g) 2 CO2 (g) + 2 H2O (l) ΔH = -1411. kJB. C2H6 (g) + 7/2 O2 (g) 2 CO2 (g) + 3 H2O (l) ΔH = -1560. kJC. H2 (g) + 1/2 O2 (g) H2O (l) ΔH = -285.8 kJ


PART 1. “” – Watch the following videos and take notes and/or answer questions as required. This will be pg. 20 of your student notes packet (but of course you use as many pages as you need to accomplish the task).

Video #1 (5 minutes)http://www.youtube.com/watch?v=c8Adft3M8mg

Take notes on the video for Hess’s Law to further your understanding Include what is meant by the term state function. (Need extra help on that?

Video #2 (5 min)http://www.youtube.com/watch?v=EAgbknIDKNo

What are the 2 types of calorimetry? Take notes on both types. Answer 4 questions on Edmodo “Quiz” (posted 10/8 at 3:45 pm)

PART 2. Get ready for the next lab (We’ll perform the lab on Wednesday)Hess’s Law Pre-Labhttp://academic.rcc.edu/freitas/1ALabReports/12HessLawV1.pdf

Write objective and materials/chemicals in your lab notebook. You’ll have to read through the entire document to find the materials. Include chemicals and equipment used (e.g. “balance” is required to take mass measurements.)

Pre-Lab: Include the following IN YOUR LAB NOTEBOOK:o Balanced chemical equations for each part (1 & 2), including states of matter.o Calculations (equations) to be used in Part 1 (see top of page 2)o Calculation (equation) to be used in Part 2 (see page 3)

Procedure written in list/bulleted format (NOT paragraph). Leave space to write down measurements such as mass, time, and temperature.

http://academic.rcc.edu/freitas/1ALabReports/12HessLawV1.pdf

http://www.youtube.com/watch?v=EAgbknIDKNo

http://www.youtube.com/watch?v=c8Adft3M8mg


ENTHALPIES OF FORMATION

Standard State _____________________________

o For a gas:

o For a Liquid or Solid:

o For a Substance in a Solution:

Standard Enthalpy Change_______

o The change in enthalpy for a process when

Standard Enthalpy of Formation ______

(a.k.a. _______________________________________________)

o For a Pure Compound:

o For a Pure Element in Its Standard State:

CALCULATING THE STANDARD ENTHALPY CHANGE FOR A REACTION

∆ H rxn° =∑ np∆ H f

° ( products )−∑ nr∆H f° ( reactants )

Keep in mind that elements in their standard states have _______________.


Examples:

1. Use the standard enthalpies of formation to determine ∆ H rxn° for the reaction:

4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g)

2. The thermite reaction, in which powdered aluminum reacts with iron oxide, is highly exothermic. Use

standard enthalpies of formation to find ∆ H rxn° for the thermite reaction.

2 Al(s) + Fe2O3(s) Al2O3(s) + 2 Fe(s)

Using Standard Enthalpies of Formation to determine stoichiometric amounts

(1) Use ∆ H f° values to calculate ∆H rxn

° for the overall reaction.(2) The thermochemical equation is the balanced chemical equation with∆ H rxn

° .(3) Convert from kJ mol grams or vice versa (as we did on page 11).

Examples:

1. A city of 100,000 people uses approximately 1.0 x 1011 kJ of energy per day. Suppose all of that energy comes from the combustion of liquid octane (C8H18) to form gaseous water and gaseous carbon dioxide. Use standard enthalpies of formation to calculate ∆H rxn

° for the combustion of octane and then determine how many kilograms of octane would be necessary to provide this amount of energy.

2. The oxidation of iron forms iron oxide according to the equation: 4 Fe(s) + 3 O2(g) 2 Fe2O3(s). Calculate ∆ H rxn

° for this reaction and calculate how much heat is produced from a hand warmer containing 15.0 g of iron powder.


12th ed Ch. 10: 72, 73, 75


Unit 5 Review Questions

1. What is thermochemistry? Why is it important?

2. What is energy? What is work? List some examples of each.

3. What is kinetic energy? What is potential energy? List some examples of each.

4. What is the law of conservation of energy? How does it relate to energy exchange between a thermodynamic system and its surroundings?

5. What is the SI Unit of energy? List some other common units of energy.

6. What is the first law of thermodynamics? What are its implications?

7. What is a state function?

8. What is internal energy? Is internal energy a state function?

9. If energy flows out of a chemical system and into the surroundings, what is the sign of ΔEsystem?

10. If the internal energy of the products of a reaction is higher than the internal energy of the reactants, what is the sign of ΔE for the reaction? In which direction does energy flow?

11. What is heat? Explain the difference between heat and energy.

12. How is the change in internal energy of a system related to heat and work (think mathematical relationship)?

13. What is heat capacity? Explain the difference between heat capacity and specific heat capacity.

14. If two objects, A and B, of different temperature come into direct contact, what is the relationship between the heat lost by one object and the heat gained by the other? What is the relationship between the temperature changes of the two objects? (Assume that the two objects do not lose any heat to anything else.)

15. What is pressure-volume work? How is it calculated?What is calorimetry? Explain the difference between a coffee-cup calorimeter and a bomb calorimeter. What is each designed to measure?

16. What is the change in enthalpy (ΔH) for a chemical reaction? How is ΔH different from ΔE?

17. Explain the difference between an exothermic and an endothermic reaction. Give the sign of ΔH for each type of reaction.

18. From a molecular viewpoint, where does the energy emitted in an exothermic reaction come from? Why does the reaction mixture undergo an increase in temperature even though energy is emitted?

19. From a molecular viewpoint, where does the energy emitted in an endothermic reaction come from? Why does the reaction mixture undergo a decrease in temperature even though energy is absorbed?

20. From a molecular viewpoint, where does the energy emitted in an exothermic reaction come from? Why does the reaction mixture undergo an increase in temperature even though energy is emitted?

21. Is the change in enthalpy for a reaction an extensive property? Explain the relationship between ΔH for a reaction and the amounts of reactants and products that undergo reaction.

22. Explain how the value of ΔH for a reaction changes upon each operation:

a. Multiplying the reaction by a factor.

b. Reversing the reaction

23. What is Hess’s Law? Why is it useful?

24. What is a standard state? What is the standard enthalpy change for a reaction?

25. What is the standard enthalpy of formation for a compound? For a pure element in its standard state?

26. How do you calculate ∆H rxn° from known standard enthalpies of formation?

chapter 4: aqueous reactions and solution...

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