© 2012 pearson education, inc. chapter 1 introduction: matter and measurement dr. subhash c. goel...

© 2012 Pearson Education, Inc.

Chapter 1

Introduction: Matter and Measurement

Dr. Subhash C. GoelSouth GA State College

Douglas, GA

Lecture Presentation


Paper Chromatography Demonstration

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Measurement© 2015 Pearson Education, Inc.

Chemistry

• Chemistry is the study of the properties and behavior of matter.

• It is central to our fundamental understanding of many science-related fields.

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Chemistry

The atom is the basic unit of chemistry and a fundamental component of matter. Chemistry is concerned with atoms and their interactions with other atoms, with particular focus on the properties of the chemical bonds formed between species.

Chemistry is also concerned with the interactions between atoms or molecules and various forms of energy (e.g. photochemical reactions, oxidation-reduction reactions, changes in phases of matter, separation of mixtures, properties of polymers, etc.)

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Matter

• is the physical material of the universe.

• has mass.

• occupies space.

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Matter

• Atoms are the building blocks of matter.

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ELEMENTS• are made up of unique atoms, the building

blocks of matter.

• Names of the elements are derived from a wide variety of sources (e.g., Latin or Greek, mythological characters, names of people or places).

For example:

Hydrogen (H) Iron (Fe)

Carbon (C) Zinc (Zn)

Oxygen (O)


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MOLECULES• are combinations of atoms held together in

specific shapes.

• Macroscopic(observable) properties of matter relate to submicroscopic realms of atoms.

• Properties relate to composition (types of atoms) and structure (arrangement of atoms) present.


Copyright © Cengage Learning. All rights reserved. 1 | 10

Compound

A substance composed of two or more elements chemically combined.

For example:

Water (H2O) Glucose

Carbon dioxide (CO2)

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Compounds and Composition

• Compounds have a definite composition. That means that the relative number of atoms of each element that makes up the compound is the same in any sample.

• This is The Law of Constant Composition (or The Law of Definite Proportions).


Mixture

A material that can be separated by physical means into two or more substances

For example:

Italian salad dressing

Saltwater


Heterogeneous MixtureA mixture that consists of physically distinct parts, each with different properties.

For example:Salt and iron filingsOil and vinegar


Homogenous Mixture

A mixture that is uniform in its properties; also called a solution.

For example:

Saltwater Sugar Solution

Air

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Methods of Classification

• State of Matter

• Composition of Matter

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Classification of Matter Based on Composition

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If you follow this scheme, you can determine how to classify any type of matter.

Homogeneous mixture

Heterogeneous mixture

Element Compound

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States of Matter

• Solids: fixed shape and volume

• Liquids: fixed volume but no shape

• Gases: No fixed shape and volume


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Matter• A Property is any characteristic that

allows us to recognize a particular type of matter and distinguish it from others.

• There are >100 substances called Elements that combines to give all kind of matter in the universe.

• We will see that properties of matter relates to both the kind of atoms in the matter contains (composition) & their arrangement in matter (structure).



Physical Change

A change in the form of matter but not in its chemical identity.

For example:

Melting

Dissolving

Density

Mass


Chemical Change = Chemical Reaction

A change in which one or more kinds of matter are transformed into a new kind of matter or several new kinds of matter.

For example:

Rusting

Burning

Reactivity with acids

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Types of Properties

• Intensive Properties…– Are independent of the amount of the

substance that is present.◦ Density, boiling point, color, etc.

• Extensive Properties…– Depend upon the amount of the substance

present.◦ Mass, volume, energy, etc.


Physical Property

A characteristic that can be observed for a material without changing its chemical identity.

For example:

Physical state

Boiling point

Color


Chemical Property

A characteristic of a material involving its chemical change.

For example:

Ability to react with oxygen

Ability to react with fluorine

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Changes in State of Matter

Converting between the three states of matter is a physical change.

When ice melts or water evaporates, there are still 2 H atoms and 1 O atom in each molecule.

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Chemical Reactions

In the course of a chemical reaction, the reacting substances are converted to new substances.

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Potassium is a soft, silvery-colored metal that melts at 64°C. It reacts vigorously with water, with oxygen, and with chlorine. Identify all of the physical properties and chemical properties given in this description.

Reacts with chlorineMelting point (64°C)

Reacts with oxygenSilvery-colored

Reacts with waterSoft

Chemical PropertyPhysical Property

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Separation of Mixtures

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Filtration

In filtration, solid substances are separated from liquids and solutions.

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Distillation

Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components.

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Chromatography

• This technique separates substances on the basis of differences in the ability of substances to adhere to the solid surface, in this case, dyes to paper.

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Numbers and Chemistry• Numbers play a major role in chemistry.

Many topics are quantitative (have a numerical value).

• Concepts of numbers in scienceUnits of measurementQuantities that are measured and calculatedUncertainty in measurementSignificant figuresDimensional analysis

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Units of Measurement

• 1960: An international agreement was reached to use metric units (called SI units) in scientific measurements.

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SI Units

• Système International d’Unités

• A different base unit is used for each quantity.

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Metric System

Prefixes convert the base units into units that are appropriate for the item being measured.

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Common Metric Conversions Centimeter / 100 = meters centimeters x 10,000 = micrometers

Centimeters x 10 = millimeters micrometers / 10,000 = centimeters

Millimeters /10 = centimeters Meter x 100 = centimeters

Grams x 1000 = milligrams nm/10,000,000 = cm

Milligrams / 1000 = grams cm x 10,000,000 = nm

Micrometers x 1000 = nanometers Grams x 1000,000 = micrograms

Nanometers / 1000 = micrometers μm = micrometers

Micrometers / 1000 = millimeters m = meters

Nanometers / 1,000,000 = millimeters mm = millimeters

Millimeters x 1,000,000 = nanometers nm = nanometers

Liters x 1000 = milliliters g = grams

Milliliters / 1000 = liters mg = milligrams

Grams / 1000 = kilograms ml = milliliters

Grams x 1000 = milligrams L = liters

Kilograms x 1000 = grams Nanometers / 10,000,000 = centimeters

Millimeter / 1000 = meters Kilometer x 1000 = meters

Millimeter x 1000 = micrometer meters / 1000 = kilometers

Meters x 1000 = millimeters meters x 1,000,000,000 = nm

Micrometers / 1,000,000 = meters nm/1,000,000,000 = meters

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Measurement

Using Prefixes1. 103 grams = 1_____grams

2. 2x10-3 meters = 2______meters

3. 2.6 x 10-1 L = 2.6______L

4. 6 x 10-2 g = 6______g

5. 1 gigajoules = 1 x _______joules

6. 9 µm (micrometer) = 9 x _______m

7. 8 cm = 8 x __________m

8. 1 ng = 1 x ________g

9. Convert 0.30 gigameters/second to a value without a prefix in scientific notation.

1. Kilo 2. milli 3. d 4. c 5. 109 6. 10-6 7. 10-2 8.10-9 9. 3.0 x 108 meters/second


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Volume• Note that volume is not a base

unit for SI; it is derived 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 is a cube 1 centimeter (cm) long on each side, also called 1 cubic centimeter (cm × cm × cm = cm3).

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Temperature

In general usage, temperature is considered the “hotness and coldness” of an object that determines the direction of heat flow.

Heat flows spontaneously from an object with a higher temperature to an object with a lower temperature.

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Temperature

• In scientific measurements, the Celsius and Kelvin scales are most often used.

• The Celsius scale is based on the properties of water.– 0 C is the freezing point of water.– 100 C is the boiling point of water.

• The kelvin is the SI unit of temperature.– It is based on the properties of gases.– There are no negative Kelvin temperatures.– The lowest possible temperature is called absolute

zero (0 K).

• K = C + 273.15

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Temperature

• The Fahrenheit scale is not used in scientific measurements, but you hear about it in weather reports!

• The equations below allow for conversion between the Fahrenheit and Celsius scales: F = 9/5(C) + 32 C = 5/9(F − 32)

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Exercise

Convert 40o F to K.


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In winter, the average low temperature in interior Alaska is -30.°F (two significant figures). What is this temperature in degrees Celsius and in kelvins?

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Density

• Density is a physical property of a substance.

• It has units that are derived from the units for mass and volume.

• The most common units are g/mL or g/cm3.

• D = m/V

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Oil of wintergreen is a colorless liquid used as a flavoring. A 28.1-g sample of oil of wintergreen has a volume of 23.7 mL. What is the density of oil of wintergreen?

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A sample of gasoline has a density of 0.718 g/mL. What is the volume of 454 g of gasoline?

V

md

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Uncertainty in Measurement

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Numbers Encountered in Science

• Exact numbers are counted or given by definition. For example, there are 12 eggs in 1 dozen.

• Inexact (or measured) numbers depend on how they were determined. Scientific instruments have limitations. Some balances measure to ±0.01 g; others measure to ±0.0001g.

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Uncertainty in Measurements• Different measuring devices have different uses and

different degrees of accuracy.

• All measured numbers have some degree of inaccuracy.

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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.

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MeasurementS. Goel, SGSC

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Accuracy versus Precision

• Accuracy refers to the proximity of a measurement to the true value of a quantity.

• Precision refers to the proximity of several measurements to each other.

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Dimensional Analysis

• We use dimensional analysis to convert one quantity to another.

• Most commonly, dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm)

1 in.

2.54 cm

2.54 cm

1 in.or

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Use the form of the conversion factor that puts the sought-for unit in the numerator:

Given unit desired unitdesired unit

given unit

Conversion factor

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Unit Conversions

1 inch = 2.54 cm cm/2.54 = inch Inch x 2.54 = cm

1 cm = 0.3937 inches

1 mile = 1.6093 Km km / 1.609 = mile Mile x 1.6 = Km

1 km = 0.6214 miles

1 mile = 5280 ft

1 gal = 3.7854 L L/3.7854 = gal gal x 3.7854=L

1 gal = 4 qrt

1 L = 1.067 qt

1 lb = 453 g g/453 = lb lb x 453 = g

1 kg = 2.20 lb kg x 2.2 = lb lb/2.2 = kg

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Example:

Convert 8.00 m to inches.

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Exercise

• How many mL are in 1.63 L?


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Exercise

A liquid helium tank has a volume of 275 L. What is the volume in m3?


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Execrcise

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


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Notation Terminology

• Fixed Decimal or Fixed Notation: numbers such as 1230.024, 236.18, or 50923.005

• Exponential Notation: very large or very small numbers are written more clearly in exponential notation. Example: instead of writing that one drop of water has 1,500,000,000,000,000,000,000 molecules, we write 1.5x1021 molecules or atomic radius of Ne atom is 0.0000000070 centimeter, we write 7.0x10-9 centimeter.

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Power of 10

106 = 1000000

103 = 1000

101 = 10

100 = 1

10-1 = 0.1

10-2 = 0.01

10-4 = 0.0001

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Multiplying/Dividing by 10, 100, 1000

• When multiplying by 10, 100, 1000, etc move the decimal to the right by the number of zeros in 10, 100, or 100072x100 = 7200 -0.0624x1000 = -62.4

• When dividing by 10, 100, 1000, etc move the decimal to the left by the number of zeros in 10, 100, or 100072/100 = 0.72 -0.624x1000 = -0.000624

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Converting Exponential Notation to Fixed Decimal Notation

• Rule 1. The sign in front never changes.• Rule 2. In Exponential notation, if the power of 10 is

positive then move decimal point to the right by the same number of places as the value of exponent.

• Rule 3. If the power of 10 is negative then move decimal point to the left.

• Examples: 2 x 102 = 200; -0.0033 x 103 = -3.3

2 x 10-2 = 0.02; 7653.8 x 10-3 = 7.6538

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Changing Exponential to Scientific Notation

• To convert an exponential to scientific notation, move the decimal so that significand is one or greater but less then ten.

• Examples:

• 0.045 x 105 = 4.5 x 103

• -0.0057 x 10-2 = -5.7 x 10-5

• -8544 x 10-7 = -8.544 x 10-4

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Converting Numbers to Scientific Notation

• Any number to the zero power is equal 1.

Examples: 20 = 1; 420 = 1; 100 = 1

• Conversions to scientific notation:

42 = 42 x 100 = 4.2 x 101

943 = 943 x 100 = 9.43 x 102

-0.00036 = -0.00036 x 100 = -3.6 x 10-4

0.0093 = 0.0093 x 100 = 9.3 x 10-3

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MULTIPLYING AND DIVIDING POWERS OF 10

• When multiplying exponential, add exponent

103x102 = 105; 10-5 x10-2 = 10-7; 105 x10-2 = 103

• When dividing exponential, subtract exponent

104/102 = 102; 10-5/102 = 10-7; 10-4/10-2 = 10-2

• When you take reciprocal of an exponential, change the sign of exponent

1/103 = 10-3; 10-3/105x10-2 = 10-6

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Multiplying and Dividing in Exponential Notation

• Do number math using number rules and exponential math using exponential rules. Then combine the two parts.

Example: (2x103)(4x1023) = 8x1026

2x4 = 8

103x1023 = 1026

Also memorize: ½ = 0.5; 1/3 = 0.33; ¼ = 0.25; 1/5 = 0.2; 2/3 = 0.67, ¾ = 0.75; 1/8 = 0.125

© 2012 pearson education, inc. chapter 1 introduction: matter and measurement dr. subhash c. goel...

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