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CHAPTER ONE: CHEMISTRY: THE SCIENCE OF CHANGE

1.1 The Study of Chemistry

A. What is Chemistry?

Chemistry is -

-

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Chemistry is the central science.

Chemistry is a quantitative science based on experimentation.

- Quantitative- Based on the of something.

- Experiment An observation of natural phenomena carried out in a controlled manner so that the results can be duplicated and rational conclusions made.

B. Five Traditional Branches of Chemistry:

1. Organic chemistry = chemistry of covalent compounds of carbon and hydrogen and

their derivatives.

The chemistry of carbon

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2. Inorganic chemistry = the study of the huge variety of substances that fall outside the

realm of organic materials.

The chemistry of everything else

3. Analytical chemistry = the identification of substances present and their amounts in a

sample.

4. Physical chemistry = applies methods of physics to the properties of matter and the

accompanying energy changes.

5. Biochemistry = the study of the chemistry of processes in living organisms.

1.2 Classification of Matter

Matter =

Broadly classified by two categories:

1. Substances = a form of matter that has a definite (constant) composition and distinct properties, such as color, smell, and taste.

a. Elements -

b. Compounds -

Properties of the individual components are lost.

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2. Mixtures

Properties of the individual components are maintained.

A mixture which is not uniform throughout is called - hetro means not the same

Mixture that has uniform properties throughout is a - homo means same - Also called a solution.

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A little quiz - categorize the following samples of matter as: A) heterogeneous mixture,

B) homogeneous mixture,

C) compound, or

D) element:

diamond / sea water (drawn from surf) / limestone / gravy / steel / table salt / sugar

Three States of Matter.

1. Solids:

2. Liquids:

3. Gases:

Solid Liquid Gas

Shape Definite Variable Variable

Volume Definite Definite Variable

Fluid No Yes Yes

Compressible No No Yes

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1.3 The Properties of Matter

Chemical properties =

Chemical change (reaction) - change in which one or more kinds of matter are transformed into one or multiple kind(s) of new matter.

Example: 2 Mg (s) + O2 (g) 2 MgO (s)

Example: Cu (s) + 4 HNO3 (aq) Cu(NO3)2 (aq) + 2 NO2 (g) + 2 H2O (l)

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Physical properties =

Can be observed without changing its chemical identity.

For example: color, density, hardness, melting point, boiling point, and electrical and thermal conductivities.

Physical change- = a change in the form of matter but not in its chemical identity.

No change in chemical composition.

Physical properties are usually altered.

Energy may be absorbed or released.

Physical properties can be further classified according to whether or not they depend on the amount of sample present.

Extensive properties -

Volume, energy, mass, etc.

Extensive properties with the same units can be added to one another.

Intensive properties

Density, temperature, etc.

Intensive properties cannot be added to one another.

All chemical properties are intensive properties.

Classification of Changes:

o Chemical Change One or more substances are used up (at least partially). One or more new substances are formed. Energy may be absorbed or released.

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Base Number

o Physical Change No change in chemical composition (no new substances formed). Physical properties are usually altered significantly. Energy may be absorbed or released. Examples =

Distillation:

1.4 Scientific Measurement

Scientific Notation o Some numbers are so big or small that they are cumbersome to write out. o For example:

Mass of earth (kg) = 5,973,600,000,000,000,000,000,000 Plancks Constant (Js) = 0.0000000000000000000000000000000006626

o In scientific notation, all numbers are written in the form of a x 10b where the coefficient a is any real number 1 |a| < 10 (referred to as the base

number) and b is an integer.

o Thus, Mass of earth (kg) = 5.9736 x 1024 Plancks Constant (Js) = 6.626 x 10-34

o Scientific notation makes these much easier to look at, compare, and use.

5.9736 x 1024 Exponent

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o Converting to scientific notation:

Find base number and set up notation. o Find base number by moving decimal place behind first non-zero digit:

Speed of light = 299792500 m/s 2.997925 x 10x

Move the decimal to get to the original number.

2.99792500 299792500

The number of moves becomes the exponent. o If moving to the right, o If moving to the left,

2.997925 x 108 m/s

o For each of the following pairs, determine which number is larger: 9.11 x 10-31 or 8.11 x 10-31 6.022 x 1023 or 6.022 x 1020 6.63 x 10-34 or 6.63 x 10-31

SI Units o International System of Units (SI) A particular choice of metric units.

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o Derived units are those resulting from a combination of a combination of any of the above base units.

o Metric and SI systems are decimal systems; prefixes used to indicate powers of ten.

Common Units o Mass- Measure of the amount of matter in a sample or object.

Weight- force exerted by an object or sample due to gravity. o Basic SI unit =

Generally use the gram (g) Conversion: 1 pound (lb) = 0.4536 kg = 453.6 g. Origin: 1 gram = mass of 1 cm3 of liquid water at 4.0 Celsius, pressure = 1

atm.

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o Length. Basic SI unit = Non SI unit = ngstrom

1 ngstrom = 0.1 nm = 1 x 10-10 m ~ size of atoms 1 in = 2.54 cm (exactly); 1 m = 39.370 inches (in); Origin: 10,000,000 meters = distance from N Pole to Equator through Paris,

France.

o Temperature Is a measure of the average kinetic energy of a substance. It can also be thought of as the An property. Basic SI unit

Called absolute temperature because 0 K is the lowest possible temperature (absolute zero).

Celsius (C) scale (named after Ander Celsius): 0C = freezing point of water at sea level. 100C = boiling point of water at sea level.

Fahrenheit (F) scale (named after Daniel Fahrenheit) 0F = freezing point of ammonium chloride/ice mixture. 100F = average human body temperature.

o Temperature conversion equations (Section 1.6):

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o Scientists predominantly use the Celsius scale.

o Since 100 Celsius degrees equal 100 Kelvins. o .the magnitude of the degree in both scales is equal. o 0C = 273.15 K, so:

o Since 100 Celsius degrees equal 180 degrees Fahrenheit o the magnitude of the Celsius degree is larger than the Fahrenheit

degree.

o 0C = 32F, so:

o Volume (Section 1.4)

A derived unit. Basic SI unit = Chemists use liters (L) or milliliters (mL) in metric system. One liter (1 L) is 1 cubic decimeter (1 dm3), or 1000 cubic centimeters (1000

cm3).

One milliliter (1 mL) is 1 cm3. Non-metric conversion: 1 L = 1.057 qt.

o Density (Section 1.4) o A derived unit. o Density = Mass per unit volume.

o For solids and liquids, usually expressed as g/mL, which is the same as g/cm3. o Water has a density of 1 g/mL at 4 C.

o 0.998 g/mL at 20 C. o Problem: 3.0 mL of liquid Hg has a mass of 40.8 g. What is its density in

g/mL?

o Problem: If reaction uses up 10.0 g ethanol, what volume of ethanol is used up? Density of ethanol is 0.789 g/mL.

d = m

V

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o Specific gravity is density of a sample relative to that of water =

Sp. Gr. = dsubstance / dwater

where dwater = 1.00 g/mL over wide range of T.

Exercise: What is the specific gravity of Hg?

1.5 Precision and Accuracy

Precision

Accuracy

Significant Figures (sig figs) those digits in a measured number (or in the result of a calculation with measured numbers) that include all certain digits plus a final

digit having some uncertainty.

The meaningful digits are in a reported number.

Good accuracy Poor precision

Good precision Poor accuracy

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Sig figs express the to which a measured number has been determined.

The more sig figs, the more precise the measurement/ number. For glassware, more (further) graduations give more precise

measurements.

o Rules for sig figs:

All non-zero digits are significant.

All embedded zeroes are significant.

Leading zeroes (to the left of the first non-zero digit) are not significant.

All terminal zeroes on the R of the decimal are significant.

Terminal zeroes at the end of the number written without a decimal point are

ambiguous (you dont know what the person intended).

340 cm you dont know whether the person actually measured length to the nearest cm and it came out exactly 0, or whether their number is an

estimate of length rounded to the nearest 10 cm.

For scientific purposes, if they put in a decimal point, the ambiguity is removed:

340. cm implies knowledge to the nearest cm.

The ambiguity can always be removed by re-expressing such numbers in scientific

notation.

30

20

22?

24

23

23.4?

24

23

23.39

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3.40 x 102 indicates knowledge to the nearest cm.

3.4 x 102 indicates knowledge only to the nearest 10 cm.

o Convert all of the following measurements to scientific notation and indicate the number of significant figures:

o 1,207 o 0.082057 o 690.

o Sig figs in calculations:

Multiplication and division The product or quotient should have the same number of sig figs as the number with the least number of sig figs.

Addition and subtraction - answer is only known to the same place as the

measurement with the least number of places.

Logarithms Answer will have as many decimal places as there were significant figures in the original measurement.

Exact numbers can be considered to have an infinite number of significant figures.

o Exact numbers are those that can be counted, or are defined as having a certain value.

When performing a series of calculations, only round values at the very end of the

calculation.

Perform the following calculations and report the answer with the proper number

of significant figures:

20.01 1.08 + 9.750 10.21

(35.683 - 35.324)

55.85

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1.6 Dimensional Analysis (Factor-Label Method)

Method of calculation in which one carries along the units for quantities Main idea: Exploit the units to your advantage to help you solve problems. Utilized conversion factors.

Conversion factor Factor equal to 1 that converts a quantity expressed in one unit to a quantity expressed in another unit.

Must always have more than one unit!

Lets use a typical example:

1 ft = 12 in.

Rearrange to give:

Use the conversion factor that allows the units to cancel. o How many inches are in 5.5 feet?

Express 1.10 in inches. 2.54 cm = 1 in.

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It is generally useful to approximate answers mentally before entering the numbers into a calculator.

o When multiplying powers of ten, the exponents are

o When dividing powers of ten, the exponents are

Dimensional Analysis Example Question: How many days have you been alive? How many hours? How many minutes? How many seconds? Express you lifespan in seconds using the most appropriate metric prefix.

Dimensional Analysis Example Question II: The distance from Cookeville to Nashville is about 80 miles.

How many minutes will it take to reach Nashville traveling at the speed limit (70 miles per hour)?

How many minutes will it take to reach Nashville traveling at 80 miles per hour?

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Dimensional Analysis Example Question III: Osmim, the most dense chemical element known, has a density of 22.59 g/mL.

What is the mass of a cube of osmium the same size as a Rubiks Cube? A standard Rubiks Cube measures about 2.25 inches per side. 2.54 = 1 in.

Dimensional Analysis Example Question IV: Express the density of mercury in lb/ft3. The density of mercury is 13.59

g/mL.

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Percentage

Per out of; cent 100 Essentially just a fraction.

% Component A = (amount A in mixture) x 100%.

(total amount of mixture)

Example: Class has 80 men and 70 women. What % are men?

Example: If the same percentage of men in the class (from the previous example problem) applies to the entire student body, and there are 12,000 students, how many

women are enrolled?

Hard problem involving Sp. Gr., Volume, and %:

A certain water/ethanol mixture is 70% ethanol by mass, and has a specific gravity of

0.95. Calculate the mass of pure ethanol in a 4.000 Liter bottle of this mixture.

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Chapter 1 Summary:

Classification of Matter Elements, substances, hetero/homogeneous mixtures. Phases of matter.

Physical/chemical processes & properties. Extensive vs. intensive properties.

Units SI units

Dimensional Analysis Unit conversions Significant figures

NOTES: