Honors Chemistry
Unit 1B
Matter, Properties, & Energy
 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)
 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?
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:
a
Chart 1: Ingredients for layered solution (you must use at least 20mL but each cup must use
different volume of water and each must differ by at least 10 mL).
Cup Number
Salt (g)
Warm water (mL)
Calculated Density
(g/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 sig figs, calculate the density for each cup and record the density in the last column of Chart
1.
4. Using the balance and 10 mL of each solution you just made, measure the density of each solution.
Record these values for density in Table 1 on the back of this sheet.
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 20 mL from each cup into the 100 mL graduated cylinder.
Observations:
Table 1: Data from the Layered Solution Activity
Cup
Number
Measured Density from
solutions made
(g/mL)
Projected order to fill
graduated cylinder
The best layers will win a prize!!!
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
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) Graph the data on the provided graph under the data table.
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.
9
Data:
Volume (mL)
Water
Mass (g)
Rubbing
Alcohol
5.00
Silicone
Oil
4.60
15.0
13.70
25.0
23.10
Slope determination of each line to calculate the density of each liquid: (You MUST show your work!
Remember to put units with your work and answer.) Circle your answer.
Water:
Rubbing alcohol:
Silicone oil:
Given your results, how will the liquids be layered if poured carefully into a graduated cylinder? Make
sketch show which liquid would form the bottom… will be in the middle, and … will be on top.
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(18.00)
= 159.62 + 90
= 249.62 g/mole
Calculate the Molar Mass for each of these compounds.
1.
KCl
2.
Li2SO4
3.
(NH4)2C2O4  H2O
4.
Potassium Hydroxide
5.
Copper (II) Bromide
6.
Magnesium Phosphate
7.
Trisilicon Heptoxide
12
Converting Moles to Grams
and
Converting Grams to Moles
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 in the space on the right. Report all answers to 2 decimal places.
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
Converting Moles to Molecules (particles)
And Back Again
1. 4.50 Moles = ? Molecules
2. 6.62 x 1024 Molecules = ? Moles
3. 91.20 Moles = ? Molecules
4. 3.01 x 1023 Molecules = ? Moles
5. 1345.9 Moles = ? Molecules
14
More Converting…
1.
2.
3.
4.
Iron 18.06x 1023 molecules = ? grams
CuSO4 1.20 x 1028 molecules = ? grams
MgCl2 380.84 grams = ? molecules
Pb(NO3)2
82.81 grams = ? molecules
Moles And Compounds Worksheet
Directions: Write the answer to the problem on the blank provided. The correct set-up
must be to the right of the problem.
______________1. .738 moles of Fe2O3 to grams
______________2. 50.5 g of FeBr3 to moles
______________3. 1.51 x 10
23
molecules of PbI2 to moles
______________4. .445 moles of CCl4 to molecules
______________5. .538 moles of Ce2(CO3)3 to grams
______________6. 150.4 g of Ce(CO3)2 to moles
______________7. 7.22 x 10
25
molecules of CuCl2 · 4 H2O to moles
______________8. 1.45 moles of Pb(C2H3O2)2 to grams
______________9. 1.22 x 10
24
molecules of CO2 to grams
_____________10. 19.3 grams of H2O to molecules
15
Heating and Cooling Curve Definitions
Specific Heat –
Solid –
LiquidGas –
Plasma –
Heating Curve Enthalpy of Fusion/ Molar heat of fusion–
Melting–
Melting Point –
Enthalpy 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!
Step B: Solid Water (Ice) Melts

By ______________energy the ice begins to _____________.
17


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
Wood = 1.76 J/goC
cal/g K or Molar C
Btu/lb F J/mol K
Substance
J/goC
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
...
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
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
24
Calorimeter
Q water = -Q system
Mass
H2O
x CH2O x ∆TH2O = Mass
sys
x Csys x ∆Tsys
Mass
H2O
x CH2O x ∆TH2O = Mass
sys
x Csys x ∆Tsys
25
Name:________________________________
Calorimetry Problems
1) A 2.8 kg sample of metal with a specific heat of 0.43 kJ/kg°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?
26
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).
27
Name ___________________________________________________________________
More Calorimetry Problems
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 17.2 ⁰C, 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?
28
Distillation Apparatus:
Separating a mixture of liquids based upon the boiling point of each liquid.
29
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?_____________________________
30