In** Heating **and **C****ooling** curves we have the representation of the amount of heat absorbed or released during phase changes.

Concept #1: If a solid substance absorbs enough thermal energy it can undergo the phase changes of fusion and vaporization.

Concept #2: If a gaseous substance evolves enough thermal energy it can undergo the phase changes of condensation and freezing.

Example #1: How much energy (kJ) is required to convert a 76.4 g acetone (molar mass = 58.08 g/mol) as a liquid at -30 ** ^{o}**C to a solid at -115.0

Practice: If 53.2kJ of heat are added to a 15.5g ice cube at - 5.00 ** ^{o}**C, what will be the resulting state and temperature of the substance?

The enthalpy of fusion of methanol (CH 3OH) is 3.16 kJ/mol. How much heat would be absorbed or released upon freezing 25.6 grams of methanol.
1. 0.253 kJ absorbed
2. 2.52 kJ released
3. 3.95 kJ absorbed
4. 2.52 absorbed
5. 3.95 kJ released
6. 0.253 kJ released

Label each letter by its phase diagram.
**extra: what phase change is Solid to Gas and Gas to Solid**

The molar heat capacity of some molecule X(l) is 10 cal/K · mol, its heat of vaporization is 5000 cal/mol and its boiling point is 75◦C. For the conversion of one mol of X(g) at 75◦C to one mol of X(l) at 60◦C,
1. 5150 cal of heat are released by X.
2. 4850 cal of heat are released by X.
3. 150 cal of heat are released by X.
4. 4850 cal of heat are absorbed by X.
5. 5150 cal of heat are absorbed by X.

How much energy is required to vaporize 98.6 g of ethanol (C 2H5OH) at its boiling point, if its ΔH vap is 40.5 kJ/mol?
A) 39.9 kJ
B) 11.5 kJ
C) 52.8 kJ
D) 86.7 kJ
E) 18.9 kJ

How much energy is required to heat 36.0 g H 2O from a liquid at 65°C to a gas at 115°C? The following physical data may be useful.
ΔHvap = 40.7 kJ/mol
Cliq = 4.18 J/g°C
Cgas = 2.01 J/g°C
Csol = 2.09 J/g°C
Tmelting = 0°C
Tboiling = 100°C
A) 52.7 kJ B) 91.7 kJ C) 87.7 kJ D) 63.5 kJ E) 10.9 kJ

The heating curve shown was generated by measuring the heat flow and temperature for a solid as it was heated. The slope of the __________ segment corresponds to the heat capacity of the liquid of the substance.

Refer to the diagram below for the following questions:

Refer to the diagram below for the following questions:

Refer to the diagram below for the following questions:

You have a 100 g sample of water at standard pressure and 70 ◦C. How much energy is required to heat this sample to steam at 120◦C?
A. 21 kJ
B. 239 kJ
C. 243 kJ
D. 17 kJ
E. 247 kJ
F. 185 kJ

You have a sample containing 15 grams of solid ethanol together with 15 g of liquid ethanol. The sample is at the freezing temperature of ethanol, -114◦C. What happens when you initially place this mixture in a deep freezer which is held at -180◦C and energy flows out of the sample in the form of heat?
A. The temperature of the sample decreases while the liquid ethanol freezes.
B. The temperature of the sample remains constant while the liquid ethanol freezes.
C. The temperature of the sample remains constant while the solid ethanol melts.
D. The temperature of the sample increases while the liquid ethanol freezes.
E. The temperature of the sample remains constant and the relative amount of liquid and solid ethanol remains constant.

Describe the (physical/chemical) phenomena exhibited in the following diagram: Where would you place heat of fusion & heat of vaporization in the figure shown?

Calculate the heat (kJ) evolved when 100.0 g of steam at 130.0 o C is converted to ice at – 5.00 o C

Name the phase transition in each of the following situations, and indicate whether it is exothermic or endothermic.Frost appears on a window on a cold winter day. Name the process.

Name the phase transition in each of the following situations, and indicate whether it is exothermic or endothermic.Droplets of water appear on a cold glass of beer. Name the process.

Suppose that 0.89 g of water condenses on a 85.0 g block of iron that is initially at 23 oC.If the heat released during condensation goes only to warming the iron block, what is the final temperature (in oC) of the iron block? (Assume a constant enthalpy of vaporization for water of 44.0 kJ/mol.)

Refer to the figure below to answer each question.
A sample of steam begins on the line segment labeled 5 on the graph.
Is heat absorbed or released in moving from the line segment labeled 5 to the line segment labeled 3?

Name the phase transition in each of the following situations, and indicate whether it is exothermic or endothermic.When ice is heated, it turns to water.
Name the process.

Name the phase transition in each of the following situations, and indicate whether it is exothermic or endothermic.Wet clothes dry on a warm summer day. Name the process.

The human body obtains 950 kJ of energy from a candy bar.If this energy were used to vaporize water at 100.0 oC, how much water (in liters) could be vaporized? (Assume the density of water is 1.00 g/mL.)

How much heat (in kJ) is required to warm 12.0 g of ice, initially at -12.0 oC, to steam at 112.0 oC? The heat capacity of ice is 2.09 J/g oC and that of steam is 2.01 J/g oC.

Ethanol (C2H5OH) melts at -114 oC and boils at 78 oC. The
enthalpy of fusion of ethanol is 5.02 kJ/mol, and its enthalpy of vaporization is 38.56 kJ/mol. The specific heat of solid and liquid ethanol are 0.97 J/gK are 2.3 J/gK respectively.How much heat is required to convert 35.5 g of ethanol at 31 oC to the vapor phase at 78 oC?

Ethanol (C2H5OH) melts at -114 oC and boils at 78 oC. The
enthalpy of fusion of ethanol is 5.02 kJ/mol, and its enthalpy of vaporization is 38.56 kJ/mol. The specific heat of solid and liquid ethanol are 0.97 J/gK are 2.3 J/gK respectively.How much heat is required to convert 35.5 g of ethanol at -154 oC to the vapor phase at 78 oC?

What is the name of the phase change that occurs when ice left at room temperature changes to liquid water?

The heat of combustion of CH4 is 890.4 kJ/mol and the heat capacity of H2O is 75.2 J/molK.Find the volume of methane measured at 298 K and 1.26 atm required to convert 1.01 L of water at 298 K to water vapor at 373 K.

Examine the heating curve for water in section 11.7 in the textbook. If heat is
added to the water at a constant rate, which of the three segments
in which temperature is rising will have the least steep slope?
Why?

Refer to the figure below to answer each question.
In moving from left to right along the line segment labeled 2 on the graph, heat is absorbed, but the temperature remains constant.
Where does the heat go?

Refer to the figure below to answer each question.
How would the graph change if it were for another substance (other than water)?

An ice cube of mass 9.0 g at temperature 0oC is added to a cup of coffee, whose temperature is 95 oC and which contains 130 g of liquid. Assume the specific heat capacity of the coffee is the same as that of water. The heat of fusion of ice (the heat associated with ice melting) is 6.0 kJ/mol.Find the temperature of the coffee after the ice melts.

You just saw that the heat capacity of ice is Cs, ice = 2.09 J/g oC and that the heat of
fusion of ice is 6.02 kJ/mol.When a small ice cube at -10oC is put into a cup of water
at room temperature, which of the following plays a greater role in cooling the liquid
water: the warming of the ice from -10oC to 0oC , or the melting of the ice?

When water boils, what are the bubbles composed of?

Heating curve for water. Changes that occur
when 1.00 mol of H2O is heated from H2O(s) at -25 to
H2O(g) at 125 at a constant pressure of 1 atm. Even though
heat is being added continuously, the system temperature does
not change during the two phase changes (red lines).What process is occurring between points C and D?

Assume that the values given here for 1 mole of methanol are constant over the relevant temperature ranges.
Melting Point
176 K
Boiling Point
338 K
Hfus
2.2 kJ/mol K
Hvap
35.2 kJ/mol K
Cs,solid
105 J/mol K
Cs,liquid
81.3 J/mol K
Cs,gas
48 J/mol K
Given the heating curve for
1 mole of methanol beginning at 170 K
and ending at 350 K
answer the following questions:
What is the value of qsolid for methanol?

Assume that the values given here for 1 mole of methanol are constant over the relevant temperature ranges.
Melting Point
176 K
Boiling Point
338 K
Hfus
2.2 kJ/mol K
Hvap
35.2 kJ/mol K
Cs,solid
105 J/mol K
Cs,liquid
81.3 J/mol K
Cs,gas
48 J/mol K
Given the heating curve for
1 mole of methanol beginning at 170 K
and ending at 350 K
answer the following questions:
What is the value of qliquid for methanol?

Assume that the values given here for 1 mole of methanol are constant over the relevant temperature ranges.
Melting Point
176 K
Boiling Point
338 K
Hfus
2.2 kJ/mol K
Hvap
35.2 kJ/mol K
Cs,solid
105 J/mol K
Cs,liquid
81.3 J/mol K
Cs,gas
48 J/mol K
Given the heating curve for
1 mole of methanol beginning at 170 K
and ending at 350 K
answer the following questions:
What is the value of qgas for methanol?

Liquid butane, C4H10, is stored in cylinders, to be used as a fuel. The normal boiling point of butane is listed as -0.5oC.How much heat must be added to vaporize 235 g of butane if its heat of vaporization is 21.3 kJ/mol?

Based on the heating curve for water, does it take more energy
to melt a mole of water or to boil a mole of water?

Examine the heating curve for water.
Explain why the curve has two segments in which heat is added to the water but the temperature does not rise.

Examine the heating curve for water.
Explain the significance of the slopes of each of the three rising segments. Why are the slopes different?

How much energy is released when 42.5 g of water freezes?

Four ice cubes at exactly 0°C with a total mass of 53.5 g are combined with 140 g of water at 85∘C in an insulated container. (ΔH°fus = 6.02 kJ/mol, water = 4.18J/g•°C) If no heat is lost to the surroundings, what is the final temperature of the mixture?

At 1 atm, how much energy is required to heat 75.0 g of H 2O(s) at –22.0 °C to H2O(g) at 145.0 °C? Helpful constants can be found below.

The specific heat of water is 1.00 cal/g C, the heat of vaporization of water is 540 cal/g, and the heat of fusion of water is 80 cal/g. How much heat would be required to convert 10 grams of ice at 0 degrees C to 10 grams of water at 75 degrees C?

What is the enthalpy change associated with heating 20g of Al from room temperature (23°C) to its melting point at 660°C and then melting it. (Specific heat of Al, Cp = 0.900J/kg °C; ΔH fus = 437 kJ/mole)

Use the information given below to calculate the energy involved in the conversion of 235 g of steam at 373 K to ice at 243 K. Be sure to use the correct sign and answer in kJ.Specific Heat, J/g.°C2.06 - ice4.18 - water2.03 - steamMolar heat of fusion for water, kJ/mol - 6.02Molar heat of vaporization for water, kJ/mol - 40.6

The specific heat of liquid ethanol, C2H5OH(l), is 2.46 J/g•°C and the heat of vaporization is 39.3 kJ/mol. The boiling point of ethanol is 78.3°C. The molecular weight of ethanol 46 g/mol. What amount of enthalpy is required to heat 50.0 g of liquid ethanol from 23.0°C to ethanol vapor at 78.3°C?1) 42.7 kJ2) 49.5 kJ3) 179 kJ4) 1970 kJ5) 6840 kJ

Ethanol (C2H5OH) melts at -114°C. The enthalpy of fusion is 5.02 kJ/mol. The specific heats of solid and liquid ethanol are 0.97 J/g•K and 2.3 J/g•K, respectively. How much heat (kJ) is needed to convert 25.0 g of solid ethanol at -135°C to liquid ethanol at -50°C? A) 207.3 kJB) -12.7 kJC) 6.91 kJD) 4192 kJE) 9.21 kJ

Consider a 1 kg block of ice at standard pressure. If it is initially at −5°C and is heated until it is steam at 109°C, how much total heat was added to the sample of water? Use the following thermodynamic values for your calculation:cice = 2.09 J/g Kcwater = 4.184 J/g Kcsteam = 2.03 J/g K∆Hvap = 2260 J/g∆Hfus = 334 J/g1. 3950 kJ2. 2710 kJ3. 2620 kJ4. 3040 kJ5. 28.7 kJ

What is the minimum mass of ice at 0 °C that must be added to the contents of a can of diet cola (340. mL) to cool the cola to 0 °C from room temperature (20.5 °C)? The heat of fusion for ice is 333 J/g. Assume the cola density is equivalent to water (1 g/mL).

An ice cube tray contains enough water at 22.0 ˚C to make 18 ice cubes that each has a mass of 30.0 g. The tray is placed in a freezer that uses CF2Cl2 as a refrigerant. The heat of vaporization of CF2Cl2 is 158 J/g. What mass of CF 2Cl2 must be vaporized in the refrigeration cycle to convert all the water at 22.0 ˚C to ice at -5.0 ˚C? The heat capacities for H2O (s) and H2O (l) are 2.03 J/g • ˚C and 4.18 J/g • ˚C, respectively, and the enthalpy of fusion for ice is 6.02 kJ/mol.

For many years drinking water has been cooled in hot climates by evaporating it from the surfaces of canvas bags or porous clay pots.How many grams of water can be cooled from 38 oC to 21 oC by the evaporation of 51 g of water? (The heat of vaporization of water in this temperature range is 2.4 kJ/g. The specific heat of water is 4.18 J/g.)

Compounds such as CCl2F2 are known as chlorofluorocarbons, or CFCs. These compounds were once widely used as refrigerants but are now being replaced by compounds that are believed to be less harmful to the environment. The heat of vaporization of CCl2F2 is 289 J/g.What mass of this substance must evaporate to freeze 210 g of water initially at 21 oC? (The heat of fusion of water is 334 J/g; the specific heat of water is 4.18 J/(g K).)

How much energy is released when 42.5 g of water freezes?

A sample of steam with a mass of 0.521 g at a temperature of 100 oC condenses into an insulated container holding 4.35 g of water at 5.0 oC. (For water, Hovap= 40.7 kJ/mol and Cwater=4.18 J/g • ˚C) Assuming that no heat is lost to the surroundings, what is the final temperature of the mixture?

In an experiment, 5.00 L of N 2 is saturated with water vapor at 22°C and then compressed to half its volume at constant T. What mass of water vapor condenses to liquid?

Find ΔH for the freezing of water at -9.50˚C. The specific heat capacity of ice is 2.04 J/g•˚C and its heat of fusion is -332 J/g.

Find ΔE for the freezing of water at -9.50˚C. The specific heat capacity of ice is 2.04 J/g•˚C and its heat of fusion is -332 J/g.

Find q for the freezing of water at -9.50˚C. The specific heat capacity of ice is 2.04 J/g•˚C and its heat of fusion is -332 J/g.

Find w for the freezing of water at -9.50˚C. The specific heat capacity of ice is 2.04 J/g•˚C and its heat of fusion is -332 J/g.

Name the phase change in each of these events: (a) Dew appears on a lawn in the morning.

Name the phase change in each of these events: (a) A diamond film forms on a surface from gaseous carbon atoms in a vacuum.

Name the phase change in each of these events: (c) Molten iron from a blast furnace is cast into ingots (“pigs”).

Liquid propane, a widely used fuel, is produced by compressing gaseous propane. During the process, approximately 15 kJ of energy is released for each mole of gas liquefied. Where does this energy come from?

How much heat is evolved in converting 1.00 mol of steam at 150.0 oC to ice at -45.0 oC? The heat capacity of steam is 2.01 J/(g) and of ice is 2.09 J/(g).

From the data below, calculate the total heat (in J) needed to convert 0.333 mol of gaseous ethanol at 300°C and 1 atm to liquid ethanol at 25.0°C and 1 atm:bp at 1 atm: 78.5°C Δ H°vap: 40.5 kJ/molcgas: 1.43 J/g•°C cliquid: 2.45 J/g•°C

Suppose that 0.46 g of water at 25 oC condenses on the surface of a 51-g block of aluminum that is initially at 25 oC. If the heat released during condensation goes only toward heating the metal, what is the final temperature (in oC) of the metal block? (The specific heat capacity of aluminum is 0.903 J/g oC and the heat of vaporization of water at 25 oC is 44.0 kJ/mol.)

Use the heating–cooling curve below to answer.What is the freezing point of the liquid?

Consider a 75.0-g sample of H 2O(g) at 125 ˚C. What phase or phases are present when 215 kJ of energy is removed from this sample?

A 0.250-g chunk of sodium metal is cautiously dropped into a mixture of 50.0 g water and 50.0 g ice, both at 0 ˚C. The reaction is2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g) ΔH = -368 kJAssuming no heat loss to the surroundings, will the ice melt? Assuming the final mixture has a specific heat capacity of 4.18 J/g • ˚C, calculate the final temperature. The enthalpy of fusion for ice is 6.02 kJ/mol.

Sulfur dioxide is produced in enormous amounts for sulfuric acid production. It melts at −73°C and boils at −10.°C. Its ΔH°fus is 8.619 kJ/mol, and its ΔH°vap is 25.73 kJ/mol. The specific heat capacities of the liquid and gas are 0.995 J/g•K and 0.622 J/g•K, respectively. How much heat is required to convert 2.500 kg of solid SO2 at the melting point to a gas at 60.°C?

Suppose that 1.08 g of rubbing alcohol (C3H8O) evaporates from a 71.0 g aluminum block. If the aluminum block is initially at 25 oC, what is the final temperature of the block after the evaporation of the alcohol? Assume that the heat required for the vaporization of the alcohol comes only from the aluminum block and that the alcohol vaporizes at 25 oC. The heat of vaporization of the alcohol at 25 oC is 45.4 kJ/mol

An 8.6-g ice cube is placed into 260 g of water. Calculate the temperature change in the water upon the complete melting of the ice. Assume that all of the energy required to melt the ice comes from the water. Express your answer in terms of the initial temperature of water, T.

How much ice (in grams) would have to melt to lower the temperature of 353 mL of water from 26 oC to 6 oC? (Assume the density of water is 1.0 g/mL.)

The fluorocarbon compound C2Cl3F3 has a normal boiling point of 47.6 oC. The specific heats of C2Cl3F3(l) and C2Cl3F3(g) are 0.91 J/g K and 0.67 J/g K, respectively. The heat of vaporization for the compound is 27.49 kJ/mol.Calculate the heat required to convert 55.5 g of C2Cl3F3 from a liquid at 12.60 oC to a gas at 87.60 oC.

If 42.0 kJ of heat is added to a 32.0-g sample of liquid methane under 1 atm of pressure at a temperature of -170oC, what is the final state of the methane once the system equilibrates? Assume no heat is lost to the surroundings. The normal boiling point of methane is -161.5 oC. The specific heats of liquid and gaseous methane are 3.48 and 2.22 J / g, respectively.

You may want to reference (Pages 450 - 453) Section 11.4 while completing this problem.If 42.0 kJ of heat is added to a 32.0-g sample of liquid methane under 1 atm of pressure at a temperature of –170˚C, what is the final temperature of the methane once the system equilibrates? Assume no heat is lost to the surroundings. The normal boiling point of methane is –161.5˚C. The specific heats of liquid and gaseous methane are 3.48 and 2.22 J/g•K, respectively. [Section 11.4]

Two heating curves, A and B , are shown. In both cases, point 1 corresponds to the crystalline solid phase.One of these graphs shows data for a liquid crystalline material. Which one?

Two heating curves, A and B , are shown. In both cases, point 1 corresponds to the crystalline solid phase.In graph A, what process does the 2-3 line segment correspond to?

Two heating curves, A and B , are shown. In both cases, point 1 corresponds to the crystalline solid phase.In graph B, what process does the 2-3 line segment correspond to?

Two heating curves, A and B , are shown. In both cases, point 1 corresponds to the crystalline solid phase.In graph A, what process does the 3-4 line segment correspond to?

Two heating curves, A and B , are shown. In both cases,
point 1 corresponds to the crystalline solid phase.In graph
B, what process does the 3-4 line segment correspond to?

A 20.0-g sample of ice at 210.0 ˚C is mixed with 100.0 g water at 80.0 ˚C. Calculate the final temperature of the mixture assuming no heat loss to the surroundings. The heat capacities of H2O(s) and H2O(l) are 2.03 and 4.18 J/g • ˚C, respectively, and the enthalpy of fusion for ice is 6.02 kJ/mol.

In regions with dry climates, evaporative coolers are used to cool air. A typical electric air conditioner is rated at 1.00 x 104 Btu/h (1 Btu, or British thermal unit = amount of energy to raise the temperature of 1 lb water by 1 ˚F). What quantity of water must be evaporated each hour to dissipate as much heat as a typical electric air conditioner?

Four ice cubes at exactly 0 oC with a total mass of 53.0 g are combined with 145 g of water at 85 oC in an insulated container. (Hfus=6.02 kJ/mol, cwater = 4.18J/g • oC)If no heat is lost to the surroundings, what is the final temperature of the mixture?

Given that the heat of fusion of water is +6.02 kJ/mol, that the heat capacity of H2O(l) is 75.2 J/mol • K and that the heat capacity of H2O(s) is 37.7 J/mol • K, calculate the heat of fusion of water at -12oC.

Some ice cubes at 0 ˚C with a total mass of 403 g are placed in a microwave oven and subjected to 750. W (750. J/s) of energy for 5.00 minutes. What is the final temperature of the water? Assume all the energy of the microwave is absorbed by the water, and assume no heat loss by the water.

A substance has a heat of vaporization of ΔHvap and heat of fusion of ΔHfus. Express the heat of sublimation in terms of ΔHvap and ΔHfus.

Refer to the figure below to answer each question.A sample of steam begins on the line segment labeled 5 on the graph. Is heat absorbed or released in moving from the line segment labeled 5 to the line segment labeled 3? What is the sign of q for this change?

Determine the amount of heat (in kJ) required to vaporize 1.55 kg of water at its boiling point. For water, ΔHvap = 40.7 kJ/mol (at 100 oC).

How much heat is required to convert 422 g of liquid H2O at 23.5 °C into steam at 150 °C?

Evaporation of sweat requires energy and thus take excess heat away from the body. Some of the water that you drink may eventually be converted into sweat and evaporate. If you drink a 20-ounce bottle of water that had been in the refrigerator at 3.8 °C, how much heat is needed to convert all of that water into sweat and then to vapor?? (Note: Your body temperature is 36.6 °C. For the purpose of solving this problem, assume that the thermal properties of sweat are the same as for water.)

Titanium tetrachloride, TiCl4, has a melting point of −23.2 °C and has a ΔHfusion = 9.37 kJ/mol.(a) How much energy is required to melt 263.1 g TiCl4?

Use the figure below to answer the following: Does it take more heat to melt 12.0 g of CH4 or 12.0 g of Hg?

Name the phase change in each of these events: (b) Icicles change into liquid water.

Use the figure below to answer the following: Does it take more heat to vaporize 12.0 g of CH4 or 12.0 g of Hg?

Name the phase change in each of these events: (c) Wet clothes dry on a summer day.

Name the phase change in each of these events: (b) Mothballs in a bureau drawer disappear over time.

Many heat-sensitive and oxygen-sensitive solids, such as camphor, are purified by warming under vacuum. The solid vaporizes directly, and the vapor crystallizes on a cool surface. What phase changes are involved in this method?

A 75.0 mL sample of water is heated to its boiling point. How much heat is required to vaporize it? (Assume a density of 1.00 g/mL.)

From the data below, calculate the total heat (in J) needed to convert 22.00 g of ice at −6.00°C to liquid water at 0.500°C: mp at 1 atm: 0.0°C Δ H°fus: 6.02 kJ/molcliquid: 4.21 J/g•°C csolid: 2.09 J/g•°C

Calculate the amount of heat required to completely sublime 56.0 g of solid dry ice (CO2 ) at its sublimation temperature. The heat of sublimation for carbon dioxide is 32.3 kJ/mol.

The ΔH°f of gaseous dimethyl ether (CH3OCH3) is −185.4 kJ/mol; the vapor pressure is 1.00 atm at −23.7°C and 0.526 atm at −37.8°C. (a) Calculate ΔH°vap of dimethyl ether.

Calculate the amount of heat (in kilojoules) required to vaporize 2.51 kg of water at its boiling point. (ΔHvap = 40.7 kJ/mol at 100oC.)

Heat is added to boiling water. Explain why the temperature of the boiling water does not change. What does change?

Substance A has the following properties.mp at 1 atm: −20.˚C bp at 1 atm: 85°CΔHfus: 180. J/g ΔH vap: 500. J/gcsolid: 1.0 J/g•°C c liquid: 2.5 J/g•°Ccgas: 0.5 J/g•°CAt 1 atm, a 25-g sample of A is heated from −40.°∆C to 100.°C at a constant rate of 450. J/min. How many minutes does it take to heat the sample to its melting point?

Assume that
the values given here for 1 mole of benzene are constant over the relevant temperature ranges.Melting point5.4 oCBoiling point80.1 oCΔHfus9.9 kJ/molΔHvap30.7 kJ/molCs,solid118 J/mol • KCs,liquid135 J/mol • KCs,gas104 J/mol • KDraw a heating curve for 1 mol
of benzene beginning at 0 °C and ending at 100 °C, and use it to answer the questions.What is the qsolid value for 1 mole of benzene?

Assume that the values given here for 1 mole of benzene are constant over the relevant temperature ranges.Melting point5.4 oCBoiling point80.1 oCΔHfus9.9 kJ/molΔHvap30.7 kJ/molCs,solid118 J/mol • KCs,liquid135 J/mol • KCs,gas104 J/mol • KDraw a heating curve for 1 mol of benzene beginning at 0 °C and ending at 100 °C, and use it to answer the questions.What is the qliquid value for 1 mole of benzene?

Assume that the values given here for 1 mole of benzene are constant over the relevant temperature ranges.Melting point5.4 oCBoiling point80.1 oCΔHfus9.9 kJ/molΔHvap30.7 kJ/molCs,solid118 J/mol • KCs,liquid135 J/mol • KCs,gas104 J/mol • KDraw a heating curve for 1 mol of benzene beginning at 0 °C and ending at 100 °C, and use it to answer the questions.What is the qgas value for 1 mole of benzene?

Substance A has the following properties.mp at 1 atm: −20.˚C bp at 1 atm: 85°CΔHfus: 180. J/g ΔH vap: 500. J/gcsolid: 1.0 J/g•°C c liquid: 2.5 J/g•°Ccgas: 0.5 J/g•°CAt 1 atm, a 25-g sample of A is heated from −40.°∆C to 100.°C at a constant rate of 450. J/min. How many minutes does it take to melt the sample?

Substance A has the following properties.mp at 1 atm: −20.˚C bp at 1 atm: 85°CΔHfus: 180. J/g ΔH vap: 500. J/gcsolid: 1.0 J/g•°C c liquid: 2.5 J/g•°Ccgas: 0.5 J/g•°CAt 1 atm, a 25-g sample of A is heated from −40.°∆C to 100.°C at a constant rate of 450. J/min. Perform any other necessary calculations, and draw a curve of temperature vs. time for the entire heating process.

A substance, X, has the following properties:Sketch a heating curve for substance X starting at 250.8C.

Use the heating–cooling curve below to answer.What is the boiling point of the liquid?

Use the heating–cooling curve below to answer.Which is greater, the heat of fusion or the heat of vaporization? Explain each term and explain how the heating–cooling curve above helps you to answer the question.

The molar heat of fusion of sodium metal is 2.60 kJ/mol, whereas its heat of vaporization is 97.0 kJ/mol. Why is the heat of vaporization so much larger than the heat of fusion?

The molar heat of fusion of sodium metal is 2.60 kJ/mol, whereas its heat of vaporization is 97.0 kJ/mol. What quantity of heat would be needed to melt 1.00 g sodium at its normal melting point?

The molar heat of fusion of sodium metal is 2.60 kJ/mol, whereas its heat of vaporization is 97.0 kJ/mol. What quantity of heat would be needed to vaporize 1.00 g sodium at its normal boiling point?

The molar heat of fusion of sodium metal is 2.60 kJ/mol, whereas its heat of vaporization is 97.0 kJ/mol. What quantity of heat would be evolved if 1.00 g sodium vapor condensed at its normal boiling point?

The molar heat of fusion of benzene (C 6H6) is 9.92 kJ/mol. Its molar heat of vaporization is 30.7 kJ/mol. Calculate the heat required to melt 8.25 g benzene at its normal melting point. Calculate the heat required to vaporize 8.25 g benzene at its normal boiling point. Why is the heat of vaporization more than three times the heat of fusion?

What quantity of energy does it take to convert 0.500 kg ice at 220. ˚C to steam at 250. ˚C? Specific heat capacities: ice, 2.03 J/g • ˚C; liquid, 4.2 J/g • ˚C; steam, 2.0 J/g • ˚C; ΔHvap = 40.7 kJ/mol; ΔHfus = 6.02 kJ/mol.

The graph below shows the heating curve of water. One can plot heating curves by measuring the increase in temperature of a given amount of ice, water or steam as a function of heat that is added at a constant rate. The first slope (from A to B) refers to the change in temperature of ice as heat is added. At the melting point of ice (B) the temperature remains constant until all of the ice is melted. The graph between C and D represents the increase in the temperature of water, until the boiling point is reached (D). At the boiling point the temperature of the water remains constant until all of the water is converted into steam (from D to E). The last portion of the curve (from E) represents the heating of steam.Determine if the following statements are True or False.Select all that are True.1. Evaporation is an endothermic process because the system is absorbing energy.2. Cooling water is an exothermic process which decreases the internal energy of the system.3. Hydrogen bonds between water molecules are broken when evaporation occurs. (From D to E)4. As water is heated, the system is absorbing energy from the surroundings.5. At the boiling point of water (1 atm and 373 K) water and steam are in equilibrium with each other.6. Heating ice (going from A to B), increases the vibrations of the atoms.7. Heating water (from C to D), decreases the kinetic energy of water.8. The freezing of water is an endothermic process because the system absorbs energy.

Calculate the heat released by cooling 54.0 g H2O from 57.0 C to -3.0 C. The transition described involves the following steps:1. Cool water from 57 C to 0 C2. Freeze water3. Cool ice from 0 C to -3 CThe specific heat values and heat of fusion is provided below: c of water = 4.18 J/g degree C c of ice = 2.09 J/g degree C Delta Hfus = 6.01 kJ/mol

he following information is given for iron at 1 atm: Tb = 2750.00°C ΔHvap (2750.00°C) = 6.338 x 103 J/gTm = 1535.00°C ΔHfus (1535.00°C) = 289.2 J/gSpecific heat solid = 0.4520 J g°C Specific heat liquid = 0.8240 J g°C A 21.20 g sample of solid iron is initially at 1511.00°C. If the sample is heated at constant pressure (P = 1 atm).________kJ of heat are needed to raise the temperature of the sample to 1849.00°C.

The following information is given for tin at 1 atm: Tb = 2270.00°C ΔHvap (2270.00°C) = 1.939 x 103 J/g Tm = 232.00°C ΔHfus (232.00°C) = 59.60 J/g Specific heat solid = 0.2260 J/g°C Specific heat liquid = 0.2430 J/g°C A 26.40 g sample of liquid tin at 497.00°C is poured into a mold and allowed to cool to 25.00°C. How many kJ of energy are released in this process? (Report the answer as a positive number.)

How many kJ of heat are needed to completely vaporize 3.30 moles of H2O? The heat of vaporization for water at the boiling point is 40.6 kJ/mole. Select one.

Enter your answer in the provided box. Sulfur dioxide is produced in enormous amounts for sulfuric acid production. It melts at -73.0°C and boils at -10.0°C. Its ΔH° is 8.619 kJ/mol, and its ΔH°vap is 25.73 kJ/mol. The specific heat capacities of the liquid and gas are 0.995 J/g • K and 0.622 J/g • K, respectively. How much heat is required to convert 8.000 kg of solid SO2 at the melting point to a gas at 60.0°C?

Be sure to answer all parts.Determine the final state and temperature of 100 g of water originally at 25.0 °C after 50.0 kJ of heat have been added to it. (a) water (b) water and steam (c) steam

Based on the thermodynamic properties provided for water (below and to the left), determine the amount of energy released for 0.6 kg of water to go from 107.7 °C to 63.2 °C.

The fluorocarbon compound C2Cl3F3 has a normal boiling point of 47.6 ºC. The specific heats of C2Cl3F3 (l) and C2Cl3F3(g) are 0.91 J/g.ºK and 0.67 J/g.ºK, respectively. The heat of vaporization for the compound is 27.49 kJ/mol. A. Calculate the heat required to convert 37.5 g of C2Cl3F3 from a liquid at 12.75 ºC to a gas at 75.10 ºC. Express your answer using two significant figures.

The following information is given for cadmium at 1 atm:boiling point = 765.0°CΔHvap(765.0°C) = 886.6 J/gmelting point = 321.0°CΔ Hfus(321.40°C) = 54.40 J/gspecific heat solid = 0.2300 J/g° Cspecific heat liquid = 0.2640 J/g°CA 27.60 g sample of solid cadmium is initially at 303.0°C. If the sample is heated at constant pressure (P = 1 atm), kJ of heat are needed to raise the temperature of the sample to 398.0°C.

After an afternoon party, a small cooler full of ice is dumped onto the hot ground and melts. If the cooler contained 7.80 kg of ice and the temperature of the ground was 38.0 ° C, calculate the energy that is required to melt all the ice at 0 ° C. The heat of fusion for water is 80.0 cal/g.

At 1 atm, how much energy is required to heat 79.0 g of H2O(s) at-22.0 degree C to H2O(g) at 119.0 ° C? Helpful constants can be found here.

To treat a burn on your hand, you decide to place an ice cube on the burned skin. The mass of the ice cube is 20.0 g, and its initial te temperature is -12.3°C. The water resulting from the melted ice reaches the temperature of your skin 29.3°C . How much heat is absorbed by the ice cube and resulting water? Assume that all the water remains in your hand. Constants may be found here.

A pure solid sample of Substance X is put into an evacuated flask. The flask is heated at a steady rate and the temperature recorded as time passes. Here is a graph of the results: Use this graph to answer the following questions:What is the melting point of X?What phase (physical state) of X would you expect to find in the flask after 9 kJ/mol of heat has been added? (check all that apply) • solid • liquid • gas

Water (2150 g) is heated until it just begins to boil. If the water absorbs 5.47 x 105 J of heat in the process, what was the initial temperature of the water?Express your answer with the appropriate units.

Ethanol (C2H3OH) melts at-114°Cand boils at 78°C. The enthalpy of fusion of ethanol is 5.02 kJ/mol and its enthalpy of vaporization is 38.56 kJ/mol. The specific heat of solid and liquid ethanol are 0 97 J/g • K are 2.3 J/g • K respectively.How much heat is required to convent 42.0 g of ethanol at -143 °C to the vapor phase at 78 °C? Express your answer using two significant figures.

An ice bag containing 294 g of ice at 0°C was used to treat sore muscles. When the bag was removed, the ice had melted and the liquid water had a temperature of 17.0°C. How many kilojoules of heat were absorbed?

A 750 gram block of ice @ -19 °Celsius absorbs 550,000 joules of energy. Describe the sample of water after it has absorbed all of that energy.

Enter your answer in the provided box.Calculate the amount of heat (in kJ) required to convert 140.1 g of water to steam at 100 °C.

Enter your answer in the provided box. From the data below, calculate the total heat (in J) needed to convert 20.00 g of ice at -4.75°C to liquid water at 0.450°C:

The heat of vaporization of water is 40.66 kJ/mol. How much heat is absorbed when 2.96 g of water boils at atmospheric pressure?

How much heat energy is required to melt 388.9 g of HBr? The molar heat of fusion of HBr is 2.41 kJ/mol.

How much heat energy is required to convert 70.7 g of solid iron at 29 °C to liquid iron at 1538 °C? The molar heat of fusion of iron is 13.8 kJ/mol. Iron has a normal melting point of 1538 °C. The specific heat capacity of solid iron is 0.449 J/g °C.

At 1 atm, how much energy is required to heat 59.0 g of H2O(s) at -120°C to H2 O(g) at 119.0°C?

How many grams of benzene, C6H6, can be melted with 39.7 kJ of heat energy? The molar heat of fusion of benzene is 9.95 kJ/mol.

How much heat energy is required to boil 93.7 g of ammonia, NH3? The molar heat of vaporization of ammonia is 23.4 kJ/mol.

Calculate the heat energy released when 14.3 g of liquid mercury at 25.00°C is converted to solid mercury at its melting point.

At 1 atm, how much energy is required to heat 39.0 g of H2O(s) at -18.0°C to H2O 125.0°C? Helpful constants can be found here.

After an afternoon party, a small cooler full of ice is dumped onto the hot ground and melts. If the cooler contained 5.60 kg of ice and the temperature of the ground was 43.0°C, calculate the energy that is required to melt all the ice at 0°C. The heat of fusion for water is 80.0 cal/g.

The fluorocarbon C2Cl3F3 has a normal boiling point of 47.6°C. The specific heats of C2Cl3F3(l) and C2Cl3F3(g) are 0.91 J/gK and 0.67 J/gK, respectively. The heat of vaporization of the compound is 27.49 kJ/mol. The heat required to convert 50.0 g of the compound from the liquid at 5.0°C to the gas at 80.0°C is ________ kJ. (i) 3031 (ii) 8.19 (iii) 10.36 (iv) 30.51 (v) 1454

Calculate the amount of heat required to completely sublime 32.0 g of solid dry ice (CO2) at its sublimation temperature. The heat of sublimation for carbon dioxide is 32.3 kJ/mol. Express your answer in kilojoules.

In a large building, oil is used in a steam boiler heating system. The combustion of 1.0-lb of oil provide: 2.4 x 107 J. How many kilograms of oil are needed to heat 170 kg of water from 22°C to 100°C?

Enter your answer in the provided box. Sulfur dioxide is produced in enormous amounts for sulfuric acid production. It melts at 73.0 °C and boils at -10.0 °C. Its ΔH°fus is 8.619 kJ/mol, and its ΔH°vap is 25.73 kJ/mol. The specific heat capacities of the liquid and gas are 0.995 J/g. K and 0.622 J/g. K, respectively. How much heat is required to convert 5.000 kg of solid SO2 at the melting point to a gas at 60.0 °C?

How much heat energy is required to convert 93.1 g of liquid sulfur dioxide, SO2, at 199.9 K to gaseous SO2 at 263.1 K if the molar heat of vaporization of SO2 is 24.9 kJ/mol, and the specific heat capacity (C) of liquid SO2 is 1.36 J/(g·°C)?

How much heat energy is required to convert 97.9 g of liquid sulfur dioxide, SO2, at 202.6 K to gaseous SO2 at 263.1 K if the molar heat of vaporization of SO2 is 24.9 kJ/mol, and the specific heat capacity (C) of liquid SO2 is 1.36 J/(g • °C)?

Calculate the heat energy released when 12.7 g of liquid mercury at 25.00°C is converted to solid mercury at its melting point.

Calculate the change in heat when 13.15 g of water vapor (steam) at 100.0 °C condenses to liquid water and then cools to 7.50 °C.

The following information is given for magnesium at 1 atm: boiling point = 1090°C ΔHvap (1090°C) = 132 kJ/molmelting point = 649°C ΔHfus (649°C) = 8.95 kJ/molspecific heat solid = 1.02 J/g°C specific heat liquid = 1.34 J/g°C ________ kJ are required to melt a 43.1 g sample of solid magnesium, Mg, at its normal melting point.

Enter your answer in the provided box.Calculate the amount of energy (in kJ) necessary to convert 457 g of liquid water from 0°C to water vapor at 172°C. The molar heat of vaporization (Hvap) of water is 40.79 kJ/mol. The specific heat for water is 4.184 J/g • °C, and for steam is 1.99 J/g • °C. (Assume that the specific heat values do not change over the range of temperatures in the problem.)

At 1 atm, how much energy is required to heat 81.0 g of H2O (s) at -24.0°C to H2O (g) at 135.0°C?

At 1 atm, how much energy is required to heat 73.0 g of H 2O(s) at - 14.0 °C to H 2O(g) at 139.0 °C?

How much energy is needed to convert 70.6 grams of ice at 0.00°C to water at 75.0°C? specific heat (ice) = 2.10 J/g°C specific heat (water) = 4.18 J/g°C heat of fusion = 333 J/g heat of vaporization = 2258 J/g

Calculate the heat energy released when 23.6 g of liquid mercury at 25.00°C is converted to solid mercury at its melting point.

Ethanol (C2H5OH) melts at-114°C and boils at 78°C. The enthalpy of fusion of ethanol is 5.02 kJ/mol, and its enthalpy of vaporization is 38.56kJ/mol. The specific heat of solid and liquid ethanol are 0.97 j/g • K are 2.3 J/g • K respectively. Part A How much heat is required to convert 24.0 g of ethanol at 21°C to the vapor phase at 78°C? Express your answer using two significant figures.

After an afternoon party, a small cooler full of ice is dumped onto the hot ground and melts. If the cooler contained 8.30 kg of ice and the temperature of the ground was 43.5°C, calculate the energy that is required to melt all the ice at 0°C. The heat of fusion for water is 80.0 cal/g.

Determine the amount of heat (in kJ) required to vaporize 1.50 kg of water at its boiling point. For water, ΔHvap = 40.7 KJ/mol (at 100°C). Express your answer to three significant figures and include the appropriate units.

Based on the thermodynamic properties provided for water, determine the energy change when the temperature of 0.850 kg of water decreased from 123°C to 19.0°C.

How much heat energy is required to convert 19.6 g of solid ethanol at-114.5°C to gasesous ethanol at 177.9°C? The molar heat of fusion of ethanol is 4.60 kJ/mol and its molar heat of vaporization is 38.56 kJ/mol. Ethanol has a normal melting point of-114.5°C and a normal boiling point of 78.4°C. The specific heat capacity of liquid ethanol is 2.45 J/g•°C and that of gaseous ethanol is 1.43 J/g·°C.

Part AHow much heat is evolved in converting 1.00 mol of steam at 145°C to ice at 55.0°C? The heat capacity of steam is 2.01 J/(g°C) and of ice is 2.09 J/(g°C). Express your answer in units of kilojoules. Assume the system is at atmospheric pressure.

How much energy would be absorbed or released by the H 2O in the process.30 grams H2O(g) at 100°C → 30 grams H2O(l) at 20°C?1. 16.2 kcal absorbed2. 16.2 kcal released3. There would be no transfer of energy due to the first law of thermodynamics4. 18.6 kcal released5. 2.4 kcal released