CHAPTER ONE
INTRODUCTION:
MATTER AND MEASUREMENT

The Study of Chemistry

Chemistry: • is the study of properties of materials and changes that they undergo. • can be applied to all aspects of life.

The Atomic and Molecular Perspective of Chemistry

Chemistry involves the study of the properties and the behavior of matter. Matter: • is the physical material of the universe. • has mass. • occupies space.

Examples of matter:

~100 elements constitute all matter.

A property is any characteristic that allows us to recognize a particular type of matter and to distinguish it from other types of matter.

Elements: • are made up of unique atoms, the building blocks of matter. • approximately 116 elements. • names of the elements are derived from a wide variety of sources.

Molecules: • are combinations of atoms held together in specific shapes. • Macroscopic (observable) propertied of matter relate to submicroscopic realms of atoms • Properties relate to composition (types of atoms present) and structure (arrangement of atoms present.

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• Combinations of atoms held together in specific shapes
CH3CH2OH HOCH2CH2OH
Ethyl alcohol ethylene glycol

Why study Chemistry?

We study chemistry because: • It has a considerable impact on society (health care, food, clothing, conservation of natural resources, environmental issues, etc.) • Chemistry serves biology, engineering, physics, marine science, etc. “Chemistry is the central science.”

Classifications of Matter

There are two ways to matter:

Physical State- solid, liquid gas

• Composition-elements, compounds, mixture

States of Matter

Solids, liquids and gases are the three forms of matter called states of matter.

Gas (vapor)-No fixed volume or shape, conforms to shape of container compressible.

Liquid-Fixed volume, no fixed shape, incompressible.
Solid-fixed volume and shape, rigid, incompressible.

Why these differences between gas, liquid and solid?

We must look at the properties on a molecular level:
Gas: Molecules far apart moving rapidly collide often.
Liquid: Molecules closer than gas, move rapidly can slide over each other.
Solid: Molecules packed closely in definite arrangements.

Pure Substances Matter that has distinct properties and composition that doe not vary from sample to sample. • Elements-cannot be decomposed into other substances.

• Compounds-composed of two or more elements.

Elements • There are 116 elements. • They vary in abundance. • Each is given a unique name and is abbreviated by a chemical symbol. • They are organized in the periodic table. • Each is given a one- or two – letter symbol derived from its name.
Common elements are found in Table 1.2 LEARN!!!!

Compounds
Compounds are combinations of elements. Elements react with each other to form compounds. (properties different from constituent elements.

Water is a combination of the elements hydrogen and oxygen.

1) H2(g) +½O2(g) → H2O(l)

H2O: 89% O2 and 11% H2

2) Fe(s) + 3/2 O2(g) Fe2O3(s)

Fe2O3: 69.9% Fe and 30.1% O

Law of Constant Composition (Law of Definite Proportions)-the elemental composition of a pure compound is always the same.
i.e. Water is always 89% O2 and 11% H2.
Mixtures
Mixture-a combination of two or more substances in which each substance retains its own chemical identity. • Each substance retains its own identity; each substance is a component of the mixture Components of Mixtures *Solvent-present in greater quantity *Solute-present in lesser quantity

• Mixtures have variable composition. Heterogeneous Mixtures-do not have the same composition, properties and appearance. Examples:

Homogeneous Mixtures- are uniform throughout, e.g. , air, sweet tea; they are called solutions. Solutions can be gas, liquid or solid. Examples:

Properties of Matter

Every substance has a unique set of properties that allow us to distinguish it from other substances.

Properties can be classified as:
Physical Properties-can be measured without changing the identity and composition of the substance.
Examples:

Chemical Properties-describe the way a substance may change or react to form other substances.
Examples:

Intensive Properties-do not depend on the amount of the sample being examined.
Examples:

Extensive Properties:-depend on the quantity of substance present.
Examples:

“Intensive properties give an idea of the composition of a substance whereas extensive properties give an indication of the quantity of a substance present.”

Physical and Chemical Changes

Physical Change: substance changes physical appearance without altering its identity.
Examples:

Chemical Change (or chemical reaction): substance transforms into a chemically different substance.
Examples:
Separation of Mixtures Key: separation techniques exploit differences in properties of the components. For example, separating a heterogeneous mixture of iron and gold filings.
How could you separate?

Another technique used to separate heterogeneous mixtures is filtration – remove solid from liquid.

Techniques used to separate homogeneous mixtures. • Distillation: based on the differences in boiling points-boil off one or more components of the mixture.

• Chromatography (the writing of colors): based on the differing abilities of substances to adhere to the surfaces of various solids such as paper and starch and exploits solubility of components.

The Scientific Method
Provides guidelines for the practice of science. • Collect data (observe, experiment, etc.) • Look for patterns, try to explain them, and develop a hypothesis (tentative explanation) • Test hypothesis, then refine. • Bring all information together into a scientific law (concise statement or equation that summarizes tested hypothesis) • Bring hypothesis and laws together into a theory. A theory should explain general principles.
Units of Measurement

Many properties of matter are quantitative, i.e., associated with numbers. The numbers must mean something- the measured quantity must have BOTH a number and a unit.

Examples:
I will be back in my office in 2.

My height is 132.08 .

The units most often used for scientific measurement is the metric system.

In 1960, an international agreement was reached specifying a particular choice of metric units for use on scientific measurements. The preferred units are called SI units, after the French “Systeme International d′Unités”.

There are seven base units- Table 1.4 LEARN!!!

Also, we must know the selected metric prefixes found in Table 1.5.

We will deal with three –length, mass and temperature.

Length • base unit-the meter (m) 1 m = 1.0936 yards 1 inch = 2.54 centimeters 1 km = 1000 m

Mass is the measure of the amount of material in an object.
Explain the difference between mass and weight. • base unit –the kilogram (kg)
` 1 kg = 1000 g 1 kg = 2.205 pounds 1 lb = 453. 6 grams

Temperature • is the measure of the hotness or coldness of an object. • Scientific studies use the Celsius and Kelvin scales Celsius Scale: water freezes at 0 ºC and boils at 100 ºC. (sea level) Kelvin Scale: water freezes at 273.15 K and boils at 373.15 K (sea level) Fahrenheit Scale (not used in science): water freezes at 32 ºF and boils at 212 ºF (sea level)

• base unit – Kelvin.

Conversions: ºF = 9/5 ºC + 32

ºC= 5/9(ºF – 32)

ºC = K-273.15

K = ºC + 273.15

Examples:
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Derived SI Units

• These are formed from the seven base units • Example: Velocity is distance traveled per unit time, so units of velocity (speed) are units of distance (m) divided by units of time (s): m/s

Volume • Units of volume = (units of length)3 = m3 (is very large)

• base unit = m3

• m3 is is very large, therefore more common used cm3 = mL (milliliter) =cc (cubic centimeter) dm3 = liter, L • cm3 and liter, L are NOT SI units.

Density • Is used to characterize substances • Density is defined as mass divided by volume. • Density = mass Volume

• Units: g/cm3 or g/mL • Densities are temperature dependent. (When reporting densities the temperature at which the density was measured must be reported.)

Note: The gram was originally defined as the mass of 1 mL of water at 4 ºC, since the density of water is 1.000 g/mL at 4 ºC.

Examples: Problem 1.27 (a), (b), and (c)

Uncertainty in Measurement

There are two types of numbers: • Exact numbers (values known exactly–known as counting or defined) Counting numbers:

Definitions:

• Inexact numbers (have some uncertainty-derived from measurement) All measurements have some degree of uncertainty in them. Examples:

Precision and Accuracy
Precision: how well measured quantities agree with each other.
Accuracy: how well measured quantities agree with the “true value”.

Precision and Accuracy

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Significant Figures

The digits that indicate the precision with which a measurement is made; all digits of a measured quantity are significant including the last digit.

A new penny has a mass of 2.54 g.

An old penny has a mass of 3.0836 g

Rules 1. Nonzero numbers and zeros between nonzero numbers are always significant. 2. Zeros before the first nonzero digit are not significant. 3. Zeros at the end of the number after a decimal point are significant. 4. Zeros at the end of a number before a decimal point are ambiguous. Exponential notation eliminates this ambiguity.

Examples:
Problem 1.37 a) 358 kg 1) 116,300

b) 0.054 s 2) 0.00296

c) 6.3050 cm 3) 8.070

d) 0.0105 L 4) 9.7750 x 10-4

e) 7.0500 x 10-3 m3 5) 12.000

Significant Figures in Calculations

• Multiplication and Division Report to the least number of significant figures.

Examples: 1) 6.221 cm x 5.2 cm = 32.3492 cm2

2) 100.00 m/10.5 s =

• Addition and Subtraction Report to the least number of decimal places. 1)20.4 g- 1.322 g = 19.078 g =

2) 12.0550 + 9.05 =

3) 257.2-19.789 =

“In multiple step calculations always retain an extra significant figure until the end to prevent rounding errors.”

Review of Rounding Off Numbers Round the following:

|Number |3 |2 |1 |0 |
| |figures |figures |figures |figures |
|76.386 | | | | |
|28.32 | | | | |
|28.62 | | | | |

Other Mathematical Operations-Scientific Notation

I. Writing Numbers in Scientific Notation Examples: 1) 1,200,000 2) 12 3) 0.000000728

II. Addition and Subtraction Examples: 1) 5.22 x 104 + 3.21 x 102 = (522 x 102) + (3.21 x 102) = 525.21 x 102 =

2) 6.25 x 10-2 - 5.77 x 10-3 = 6.25 x 10-2 - 0.577 x 10-2 = 5.67 x 10-2

3) 5.00 x 10-2 + 4.7 x 10-3 =

III. Multiplication and Division Examples: 1) (5.4 x 102)(2.1 x 103) = (5.4)(2.1) x 10 2+3 11.34 x 105 =

2) (1.2 x 105)(3.22 x 10-3) = (1.2)(3.22) x 10 5-3 3.864 x 102 =

3) (3.2 x 105)/(6.5 x 102) = 3.2/6.5 x 10 5-2 = 0.492 x 103 =

4) (5.7 x 107)/(8.5 x 10-2) = 0.671 x 10 7- (-2) =

IV. Powers and Roots Examples:

1) (1.2 x 105)3 = (1.2)3 x 10 5x3 = 1.728 x 10 15

2) 3√(2.5 x 106) = 3√(2.5 x 10 6/3

Examples: 1) (5.98 x 12)(2.77 x 10-5)

2) 4√(1.75 x 10-12)

3) (6.67 x 10-8)3 =
Dimensional Analysis

• A method of problem solving in which units are carried through all calculations.

• Dimensional analysis ensures that the final answer of a calculation has the desired unit.

• Given units can be multiplied or divided to give the desired unit. Given unit x desired unit = desired unit given unit

• Conversion Factors are simple ratios.

Conversion factor = (desired unit)/(given unit) These are fractions whose numerator and denominator are the same quantity expressed in different units.

• Multiplication by a conversion factor is equivalent to multiplying be a factor of one.

Examples:
Practice Exercise 1.9 on page 26.

Problem 1.47 (a), (c), (d), (f)

Practice Exercise 1.12 page 29

Problem 1.49 (b)