Measuring




Volume
Temperature
Mass
Distance
Reading the Meniscus
Always read volume from
the bottom of the
meniscus. The meniscus
is the curved surface of
a liquid in a narrow
cylindrical container.
Try to avoid parallax errors.
Parallax errors arise when a meniscus or needle is
viewed from an angle rather than from straight-on
at eye level.
Incorrect: viewing the
meniscus
from an angle
Correct: Viewing the
meniscus
at eye level
Graduated Cylinders
The glass cylinder
has etched marks to
indicate volumes, a
pouring lip, and quite
often, a plastic
bumper to prevent
breakage.
Measuring Volume
 Determine the volume contained in a
graduated cylinder by reading the bottom
of the meniscus at eye level.
 Read the volume using all certain digits
and one uncertain digit.
 Certain digits are determined
from the calibration marks on
the cylinder.
The uncertain digit (the last digit
of the reading) is estimated.
Use the graduations to find all certain
digits
There are two
unlabeled graduations
below the meniscus,
and each graduation
represents 1 mL, so
the certain digits of
the reading are…
52 mL.
Estimate the uncertain digit and take
a reading
The meniscus is
about eight tenths
of the way to the
next graduation, so
the final digit in
the reading is 0.8 mL
.
The volume in the graduated cylinder is 52.8 mL.
10 mL Graduate
What is the volume of liquid in the graduate?
6
_ .6
_ 2
_ mL
25mL graduated cylinder
What is the volume of liquid in the graduate?
1
5 mL
_1
_ . _
100mL graduated cylinder
What is the volume of liquid in the graduate?
5
7 mL
_2
_ . _
Self Test
Examine the meniscus below and determine
the volume of liquid contained in the
graduated cylinder.
The cylinder contains:
7
_6
_ . 0
_ mL
The Thermometer
o Determine the
temperature by reading
the scale on the
thermometer at eye
level.
o Read the temperature
by using all certain
digits and one uncertain
digit.
o Certain digits are determined from the calibration
marks on the thermometer.
o The uncertain digit (the last digit of the reading) is
estimated.
o On most thermometers encountered in a general
chemistry lab, the tenths place is the uncertain digit.
Do not allow the tip to touch the
walls or the bottom of the flask.
If the thermometer bulb
touches the flask, the
temperature of the glass
will be measured instead
of the temperature of
the solution. Readings
may be incorrect,
particularly if the flask
is on a hotplate or in an
ice bath.
Reading the Thermometer
Determine the readings as shown below on
Celsius thermometers:
8 _
7. _
4 C
_
3
0 C
_5
_ . _
Measuring Mass – Electronic balance
Our balances read to the hundredths place,the
uncertain digit is the hundredths place ( _ _ _ . _
X)
Balance Rules
In order to protect the balances and ensure accurate
results, a number of rules should be followed:
 Always check that the machine is zeroed
before adding any substance on it. This is done
by pushing the ZERO button before weighing any
substance.
 Never weigh directly on the balance pan.
Always use a piece of weighing paper, weigh
boat, or beaker to protect it.
 Do not weigh hot or cold objects.
 Clean up any spills around the balance
immediately.
Determining Mass
1. Place
container
(weigh boat
or beaker)
on the pan
2. Zero the
balance by
hitting the
zero bottom
3. Place object
on pan
4. Read the
digits to the
last digit
Metric Ruler
• Rulers are used to measure distance
• Can be used to measure the volume of a
cube (L x W x H) length x width x height
• Determine the temperature by reading the
scale on the thermometer at eye level.
• Read the ruler by using all certain digits
and one uncertain digit.
Reading a Ruler
• Determine the readings as shown below on
centimeter ruler:
_ _ . _ cm
_ _ . _ cm
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The SI System of Measurement
The Nature of Measurement
A Measurement is a quantitative
observation consisting of TWO parts
• Part 1 - number
• Part 2 - scale (unit)
•
Examples:
• 20 grams
• 6.63 x 10-34 Joule·seconds
The Fundamental SI Units
(le Système International, SI)
Base Quantity
Name of unit
Symbol
Length
Meter
m
Mass
Kilogram
kg
Time
Second
s
Temperature
Kelvin
K
Amount of Substance
Mole
mol
SI Prefixes
Common to Chemistry
Prefix
Kilo
Deci
Centi
Milli
Micro
Unit Abbr.
k
d
c
m

Exponent
103
10-1
10-2
10-3
10-6
Metric Acronym
•
•
•
•
•
•
•
K
H
D
M
D
C
M
King
Henry
Died
Monday
Drinking
Chocolate
Milk
Metric Conversions
103
kilo
102
101
hecto deka
g
m
L
Base
unit
10-1
deci
10-2
centi
10-3
milli
Conversions in the metric system are merely a
matter of moving a decimal point. The “base
unit” means the you have a quantity (grams,
meters, Liters, etc without a prefix.
Metric Conversions
103
kilo
102
101
hecto deka
g
m
L
Base
unit
18 L
10-1
10-2
deci
1
10-3
centi
2
milli
3
18 liters = 18 000 milliliters
Example #1: Convert 18 liters to milliliters
Metric Conversions
103
kilo
102
101
hecto deka
450 mg = 0.450 g
g
m
L
10-1
Base
unit
10-2
deci
3
10-3
centi
2
milli
1
Example #2: Convert 450 milligrams to grams
450 mg
Blank
Metric Conversions
Ladder Method
T. Trimpe 2008 http://sciencespot.net/
Ladder Method
1
2
KILO
1000
Units
3
HECTO
100
Units
DEKA
10
Units
DECI
0.1
Unit
Meters
Liters
Grams
How do you use the “ladder” method?
CENTI
0.01
Unit
MILLI
0.001
Unit
4 km = _________ m
1st – Determine your starting point.
Starting Point
2nd – Count the “jumps” to your ending point.
How many jumps does it take?
3rd – Move the decimal the same number of
jumps in the same direction.
Ending Point
4. __. __. __. = 4000 m
1
2
3
Conversion Practice
Try these conversions using the ladder method.
1000 mg = _______ g
1 L = _______ mL
160 cm = _______ mm
14 km = _______ m
109 g = _______ kg
250 m = _______ km
Compare using <, >, or =.
56 cm
6m
7g
698 mg
Metric Conversion Challenge
Write the correct abbreviation for each metric unit.
1) Kilogram _____
4) Milliliter _____
7) Kilometer _____
2) Meter _____
5) Millimeter _____
8) Centimeter _____
3) Gram _____
6) Liter _____
9) Milligram _____
Try these conversions, using the ladder method.
10) 2000 mg = _______ g
15) 5 L = _______ mL
20) 16 cm = _______ mm
11) 104 km = _______ m
16) 198 g = _______ kg
21) 2500 m = _______ km
12) 480 cm = _____ m
17) 75 mL = _____ L
22) 65 g = _____ mg
13) 5.6 kg = _____ g
18) 50 cm = _____ m
23) 6.3 cm = _____ mm
14) 8 mm = _____ cm
19) 5.6 m = _____ cm
24) 120 mg = _____ g
Compare using <, >, or =.
25) 63 cm
26) 536 cm
6m
53.6 dm
27) 5 g
28) 43 mg
508 mg
5g
29) 1,500 mL
30) 3.6 m
1.5 L
36 cm
Blank
Metric Dimensional
Analysis Practice
Problem #1
Convert 400 mL to Liters
400 mL
1
L
1 000 mL
= .400 L
= 0.4 L
= 4x10-1 L
Problem #2
Convert 10 meters to mm
10 m
1 000 mm
1
m
= 10 000 mm
= 1x104 mm
Problem #3
Convert 73 grams to kg
73 g
1 kg
1 000 g
= 0.073 kg
= 7.3x10-2 kg
Problem #4
Convert 0.02 kilometers to m
0.02 km 1 000 m
1
km
= 20 m
= 2x101 m
Problem #5
Convert 20 centimeters to m
20 cm
1 m
100 cm
= 0.20 m
= 2x10-1 m
Problem #6
Convert 450 milliliters to dL
450 mL
1 dL
100 mL
= 4.5 dL
Problem #7
Convert 10 kilograms to grams
10 kg
1 000 g
1
kg
= 10 000 g
= 1x104 g
Problem #8
Convert 935 mg to cg
935 mg
1
cg
10 mg
= 93.5 cg
= 9.35x101 cg
Problem #9
Convert 5.2 kg to mg
5.2 kg
1 000 000 mg
1 kg
= mg
= 5 200 000 mg
= 5.2x106 mg
Problem #10
Convert 175 mL to cm3
175 mL
1
cm3
1 mL
3
175
cm
=
= 1.75x102 cm3
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Uncertainty and Significant Figures
Cartoon courtesy of Lab-initio.com
Uncertainty in Measurement
•
A digit that must be estimated is
called uncertain. A measurement
always has some degree of
uncertainty.
Why Is there Uncertainty?
 Measurements are performed with
instruments
 No instrument can read to an infinite number
of decimal places
Which of these balances has the greatest
uncertainty in measurement?
Precision and Accuracy
•
Accuracy refers to the agreement of a
particular value with the true value.
•
Precision refers to the degree of
agreement among several measurements
made in the same manner.
Neither
accurate nor
precise
Precise but not
accurate
Precise AND
accurate
Types of Error
•
Random Error (Indeterminate Error) measurement has an equal probability of being
high or low.
•
Systematic Error (Determinate Error) - Occurs in
the same direction each time (high or low), often
resulting from poor technique or incorrect
calibration.
Rules for Counting Significant Figures Details
• Nonzero integers always count as significant
figures.
• 3456 has
• 4 significant figures
Rules for Counting Significant Figures Details
•
•
Zeros
Leading zeros do not count as significant
figures.
• 0.0486 has
• 3 significant figures
Rules for Counting Significant Figures Details
•
•
Zeros
Captive zeros always count as significant
figures.
• 16.07 has
• 4 significant figures
Rules for Counting Significant Figures Details
•
•
Zeros
Trailing zeros are significant only if the
number contains a decimal point.
• 9.300 has
• 4 significant figures
Rules for Counting Significant Figures Details
•
Exact numbers have an infinite number
of significant figures.
• 1 inch = 2.54 cm, exactly
Sig Fig Practice #1
How many significant figures in each of the
following?
5 sig figs
1.0070 m 
17.10 kg 
4 sig figs
100,890 L 
5 sig figs
3.29 x 103 s 
3 sig figs
0.0054 cm 
2 sig figs
3,200,000 
2 sig figs
Rules for Significant Figures in Mathematical
Operations
•
Multiplication and Division: # sig figs in the
result equals the number in the least precise
measurement used in the calculation.
• 6.38 x 2.0 =
• 12.76  13 (2 sig figs)
Sig Fig Practice #2
Calculation
Calculator says:
Answer
3.24 m x 7.0 m
22.68 m2
100.0 g ÷ 23.7 cm3
4.219409283 g/cm3
0.02 cm x 2.371 cm
0.04742 cm2
0.05 cm2
710 m ÷ 3.0 s
236.6666667 m/s
240 m/s
1818.2 lb x 3.23 ft
5872.786 lb·ft
5870 lb·ft
1.030 g ÷ 2.87 mL
2.9561 g/mL
2.96 g/mL
23 m2
4.22 g/cm3
Rules for Significant Figures in
Mathematical Operations
•
Addition and Subtraction: The number of
decimal places in the result equals the number of
decimal places in the least precise measurement.
• 6.8 + 11.934 =
• 18.734  18.7 (3 sig figs)
Sig Fig Practice #3
Calculation
Calculator says:
Answer
3.24 m + 7.0 m
10.24 m
10.2 m
100.0 g - 23.73 g
76.27 g
76.3 g
0.02 cm + 2.371 cm
2.391 cm
2.39 cm
713.1 L - 3.872 L
709.228 L
709.2 L
1818.2 lb + 3.37 lb
1821.57 lb
1821.6 lb
2.030 mL - 1.870 mL
0.16 mL
0.160 mL
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Scientific
Notation
Scientific Notation
In science, we deal with some very
LARGE numbers:
1 mole = 602000000000000000000000
In science, we deal with some very
SMALL numbers:
Mass of an electron =
0.000000000000000000000000000000091 kg
Imagine the difficulty of calculating
the mass of 1 mole of electrons!
0.000000000000000000000000000000091 kg
x 602000000000000000000000
???????????????????????????????????
Scientific Notation:
A method of representing very large or
very small numbers in the form:
M x 10n
 M is a number between 1 and 9
 n is an integer
.
2 500 000 000
9 8 7 6 5 4 3 2 1
Step #1: Insert an understood decimal point
Step #2: Decide where the decimal must end
up so that one number is to its left
Step #3: Count how many places you bounce
the decimal point
Step #4: Re-write in the form M x 10n
2.5 x
9
10
The exponent is the
number of places we
moved the decimal.
0.0000579
1
2 3
4
5
Step #2: Decide where the decimal must end
up so that one number is to its left
Step #3: Count how many places you bounce
the decimal point
Step #4: Re-write in the form M x 10n
5.79 x
-5
10
The exponent is negative
because the number we
started with was less
than 1.
PERFORMING
CALCULATIONS
IN SCIENTIFIC
NOTATION
ADDITION AND SUBTRACTION
Review:
Scientific notation expresses a
number in the form:
M x
1  M  10
n
10
n is an
integer
4 x 106
6
+ 3 x 10
7 x 106
IF the exponents are
the same, we simply
add or subtract the
numbers in front and
bring the exponent
down unchanged.
106
6
10
4 x
- 3 x
6
1 x 10
The same holds true
for subtraction in
scientific notation.
106
4 x
+ 3 x 105
If the exponents are
NOT the same, we
must move a decimal
to make them the
same.
6
10
4.00 x
4.00 x
6
5
+ .30 x 10
+ 3.00 x 10
6
4.30 x 10
Move the
decimal on
the smaller
number!
6
10
A Problem for you…
-6
10
2.37 x
-4
+ 3.48 x 10
Solution…
-6
002.37 x 10
-4
+ 3.48 x 10
Solution…
-4
0.0237 x 10
-4
+ 3.48
x 10
-4
3.5037 x 10
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Density
Density
Density is an important intensive property,
which can be used to help determine the
identity of an unknown substance.
While the mass or the volume of a
substance will vary from sample to
sample, the density will remain the same
at a given temperature.
As you know, the density of a substance is
a measure of how much mass is present
in a given unit of volume.
Density is the measure of the
“compactness” of a material
 How close the atoms or molecules are to each
other
 More than “heaviness” - density includes how
much space an object takes up!!
 All substances have density including liquids,
solids, and gases
Density is the measure of the
“compactness” of a material
 How close the atoms or molecules are to each
other
 More than “heaviness” - density includes how
much space an object takes up!!
 All substances have density including liquids,
solids, and gases
Density
D = m/v
m = mass (g)
Which is the amount of matter in a
substance
DO NOT confuse mass with weight
Weight – is the force of gravity
exerted on an object
v = volume (ml) or (cm3)
Is measured by water displacement or
using a ruler
Question
Which weighs more?
50 kilograms of iron
Or
50 kilograms of feathers
Question
Which has a greater
density?
iron
Or
feathers
Water
• Waters density = 1.000 g/ml at 4oC
• Ice floats on water
Densities of Common Substances
If water has a density of
1.00 g/cm3 , which of
the following will float or
sink?
Factors that Affect Density
Temperature
• For most substances, as temperature increases
the volume increases and as a result the
density decreases.
Pressure
Dissolved solids in Liquids
2 Ways to Measure the Volume of a
Solid
1. Objects with regular sides calculate L x W x H
using a ruler
2. Water displacement. Place irregularly shaped
solid into a volumetric cylinder half filled with
water and place the solid into the cylinder and
subtract the volumes
Useful Conversion Factors
1 cm3 = 1 ml
1 dm3 = 1 L
Manipulating the Density Formula
𝑚
𝐷=
𝑣
Sample Problem 1
A student determines that a piece of an unknown
material has a mass of 5.854 g and a volume of 7.57
cm3. What is the density of the material, rounded to the
correct number of significant digits?
Sample Problem 2
Iron has a known density of 7.87 g/cm3. What would be
the mass of a 2.5 cm3 piece of iron?
Sample Problem 3
Mercury has a density of 13.5 g/cm3. How much space
would 50.0 g of mercury occupy?
Blank
Percent Error
Percent Error
Students often assume that each measurement that
they make in the laboratory is true and
accurate. Likewise, they often assume that the
values that they derive through experimentation are
very accurate. However, sources of error often
prevent students from being as accurate as they
would like. Percent error calculations are used to
determine how close to the true values, or how
accurate, their experimental values really are.
Percent Error
The value that the student comes up with is
usually called the observed value, or the
experimental value. A value that can be found in
reference tables is usually called the true value, or
the accepted value. The percent error can be
determined when the true value is compared to
the observed value according to the equation
below:
Sample Problem 1
A student measures the mass and volume of a piece of
copper in the laboratory and uses his data to calculate the
density o the metal. According to his results, the copper
has a density of 8.37 g/cm3. Curious about the accuracy of
his results, the student consults a reference table and finds
that the accepted value for the density of copper is 8.92
g/cm3. What would be the student's percent error?
Sample Problem 2
A student experimentally determines the specific heat of
water to be 4.29 J/g x Co. He then looks up the specific
heat of water on a reference table and finds that is 4.18 J/g
x Co. What is his percent error?
Sample Problem 3
A student takes an object with an accepted mass of 200.00
grams and masses it on his own balance. He records the
mass of the object as 196.5 g. What is his percent error?
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Classification of Matter
Composition of Matter Flowchart
MATTER
yes
MIXTURE
yes
Is the composition
uniform?
Homogeneous
Mixture
(solution)
no
Can it be physically
separated?
PURE SUBSTANCE
no
Heterogeneous
Mixture
yes
Can it be chemically
decomposed?
Compound
no
Element
Classifying Matter by Composition
Elements- simplest kind of matter, made
of one type of atom
An atom is the smallest unit of an element
that maintains the properties of that
element.
Cannot be broken down into simpler
substances by ordinary chemical means
Ex. gold, copper, oxygen (on the periodic
table)
Element
Classifying Matter by Composition
Compounds – matter composed of the atoms
of two or more elements chemically
bonded
Compounds can be broken down by chemical
methods
When they are broken down, the
components have completely different
properties than the compound.
Ex. Sugar, salt, water, carbon dioxide
Compound
Classifying Matter by Composition
A mixture is a blend of two or more kinds
of matter, each of which retains its own
identity and properties.
A mixture is mixed together physically.
Variable composition, often expressed by a
percent composition by mass or volume
(Ex. 5% salt and 95% water)
Mixture
Classifying Matter by Composition
Homogeneous Mixture– matter
with a uniform composition
Homogeneous mixtures are also
called solutions.
Ex. Salt water and Kool –aid
Homogeneous Mixture
Classifying Matter by Composition
A heterogeneous mixture is not
the same throughout (not
uniform).
Examples: M & M’s, Chocolate
chip cookie, gravel, soil, rocks
such as granite, blood, milk,
salad, ocean water, etc.
Heterogeneous Mixture
Classify It
copper wire, aluminum foil
Classify It
» EX: table salt (NaCl)
Classify It
Granite
Apple Juice
Classify It
Examples:
» magnesium
» Pizza
» Calcium chloride
» Orange juice
» Club soda
Classify It
Examples:
» magnesium
» pizza
» Calcium chloride
» Orange juice
» Club soda
element
hetero. mixture
compound
hetero. mixture
Homo. (solution)
States of Matter
Solid- matter that can not flow and has
definite volume and shape
Liquid- definite volume but no definite
shape and can flow
Gas- a substance without definite volume
or shape and can flow.
Plasma- a substance that is similar to a
gas, but loses electrons due to its high
temperature
States of Matter
Definite
Volume?
Definite Particle position
Shape? and movement
Solid
YES
YES
Liquid
YES
NO
Gas
NO
NO
Packed tightly,
vibrate about fixed
point
Close together, can
move past each other
- flow
Far apart, move
rapidly - flow
States of Matter
Separating Mixtures
Mixtures are separated by their
physical properties.
Primary methods of separating
mixtures are:
filtration
distillation
centrifuge
chromatography
Separating Mixtures
Filtration is a method used to
separate the components of
mixtures that contain an
insoluble solid and a liquid.
Example: sand and water
Filtration
Separating Mixtures
Distillation is a method of
separating substances in a
mixture by evaporation of a
liquid and subsequent
condensation of its vapor.
Example: desalination of salt
water
Separating Mixtures
Centrifuge
Used to separate solid-liquid
mixtures such as those in
blood. The centrifuge spins
rapidly and causes the solid
to settle to the bottom.
Ex. Separating blood
Separating Mixtures
Chromatography is a method of
separating mixtures that uses a
stationary phase and a mobile
phase. Paper chromatography
can be used to separate pigments
because they move at different
rates on the paper.
Chromatography
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Properties & Changes of
Matter
Properties of Matter
Physical Property- a property that can be
observed and measured without changing the
identity of the substance.
Examples? Mass, Density, Melting and Boiling
Points
Chemical Property-relates to a substance’s ability
to undergo changes that transform it into
different substances.
Examples? Reactivity, Toxicity, and Chemical
stability
Properties of Matter
Chemists use properties to identify and separate
matter. More than one property must be used
for identification.
Intensive Properties – do not depend on the amount
of matter present
Ex. Melting pt., boiling pt, density, conduct
electricity
Extensive Properties – depend on the amount of
matter present
Ex. Volume, mass
Changes in Matter
A physical change does not change
the composition or identity of the
substance.
• Examples?
• Boiled water is still water.
• All phase changes are physical
changes
Freeze
Condense
Melt
Evaporate
Solid
Liquid
Gas
Changes in Matter
Sublimation is a process in which a solid
changes directly to a gas without going
through the liquid phase.
Ex: dry ice  CO2
Deposition is a process in which a gas
changes directly to a liquid without going
through the liquid phase
Ex: liquid vapor to ice (frost on windshields)
Changes in Matter
A chemical change occurs when
one or more substances are
changed into new substances.
Reactants- substances that react
Products- substances that form
Products have NEW PROPERTIES
2H2 + O2 → 2H2O
reactants
product
Law of Conservation of mass – Also
known as Conservation of Matter.
Matter can be neither created nor
destroyed, though it can be
rearranged. Mass remains constant in
an ordinary chemical change.
Indications of chemical
change
1. Production of energy in the
forms of heat, light, sound,
or electricity
2. Production of a gas
3. Formation of a precipitate
4. A change in color
5. A change in odor
What is Energy?
Energy - is the ability to or capacity to do
work or to produce change
Conservation of Energy (1st Law of
Thermodynamics) - Energy can be neither
created nor destroyed; the energy of the
universe is constant.
Two types:
Kinetic energy – energy of motion
Potential energy – is the energy due to
position of object
Chemical Energy
Chemical energy - is a special kind of
potential energy
• Is the energy involved in chemical
reactions
Energy Changes
• Some changes in matter release energy.
• For example, the explosion that occurs
when hydrogen and oxygen react to form
water is a release of energy.
• Heat energy and light energy are
released as the reaction takes place.
Energy Changes
• A change in matter in which energy is absorbed
from the surroundings is an endothermic process
(heat enters).
•EXAMPLES: melting ice & boiling water
•When barium hydroxide reacts ammonium nitrate
are mixed the test-tube feels cold to touch
because energy has been absorbed
Energy Changes
• A change in matter in which energy is
released is an exothermic process (heat
exits).
• Examples: freezing water & condensation
• Burning of paper gives off heat to the
surroundings.
Blank
Phase Changes
Courtesy www.lab-initio.com
What do you notice about the temperature
during a phase change?
Water phase
changes
Phase Diagram
 Represents phases as a function of
temperature and pressure.
 Critical temperature: temperature above
which the vapor can not be liquefied.
 Critical pressure: pressure required to
liquefy AT the critical temperature.
 Critical point: critical temperature and
pressure (for water, Tc = 374°C and 218 atm).
Phase changes by Name
Water
Carbon dioxide
Phase
Diagram
for Carbon
dioxide
Phase
Diagram
for Carbon
Carbon
Intro to
Thermochemistry
Heat and
Temperature
Thermochemistry
Thermochemistry – is the study of the
transfer of energy as heat that
accompany chemical reactions and
physical changes
Heat and Temperature
Heat - is energy transferred between a
system and its surroundings
•
Heat absorbed or released is
measured by a calorimeter
Heat
• Energy that flows from something warm to
something cooler
• A hotter substance gives KE to a cooler one
• When heat is transferred (lost or gained), there
is a change in the energy within the substance
Questions
A. When you touch ice, heat is transferred
from
1) your hand to the ice
2) the ice to your hand
B. When you drink a hot cup of coffee,
heat
is transferred from
1) your mouth to the coffee
2) the coffee to your mouth
Answer
A. When you touch ice, heat is transferred
from
1) your hand to the ice
B. When you drink a hot cup of coffee, heat
is transferred from
2) the coffee to your mouth
Question
C. When you heat 200 g of water for 1 minute,
the water temperature rises from 10°C to
18°C.
400 g
200 g
If you heat 400 g of water at 10°C in the same
pan with the same amount of heat for 1 minute,
what would you expect the final temperature to
be?
1) 10 °C
2) 14°C
3) 18°C
Answer
2)14°C
Heating twice the mass of water using
the same amount of heat will raise the
temperature only half as much.
200 g
400 g
Temperature
Temperature – is a measure of the average kinetic
energy of particles
•
Measured in Kelvin
•
Common Temperatures you need to KNOW
Celsius Scale (0C)
– 00C water freezes, 1000C water boils
Fahrenheit Scale (0F)
– 320F water freezes, 2120F water boils
Kelvin Scale (K)
– 273 K water freezes, 373 K water boils
– ABSOLUTE Zero 0 K
Temperature
•
Measured in Kelvin
K = 273 + oC
Convert the following
K
oC
50
357
32
298
Some Equalities for Heat
Heat is measured in calories or joules
• 1 kcal = 1000 cal
• 1 calorie = 4.18 J (Joules)
• 1 kJ = 1000 J
Specific Heat
• Why do some foods stay hot longer than
others?
• Why is the beach sand hot, but the water is
cool on the same hot day?
Specific Heat
Different substances have different
capacities for storing energy
It may take 20 minutes to heat water to
75°C. However, the same mass of
aluminum might require 5 minutes and the
same amount of copper may take only 2
minutes to reach the same temperature.
Specific Heat Values
Specific heat - is the amount of heat needed
to raise the temperature of 1 g of a substance
by 1°C
cal/g°C
J/g°C
water
aluminum
copper
silver
gold
1.00
0.22
0.093
0.057
0.031
4.18
0.90
0.39
0.24
0.13
Questions
A. A substance with a large specific heat
1) heats up quickly
2) heats up slowly
B. When ocean water cools, the surrounding air
1) cools 2) warms
3) stays the same
C. Sand in the desert is hot in the day, and cool
at night. Sand must have a
1) high specific heat 2) low specific heat
Answers
A. A substance with a large specific heat
2) heats up slowly
B. When ocean water cools, the surrounding
air
2) warms
C. Sand in the desert is hot in the day, and cool
at night. Sand must have a
2) low specific heat
Measuring Heat
Requires
•
Grams of substance
•
Temperature change T
T = Tf - Ti
•
Specific heat of the substance
Calculating Heat
Energy lost
or gained
=
mass x temp. change x specific heat
q = m x ∆T x Cp
A few key ideas:
•
•
If a substance receives heat and
experiences an increase in temperature
then Q is a positive number and ∆T is a
positive number.
If a substance loses heat and experiences a
decrease in temperature then Q is a negative
number and ∆T is a negative number.
A few key ideas:
•
•
•
Q, heat energy, can be measured in either
Joules or calories. Just make sure that your
units for c are consistent with your units
for Q.
∆T, change in temperature, can be measured in
K, C, or F. Just make sure that your units for
c are consistent with your units for ∆T.
Always start a problem by listing the given
information (with units) and writing down the
specific heat capacity equation.
A few key ideas:
• You must ALWAYS show all work and make
sure you have consistent units on your final
answer.
• The First Law of Thermodynamics states that
if two substances exchange heat, the quantity
of heat gained by one substance is exactly
equal and opposite to the quantity of heat lost
by the other substance.
Problem
The element hydrogen has the highest specific
heat of all elements. At a temperature of 25C,
hydrogen’s specific heat capacity is 14300J/(kg ∙
K). If the temperature of a 0.34 kg sample of
hydrogen is to be raised by 25 K, how much heat
will have to be transferred to the hydrogen?
Problem
At 25C, radon’s specific heat capacity is 94J/(kg
∙K). If the temperature of a 0.34kg sample of
radon is to be raised by 25K, how much heat will
have to be transferred to the radon?
Problem
A 0.59 kg brass candlestick has an initial
temperature of 98.0C. If 21,100J of heat is
removed from the candlestick to lower its
temperature to 6.8C, what is the specific heat
capacity of brass?
Problem
A 0.38kg drinking glass is filled with a hot liquid.
The liquid transfers 7032J of heat to the glass.
If the temperature of the glass increases by
22K, what is the specific heat capacity of the
glass?
End Of Unit 1