Honors Chemistry
Unit 1B
Matter, Properties, & Energy
(2016-2017)
 Describe states of matter and common properties.
o Molar Mass
o Moles, Molecules, and Grams Conversions
 Label a heating / cooling curve
o Solid, liquid, gas
o Evaporation, condensation, freezing, melting
o Enthalpy of fusion, enthalpy of vaporization
o Specific heat
o Boiling point, melting point
 Separate mixtures based on physical properties.
o boiling point (distillation), magnetism, density etc.
 Evaluate energy changes of matter (Specific Heat & Calorimetry)
o Calculation of specific heat (q = mCΔT)
o Calorimetry of various systems involving exothermic and
endothermic heat exchange.
o The calculation of energy released from a food substance using
calorimetry.
1
We are looking for:
1a. Calculate the molar mass of a compound/element using a periodic table.
1b. Using molar mass and unit analysis, convert moles of a given compound to grams of that compound and vice versa.
1c. Using Avogadro’s number and unit analysis, convert atoms/molecules of a compound to moles
of that compound.
2a. Physical properties such as boiling point, magnetism, density etc.
2b. Measure the density of various samples and use the density to identify the material.
3a. Identification of all phase changes and energy change values
3b. Evaporation, condensation, freezing, melting
3c. Enthalpy of fusion, enthalpy of vaporization
3d. Specific heat
3e. Boiling point, melting point
3f. Solid, liquid, gas
4a. Calculations of energy released/gained using specific heat (q = mCΔT)
4b. Calorimetry of various systems involving exothermic and endothermic heat exchange.
4c. The calculation of energy released from a food substance using calorimetry.
What’s the MATTER
Matter:
 Anything that has _______ and takes up __________.
Matter is made up of building blocks:
_________
– smallest unit of an element.
_________
– a pure substance made of only one kind of atom.
_________
– made of two or more atoms that are chemically
combined.
 90% of the Earth’s crust is made up of only 5 elements:
Oxygen
49.2%
Silicon
25.7 %
Aluminum
7.5%
Iron
4.7%
Calcium
3.4%
States of Matter
 Solid



Definite __________ and ___________
Particles are __________ packed
Slight expansion when ___________
Incompressible
2
 Liquid



Has definite __________, but no definite __________ (assumes the shape of the container)
Particles are ____________ packed (can flow)
Easily expand when ___________
Considered incompressible
 Gas



No definite ___________or ___________
___________ to fill the container
Particles are spaced _______ apart
Compressible
 Plasma



Consists of _______________ charged particles
It’s an ionized _________
Common in __________, but very rare on ____________
Found in lightning, fluorescent lights and neon signs
Energy Amounts in States of Matter




Solid- little energy, particles vibrate and rotate
Liquid- more energy, they move freely
Gas- even more energy, move quickly
Plasma- most energy, move extremely fast
Names of Phase Changes






Solid to Liquid
Liquid to Gas
Gas to Liquid
Liquid to Solid
Solid to Gas
Gas to Solid
= ______________
= Boiling / Evaporation
= ______________
= Freezing
= ______________
= Deposition
Types of Matter
 Pure Substance Matter with a ________ composition
 It has distinct properties
 Examples =
elements
compounds
 Mixtures Most matter is a mixture
 The composition is not fixed (changes from sample to sample)
 Solute:
 Solvent:
 Two Types –
____________________
____________________
3
Homogeneous Mixtures:
 Composition is ______________ throughout
 Solution
 Particle size = 0.01 – 1 nm
 Doesn’t settle out upon ________________
 Can’t be separated by __________________
 Doesn’t scatter _____________
 Example = distilled water
Heterogeneous Mixtures
 Composition is ______________ throughout
 Suspension
 The sample varies in composition, properties and _______________
 No ______________
 Particle size is greater than 1000 nm
 Particles ____________________ upon standing
 Can be separated by filtration
 Might scatter light
 Examples = soil, trail mix, _________________
 Colloid
 Particle size = 1 – 1000 nm
 Doesn’t settle out upon standing
 Can’t be separated by filtering
 Scatters light ( ____________________ )
 Examples = milk, gelatin, smoke
Physical vs. Chemical Properties
 Every substance has a unique set of properties (characteristics that identify that substance)
Physical Change A change in matter from one form to another without changing its ______________________________ (most can
be reversed)
 Examples =
 Change in state
 Dissolving
 Compressing
 Physical Properties Properties that can be measured without changing the identity and composition of the substance
 Physical Property Examples Color
 Odor
 Density
 Melting Point
 Boiling Point
 Hardness
 Solubility
4
Chemical Change A change in matter from one form to another by changing its ______________________________ (most cannot be
reversed)
 Examples of what to look for = Chemists Get Practice Trying Labs
 C__________
 G__________
 P__________
 T__________
 L__________
 Examples =
 Combustion
 Electrolysis of water
 Any reaction that produces a new product like water, a gas, a solid (precipitate in solution)
 Chemical Properties Properties that describe the way a substance may change to form other substances
 Only observed when a _______________________takes place
 Chemical Property Examples




Heating to combustion
Reactivity with water or acid
Flammability
Corrosion
Decomposition
Law of Conservation of Mass = In a physical change or a chemical reaction, mass is neither created nor destroyed (Antoine
Lavoisier)
5
Density
Density is a measure of mass per volume.
D=
Answer the following questions on density and experimental error. Show all your work with units and round your
answer to the correct number of sig. figs.!
1. What is the density of a cardboard if 6.2 g occupy 8.56 cubic centimeters?
2. What is the density of a gold nugget having a volume of 2.39 cubic centimeters and a mass of 45.58 grams?
3. What is the mass of a piece of aluminum having a volume of 15.12 cubic centimeters and a density of 2.70 grams per
cubic centimeter.
4. Cerium sulfate has a density of 3.17 grams per cubic centimeter. What is the volume of .54 grams of this substance?
5 What is the density of a brick if 51.21 g occupy 31.32 cubic centimeters?
6. Cerium sulfate has a density of 3.17 grams per cubic centimeter. What is the volume of 1.25 grams of this substance?
7. Tin has a density of 7.28 grams per cubic centimeter. What is the volume of 11.2 grams of this substance?
6
Percent Error (Experimental Error)
Density and Percent Error Practice
1) You measure the density of substance in the lab as 4.25 g/mL. The true value of the density of the substance is 4.32
g/mL. Calculate your percent error.
2) Your measurement of the volume of a sample of tap water is 9.6 mL, 9.52 mL and 9.553 mL using 3 different
graduates. What is the average volume of the tap water?
3)
The average mass of this tap water is 9.2 g, what is the density?
4) The standard density of water is 1.00 g/mL, What is your percentage error for the above problem?
5) In the lab, you determine the density of a metal to be 2.57 g/cm3. The accepted or theoretical value of the density
for this metal is 2.72 g/cm3. What is your % error for your lab?
7
Layered Solutions Activity
Objective:
To create a column of distinct layers of different solutions.
Procedure:
1. In five different cups, place the following ingredients:
Chart 1: Ingredients for layered solution (you must use at least 20mL but each cup must use
a different volume of water and each must differ by at least 10 mL).
Cup Number/color
Salt (g)
Warm water (mL)
2. Once you have added all of the necessary materials, stir each one with a spoon for at least one
minute until all of the salt is completely dissolved.
3. Using the balance and 10 mL of each solution you just made, measure the mass of each solution.
Record these values in Table 1 below.
4. Calculate the density of each cup’s solution and show your work in the appropriate column. Round
your answer to the correct number of sig. figs.
5. Determine the order to put the liquids into the graduated cylinder. Which one should go in first,
the least dense or the most dense?
6. Add food coloring to create a rainbow effect with the different layers.
7. Using your pipette, carefully transfer 10 mL from each cup into the 50 mL graduated cylinder. (or
use about 2 ml of each and put into a 10 mL cylinder)
Observations:
Table 1: Data from the Layered Solution Activity
Cup
Color
Sample
Volume (mL)
Sample Mass
(g)
Work for density
calculation
Density of
solutions
made
(g/mL)
Projected
order to
fill
graduate
d
cylinder
8
Lab: Graphing and Density
Name:________________________
Class Period:_____
Purpose: -Determine the density of a liquid from a graph of mass and volume.
-Determine the layering order of three liquids if poured together into a
graduated cylinder.
Problem: Are density and the layering order of liquids in a graduated cylinder
related?
Hypothesis: (If…,then…)
Experiment:
Materials:
25mL graduated cylinder
balance
rubbing alcohol
water
dropper
calculator
Leogos
Procedure:
1) Determine the mass of 5mL of water and record this in the data table.
a. Place the empty graduated cylinder on the balance and record this mass below:
i. Mass of empty cylinder:_____________g
b. Place the 5mL of water in the cylinder and carefully place it on the balance.
c. Subtract the mass of the empty cylinder.
d. Record the mass in the data table.
2) Repeat step 1 for 15mL and 25mL of water.
3) Repeat step 1 for 5mL, 15mL, and 25mL of rubbing alcohol.
a. Use the beaker of rubbing alcohol that is at your table. Return the alcohol to the beaker
when finished for the next class to use!!
4) We will not be measuring the values for silicone oil because it is too messy. The data has already
been given to you in the data table.
5) Measure each dimension (L, W, H) of each lego block sample and record in the data table.
6) Measure the mass of each lego sample and record in the data table.
7) Graph the data for each substance on the same sheet of graph paper.
a. Use a different colored pencil for each line.
b. Provide a key to identify each color.
c. Make a title for the graph.
d. Make a best-fit line for each color. The line must go through 0,0.
e. Determine the slope of each line and show your work on the graph. Make sure you put
units on your slopes.
9
Liquids Data:
Volume (mL)
Mass (g)
Water
Rubbing
Alcohol
Silicone
Oil
5.00
4.60
15.0
13.70
25.0
23.10
Lego Data:
Length
(cm)
Width
(cm)
Height
(cm)
Calculated volume,
rounded for sig.
figs. (cm3)
Mass (g)
10
Stoichiometry
From 2 greek words:
Stoicheion = element
1792 – German Chemist
Metron = measure
Jeremias Benjamin Richter
Is concerned with the amount of substances involved in a reaction
Composition stoichiometry = mass relationships between elements in compounds
Ex: Na2SO4
2 Na / 1 SO4
Avogadro’s Number = Number of particles in a mole
6.02 X 1023
602,000,000,000,000,000,000,000
(Named after Amadeo Avogadro – 1776-1856 Italian chemist and physicist)
Molar Mass
Mass in grams of one mole of an element or compound
Numerically equal to the atomic weight of the element
or
the sum of all the atomic weights in the formula
11
Molar Mass Examples
1. NaCl = 22.99 + 35.45 = 58.44 g/mole
2. CuSO4  5H2O = 63.55 + 32.07 + 4(16.00) + 5 [(2 x 1.01)+ (16.00)]
= 159.62 + 90.10
= 249.72 g/mole
Calculate the Molar Mass for each of these compounds.
1.
KCl
2.
Li2SO4
3.
(NH4)2C2O4  H2O
4.
KOH
5.
CuBr2
6.
Mg3(PO4)2
7.
Si3O7
12
Using Molar Mass
Converting Moles to Grams
and
Converting Grams to Moles
Use the molar mass to convert the given moles into grams or given grams to moles. Write the answer
to the problem on the line provided. Show all of your work using factor label method. Report all
answers to appropriate sig. figs. using the given moles or mass in the problem.
Work Space
__________ 1. 7.8 moles of Fe2O3
__________ 2. 100.2 moles of Pb(NO3)2
___________3. 1.22 moles of CO2 to grams
__________ 4. 120.8 grams of K2SO4 to moles
__________ 5. 4.6 grams of MgCl2 to moles
__________ 6. 2.3 grams of Ba3(PO4)2 to moles
13
Using Avogadro’s Number (6.02x1023 )
Converting Moles to Molecules (or particles or atoms)
And Back Again
1) 4.50 Moles NaCl = ? Molecules of NaCl
2) 6.62 x 1024 Molecules of H2O = ? Moles of H2O
3) 91.20 Moles of CO2 = ? Molecules of CO2
4) 3.01 x 1023 Molecules CuCl2 = ? Moles of CuCl2
5) 1345.9 Moles MgO = ? Molecules of MgO
6) 4.87 x 1035 molecules of H2O contains how many atoms of hydrogen?
7) 2.36 moles of FeCl3 contains how many atoms of chlorine?
14
More Converting…
Using Avogadro’s Number and Molar Mass
Work out the following conversions on a separate sheet of paper. Show all your work with units of
measurement. Round your answer to sig. figs.
1) 1.806x 1024 atoms of iron = ? grams iron
2) 1.20 x 1028 molecules CuSO4 = ? grams CuSO4
3) 380.84 grams MgCl2 = ? molecules MgCl2
4) 82.81 grams Pb(NO3)2 = ? molecules Pb(NO3)2
5) 37.6 grams H2SO4 = ? atoms of oxygen
6) 75.26 grams Al2(SO4)3 = ? atoms of sulfur
Moles, Mass, and Molecules (or atoms) of Compounds (or elements)
Work out the following conversions on a separate sheet of paper. Show all your work with units of
measurement. Round your answer to sig. figs.
1) 0.738 moles of Fe2O3 to grams.
2) 50.5 g of FeBr3 to moles.
23
3) 1.51 x 10 molecules of PbI2 to moles.
4) 0.445 moles of CCl4 to molecules.
5) 0.538 moles of Ce2(CO3)3 to grams.
6) 150.4 g of Ce(CO3)2 to moles.
25
7) 7.22 x 10 molecules of CuCl2 · 4 H2O.
8) 1.45 moles of Pb(C2H3O2)2 to grams.
24
9) 1.22 x 10 molecules of CO2 to grams.
10) 19.3 grams of Mg to atoms.
15
Heating and Cooling Curve Terms/Definitions
TemperatureSpecific Heat –
Solid –
LiquidGas –
Plasma –
Heating Curve
Enthalpy (heat) of Fusion/ Molar heat of fusion–
Melting–
Melting Point –
Enthalpy (heat) of Vaporization/Molar heat of Vaporization–
Evaporation–
Boiling Point –
Sublimation-
Cooling Curve
Condensation –
Condensation Point –
Freezing –
Freezing Point –
Deposition 16
The graph below shows the relationship between heat (energy) added, in calories (cal), and
temperature for 1 g of water. A student applied heat to 1 g of ice that had been cooled to -40⁰C and
measured the rise in temperature.
Read and fill-in the notes below and on the following pages and label the steps/regions A, B, C, D, E on
the graph.
Step A:Solid Water (Ice) Rises in Temperature (Keep in mind the graph is for
water!)

If the __________________ is not at 0oC, it will rise as heat is ____________to
get there. (Kinetic energy is _________________)

Each gram of water requires a constant amount of energy to increase 1o = specific
heat

IMPORTANT – the ice has not________________ yet!
17
Step B: Solid Water (Ice) Melts

By ______________energy the ice begins to _____________.

Temperature does not ___________ as more energy is being ______________
(Kinetic energy is _____________________ but potential energy is
____________)
Each mole of water requires a given amount of energy to melt = molar heat of fusion
(∆ Hfus) in kJ / mole.


Energy is overcoming water molecules attraction for each other so it can be converted
from a solid to liquid.

How many calories of energy did it take to completely change the 1 gram of solid
water (ice) at 0⁰C to liquid water?________________________
Step C: Liquid Water Rises in Temperature

Now the ice is completely _________ and the water temperature begins to
_________________ as heat is ________________. (specific heat)

Kinetic energy is ______________________.

The water has not started to____________ yet.

How many calories of energy did it take to make the 1 gram of liquid water to change
temperature from 0⁰C to 100⁰C (just beginning to boil)?____________
Step D: Liquid Water Boils

As we __________ energy the temperature does not change.

Each mole of water will require a constant amount of energy to boil = molar heat of
vaporization (∆Hvap) KJ/mole.

The energy is being used to overcome water's attraction to each other to convert the
liquid to a gas (kinetic energy _________________ but potential energy is
_________________).
How many calories of energy did it take to make the 1 gram of liquid water to
completely turn to steam once it hit 100⁰C?________________________

18
Step E: Steam Rises in Temperature

Temperature ___________ again when all water is turned to steam

Each gram of water requires a constant amount of energy to rise 1o = specific heat.
Specific Heat Capacity “C”
The amount of energy required to be absorbed to warm 1 gram of a substance by 1 oC (or
1 K) or the amount of energy required to be released to cool 1 gram of a substance by 1
o
C (or 1 K).
-orHow easily things warm up & cool down.
Energy Calculations Involving Specific Heat:
q = mC∆T
where:
q = Heat Energy
+ q means heat/energy is being absorbed (endothermic process)
- q means heat/energy is being released (exothermic process).
m = mass in grams
c = specific heat capacity (also “s”)
∆T = change in temperature (temperature final – temperature initial)
Energy Units:
Heat energy (q) is in joules(J), kilojoules (kJ) or calories (cal).
1 calorie = 4.184 joules
Mass (m) is in grams or kilograms
Specific heat capacity, c, is in J/g oC or kJ/kgoC
Water (L) = 4.184 J/goC Water (s) = 2.03 J/goC
Water (g) = 2.0 J/goC
Temperature , T, is usually in oC (temperature can be in K)
19
Metals have low specific heat values
Aluminum
0.900 J/goC
Iron
0.450 J/goC
Gold
0.126 J/goC
Doesn’t take much heat to heat them up and they don’t hold the heat well!!! (better
conductors of heat/energy)
Water and organic materials hold heat much better – have higher specific heats also
takes more energy to heat them up. (better insulators of heat/energy)
Water = 4.184 J/goC
Substance
J/goC cal/g K or Molar C
Wood = 1.76 J/goC
Btu/lb F
J/mol K
Aluminum
0.900
0.215
24.3
Bismuth
0.123
0.0294
25.7
Copper
0.386
0.0923
24.5
Brass
0.380
0.092
...
Gold
0.126
0.0301
25.6
Lead
0.128
0.0305
26.4
Silver
0.233
0.0558
24.9
Tungsten
0.134
0.0321
24.8
Zinc
0.387
0.0925
25.2
Mercury
0.140
0.033
28.3
2.4
0.58
111
Water
4.184
1.00
75.2
Ice (-10 C)
2.05
0.49
36.9
Granite
.790
0.19
...
Glass
.84
0.20
...
Iron
.450
Alcohol(ethyl)
20
Name _____________________________________
Energy & Specific Heat Problems
1. How much heat energy does a copper sample absorb if its specific heat is 0.386 J/g oC, its
mass is 12.5 g and it is heated from 25.0 oC to 40.0 oC?
2. How much heat energy is released by 10.0 g of gold, when it is cooled from 35.0 oC to 25.0
o
C? The specific heat of gold is 0.129 J/g oC.
3. A 4.00 kg sample of iron was heated from 0.0 oC to 20.0 oC. It absorbed 35.2 kJ of energy
as heat. What is the specific heat of this piece of iron?
4. 42.6 J of energy is needed to heat 2.00 grams of carbon from 50.0 oC to what final
temperature? The specific heat of carbon is 0.790 J/g oC.
21
Name _____________________________________________
Energy & Specific Heat Problems2
1. What amount of heat is required to raise the temperature of 85.9 g of water by 7.0C?
2. When 1045 joules are absorbed by a certain mass of water, the temperature of the water
increases from 45.0 ºC to 50.0 ºC. What is the mass of the water sample?
3. How many joules are required to heat 38.0 grams of gold from 60.0 ºC to 260.0 ºC? The
specific heat of gold is 0.126 J/(g·ºC).
4. Iron has a specific heat of 0.450 J/(g·ºC). If 1400. joules are absorbed by a chunk of iron
that weighs 40.0 grams, how much does the temperature of the iron increase?
22
Name _____________________________________
More Energy & Specific Heat Problems
**Pay
attention to units AND sig figs**
1. What is the specific heat value of a sample of unknown material, if it weighs 36.359 grams
and 59.912 J of heat raise its temperature 152.0 oC?
2. What would be the final temperature of a 73.174 g sample of cobalt with an initial
temperature of 102.0 oC, after it loses 800 J? (The specific heat of cobalt is 0.4210
J/goC)
3. What mass of iron would release 0.1854 kJ when its temperature changed from 1550.0 oC
to 75.0 oC? (The specific heat of iron is 0.450 J/g oC)
4. The specific heat of mercury is 0.0335 cal/g oC. If 152.00 g of mercury at 75.0 oC are
cooled to 23.5 oC, what is the value of q in Joules?
5. Kelly has 2.00 kg of water at 80.0 oC and wants it to cool to 45 oC. If the water releases
20.9 kJ of energy every minute, how long will it take to cool?
23
Calorimetry
Students made a calorimeter from 2 foam cups,
a lid, and a thermometer. One cup was nested inside the
other cup. The thermometer was fitted through a hole in
the lid (see Figure 1). Liquids could be added by
temporarily lifting up a flap in the lid.
Experiment 3
The procedure for Experiment 1 was repeated
except that, for each trial, a different mass of solid H2O
(ice cubes from a freezer), instead of RT H2O, was added
to the boiling H2O in the calorimeter (see Table 3).
Table 3
Trial
11
12
13
14
Experiment 1
The students placed 100 g of boiling H2O in the
calorimeter. Then 20 g of H2O at room temperature, RT
(which was a constant 23⁰C), was added to the
calorimeter. The calorimeter was swirled until the
temperature of its contents became constant, and this
final temperature, Tf, was recorded. The procedure was
repeated except that, for each trial, a different mass of
RT H2O was added to the 100 g of boiling H2O in the
calorimeter (see Table 1).
Mass of RT
H2O added (g)
1
2
3
4
5
20
40
60
80
100
2.
In which of the following trials was the final
temperature in the calorimeter highest?
A. Trial 1
B. Trial 5
C. Trial 10
D. Trial 11
3.
Which of the following statements best
describes what happened after the second
addition of H2O to the calorimeter in Trial 4?
The H2O added in the second addition was:
A. Hotter than the H2O in the calorimeter,
so the temperature in the calorimeter
decreased to 66⁰C.
B. Hotter than the H2O in the calorimeter,
so the temperature in the calorimeter
increased to 66⁰C.
C. Cooler that the H2O in the calorimeter,
so the temperature in the calorimeter
decreased to 66⁰C.
D. Cooler than the H2O in the calorimeter,
so the temperature in the calorimeter
increased to 66⁰C.
4.
If an additional trial had been done in
Experiment 2 in which Tf was 76⁰C, Ta would
most likely have been:
A. Less than 45⁰C.
B. Between 45⁰C and 58⁰C.
C. Between 71⁰C and 88⁰C.
D. Greater than 88⁰C.
Experiment 2
The procedure for Trial 5 was repeated except
that, for each trial, the 100 g of H2O added to the 100 g
of boiling H2O was at a different temperature, Ta (see
Table 2).
Trial
Ta (⁰ C)
6
7
8
9
10
34
45
58
71
88
Tf
(⁰ C)
67
73
79
86
94
91
82
70
46
In which trial, if any, did the first amount of H2O
placed in the calorimeter have the same initial
temperature as did the second amount of H2O
added to the calorimeter?
A. Trial 4
B. Trial 8
C. Trial 11
D. None of the trials
Tf
(⁰ C)
87
78
71
66
62
Table 2
Tf (⁰ C)
1.
Table 1
Trial
Mass of ice
added (g)
5
11
20
42
24
5.
Which of the following graphs best shows how
Tf varied with the mass of RT H2O that was
added in Experiment 1?
A.
C.
B.
D.
6.
Which of the following statements gives the
most likely reason that 2 foam cups were used
to make the calorimeter instead of just 1 cup?
Using the second cup:
A. Kept the H2O from leaking out of the
hole made in the inner cup to hold the
thermometer.
B. Kept the 100⁰C H2O from contacting
the added RT H2O.
C. Decreased the rate at which heat was
lost to the atmosphere.
D. Increased the rate at which heat was
lost to the atmosphere.
7. In which experiment, if any, was H2O present in
3 different states (phases) within the calorimeter?
A. Experiment 1
B. Experiment 2
C. Experiment 3
D. None of the experiments
25
Calorimetry
From the point of view of the system
Endothermic
Exothermic
Feels cold
Feels hot
Surroundings lose
heat (energy)
Surroundings gain
heat (energy)
System gains energy
System loses energy
(+) Energy term
Energy is absorbed
(-) Energy term
Energy is released
Measured in Joules
Measured in Joules
To convert between Joules and Calories:
1 calorie = 4.184 Joules
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Calorimeter
Q water = -Q system
Mass
H2O
x CH2O x ∆TH2O = Mass
sys
x Csys x ∆Tsys
(Tf-Ti)
(Tf-Ti)
Tf is the same value for both
Mass
H2O
x CH2O x ∆TH2O = Mass
sys
x Csys x ∆Tsys
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Name:________________________________
Calorimetry Problems
1) A 2.80 kg sample of metal with a specific heat of 0.43 J/g°C is heated to 100.0°C and then placed
in a 50.0g sample of water at 30.0°C. What is the final temperature of the water and the metal?
2) The specific heat of mercury is 0.033 cal/g°C. If 152g of mercury at 75.0°C is placed in 145g of
water at 23.5°C, what will be the final temperature of the water?
3) A 37.7 g piece of metal is heated to 100.0C and placed into 75.0 g of water in a coffee-
cup calorimeter. Initially, the temperature of the water in the calorimeter was 23.1C.
After the metal was added to the water the temperature of the water increased until its
temperature and the temperature of the metal were 30.6C.
a. What is the specific heat of the metal?
b. What kind of metal was added to the water in the calorimeter?
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4) A 440.00 g sample of mercury (specific heat = 0.140 J/goC, initial temperature of 22.00oC)
is placed into 134.00 g of water (initial temperature of 35.00oC). Find the final
temperature of the system.
5) Abbey is testing her baby’s bath water and finds that it is too cool, so she adds some hot water
from kettle on the stove. If Abbey adds 2.00 kg of water at 80.0°C to 20.0 kg of water at 27.0°C,
what is the final temperature of the bath water?
6) Jason is emptying the dishwasher. He removes a 0.200 kg glass that has a temperature of 30.0°C.
Into the glass, he pours 0.100 kg of diet soda (mostly water) which comes out of the refrigerator
with a temperature of 5.00°C. Assuming no external heat loss, what will be the final equilibrium
temperature of the glass of diet soda (no ice was added)? (c for glass =0.84 J/g°C).
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Name ___________________________________________________________________
More Calorimetry Problems (1 cal = 4.184 J)
1) 45.3 g of a shiny metal, with a specific heat of 0.561 cal/g ⁰C, is placed into a water bath that has a temperature of
99.7 ⁰C. It is then placed into a calorimeter that has 54.7 mL of water. If the water and the metal end up with a
temperature of 48.0⁰C, what was the initial temperature of the water in the calorimeter?
2) A metal with a mass of 97.4 g is heated to a temperature of 81.4 ⁰C. It is then placed into a calorimeter containing
0.246 kg of benzene, which has a specific heat of 1.74 J/g ⁰C. The temperature of the benzene rises from 15.5 ⁰C
to 32.5 ⁰C. What is the specific heat of the metal in calories?
3) A metal with a specific heat of 0.126 cal/g ⁰C is placed into a water bath with a temperature of 94.5 ⁰C. The metal
is then placed into a calorimeter containing 86.5 g of acetic acid at a temperature of 20.6 ⁰C. The acetic acid and
metal have a final temperature of 35.5 ⁰C. The acetic acid has a specific heat of 2.05 J/g ⁰C. What is the mass of
the metal?
4) A metal with a specific heat of 2.03 J/g ⁰C and a mass of 68.5 g is placed into a hot water bath with a temperature
of 74.5 ⁰C. The metal is then placed into a calorimeter containing acetic acid at a temperature of 14.5 ⁰C. The final
temperature of the acetic acid and metal is 45.5 ⁰C. The density of acetic acid is 1.04 g/mL and a specific heat of
0.49 cal/g ⁰C. What is the volume of acetic acid in the calorimeter?
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Distillation Apparatus:
Separating a mixture of liquids based upon the boiling point of each liquid.
Using the graphed data and the table of compounds and their boiling point temperatures, which
compound(s) is definitely not in the mixture?_______________________________
Which compound(s) are definitely in the mixture?_____________________________
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Distillation
To separate a mixture of liquids, the liquid can be heated
to force components, which have different boiling points,
into the gas phase. The gas is then condensed back into
liquid form and collected.
impure liquid
(mixture)
Distilled liquid
(distillate)
Separation of Unknown with
Fractional Distillation
120
Temperature (⁰C)
100
Solvent Boiling Point °C
Acetone = 56.5
Methanol = 64.7
Hexane = 68.8
Ethyl Acetate = 77
2-Methyl-2-propanol = 82.2
Water = 100.0
Toluene = 110.6
1-Butanol = 117.2
80
60
40
20
0
0
1
2
3
4
5
6
7
8
9
10 11 12 13 14 15 16 17 18 19 20
Time (min.)
Identify the material present in the mixture that has the lower boiling
point.________________________
Identify the material present in the mixture that has the higher boiling
point._______________________
What material(s) could still be present in the mixture that is still being heated?
____________________
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Distillation Results
Temp. ( °C)
Solvent Boiling Point °C
Acetone = 56.5
Methanol = 64.7
Hexane = 68.8
Ethyl Acetate = 77
2-Methyl-2-propanol = 82.2
Water = 100.0
Toluene = 110.6
1-Butanol = 117.2
Time (min.)
Based on the above graph, how many different components are definitely in this mixture?? ___________
What is the lowest distillation temperature on this graph? ____________________
What material boils at this temperature? ______________________
What is temperature when the next component begins to distill? _____________________
Identify this material based on the given table _________________________________
At what temperature does the last component boil? _______________________
This component‘s chemical name is _____________________________________
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