unit 11 kinetic molecular theory packet page 1 of 13 ... 11 packet.pdfchemistry—unit 11 kinetic...

Unit 11 Kinetic molecular theory packet of 13

Chemistry—Unit 11 Kinetic Theory

Unit Quiz: Test

Objectives

Be able to define pressure and memorize the basic pressure units.

Be able to convert to/from: atms, kPa, Pa, mm Hg, torrs.

Differentiate between the three states of matter.

Be able to identify and name the changes of state.

Be able to read a phase diagram and identify the triple point, phases and phases changes

Be able to describe surface tension

Understand the concept of atmospheric pressure and its effect on boiling point.

Understand the properties of water due to hydrogen bonding.

Be able to label a boiling point/freezing point graph

Be able to do a heat of fusion and heat of vaporization problem.

Be able to describe in detail properties of solids, liquids and gases

Be able to complete specific heat problems.

Understand how to measure temperature and heat (including units)

Vocabulary –

Sublimation Vapor Melting Condensation Evaporation Triple point

Surface tension Kinetic Theory Kelvin Absolute Zero Fluids ideal gas

volatile molar heat of vaporization Molar heat of fusion heat

exothermic endothermic Joule Specific heat temperature Viscosity

Phase Change Diagram

Rank Solids, liquids and gases from weakest attractive forces to strongest:

________________ _______________ ______________

Boiling point is the temperature at which the vapor pressure of the liquid equals the atmospheric pressure. For

vapor to form, the molecules of the liquid must overcome the forces of attraction between them.

___________________________________________________. If the atmospheric pressure is lower the boiling point will be

lower; if the atmospheric pressure is higher the boiling point will be higher.

**** ****

Boiling point is ___________________________ Boiling point is __________________________________


VAPOR PRESSURE GRAPHS Use the graph to answer

the following questions.

1. What is the vapor pressure of CHCl3 at

50C?

2. What is the boiling point of H2O when

theexternal pressure is 30 kPa?

3. What is the normal boiling point of CCl4?

Which substance has the weakest IMF?

__________________

Kinetic theory of Matter class notes:

Fill in the following assumptions of the kinetic theory:

A. Gases are made up of tiny particles that are _____

apart.

B. Collisions between particles are elastic meaning no

energy is lost or gained overall.

C. Gases are in constant motion - they possess

_______ energy – energy of motion

D. There are ______ forces of attraction between gas particles.

E. The average kinetic energy of the particles depends on the ___________ of the gas.

Higher temp = _____________ KE


PHASE DIAGRAM EXPLANATION (water)

Much of the information we have discussed can be shown in a phase diagram, which shows how the states of

matter in a system are affected by changes in temperature and pressure. Below is a phase diagram for water.

Each area represents all of the conditions under which water in that phase can exist. Together we will label

the phases of matter. Point A is called the triple point because all three states are in equilibrium at this

temperature and pressure (0.01˚C and 0.006 atm). Point C represents the critical point; above this point, there

is no vapor pressure curve. Only the gaseous state exists at pressures and temperatures above this point.

G is the normal melting point, at 1 atm (standard(normal) atmospheric pressure where the state changes from

solid to liquid. F, the normal boiling point, is that temperature at which the liquid-vapor equilibrium curve is

cut by the pressure line of 1atm.


Answer the following Q’s for the diagram

on the right:

______1. What is the normal melting point?

___2. What is the normal condensing

point?

______3. What is the normal freezing point?

___4. What is the normal boiling point ?

______5. What is the temperature and

pressure at the triple point?


Heating and cooling curves represent how Temperature (or energy) affects the states of matter; we will fill in

the heating curve graph that is below together!


Practice Sheet Heating and Cooling curves!

For each of the above label the states of matter, the BPt, MPt cond Pt. and Frz pt.

For Graph A – cooling of a gas 1.What is the freezing point temp? _________ The boiling point temp? _________

2.Where is Kinetic energy changing? _______ Potential energy? ________

For Graph B – water heating 1.What is the melting point temp? _________ The boiling point temp? _________

2.Where is Kinetic energy changing? _______ Potential energy? ________

3.Why is the temperature staying constant between points 2 and 3 even though heat is being added to the

water?

Energy and specific heat notes

Temperature: ____________________________________________________________________________________________________

Units for temperature:

Conversions: 0C = 5/9 (0F -32) K = 0C + 273 In packet You will not use 0F in this class.

Scale comparisons:

Water K 0C 0F

______________________ 373 100 212

______________________ 273 0 32

______________________ 0 -273 -460 also known as absolute zero, no movement

Heat: _______________________________________________________________________________________________________________

Heat/Energy – _____________________________________________________________________________________________________

___________________________________________________________________________________________ (T).

Unit = J (joule) or cal (calorie) 1 calorie (cal) = 4.184 J

Food calories are actually kilocalories 1 food calorie = 1000 “chemical” calories

Calorimeter: device used to measure energy changes – see picture on pg 519

(a thermometer measures temperature changes not energy!)


Law of conservation of Energy:

Heat lost in a calorimeter is equal to the heat gained by the water in the calorimeter.

qlost = (-qgain) Therefore you can use a calorimeter to calculate the heat or the specific heat for a substance

burned in a calorimeter.

specific heat: C or Cp : _______________________________________________________________________________________

_________________________________________________________________________________________________________________

because the sizes of the degree divisions on both scales are equal.

**Specific heat is usually measured under constant pressure conditions – the subscript p (Cp) – is used as a

reminder.

Specific heat is a constant for a substance. ***** _______________________________________________________________

__________________________________________________________________________________________________________________

***_______________________________________________________________________________________________________________

___________________________________________________________________________________________________________________

Specific heat for liquid water is 4.18 J/g0C (given in packet!)

Formula for specific heat: ________________________________________________________________________________________

Equation can be rearranged to

Example copy work!

A 4.0 g sample of glass was heated from 0 0C to 41 0C, and was found to absorb 32 J of heat. What is the

specific heat of the glass?

Example 2: How much heat energy is needed to raise the temperature of a 55 g sample of aluminum from

22.4 0C to 94.6 0C? The specific heat for Al is 0.897 J/g-C


You try: How much heat energy is needed to raise the temperature of a 25 g sample of water from 32 0C to 55 0C (use your packet to find the specific heat of liquid water)

Example 3 A 32.70 gram piece of metal is heated to 98.7 0C. It is then added to a calorimeter containing 150g

of water at 23.5°C. The final temperature of the water and metal is 25.2 0C. What is the specific heat of this

metal? What’s the ID of the metal?

You try: A 180 gram piece of metal is heated to 65 0C. It is then added to a calorimeter containing 100 g of

water at 25 0C. The final temperature of the water and metal is 42 0C. What is the specific heat of this metal?

To calculate the heat required for changes of state—Use this when temperature stays constant

q = mHv for boiling/condensing and q = mHf for freezing/melting

for water: Hv = 2,260 J/g and Hf = 334J/g (all of this is in the packet!)

Example: Calculate the amount of heat required to melt 55 g of ice (at 0 0C)

You try: Calculate the amount of heat required to boil 42 g of water (at 100 0C)

To help with calculations – remember the heating curve

changes in temp (kinetic energy changes ) = use specific heat ( no phase change)

changes in phase (change in potential energy ) = use molar heat (no temp change )


Practice Heat and It’s measurement: Problems:

1. How many joules of heat are given off when 5.0g of water cools from 75 0C to 250C?

(specific heat of water is in the packet – look it up!)

2. How many joules does it take to melt 35 g of ice at 0 0C ?

3. How much heat energy is given off when 85 g of steam condenses?

4. How many joules of heat are necessary to raise the temperature of 25 g of water from 10 0C to 60 0C ?

5.What mass of a substance (with a specific heat of 0.071 J/g-K) will absorb 45 J as it is heated from 293 K to

313 K?

6.When 80.0 grams of a certain metal at 90.0 °C was mixed with 100.0 grams of water at 30.0 °C, the final

equilibrium temperature of the mixture was 36.0 °C. What is the specific heat of the metal?

Heating Curve/Calorimetry

1) What is happening to the average kinetic energy of the

molecules in the sample during section

2?__________________________

2) As a substance goes through section (2), what happens

to the distance between the particles?________________________

3) What is the name of the process happening during

section (4)?

4) What would be the name of the process happening

during section (4) if time were going the other way?

_________________

5) What is the melting point of this

substance?__________________________________

6) At what temperature would this sample finish boiling?

_________

7) When this substance is melting, the temperature of the

ice-water mixture remains constant because:

a. Heat is not being absorbed c. Heat is being converted to potential energy

b The ice is colder that the water d. Heat is being converted to kinetic energy


8) When a given quantity of water is heated at a constant rate, the phase change from liquid to gas takes

longer than the phase change from solid to liquid because

a. The heat of vaporization is greater than the heat of fusion

b. The heat of fusion is greater than the heat of vaporization

c. The average kinetic energy of the molecules is greater in steam than in water

d. Ice absorbs energy more rapidly than water does

9) The temp. when a substance in the liquid state freezes is the same as the temp at which the substance

a. Melts b. Sublimes c Boils d.Condenses

10) If 150 grams of water is heated from 20oC to 30oC, the number of joules of heat energy absorbed is:

11) If a 2.0 g sample of water at 5.0oC absorbs 21.8 J of heat energy, the temperature of the sample will be

raised by?

12) How much heat is required to boil 4.5 g of water (at 100 C)?

13) A sample of water is heated from 10.oC to 15oC by the addition of 130 joules of heat. What is the mass of

the water?

14) How much energy is released as 30.0 grams of water (at 0 oC) is frozen?

15) A 2.70 gram piece of metal is heated to 98.7 C. It is then added to a beaker containing 150 mL (=150 g) of

water at 23.5 C. The final temperature of the water and metal is 25.2 C. What is the specific heat of this metal

Review Kinetic Theory

Convert the following

1. 1.25 atm to mm Hg 2. 5.38 kPa to torr 3. 324 KPa to atm

4. How would the boiling point of water be different in Death Valley (below sea level)? Why?

5. How much energy would be required for each of the following?

A. melting 45 g of ice B. boiling 45.0 g of water

6. a 250 g sample of metal is heated to 55 0C. It is then added to a calorimeter holding 100 g of water at 30 0C. The final temperature of the water and metal is 43 0C. What is the specific heat of the metal


Specific heat lab

Background: One physical property of a pure substance is the amount of heat energy it will absorb per unit

mass. This property is called specific heat.

We will use a calorimeter in lab to measure specific heat. Our calorimeter will be a simple one made out of a

styrofoam cup.

To calculate specific heat you will use two formulas:

Heat gained by the water = (heat lost by the metal) = mass of water x T x specific heat of water (4.18 J/g-C)

Specific heat of metal = heat gained by the water

Mass of metal x T

Objective: To use lab data to calculate the specific heat for a metal.

To determine your percent error in calculating specific heat.

Equipment: Large beaker styrofoam cup thermometer test tube metals

Procedure:

1. Fill a large beaker about half full of tap water.

2. Place the beaker on a hot plate and heat to boiling.

3. Measure the mass of an empty and dry test tube and record.

4. Add the sample metal pieces until the test tube is half full. Record the mass of the test tube and metal and

record the identity of the metal.

5. Place the test tube gently into the beaker of water and continue heating the water to the boiling point –

use the gold test tube tongs.

6. Obtain a styrofoam cup (used as a calorimeter) & measure mass carefully. Record.

7. Fill the cup about half full with water and record the temperature of the water and the new mass of the cup

(do not get water on the balance!)

8. Allow the water in the beaker to boil for several minutes – (It will be assumed that the temperature of the

metal is the same as the boiling water – 100 0C)

9. Remove the test tube from the boiling water and dump the metal into the styrofoam cup – try to get as

little water into the cup as possible.

10. Stir the water in the cup slowly and carefully and then record the highest temperature reached.

11. Recover the metal by carefully pouring the water off (decanting) and then spreading the solid on a paper

towel. Do not allow any metal to get in the sink

Data Table: Include units!

A. Mass of empty test tube _____________

B. Mass of test tube & metal ___________

C. Mass of metal_______________ (B-A)

D. Mass of cup________

E. Mass of cup and water:__________

F. Mass of water____________(E-D)

G. Temperature of water in cup_________

H. Highest temperature of water and hot metal in cup________

I. Change in water temperature (∆T) caused by hot metal = ___________ (H-G)


Names: ______________________________________ Pd: __________________

Calculations: Include units on final answers!

1. Find the heat transferred from the hot metal to the water in the calorimeter.

q = m x C x Δ T

C = 4.18 J/g-0C m = mass of water in calorimeter (line F) Δ T = change of temp of water (line I)

q = ________________________

2. Find the specific heat of your metal by using the specific heat formula.

q = m x C x Δ T

q = answer found above m = mass of metal (line C) ΔT = temp change of metal (100 0 C – line H)

C = _______________________

3. Match the C, specific heat, you calculated in #2 to the closest metal from your reference packet.

True value for ______________________(metal name) ‘s Specific heat is __________________ ( C in RefPckt)

(Write the name of your metal above) (specific heat w/ units)

4. Find your percent error (show work!)

% error:___________________

5. Out of aluminum, copper, zinc and nickel which would be the best:

A. Heat conductor:_______________ Why?

B. Insulator:___________________ Why?

5)Test review questions A. How much heat is required to raise the temp of 50 g of water from 22 0 C to 88 0 C?

_____________________

B. What is the specific heat of 25 g of a metal that releases 120 J of heat when cooled from 75 0C to 30 0C?

_____________________

unit 11 kinetic molecular theory packet page 1 of 13 ... 11 packet.pdfchemistry—unit 11 kinetic...

Documents