CHAPTER 1Matter & Measurement

General, Organic, & Biological Chemistry

Janice Gorzynski Smith

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CHAPTER 1: Matter & Measurement

Smith. General Organic & Biolocial Chemistry 2nd Ed.

Learning Objectives: Definition of matter

Solids, liquids, and gases

Physical vs chemical properties and changes

Pure substances: Elements & Compounds

Mixtures: Heterogeneous vs Homogeneous

Units of the metric system & common prefixes

Measured vs exact numbers

Significant figures: identify & use in calculations

Scientific Notation

Conversion factors for calculations to cancel units

The three temperature scales

Density and Specific Gravity

Matter Definition

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Matter is anything that has mass and takes up volume.

Naturally occurring:• cotton• sand• digoxin, a cardiac drug

Synthetic (human-made):• nylon• Styrofoam• ibuprofen

Matter Solids, Liquids, Gases

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The Solid State:• A solid has a definite

volume.

• It maintains its shape regardless of its container.

• Solid particles lie close together in a regular pattern.

The Liquid State:• A liquid has a

definite volume.

• It takes the shape of its container.

• Liquid particles are close together but can move past one another.

The Gas State:

• A gas has no definite shape; it assumes the shape of its container.

• It has no definite volume; it assumes the volume of its container.

• Gas particles are very far apart and move around randomly.

Matter Physical Properties

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Physical properties can be observed or measured without changing the composition of the material.

•boiling point

•melting point

•solubility

•color

•odor

•state of matter

Matter Chemical Properties

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Chemical properties determine how a substance can be converted into another substance.

Chemical change is the chemical reaction that converts one substance into another (Chapters 5 and 6).

Matter Pure Substances: Elements

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• A pure substance is composed of only a single component (atom or molecule).

• It has a constant composition, regardless ofsample size or origin of sample.

• It cannot be broken down to other pure substances by a physical change.

Pure Substances

An element is a pure substance that cannot be broken down by a chemical change.

aluminum metal (Al)

Matter Pure Substances: Compounds

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A compound is a pure substance formed by chemically joining two or more elements.

table salt (NaCl)

Matter Mixtures

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Mixtures

• Mixtures are composed of more than onecomponent.

• They can have varying composition (anycombination of solid, liquid, and gas).

• Mixtures can be separated into their componentsby a physical process.

All matter can be classified as either a pure substanceor a mixture.

Matter Mixtures: Heterogeneous & Homogeneous

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Homogeneous Mixture

Example: simple syrup

HeterogeneousMixture

Example: vinaigrette

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Matter Definition

http://ridenourmhs.wikispaces.com/ESUnit2

Measurements Metric System

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Each type of measurement has a base unit in themetric system.

Measurements Common Prefixes

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The prefix of the unit name indicates if the unit is larger or smaller than the base unit.

Measurements Common Prefixes

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1 kilometer (km) = 1,000 meters (m) 1 km = 1,000 m

1 millimeter (mm) = 0.001 meters (m)1 mm = 0.001 m

1 centimeter (cm) = 0.01 meters (m)1 cm = 0.01 m

The base unit of length is the meter (m).

Measurements Common Prefixes

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Mass is a measure of the amount of matter in an object.

Weight is the force that matter feels due to gravity.

1 kilogram (kg) = 1,000 grams (g)1 kg = 1,000 g

1 milligram (mg) = 0.001 grams (g)1 mg = 0.001 g

The base unit of mass is the gram (g).

Measurements Common Prefixes

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1 kiloliter (kL) = 1,000 liters (L) 1 kL = 1,000 L

1 milliliter (mL) = 0.001 liters (L) 1 mL = 0.001 L

Volume = Length x Width x Height= cm x cm x cm= cm3

1 mL = 1 cm3 = 1 cc

The base unit of volume is the liter (L).

Measurements Units

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Measurements Exact Numbers

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An exact number results from counting objects or is part of a definition.

•10 fingers •10 toes•1 meter = 100 centimeters

An inexact number results from a measurement or observation and contains some uncertainty.

•15.3 cm•1000.8 g•0.0034 mL

Measurements Significant Figures

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Significant figures are all the digits in a measured number including one estimated digit.

All nonzero digits are always significant.

3 sig. figures 6 sig. figures

65.2 g 255.345 g 65.2 g 255.345 g

Measurements Significant Figures

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3 sig. figures

Rule 1: A zero counts as a significant figure when it occurs:

•between two nonzero digits

5 sig. figures

•at the end of a number with a decimal place

29.05 g4 sig. figures

1.0087 mL29.05 g 1.0087 mL

3.7500 cm

5 sig. figures

620. lb3.7500 cm 620. lb

Measurements Significant Figures

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Rule 2: A zero does not count as a significant figure when it occurs:

5 sig. figures

•at the beginning of a number

1 sig. figure

•at the end of a number that does not have a decimal

0.00245 mg 0.008 mL3 sig. figures

0.00245 mg 0.008 mL

2570 m 1245500 m

3 sig. figures

2570 m 1245500 m

Measurements Significant Figures: Multiplication & Division

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351.2 miles

5.5 hour= 63.854545 miles

hour351.2 miles

5.5 hour

4 sig. figures

2 sig. figures Answer must have2 sig. figures.

Multiplication/Division Rules: The answer has the same number of significant figures as the original number with the fewest significant figures.

Measurements Significant Figures: Multiplication & Division

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63.854545 miles

to be retained to be dropped

first digit to be dropped

hour63.854545

If the first digit to be dropped is: Then:

•between 0 and 4 •drop it and all remaining digits

•between 5 and 9 •round up the last digit to be retained by adding 1

=

2 sig. figuresAnswer

64 mileshour

Measurements Significant Figures:Addition & Subtraction

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Addition/Subtraction Rules: The answer has the same number of decimal places as the original number with the fewest decimal places.

10.11 kg

3.6 kg

6.51 kg

10.11 kg

3.6 kg

2 decimal places

1 decimal place

answer must have1 decimal place

= 6.5 kg final answer 1 decimal place

Measurements Scientific Notation

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Coefficient:A number between1 and 10.

y x 10x Exponent:Any positive or negativewhole number.

In scientific notation, a number is written as:

Measurements Scientific Notation

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•When the exponent x is positive, move the decimal point x places to the right.

2.80 x 10–2 =

2.800 x 102 =

•When the exponent x is negative, move the decimal point x places to the left.

280.0

0.0280

Measurements Conversion Factors

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• Conversion factor: A term that converts a quantity in one unit to a quantity in another unit.

• Conversion factors are usually written as equalities.

2.21 lb = 1 kg

• To use them, they must be written as fractions.

original quantity

conversion factordesired quantityx =

2.21 lb1 kg

or 1 kg2.21 lb

Measurements Conversion Factors

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•make sure all unwanted units cancel

Factor-label method: Using conversion factors toconvert a quantity in one unit to a quantity in another unit.

•units are treated like numbers

To convert 130 lb into kilograms:

130 lb x conversion factor = ? kg

original quantity

desired quantity

Measurements Conversion Factors

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or

2.21 lb1 kg

1 kg2.21 lb

130 lb x

= 59 kg

Answer2 sig. figures

•The bottom conversion factor has the original unit in the denominator.

•The unwanted unit lb cancels.

•The desired unit kg does not cancel.

Measurements Temperature

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Temperature is a measure of how hot or cold an object is.

1. Degrees Fahrenheit (oF)2. Degrees Celsius (oC)3. Kelvin (K)

• Three temperature scales are used:

To convert from oC to oF: To convert from oF to oC:

oF = 1.8(oC) + 32oC = oF − 32 1.8

To convert from oC to K:

K = oC + 273 oC = K − 273

To convert from K to oC:

Measurements Temperature

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Measurements Density

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density =mass (g)

volume (mL or cc)

Density: A physical property that relates the mass of a substance to its volume.

To convert volume (mL) to mass (g):

To convert mass (g) to volume (mL):

mL xg

mL= g g x

mLg

= mL

density inverse of density

Measurements Specific Gravity

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Specific gravity: A quantity that compares the density of a substance with the density of water at thesame temperature.

specific gravity = density of a substance (g/mL)density of water (g/mL)

•The units of the numerator (g/mL) cancel the units of the denominator (g/mL).

•The specific gravity of a substance is equal to its density, but contains no units.