The document discusses measurements and units used in chemistry. It begins by discussing measurements made by the Mars rover Spirit and the importance of accuracy and precision in measurements. It then discusses scientific notation and defines accuracy and precision when evaluating measurements. Significant figures and proper reporting of measurements are also covered. Finally, it discusses the International System of Units (SI units) including common units for length, volume, mass, temperature, energy and converting between units.

1. Chemistry
Toombs County High School

2. 3.1
 On January 4, 2004, the
Mars Exploration Rover
Spirit landed on Mars. Each
day of its mission, Spirit
recorded measurements for
analysis. In the chemistry
laboratory, you must strive
for accuracy and precision in
your measurements.

3. Using and Expressing
Measurements
How do measurements
relate to science?

4. A measurement is a quantity that
has both a number and a unit.
 Measurements are fundamental to
the experimental sciences. For that
reason, it is important to be able to
make measurements and to decide
whether a measurement is correct.

5.  In scientific notation, a
given number is written as
the product of two
numbers: a coefficient
and 10 raised to a power.
 The number of stars in a
galaxy is an example of an
estimate that should be
expressed in scientific
notation.

6. Accuracy, Precision, and
Error
How do you evaluate accuracy
and precision?

7. Accuracy and Precision
▪Accuracy is a measure of how close a
measurement comes to the actual or
true value of whatever is measured.
▪Precision is a measure of how close a
series of measurements are to one
another.

8. To evaluate the accuracy of a
measurement, the measured value must
be compared to the correct value. To
evaluate the precision of a
measurement, you must compare the
values of two or more repeated
measurements.

10. Determining Error
▪ The accepted value is the correct value based
on reliable references.
▪ The experimental value is the value measured
in the lab.
▪ The difference between the experimental value
and the accepted value is called the error.

11. ▪ The percent error is the absolute value of the error
divided by the accepted value, multiplied by 100%.

13.  Just because a measuring device works, you
cannot assume it is accurate. The scale below has
not been properly zeroed, so the reading obtained
for the person’s weight is inaccurate.

14. Significant Figures in
Measurements
Why must measurements be
reported to the correct number of
significant figures?

15. ▪ Suppose you estimate a weight that is between
2.4 lb and 2.5 lb to be 2.46 lb. The first two digits
(2 and 4) are known. The last digit (6) is an
estimate and involves some uncertainty. All
three digits convey useful information, however,
and are called significant figures.
▪ The significant figures in a measurement
include all of the digits that are known, plus a
last digit that is estimated.

16. Measurements must always be reported
to the correct number of significant
figures because calculated answers
often depend on the number of
significant figures in the values used in
the calculation.

17. 3.1

18. 3.1

22.  Significant Figures in Calculations
 How does the precision of a calculated answer
compare to the precision of the measurements used
to obtain it?

23.  In general, a calculated answer cannot be
more precise than the least precise
measurement from which it was calculated.
The calculated value must be rounded to
make it consistent with the
measurements from which it was
calculated.

24. Rounding
 To round a number, you must first
decide how many significant figures
your answer should have. The answer
depends on the given measurements
and on the mathematical process used
to arrive at the answer.

27. Addition and Subtraction
 The answer to an addition or
subtraction calculation should be
rounded to the same number of
decimal places (not digits) as the
measurement with the least number of
decimal places.

30. Multiplication and Division
▪ In calculations involving multiplication and
division, you need to round the answer to the
same number of significant figures as the
measurement with the least number of
significant figures.
▪ The position of the decimal point has nothing to
do with the rounding process when multiplying
and dividing measurements.

34.  In the signs shown here, the
distances are listed as numbers
with no units attached.
Without the units, it is
impossible to communicate the
measurement to others. When
you make a measurement, you
must assign the correct units to
the numerical value.

35.  Measuring with SI Units
 Which five SI base units do chemists commonly
use?

36.  All measurements depend on units that serve
as reference standards. The standards of
measurement used in science are those of the
metric system.
 The International System of Units
(abbreviated SI, after the French name, Le
Système International d’Unités) is a revised
version of the metric system.

37. The five SI base units commonly used by
chemists are the meter, the kilogram, the
kelvin, the second, and the mole.

38. Units and Quantities
 What metric units are commonly used
to measure length, volume, mass,
temperature and energy?

39.  Units of Length
 In SI, the basic unit of length, or linear measure, is the
meter (m). For very large or and very small lengths, it may
be more convenient to use a unit of length that has a
prefix.

40.  Common metric units of length include the
centimeter, meter, and kilometer.

41.  Units of Volume
 The SI unit of volume is the amount of space
occupied by a cube that is 1 m along each edge.
This volume is the cubic meter (m)3
. A more
convenient unit of volume for everyday use is the
liter, a non-SI unit.
 A liter (L) is the volume of a cube that is 10
centimeters (10 cm) along each edge (10 cm × 10
cm × 10 cm = 1000 cm3
= 1 L).

42.  Common metric units of volume include
the liter, milliliter, cubic centimeter, and
microliter.

43.  The volume of 20 drops of liquid from a
medicine dropper is approximately 1 mL.

44. A sugar cube has a volume of 1 cm3
. 1 mL
is the same as 1 cm3
.

45.  A gallon of milk has about twice the volume
of a 2-L bottle of soda.

46. Units of Mass
 The mass of an object is measured in
comparison to a standard mass of 1
kilogram (kg), which is the basic SI unit of
mass.
 A gram (g) is 1/1000 of a kilogram; the mass
of 1 cm3
of water at 4°C is 1 g.

47.  Common metric units of mass include
kilogram, gram, milligram, and
microgram.

48.  Weight is a force that measures the pull on a
given mass by gravity.
 The astronaut shown on the surface of the
moon weighs one sixth of what he weighs on
Earth.

49.  Units of Temperature
 Temperature is a measure of
how hot or cold an object is.
 Thermometers are used to
measure temperature.

50. ▪
Scientists commonly use two
equivalent units of temperature,
the degree Celsius and the
kelvin.

51. ▪ On the Celsius scale, the freezing point of
water is 0°C and the boiling point is 100°C.
▪ On the Kelvin scale, the freezing point of
water is 273.15 kelvins (K), and the boiling
point is 373.15 K.
▪ The zero point on the Kelvin scale, 0 K, or
absolute zero, is equal to −273.15 °C.

52.  Because one degree on the Celsius scale is
equivalent to one kelvin on the Kelvin scale,
converting from one temperature to another is
easy. You simply add or subtract 273, as shown in
the following equations.

53. 3.2
▪ Conversions Between the Celsius and Kelvin Scales

56.  Units of Energy
Energy is the capacity to do work
or to produce heat.
▪ The joule and the calorie are
common units of energy.

57.  The joule (J) is the SI unit of energy.
 One calorie (cal) is the quantity of heat that
raises the temperature of 1 g of pure water by
1°C.

58.  This house is equipped with solar panels. The
solar panels convert the radiant energy from
the sun into electrical energy that can be
used to heat water and power appliances.

60.  Because each country’s
currency compares differently
with the U.S. dollar, knowing
how to convert currency units
correctly is very important.
Conversion problems are
readily solved by a problem-
solving approach called
dimensional analysis.

61. Conversion Factors
What happens when a
measurement is multiplied by a
conversion factor?

62. ▪ A conversion factor is a ratio of equivalent
measurements.
▪ The ratios 100 cm/1 m and 1 m/100 cm are examples of
conversion factors.

63. When a measurement is multiplied by a
conversion factor, the numerical value is
generally changed, but the actual size of
the quantity measured remains the
same.

64.  The scale of the micrograph is in nanometers. Using the
relationship 109
nm = 1 m, you can write the following
conversion factors.

65.  Dimensional Analysis
 Why is dimensional analysis useful?

66.  Dimensional analysis is a way to analyze and
solve problems using the units, or
dimensions, of the measurements.
 Dimensional analysis provides you with an
alternative approach to problem solving.

71. Converting Between Units
 What types of problems are easily solved by
using dimensional analysis?

72. Problems in which a
measurement with one unit is
converted to an equivalent
measurement with another unit
are easily solved using
dimensional analysis.

75.  Multistep Problems
 When converting between units, it is often
necessary to use more than one conversion
factor. Sample problem 3.8 illustrates the
use of multiple conversion factors.

78.  Converting Complex Units
 Many common measurements are
expressed as a ratio of two units. If you use
dimensional analysis, converting these
complex units is just as easy as converting
single units. It will just take multiple steps
to arrive at an answer.

83.  If you think that these lily
pads float because they are
lightweight, you are only
partially correct. The ratio
of the mass of an object to
its volume can be used to
determine whether an
object floats or sinks in
water.

84. Determining Density
What determines the density of a
substance?

85. Density is the ratio of the mass of an
object to its volume.

86.  Each of these 10-g samples has a different volume
because the densities vary.

87.  Density is an intensive property that
depends only on the composition of a
substance, not on the size of the sample.

88. The density of corn
oil is less than the
density of corn
syrup. For that
reason, the oil
floats on top of the
syrup.

89. Density and Temperature
How does a change in
temperature affect density?

90. Experiments show that the volume of
most substances increases as the
temperature increases. Meanwhile, the
mass remains the same. Thus, the
density must change.
 The density of a substance generally
decreases as its temperature increases.

98.  You could measure the
amount of sand in a sand
sculpture by counting each
grain of sand, but it would
be much easier to weigh the
sand. You’ll discover how
chemists measure the
amount of a substance using
a unit called a mole, which
relates the number of
particles to the mass.
10.1

99. Measuring Matter
What are three methods for
measuring the amount of
something?
10.1

100. ▪ You often measure the amount of something by one
of three different methods—by count, by mass, and
by volume.
10.1

103. What Is a Mole?
How is Avogadro’s number related
to a mole of any substance?
10.1

104. A mole of any substance contains
Avogadro’s number of representative
particles, or 6.02 × 1023
representative
particles.
 The term representative particle refers to
the species present in a substance: usually
atoms, molecules, or formula units.
10.1

105.  Converting Number of Particles to Moles
 One mole (mol) of a substance is 6.02 × 1023
representative particles of that substance and is
the SI unit for measuring the amount of a
substance.
 The number of representative particles in a mole,
6.02 × 1023
, is called Avogadro’s number.
10.1

106. 10.1

110.  Converting Moles to Number of Particles
10.1

113. The Mass of a Mole of an Element
How is the atomic mass of an
element related to the molar mass
of an element?
10.1

114. The atomic mass of an element
expressed in grams is the mass of a
mole of the element.
 The mass of a mole of an element is its
molar mass.
10.1

115. One molar mass of carbon, sulfur,
mercury, and iron are shown.
10.1

116. 10.1

117. The Mass of a Mole of a Compound
 How is the mass of a mole of a
compound calculated?
10.1

118.  To calculate the molar mass of a
compound, find the number of grams
of each element in one mole of the
compound. Then add the masses of
the elements in the compound.
10.1

119. ▪ Substitute the unit grams for atomic mass units. Thus 1
mol of SO3 has a mass of 80.1 g.
10.1

120. ▪ Molar Masses of Glucose, Water, and
Paradichlorobenzene
10.1

125.  How can you guess the number of
jelly beans in a jar? You estimate the
size of a jelly bean and then estimate
the dimensions of the container to
obtain its volume. In a similar way,
chemists use the relationships
between the mole and quantities
such as mass, volume, and number of
particles to solve chemistry
problems.
10.2

126. The Mole–Mass Relationship
 How do you convert the mass of a
substance to the number of moles of
the substance?
10.2

127. ▪ Use the molar mass of an element or compound to
convert between the mass of a substance and the
moles of a substance.
10.2

134. The Mole–Volume Relationship
 What is the volume of a gas at STP?
10.2

135.  Avogadro’s hypothesis states that equal
volumes of gases at the same temperature
and pressure contain equal numbers of
particles.
10.2

136.  The volume of a gas varies with temperature
and pressure. Because of these variations,
the volume of a gas is usually measured at a
standard temperature and pressure.
 Standard temperature and pressure (STP)
means a temperature of 0°C and a pressure
of 101.3 kPa, or 1 atmosphere (atm).
10.2

137. ▪At STP, 1 mol or, 6.02 × 1023
representative particles, of any gas
occupies a volume of 22.4 L.
▪The quantity 22.4 L is called the molar
volume of a gas.
10.2

138.  Calculating Volume at STP
10.2

142.  Calculating Molar Mass from Density
10.2

147. 10.2

148. 10.2

149. 10.2

150. 10.2

152. Particles Move - KINETIC

153.  The skunk releases its spray!
Within seconds you smell
that all-too-familiar foul
odor. You will discover some
general characteristics of
gases that help explain how
odors travel through the air,
even on a windless day.
13.1

154. Kinetic Theory and a Model
for Gases
What are the three
assumptions of the kinetic
theory as it applies to gases?
13.1

155. The word kinetic refers to motion.
 The energy an object has because of
its motion is called kinetic energy.
 According to the kinetic theory, all
matter consists of tiny particles that
are in constant motion.
13.1

156. According to kinetic theory:
▪The particles in a gas are considered
to be small, hard spheres with an
insignificant volume.
▪The motion of the particles in a gas
is rapid, constant, and random.
▪All collisions between particles in a
gas are perfectly elastic.
13.1

157. ▪ Particles in a gas are in rapid, constant motion.
13.1

158. ▪ Gas particles travel in straight-line paths.
13.1

159. ▪ The gas fills the container.
13.1

160. Gas Pressure
 How does kinetic theory explain
gas pressure?
13.1

161.  Gas pressure results from the force exerted
by a gas per unit surface area of an object.
 An empty space with no particles and no pressure
is called a vacuum.
 Atmospheric pressure results from the collisions
of atoms and molecules in air with objects.
13.1

162. Gas pressure is the result of
simultaneous collisions of billions of
rapidly moving particles in a gas
with an object.
13.1

163.  A barometer is a device that is used to
measure atmospheric pressure.
13.1

164.  The SI unit of pressure is the pascal (Pa).
 One standard atmosphere (atm) is the
pressure required to support 760 mm of
mercury in a mercury barometer at 25°C.
13.1

170. Kinetic Energy and Temperature
What is the relationship between
the temperature in kelvins and the
average kinetic energy of
particles?
13.1

171. Average Kinetic Energy
 The particles in any collection of atoms or
molecules at a given temperature have a
wide range of kinetic energies. Most of the
particles have kinetic energies somewhere
in the middle of this range.
13.1

172. 13.1

173.  Absolute zero (0 K, or –273.15°C) is the
temperature at which the motion of
particles theoretically ceases.
 Particles would have no kinetic energy at
absolute zero.
 Absolute zero has never been produced in the
laboratory.
13.1

174. Average Kinetic Energy and
Kelvin Temperature
The Kelvin temperature of a
substance is directly proportional to
the average kinetic energy of the
particles of the substance.
13.1

175. ▪ In this vacuum chamber, scientists cooled sodium vapor
to nearly absolute zero.
13.1

176.  END OF SHOW

178.  In organized soccer, a ball
that is properly inflated
will rebound faster and
travel farther than a ball
that is under-inflated. If
the pressure is too high,
the ball may burst when it
is kicked. You will study
variables that affect the
pressure of a gas.

179.  Compressibility
 Why are gases easier to compress than solids or
liquids are?

180.  Compressibility is a measure of how much the
volume of matter decreases under pressure.
When a person collides with an inflated airbag,
the compression of the gas absorbs the energy
of the impact.

181.  Gases are easily compressed because of the
space between the particles in a gas.
▪ The distance between particles in a gas is
much greater than the distance between
particles in a liquid or solid.
▪ Under pressure, the particles in a gas are
forced closer together.

182. ▪ At room temperature, the distance between particles in
an enclosed gas is about 10 times the diameter of a
particle.

183. Factors Affecting Gas Pressure
 What are the three factors that affect
gas pressure?

184. The amount of gas, volume, and
temperature are factors that affect
gas pressure.

185. Four variables are generally used to
describe a gas. The variables and
their common units are
 pressure (P) in kilopascals
 volume (V) in liters
 temperature (T) in kelvins
 the number of moles (n).

186. Amount of Gas
 You can use kinetic theory to predict and
explain how gases will respond to a change
of conditions. If you inflate an air raft, for
example, the pressure inside the raft will
increase.

187.  Collisions of particles with the inside walls of
the raft result in the pressure that is exerted by
the enclosed gas. Increasing the number of
particles increases the number of collisions,
which is why the gas pressure increases.

188.  If the gas pressure increases until it exceeds the
strength of an enclosed, rigid container, the
container will burst.

189.  Aerosol Spray Paint

190. Volume
 You can raise the pressure exerted by a
contained gas by reducing its volume.
The more a gas is compressed, the
greater is the pressure that the gas
exerts inside the container.

191.  When the volume of the container is halved, the
pressure the gas exerts is doubled.

192. Temperature
 An increase in the temperature of an
enclosed gas causes an increase in its
pressure.
 As a gas is heated, the average kinetic
energy of the particles in the gas increases.
Faster-moving particles strike the walls of
their container with more energy.

193.  When the Kelvin temperature of the enclosed gas
doubles, the pressure of the enclosed gas doubles.

195.  Hot lava oozes and flows,
scorching everything in its
path, and occasionally
overrunning nearby houses.
When the lava cools, it
solidifies into rock. The
properties of liquids are related
to intermolecular interactions.
You will learn about some of
the properties of liquids.
13.2

196. A Model for Liquids
What factors determine the
physical properties of a liquid?
13.2

197. ▪ Substances that can flow are referred to as fluids. Both
liquids and gases are fluids.
13.2

198. The interplay between the disruptive
motions of particles in a liquid and the
attractions among the particles
determines the physical properties of
liquids.
13.2

199. Evaporation
What is the relationship between
evaporation and kinetic energy?
13.2

200.  The conversion of a liquid to a gas or vapor is
called vaporization.
 When such a conversion occurs at the surface
of a liquid that is not boiling, the process is
called evaporation.
13.2

201. 13.2
 In an open container, molecules that
evaporate can escape from the container.

202.  In a closed container, the molecules cannot
escape. They collect as a vapor above the
liquid. Some molecules condense back into a
liquid.
13.2

203. During evaporation, only those
molecules with a certain minimum
kinetic energy can escape from the
surface of the liquid.
13.2

204. Vapor Pressure
 When can a dynamic equilibrium exist
between a liquid and its vapor?
13.2

205. Vapor pressure is a measure
of the force exerted by a gas
above a liquid.
13.2

206.  In a system at constant vapor pressure, a
dynamic equilibrium exists between the
vapor and the liquid. The system is in
equilibrium because the rate of evaporation
of liquid equals the rate of condensation of
vapor.
13.2

207. Vapor Pressure and Temperature
Change
▪ An increase in the temperature of a contained
liquid increases the vapor pressure.
▪ The particles in the warmed liquid have
increased kinetic energy. As a result, more of
the particles will have the minimum kinetic
energy necessary to escape the surface of the
liquid.
13.2

208. 13.2

209.  Vapor Pressure Measurements
▪ The vapor pressure of a liquid can be
determined with a device called a
manometer.
13.2

210. ▪ Manometer
13.2

211.  Boiling Point
 Under what conditions does boiling occur?
13.2

212. When a liquid is heated to a
temperature at which particles
throughout the liquid have enough
kinetic energy to vaporize, the liquid
begins to boil.
13.2

213. The temperature at which the vapor
pressure of the liquid is just equal to the
external pressure on the liquid is the
boiling point (bp).
13.2

214.  Boiling Point and Pressure Changes
▪ Because a liquid boils when its vapor pressure
is equal to the external pressure, liquids don’t
always boil at the same temperature.
▪ At a lower external pressure, the boiling point
decreases.
▪ At a higher external pressure, the boiling point
increases.
13.2

215. ▪ Altitude and Boiling Point
13.2

216. 13.2

217.  Normal Boiling Point
 Because a liquid can have various boiling points
depending on pressure, the normal boiling point
is defined as the boiling point of a liquid at a
pressure of 101.3 kPa.
13.2

218. 13.2

219.  END OF SHOW

220.  In 1985, scientists
discovered a new form of
carbon. They called this
form of carbon
buckminsterfullerene, or
buckyball for short. You will
learn how the arrangement
of particles in solids
determines some general
properties of solids.
13.3

221. A Model for Solids
 How are the structure and properties
of solids related?
13.3

222.  The general properties of solids reflect
the orderly arrangement of their
particles and the fixed locations of
their particles.
13.3

223.  The melting point (mp) is the temperature at
which a solid changes into a liquid.
13.3

224.  Crystal Structure and Unit Cells
 What determines the shape of a crystal?
13.3

225. ▪ In a crystal, the particles are arranged in an orderly,
repeating, three-dimensional pattern called a crystal
lattice.
13.3

226. Allotropes
 Allotropes are two or more different
molecular forms of the same element
in the same physical state.
▪ Allotropes have different properties
because their structures are different.
▪ Only a few elements have allotropes.
13.3

227. ▪ Carbon Allotropes
13.3

228. Non-Crystalline Solids
 An amorphous solid lacks an ordered
internal structure.
▪ Rubber, plastic, asphalt, and glass are
amorphous solids.
▪ A glass is a transparent fusion product of
inorganic substances that have cooled to a
rigid state without crystallizing.
13.3

229.  END OF SHOW

230. SC .

231.  Familiar weather events can
remind you that water exists on
Earth as a liquid, a solid, and a
vapor. As water cycles through
the atmosphere, the oceans,
and Earth’s crust, it undergoes
repeated changes of state. You
will learn what conditions can
control the state of a substance.
13.4

232.  Sublimation
 When can sublimation occur?
13.4

233. The change of a substance from a
solid to a vapor without passing
through the liquid state is called
sublimation.
▪ Sublimation occurs in solids with vapor
pressures that exceed atmospheric pressure at
or near room temperature.
13.4

234.  When solid iodine is
heated, the crystals
sublime, going
directly from the solid
to the gaseous state.
When the vapor cools,
it goes directly from
the gaseous to the
solid state.
13.4

235.  Phase Diagrams
 How are the conditions at which phases are in
equilibrium represented on a phase diagram?
13.4

236. A phase diagram is a graph that gives
the conditions of temperature and
pressure at which a substance exists as
solid, liquid, and gas (vapor).
13.4

237. The conditions of pressure and
temperature at which two phases exist
in equilibrium are indicated on a phase
diagram by a line separating the phases.
13.4

238. 13.4

239.  The triple point
describes the only set of
conditions at which all
three phases can exist in
equilibrium with one
another.
13.4

240.  END OF SHOW