Course: Instructor: Office location: Office phone: Office hours:

General Chemistry I (CHEM 140, Section: 02) Cevdet Akbay LS 227 672-1943 MTWRF 10:00 AM-12:00 PM or by appointments

Chemistry:

The Study of Change

Chapter 1

Why ice melts and water evaporates?

Why keeping foods cold slows their spoilage?

How our bodies use food to maintain life? Why leaves turn colors in the fall?

How a battery generates electricity? How Blah blah blah

Chemistry

supplies answers to these questions and countless others like them. Then, how we define Chemistry? It depends:

Chemistry

is the study of the properties of materials and the changes that materials undergo.

Chemistry

is the scientific discipline that treats the composition, properties, and transformations of matter.

One of the joys of learning

chemistry

is seeing how chemical principles operate in all aspects of our lives, from lighting a match to the development of drugs to cure cancer. You are just beginning the journey of learning

chemistry

. This first chapter provides an overview of what

chemistry

is about and deals with some fundamental concepts of matter and scientific measurements.

Chemistry

is the study of matter and the changes it undergoes

(another definition)

1. Matter

is anything that occupies space and has mass.

2.

A

substance

is a form of matter that has a definite composition and distinct properties.

water, ammonia, sucrose, gold, oxygen

A

mixture

is a combination of two or more substances in which the substances retain their distinct identities.

1.

Homogenous mixture

– composition of the mixture is the same throughout.

air, soft drink, milk, solder 2.

Heterogeneous mixture

– composition is not uniform throughout.

cement, iron filings in sand

Physical means

can be used to separate a mixture into its pure components.

distillation magnet

An

element

is a substance that

cannot

be separated into simpler substances by

chemical

means.

• 115 elements have been identified • 83 elements occur naturally on Earth gold, aluminum, lead, oxygen, carbon • 32 elements have been created by scientists technetium, americium, seaborgium

These elements vary widely in their abundance, as shown in Figure 1.6.

Figure 1.6

Elements in percent by mass in (a) Earth's crust (including oceans and atmosphere) and (b) the human body.

Some of the more familiar

elements

are listed in Table 1.2, along with the chemical abbreviations —or chemical

symbols

—used to denote them. The symbol for each element consists of

one or two letters

, with the first letter capitalized. These symbols are often derived from the English name for the element, but sometimes they are derived from a foreign name instead (last column in Table 1.2). All the known elements and their symbols are listed on the front inside cover of this text. The table in which the symbol for each element is enclosed in a box is called the

periodic table

.

A

compound

is a substance composed of atoms of two or more elements chemically united in fixed proportions.

Compounds can only be separated into their pure components (elements) by

chemical

means.

Water (H 2 O) Glucose (C 6 H 12 O 6 ) Ammonia (NH 3 )

Most elements can interact with other elements to form

compounds

. Hydrogen gas, for example, burns in oxygen gas to form water.

Figure 1.10

Conversely, water can be decomposed into its component

elements

(hydrogen and oxygen) by passing an electrical current through it, as shown in Figure 1.7.

Figure 1.7

Water decomposes into its component elements, hydrogen and oxygen, when a direct electrical current is passed through it. The volume of hydrogen (on the right) is twice the volume of oxygen (on the left).

Pure

water

, regardless of its source, consists of 11% hydrogen and 89% oxygen by mass. This macroscopic composition corresponds to the molecular composition, which consists of

two hydrogen atoms

combined with

one oxygen atom

.

As seen in Table 1.3, the properties of water bear no resemblance to the properties of its component elements.

Hydrogen

,

oxygen

, and

water

are each unique substances.

Law of constant composition:

A law that states that the elemental composition of a pure compound is always the same, regardless of its source. This law is also known as

law of definite proportions.

It was first put forth by the French chemist Joseph Louis Proust in 1800s.

Each of the followings can be classified as a heterogeneous mixture, homogeneous mixture, compound, or element. How would you classify each? 1) Iced tea

a) heterogeneous mixture

2) Ethyl alcohol

a) heterogeneous mixture b) homogeneous mixture c) compound b) homogeneous mixture c) compound d) element d) element

3) Ozone

a) compound b) homogeneous mixture c) heterogeneous mixture d) element

4) Air in the classroom

a) heterogeneous mixture b) homogeneous mixture c) compound d) element

5) Table salt

a) heterogeneous mixture b) c) d) homogeneous mixture compound element

6) A salt and sand mixture

a) heterogeneous mixture b) c) d) homogeneous mixture compound element

1) b 2) c 3) a 4) b 5) c 6) a 1)

a) b) c) d)

Which of the following is not matter?

elemental phosphorus light dust pizza

2)

a) b) c) d)

Which of the following is a pure substance?

concrete nitrogen blue-cheese salad dressing air

4)

a) b) c) d)

Which of the following is not a state of matter?

Gas Vacuum Solid Liquid

3)

a) b) c) d)

Passing an electric current through a certain substance produces oxygen and sulfur. This substance cannot be a(n)

Compound Element Mixture Pure substance

1) b 2) b 3) b 4) b

Three States of Matter

Matter can exist in one of three

states of matter

: a

gas

, a

liquid

, or a

solid

. A

gas

is highly compressible and will assume both the shape and the volume of its container.

A

liquid

is not compressible and will assume the shape but not the volume of its container.

A

solid

also is not compressible, and it has a fixed volume and shape of its own.

Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior gas assumes the shape and volume of its container liquid assumes the shape of the part of the container which it occupies solid retains a fixed volume and shape

particles can move past one another particles can move/slide past one another

compressible not easily compressible

lots of free space between particles little free space between particles

flows easily flows easily

rigid - particles locked into place

not easily compressible

little free space between particles

does not flow easily

particles can move past one another particles can move/slide past one another rigid - particles cannot move/slide past one another

Properties of Matter

Every substance has a unique set of

properties

(or characteristics) that allow us to recognize it and to distinguish it from other substances. For example, the properties listed in Table 1.3 allow us to distinguish hydrogen, oxygen, and water from one another. The properties of matter can be categorized as

physical properties

or

chemical properties. Physical properties

can be measured without changing the identity and composition of the substance. These properties include color, odor, density, melting point, boiling point, and hardness.

Chemical properties

describe the way a substance may change or

react

to form other substances. A common chemical property is flammability, the ability of a substance to burn in the presence of oxygen.

Some properties —such as temperature, melting point, and density —

do not depend on

the amount of the sample (matter) being examined. These properties, called

intensive properties

, are particularly useful in chemistry because many can be used to

identify

substances.

Extensive properties

of substances

depend on

the quantity of the sample and include measurements of mass and volume. Extensive properties relate to the

amount

of substance present.

Physical or Chemical?

A

physical change

does not alter the composition or identity of a substance.

ice melting sugar dissolving in water A

chemical change

alters the composition or identity of the substance(s) involved.

hydrogen gas burns in oxygen gas to form water

Which of the following is a chemical process?

a) dissolving sugar in water b) crushing of stone c) tarnishing of silver d) melting of lead Which of the following is an intensive quantity?

a) heat content of a substance b) mass of a substance c) density of a substance d) volume of a substance Which one of the following is a physical process?

a) the rusting of iron b) the explosion of nitroglycerine c) the baking of a potato d) the condensation of water vapor e) the formation of polyethylene from ethylene c c d

Matter - anything that occupies space and has

mass

.

mass

– measure of the quantity of matter SI unit of mass is the

kilogram

(kg) 1 kg = 1000 g = 1 x 10 3 g

weight

– force that gravity exerts on an object weight =

g

x mass on earth,

g

= 9.8 m/s 2 on moon,

g=

1.63 m/s 2 A 60 kg bar will weigh ~600 N on earth ~100 N on moon

Table 1.2 SI Base Units

Base Quantity Name of Unit Length Mass Time Current Temperature Amount of substance Luminous intensity meter kilogram second ampere kelvin mole candela Symbol m kg s A K mol cd

Unit Ex. meter Table 1.3 Prefixes Used with SI Units Prefix Tera Giga Mega Kilo Deci Centi Milli Micro Nano Pico Symbol T G M k d c m m n p Meaning 10 12 10 9 10 6 10 3 10 -1 10 -2 10 -3 10 -6 10 -9 10 -12

Volume

– SI derived unit for volume is cubic meter (m 3 ) 1 cm 3 = (1 x 10 -2 m) 3 = 1 x 10 -6 m 3 1 dm 3 = (1 x 10 -1 m) 3 = 1 x 10 -3 m 3 1 L = 1000 mL = 1000 cm 3 = 1 dm 3 1 mL = 1 cm 3

Density

– SI derived unit for density is kg/m 3 1 g/cm 3 = 1 g/mL = 1000 kg/m 3 density = mass volume

d

=

m V

A piece of platinum metal with a density of 21.5 g/cm 3 has a volume of 4.49 cm 3 . What is its mass?

d

=

m V m

=

d

x

V

= 21.5 g/cm 3 x 4.49 cm 3 = 96.5 g

K = 0 C + 273.15

273 K = 0 0 C 373 K = 100 0 C 0 9 5 0 C + 32 32 0 F = 0 0 C 212 0 F = 100 0 C 0 5 9 0 F – 32) 0 0 C = 32 0 F 100 0 C = 212 0 F

Convert 172.9 0 F to degrees Celsius.

0 9 5 0 C + 32 0 F 9 5 0 C 5 9 x ( 0 F – 32) = 0 C 0 5 9 0 F – 32) 0 5 9 – 32) = 78.3

Scientific Notation The number of atoms in 12 g of carbon: 602,200,000,000,000,000,000,000 6.022 x 10 23 The mass of a single carbon atom in grams: 0.0000000000000000000000199

1.99 x 10 -23 N x 10

n

N is a number between 1 and 10

n

is a positive or negative integer

Scientific Notation

568.762

move decimal left

n > 0

568.762 = 5.68762 x 10

2

0.00000772

move decimal right

n < 0

0.00000772 = 7.72 x 10

-6 Addition or Subtraction

1. Write each quantity with the same exponent

n

2. Combine N 1 and N 2 3. The exponent,

n

, remains the same 4.31

x 10

4

+ 3.9

x 10

3

= 4.31

x 10

4

+ 0.39

x 10

4

= 4.70 x 10

4

Scientific Notation

Multiplication

1. Multiply N 1 and N 2 2. Add exponents

n 1

and

n 2

( 4.0

( x 10 -5 ) x ( 7.0

4.0

x 10 3 ) = x 7.0

) x (10 -5

+

3 ) = 28 x 10 -2

2.8 x 10 -1

=

Division

1. Divide N 1 and N 2 2. Subtract exponents

n 1

and

n 2

8.5

( x 10 4 8.5

÷ ÷ 5.0

x 10 5.0

) x 10 4

-

9 9

1.7 x 10 -5

= =

Significant Figures

• Any digit that is not zero is significant 1.234 kg

4

significant figures • Zeros between nonzero digits are significant 606 m

3

significant figures • Zeros to the left of the first nonzero digit are

not

significant 0.08 L

1

significant figure • If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg

2

significant figures • If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant 0.00420 g

3

significant figures

• When a number ends in zeros but contains no decimal point, the zeros may or may not be significant. 5000 ( one , two, three , or four significant figures) 10,500 (three, four, or five significant figures) The use of exponential notation eliminates the potential ambiguity of whether the zeros at the end of a number are significant 1.03

x 10 4 1.030

x 10 4 1.0300

x 10 4 Three Four Five significant figures significant figures significant figures

How many significant figures are in each of the following measurements?

24 mL 2 significant figures 4 significant figures 3001 g 0.0320 m 3 6.4 x 10 4 molecules 560 kg 560. kg 3 significant figures 2 significant figures 2 or 3 significant figures 3 significant figures

Significant Figures

Addition or Subtraction

The answer

cannot

have

more digits

to the right of the decimal point than any of the original numbers.

89.332

+ 1.

1 90.432

one significant figure after decimal point round off to 90.

4 3.

70 -2.9133

0.7867

two significant figures after decimal point round off to 0.

79

Significant Figures

Multiplication or Division

The number of significant figures in the result is set by the original number that has the

smallest

number of significant figures 4.51 x 3.6666 = 16.536366

=

16.5

2 3 sig figs round to 3 sig figs 6.8 ÷ 112.04 = 0.0606926 =

0.061

sig figs round to 2 sig figs

Accuracy

– how close a measurement is to the

true

value

Precision

– how close a set of measurements are to each other accurate & precise precise but not accurate not accurate & not precise

Dimensional Analysis

1. Determine which unit conversion factor(s) are needed 2. Carry units through calculation 3. If all units cancel except for the desired unit(s), then the problem was solved correctly.

Given unit x Desired unit Given unit = Desired unit

How many mL are in 1.63 L?

1 L = 1000 mL 1.63 L x 1000 mL 1L = 1630 mL 1.63 L x 1L 1000 mL = 0.001630

L 2 mL

We know that 1 inch is the same length as 2.54 centimeters. We’re told that the new start of the Indians basketball team is 64 cm tall and that he’s going to be the starting center. Based on height alone we’ll be able to tell if he’ll help the team but we Americans think in INCHES not CENTIMETERS so we need to convert units.

Remember Given unit x Desired unit = Desired unit Given unit

Desired unit

1 in 64 cm X 2.54 cm

Given unit

64 2.54

in

= 25.4 in

The speed of sound in air is about 343 m/s. What is this speed in miles per hour?

meters to miles seconds to hours 1 mi = 1609 m 1 min = 60 s 1 hour = 60 min 343 m s 1 mi x 1609 m x 60 s 1 min x 60 min 1 hour mi = 767 hour

A car travels 28 mi per gallon of gasoline. How many kilometer per liter will it go?

1.0 km = 0.62137 mi or 0.62137 mi 1.0 km or 1.0 km 0.62137 mi 1.0 gal = 3.7854 L or 3.7854 L 1.0 gal or 1.0 gal 3.7854 L

TWO

steps 28 mi gal x 1.0 km 0.62137 mi = 45 km gal

Step 1

45 km gal x 1 gal 3.7854 L ONE step = 12 km L

Step 2

28 mi gal x 1 km 0.62137 mi x 1 gal 3.7854 L = 12 km L Watch for significant figures. The real number is 11.9 km/L, however, here we have three sig. figures. We have started with a two sig. figure number (28). Thus, the final answer has to have two significant figures!