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Chemistry is the study of matter, its properties, composition, and structure, and the changes it undergoes.

1.
Chapter 1

Introduction: Matter and Measurement

Learning outcomes:

Distinguish between the macroscopic world and the submicroscopic realm of atoms

and molecules.

Categorize matter into its different states.

Differentiate among pure substances, homogeneous mixtures, and heterogeneous

mixtures.

Categorize the properties of pure substances.

Contrast the different techniques used to separate mixtures of substances.

Differentiate the different forms of energy.

Use the metric system to express scientific quantities quantitatively.

Differentiate between metric base units and derived units, such as volume, density,

and energy.

Determine the uncertainty of a given quantity in scientific calculations.

Analyze calculations to determine the correct number of significant figures for the

answer.

Apply dimensional analysis to track the conversion from given units to desired units.

• Chemistry is the study of matter, its properties,

composition, and structure and the changes it undergoes.

• It is central to our fundamental understanding of many

science-related fields.

1

Introduction: Matter and Measurement

Learning outcomes:

Distinguish between the macroscopic world and the submicroscopic realm of atoms

and molecules.

Categorize matter into its different states.

Differentiate among pure substances, homogeneous mixtures, and heterogeneous

mixtures.

Categorize the properties of pure substances.

Contrast the different techniques used to separate mixtures of substances.

Differentiate the different forms of energy.

Use the metric system to express scientific quantities quantitatively.

Differentiate between metric base units and derived units, such as volume, density,

and energy.

Determine the uncertainty of a given quantity in scientific calculations.

Analyze calculations to determine the correct number of significant figures for the

answer.

Apply dimensional analysis to track the conversion from given units to desired units.

• Chemistry is the study of matter, its properties,

composition, and structure and the changes it undergoes.

• It is central to our fundamental understanding of many

science-related fields.

1

2.
Atomic and Molecular Perspective

Matter – Anything that has

mass and occupies space.

Atom – The smallest stable

building block of matter.

Molecule – Groups of atoms

held together with a specific

connectivity and shape.

Composition - the types of atoms that are present in a

compound and the ratio of these atoms (for example H2O, C2H6O).

Structure - how atoms are connected (bonded) to each other,

how far apart they are, and the shape of the molecule.

Methods of Classification of Matter

State of Matter - physical state is gas, liquid, or solid.

Composition of Matter - element, compound, or mixture

States of Matter

1) Gas (vapor) – has no fixed

volume or shape, uniformly

expands to fill its container,

compressible, flows readily,

diffusion occurs rapidly.

2) Liquid - has a distinct

volume independent of its

container, assumes the shape

of the portion of the

container it occupies, not

significantly compressible,

diffusion occurs but slower

than a gas.

3) Solid - has both a definite

shape and definite volume,

not significantly

compressible, diffusion

occurs extremely slowly.

2

Matter – Anything that has

mass and occupies space.

Atom – The smallest stable

building block of matter.

Molecule – Groups of atoms

held together with a specific

connectivity and shape.

Composition - the types of atoms that are present in a

compound and the ratio of these atoms (for example H2O, C2H6O).

Structure - how atoms are connected (bonded) to each other,

how far apart they are, and the shape of the molecule.

Methods of Classification of Matter

State of Matter - physical state is gas, liquid, or solid.

Composition of Matter - element, compound, or mixture

States of Matter

1) Gas (vapor) – has no fixed

volume or shape, uniformly

expands to fill its container,

compressible, flows readily,

diffusion occurs rapidly.

2) Liquid - has a distinct

volume independent of its

container, assumes the shape

of the portion of the

container it occupies, not

significantly compressible,

diffusion occurs but slower

than a gas.

3) Solid - has both a definite

shape and definite volume,

not significantly

compressible, diffusion

occurs extremely slowly.

2

3.
Elements, Compounds & Mixtures

Pure Substance Matter that has a fixed composition and distinct properties. All

pure substances are either elements or compounds.

Elements All atoms are the same kind, elements have only one type of atom. e.g.

oxygen (O2), gold (Au), silicon (Si) and diamond (C).

Compounds Contains more than one type of atom, but all molecules (or repeat

units) are the same, e.g. water (H2O), ethanol (C2H6O), quartz (SiO2), sodium

chloride (NaCl).

Mixture Have variable composition and can be separated into component parts by

physical methods. Mixtures contain more than one kind of molecule, and their

properties depend on the relative amount of each component present in the mixture.

Homogeneous & Heterogeneous Mixtures

The composition is variable for both

heterogeneous and homogeneous mixtures.

Heterogeneous Mixture - non-uniform.

Chocolate Chip Cookie – Chocolate, dough, etc.

Concrete – Cement, rocks, etc.

Nachos – Chips, cheese, jalapeños, salsa, etc.

Homogeneous Mixture – uniform throughout,

also called a solution.

Air – principle components include O2, N2 & CO2

Vodka – principle components are ethanol and water

Brass – solid solution of Cu and Zn

Ruby – solid solution of Al2O3 and Cr2O3

3

Pure Substance Matter that has a fixed composition and distinct properties. All

pure substances are either elements or compounds.

Elements All atoms are the same kind, elements have only one type of atom. e.g.

oxygen (O2), gold (Au), silicon (Si) and diamond (C).

Compounds Contains more than one type of atom, but all molecules (or repeat

units) are the same, e.g. water (H2O), ethanol (C2H6O), quartz (SiO2), sodium

chloride (NaCl).

Mixture Have variable composition and can be separated into component parts by

physical methods. Mixtures contain more than one kind of molecule, and their

properties depend on the relative amount of each component present in the mixture.

Homogeneous & Heterogeneous Mixtures

The composition is variable for both

heterogeneous and homogeneous mixtures.

Heterogeneous Mixture - non-uniform.

Chocolate Chip Cookie – Chocolate, dough, etc.

Concrete – Cement, rocks, etc.

Nachos – Chips, cheese, jalapeños, salsa, etc.

Homogeneous Mixture – uniform throughout,

also called a solution.

Air – principle components include O2, N2 & CO2

Vodka – principle components are ethanol and water

Brass – solid solution of Cu and Zn

Ruby – solid solution of Al2O3 and Cr2O3

3

4.
Periodic Table Relative abundances of

elements in the Earth’s

crust and human body.

Elements are represented as symbols with one

or two letters; the first is always capitalized.

4

elements in the Earth’s

crust and human body.

Elements are represented as symbols with one

or two letters; the first is always capitalized.

4

5.
Compound Elements

Elements can interact with other elements to form compounds,

and compounds can be decomposed into elements.

Law of Constant Composition

The elemental composition of a compound is always

the same, which is known as the Law of Constant

Composition (or Law of Definite Proportions).

Example: According to the law of definite proportions,

if a sample of compound A contains 12 g of sulfur and

6 g of oxygen, another sample of A that contains 50

grams of sulfur must contain ________ g oxygen.

5

Elements can interact with other elements to form compounds,

and compounds can be decomposed into elements.

Law of Constant Composition

The elemental composition of a compound is always

the same, which is known as the Law of Constant

Composition (or Law of Definite Proportions).

Example: According to the law of definite proportions,

if a sample of compound A contains 12 g of sulfur and

6 g of oxygen, another sample of A that contains 50

grams of sulfur must contain ________ g oxygen.

5

6.
Chemical and Physical Properties

Physical Properties Some properties can be readily measured

with our senses, e.g. odor and color, instruments are needed to

measure other properties, such as electrical resistivity, hardness,

melting point, boiling point, density, mass, volume, etc.

Chemical Properties Describe the reactivity of a substance

toward other substances. Examples include:

Ethanol burns in air (reacts with oxygen)

Sodium reacts vigorously with water,

Corrosion of metal parts (rust),

Trinitrotoluene (TNT) is explosive.

Physical changes are changes in matter that

do not change the composition of a substance.

– Examples include changes of state,

temperature, and volume.

Chemical changes result in new substances.

– Examples include reactions: combustion,

oxidation, tarnishing, and decomposition.

Properties of Matter

• Intensive Properties:

□ Independent of the amount of the

substance that is present.

• Density, temperature, melting point,

boiling point, hardness, color, etc.

• Extensive Properties:

□ Dependent upon the amount of the

substance present.

• Mass, volume, energy, etc.

6

Physical Properties Some properties can be readily measured

with our senses, e.g. odor and color, instruments are needed to

measure other properties, such as electrical resistivity, hardness,

melting point, boiling point, density, mass, volume, etc.

Chemical Properties Describe the reactivity of a substance

toward other substances. Examples include:

Ethanol burns in air (reacts with oxygen)

Sodium reacts vigorously with water,

Corrosion of metal parts (rust),

Trinitrotoluene (TNT) is explosive.

Physical changes are changes in matter that

do not change the composition of a substance.

– Examples include changes of state,

temperature, and volume.

Chemical changes result in new substances.

– Examples include reactions: combustion,

oxidation, tarnishing, and decomposition.

Properties of Matter

• Intensive Properties:

□ Independent of the amount of the

substance that is present.

• Density, temperature, melting point,

boiling point, hardness, color, etc.

• Extensive Properties:

□ Dependent upon the amount of the

substance present.

• Mass, volume, energy, etc.

6

7.
Separation of Mixtures

Filtration Distillation

Chromatography

Energy

• Energy is the capacity to

do work or transfer heat.

• Work is the energy

transferred when a force

exerted on an object

causes a displacement of

that object.

• Heat is the energy used to

cause the temperature of

an object to increase.

• Force is any push or pull

on an object.

7

Filtration Distillation

Chromatography

Energy

• Energy is the capacity to

do work or transfer heat.

• Work is the energy

transferred when a force

exerted on an object

causes a displacement of

that object.

• Heat is the energy used to

cause the temperature of

an object to increase.

• Force is any push or pull

on an object.

7

8.
Fundamental Forms of Energy

Kinetic energy is

the energy of motion.

– Its magnitude

depends on the

object’s mass

and its velocity:

KE = ½mv2

• Potential energy of an object depends on its relative position

compared to other objects.

• Potential energy also refers to the composition of an object,

including the chemical energy stored in chemical bonds.

One of the goals in chemistry is to related the energy changes in

the macroscopic world to the kinetic or potential energy of

substances at the molecular level.

Numbers and Units in Chemistry

Major role in quantifying:

– Units of measurement

– Quantities that are measured and calculated

– Uncertainty in measurement

– Significant figures

– Dimensional analysis

(e.g. 1 inch = 2.54 cm)

Energy is a derived unit, mass times velocity squared

Joule (J) = kg·m2/s2

8

Kinetic energy is

the energy of motion.

– Its magnitude

depends on the

object’s mass

and its velocity:

KE = ½mv2

• Potential energy of an object depends on its relative position

compared to other objects.

• Potential energy also refers to the composition of an object,

including the chemical energy stored in chemical bonds.

One of the goals in chemistry is to related the energy changes in

the macroscopic world to the kinetic or potential energy of

substances at the molecular level.

Numbers and Units in Chemistry

Major role in quantifying:

– Units of measurement

– Quantities that are measured and calculated

– Uncertainty in measurement

– Significant figures

– Dimensional analysis

(e.g. 1 inch = 2.54 cm)

Energy is a derived unit, mass times velocity squared

Joule (J) = kg·m2/s2

8

9.
Metric System Prefixes

Example

Example 1: How many grams are in 10 kg?

10 kg 1×103 g

× =10×103 g

1 kg

In scientific notation, 1.0x104 g.

Example 2: Use appropriate metric prefixes to write

the following measurement without use of exponents,

7.6×10−10 g.

Approach: Use the closest prefix to the exponent (-10),

choose 1 ng = 1×10−9 g.

7.6×10−10 g 1 ng

× = 0.76 ng

1×10−9 g

9

Example

Example 1: How many grams are in 10 kg?

10 kg 1×103 g

× =10×103 g

1 kg

In scientific notation, 1.0x104 g.

Example 2: Use appropriate metric prefixes to write

the following measurement without use of exponents,

7.6×10−10 g.

Approach: Use the closest prefix to the exponent (-10),

choose 1 ng = 1×10−9 g.

7.6×10−10 g 1 ng

× = 0.76 ng

1×10−9 g

9

10.
Volume

Volume is not a base unit for

SI; it is a derived unit from

length (m × m × m = m3).

The most commonly used

metric units for volume are the

liter (L) and the milliliter (mL).

– A liter is a cube 1

decimeter (dm) long on

each side.

– A milliliter (mL) is a cube

1 centimeter (cm) long on

each side, also called 1

cubic centimeter

(cm × cm × cm = cm3).

Glassware for Measuring Volume

Uncertainty in Measurements - Different measuring devices

have different uses and different degrees of precision.

10

Volume is not a base unit for

SI; it is a derived unit from

length (m × m × m = m3).

The most commonly used

metric units for volume are the

liter (L) and the milliliter (mL).

– A liter is a cube 1

decimeter (dm) long on

each side.

– A milliliter (mL) is a cube

1 centimeter (cm) long on

each side, also called 1

cubic centimeter

(cm × cm × cm = cm3).

Glassware for Measuring Volume

Uncertainty in Measurements - Different measuring devices

have different uses and different degrees of precision.

10

11.
Temperature Scales

Temperature – the “hotness and coldness” of an object.

Heat flows spontaneously from an object with a higher

temperature to an object with a lower temperature.

Temperature

• The Kelvin is the SI

unit of temperature.

• It is based on the

properties of gases.

• There are no

negative Kelvin

temperatures.

• K = C + 273.15

Example: Express 25.00 °C in K.

11

Temperature – the “hotness and coldness” of an object.

Heat flows spontaneously from an object with a higher

temperature to an object with a lower temperature.

Temperature

• The Kelvin is the SI

unit of temperature.

• It is based on the

properties of gases.

• There are no

negative Kelvin

temperatures.

• K = C + 273.15

Example: Express 25.00 °C in K.

11

12.
Temperature

•The Fahrenheit scale is not

used in scientific

measurements.

F = 9/5(C) + 32

C = 5/9(F − 32)

The ‘9/5’, ‘5/9’, and ’32’ are

exact numbers and do not

influence significant figures.

Examples:

Density

mass m Example: A piece of unknown metal with a

Density = d = =

volume V right rectangular prism shape has a width

of 3.2 cm, a length of 17.1 cm, and height

of 4.0 cm. Its mass is 1.5 kg. Calculate the

density of the metal in g/cm3.

12

•The Fahrenheit scale is not

used in scientific

measurements.

F = 9/5(C) + 32

C = 5/9(F − 32)

The ‘9/5’, ‘5/9’, and ’32’ are

exact numbers and do not

influence significant figures.

Examples:

Density

mass m Example: A piece of unknown metal with a

Density = d = =

volume V right rectangular prism shape has a width

of 3.2 cm, a length of 17.1 cm, and height

of 4.0 cm. Its mass is 1.5 kg. Calculate the

density of the metal in g/cm3.

12

13.
Example

The world’s largest gold bar in the Toi Gold Museum in

Japan can be seen and touched by visitors. It has an

irregular shape with dimensions about 17.9 in. by 8.9 in.

by 6.7 inches, with a volume of 12.94 L. Determine the

mass and weight (1 kg = 2.205 lb).

Numbers in Chemistry

• Exact numbers are counted or given by definition. For

example, there are 12 eggs in 1 dozen and 3 feet in 1 yard.

• Inexact (or measured) numbers depend on how they were

determined. Scientific instruments have limitations

(equipment errors) and individuals can read some

instrumentation differently (human errors).

Digital Reading Scale read by eye

The last digit measured is considered reliable, but not exact.

13

The world’s largest gold bar in the Toi Gold Museum in

Japan can be seen and touched by visitors. It has an

irregular shape with dimensions about 17.9 in. by 8.9 in.

by 6.7 inches, with a volume of 12.94 L. Determine the

mass and weight (1 kg = 2.205 lb).

Numbers in Chemistry

• Exact numbers are counted or given by definition. For

example, there are 12 eggs in 1 dozen and 3 feet in 1 yard.

• Inexact (or measured) numbers depend on how they were

determined. Scientific instruments have limitations

(equipment errors) and individuals can read some

instrumentation differently (human errors).

Digital Reading Scale read by eye

The last digit measured is considered reliable, but not exact.

13

14.
Precision and Accuracy

• Precision is a measure of how closely individual

measurements to agree with one another.

• Accuracy refers to how closely individual

measurements agree with the correct “true” value.

Significant Figures

• The term significant figures refers to digits

that were measured.

• When rounding calculated numbers, we pay

attention to significant figures so we do not

overstate the accuracy of our answers.

14

• Precision is a measure of how closely individual

measurements to agree with one another.

• Accuracy refers to how closely individual

measurements agree with the correct “true” value.

Significant Figures

• The term significant figures refers to digits

that were measured.

• When rounding calculated numbers, we pay

attention to significant figures so we do not

overstate the accuracy of our answers.

14

15.
Significant Figures

1. Zeros between non-zero numbers are always

significant.

2. Zeros at the beginning of a number are never

significant, merely indicate the position of the

decimal point.

3. Zeros at the end of the number after a decimal

place are significant if the number contains a

decimal point.

4. Zeros at the end of a number before a decimal place

are ambiguous (e.g. 23,800 g), unless a decimal

point is written at the end (i.e. 23,800. g). Assume

the zeros are insignificant, unless there is a decimal

point. Avoid ambiguity by using scientific notation.

Significant Figures

How many significant figures are present in

each of the measured quantities?

0.0012

108

900.0

3.0012

0.002070

4.80 10-3

4.800 10-3

15

1. Zeros between non-zero numbers are always

significant.

2. Zeros at the beginning of a number are never

significant, merely indicate the position of the

decimal point.

3. Zeros at the end of the number after a decimal

place are significant if the number contains a

decimal point.

4. Zeros at the end of a number before a decimal place

are ambiguous (e.g. 23,800 g), unless a decimal

point is written at the end (i.e. 23,800. g). Assume

the zeros are insignificant, unless there is a decimal

point. Avoid ambiguity by using scientific notation.

Significant Figures

How many significant figures are present in

each of the measured quantities?

0.0012

108

900.0

3.0012

0.002070

4.80 10-3

4.800 10-3

15

16.
Rounding*

After determining the appropriate number of

significant figures, round off your final answer.

1. If the first digit you drop is greater than 5, add 1

to the last digit you keep. You are rounding up.

2. If the first digit you drop is less than 5, do

nothing to the digits you keep. You are rounding

down.

3. If the digit you drop is 5, and there are no

following digits, round down. If there are digits

following the 5, round up.

*You may receive a different rule #3 from your lab instructor.

Significant Figures & Calculations

Addition and Subtraction

Line up the numbers at the decimal point and the

answer cannot have more decimal places than the

measurement with the fewest number of decimal

places.

16

After determining the appropriate number of

significant figures, round off your final answer.

1. If the first digit you drop is greater than 5, add 1

to the last digit you keep. You are rounding up.

2. If the first digit you drop is less than 5, do

nothing to the digits you keep. You are rounding

down.

3. If the digit you drop is 5, and there are no

following digits, round down. If there are digits

following the 5, round up.

*You may receive a different rule #3 from your lab instructor.

Significant Figures & Calculations

Addition and Subtraction

Line up the numbers at the decimal point and the

answer cannot have more decimal places than the

measurement with the fewest number of decimal

places.

16

17.
Addition and Subtraction

• The absolute uncertainty can be no

smaller than the least accurate number.

• 12.04

- 10.4

1.64 1.6

• The answer should have no more decimal

places than the least accurate number.

Multiplication and Division

The answer cannot have more significant

figures than the measurement with the

fewest number of significant figures.

3121 12 = 37452 = 3.7 104

# sig. digits 4 2 2

Know the number of appropriate digits throughout, round at the end.

17

• The absolute uncertainty can be no

smaller than the least accurate number.

• 12.04

- 10.4

1.64 1.6

• The answer should have no more decimal

places than the least accurate number.

Multiplication and Division

The answer cannot have more significant

figures than the measurement with the

fewest number of significant figures.

3121 12 = 37452 = 3.7 104

# sig. digits 4 2 2

Know the number of appropriate digits throughout, round at the end.

17

18.
Mixed Operations

Determine accuracy in the same order as the

mathematical operations, # of significant digits in blue

•Retain at least one additional digit past the significant

figures in combined operations, so rounding doesn’t

affect results…

-keep track of the proper significant figures to use at

the end.

3 3

m 2.79 g 2.79 g

d= = =

v 8.34 mL - 7.58 mL 0.76 mL

3 3 2

d = 3.7 g/mL

2

Evaluate the expression to the correct number of

significant figures. How many sig. figs. in the answer?

4.184 × 100.620 × (25.27 - 24.16) = _____

Using order of operations, evaluate subtraction first.

Then multiply each number and evaluate.

4.184 × 100.620 × (1.11) = 467.3034288

(4 s.f.) (6 s.f.) (3 s.f.) 467

Multiplication uses fewest number of sig. figures.

The answer should have three significant figures.

Retain at least one additional digit past the significant figures in

combined operations, so rounding doesn’t affect result…

-keep track of the proper significant figures for the final answer.

18

Determine accuracy in the same order as the

mathematical operations, # of significant digits in blue

•Retain at least one additional digit past the significant

figures in combined operations, so rounding doesn’t

affect results…

-keep track of the proper significant figures to use at

the end.

3 3

m 2.79 g 2.79 g

d= = =

v 8.34 mL - 7.58 mL 0.76 mL

3 3 2

d = 3.7 g/mL

2

Evaluate the expression to the correct number of

significant figures. How many sig. figs. in the answer?

4.184 × 100.620 × (25.27 - 24.16) = _____

Using order of operations, evaluate subtraction first.

Then multiply each number and evaluate.

4.184 × 100.620 × (1.11) = 467.3034288

(4 s.f.) (6 s.f.) (3 s.f.) 467

Multiplication uses fewest number of sig. figures.

The answer should have three significant figures.

Retain at least one additional digit past the significant figures in

combined operations, so rounding doesn’t affect result…

-keep track of the proper significant figures for the final answer.

18

19.
Evaluate the expression to the correct number of

significant figures. How many sig. figs. in the answer?

9.6 x 100.65

+ 4.026 =

8.321

Evaluate the expression to the correct number of

significant figures. How many sig. figs. in the answer?

320.75 - (6102.1 / 3.1) =

Evaluate the expression to the correct number of

significant figures. How many sig. figs. in the answer?

(14.20000 0.7288) + (12.00536 0.0201) =

19

significant figures. How many sig. figs. in the answer?

9.6 x 100.65

+ 4.026 =

8.321

Evaluate the expression to the correct number of

significant figures. How many sig. figs. in the answer?

320.75 - (6102.1 / 3.1) =

Evaluate the expression to the correct number of

significant figures. How many sig. figs. in the answer?

(14.20000 0.7288) + (12.00536 0.0201) =

19

20.
Dimensional Analysis

Units are multiplied together or divided into

each other along with the numerical values.

• Keep track of both numerical values and units.

http://www.wired.com/2010/11/1110mars

-climate-observer-report/

Conversions: Two or More Factors

What is the mass in g, of 1.00 gal of H2O? The density

of water is 1.00 g/mL. 1 L = 1.057 qt, 1 gal = 4 qt

20

Units are multiplied together or divided into

each other along with the numerical values.

• Keep track of both numerical values and units.

http://www.wired.com/2010/11/1110mars

-climate-observer-report/

Conversions: Two or More Factors

What is the mass in g, of 1.00 gal of H2O? The density

of water is 1.00 g/mL. 1 L = 1.057 qt, 1 gal = 4 qt

20

21.
Conversions Involving Volume

Express a volume of 1.250 L in mL and cm3

1 mL

(1.250 L)× =1,250. mL

1 × 10 L

1000 mL

(1.250 L)× =1,250. mL

1L

1000 cm3

(1.250 L)× =1,250. cm3

1L

Express a volume of 1,250. cm3 in m3.

The prefix centi is 10-2, 1 cm = 110-2 m for length.

Volume involves cubed units, create a conversion:

110−2 m 110−2 m 110−2 m 110−6 m3

=

1 cm 1 cm 1 cm 1 cm3

Use the conversion to express the volume in m3:

110−6 m3

(1250. cm3 ) × = 1.250 × 10−3 m3

1 cm3

21

Express a volume of 1.250 L in mL and cm3

1 mL

(1.250 L)× =1,250. mL

1 × 10 L

1000 mL

(1.250 L)× =1,250. mL

1L

1000 cm3

(1.250 L)× =1,250. cm3

1L

Express a volume of 1,250. cm3 in m3.

The prefix centi is 10-2, 1 cm = 110-2 m for length.

Volume involves cubed units, create a conversion:

110−2 m 110−2 m 110−2 m 110−6 m3

=

1 cm 1 cm 1 cm 1 cm3

Use the conversion to express the volume in m3:

110−6 m3

(1250. cm3 ) × = 1.250 × 10−3 m3

1 cm3

21