Chapter 1

• Reason for studying chemistry:

• Essentially effects all aspects of our lives.

• Study helps us to understand this and make

rational decisions for today and for the future.

• We will learn some basic chemistry principles

so that we can talk with knowledge later in the

semester about relevant topics for our work

and our lives

Basic Research: The concepts that form the foundation of our scientific knowledge. This is essentially discovered through pure research, usually without a specific end goal in mind.

Applied Research: Research towards the solution of a particular problem. In industry today, most research is applied research.

Technology: Using basic science and/or applied science to produce products of practical application, frequently for mass production.

Scientific Method:(summarizes how scientists work)

1. Observations

2.Hypotheses

3. Experimental Testing

4. Theories

2a. Laws

• Chemistry is the study of the structure and behavior of

matter.

• Matter is anything that occupies space and has mass.

Mass does not change with location. Weight depends

on gravity.

• Chemists study matter in quantities large enough to be

seen and handled in an ordinary manner (which we call

macroscopic), or so small to be only seen through a

microscope (which we call microscopic).

• Finally, chemistry did not become a modern science until chemists realized there were particles that could not even be seen through microscopes (submicroscopic), and studied their effects on how matter behaves.

Chemists study matter by studying its identity and behavior. In doing so, they look at the properties of matter. 2 Main Classifications:

1. Physical – properties (or characteristics)

that can be studied without changing the

substance into a new substance or trying to

do so. Examples are color, physical state

(solid, liquid or gas) at specific temperatures,

and something called density

• Density is defined as the mass (or weight) in a

specific volume of space for that substance.

• For example, water has a mass of 1.0 grams in

every 1.0 milliliter. This is a much better way to

identify a substance than using mass or volume

separately. We will look at this in one of our

experiments.

• Usually reported as grams/ milliliter.

Density is defined as Mass divided by the volume of a sample. Mathematically that is:

D = M / V

Besides helping identify substances, density tells us whether something will float or sink in a liquid.

1

1.5

2

D=1 g/mL

D=1.5 g/mL

D=2 g/mL

Also, we can use that idea to determine the density of any solid object. What are the approximate densities of the green and red object?

When we study physical properties we frequently do

change the substance, but only its appearance not its

identity. For example, the temperature a substance freezes

at is called its Freezing Point or FP. To measure FP, we

simply have to freeze a liquid sample of the substance and

measure the temperature at which this occurs. We have

only changed the state from liquid to solid, but the

substance is the same. (Think of ice and water). These

types of changes are called Physical Changes.

2. Chemical: Properties that can only be studied by changing or trying to change the substance into a new, different substance.

• One example is flammability. Even if a substance does not burn, this is still a chemical property. In studying chemical properties we perform Chemical Changes on the substances.

Reactants Products for example:

Hydrogen + Oxygen water

There are 3 forms or phases of Matter:

1. Solid – Holds its shape and size no matter what container it is in.

2. Liquid – Holds its size but takes shape of its container.

3. Gas – Takes shape & size of its container.

Any substance can be (under the proper conditions) a solid, liquid or gas. Think of ice, water and steam. They are all different states for the substance water.

At the macroscopic level: 3 Types of matter:

1. Elements – Basic substances in nature. Cannot be decomposed into simpler substances.

2. Compounds – 2 or more elements, chemically combined together in a specific way, and always the same way. Can only be chemically decomposed into other smaller compounds or elements. Compounds have completely different properties than the elements that they contain.

a) Pure Substance – (your book simply uses the term substance) Any pure element or compound

There are 118 known elements (26 of which are so called man-made elements. Each element has a 1 or 2 letter symbol to represent it. A chemist needs to know the names that go with the symbols and vice-versa. We will learn some of them in this course.

Compounds also have symbols based on the symbols

of the elements composing the compound. Examples:

NaCl, NH4NO3, Mg(OH)2. These are called molecular

formulas. Later in the semester we will learn about

other ways to represent the formulas of compounds.

3. Mixtures – 2 or more pure substances physically

mixed together in any way and not always the same

way. Frequently easy to separate into its components.

The components retain some or all of their own

properties.

At the submicroscopic level: 2 Particles: (for now)

1. Atoms – Smallest particle identifiable as an element.

2. Molecules – 2 or more atoms combined together chemically. Smallest particle that can be identified as a compound.

Mixtures can further be divided into 2 types:

1. Homogeneous – All parts of the mixture are identical, even at the microscopic level. All homogeneous mixtures are also called solutions. Some common examples are clean air, sugar water, brass, soda water, Scotch.

a) Also, all elements and compounds, by definition, are homogeneous.

2. Heterogeneous – Different parts have different

composition. In many cases these different parts can be

obviously seen ( cinnamon & sugar, basket of different

fruit, metal ores), and sometimes only under special

conditions (dusty air, milk, blood)

• All Chemical & some physical changes also involve energy changes. Energy can either be absorbed or released during these processes.

• If it is released, this is called an exothermic reaction.

• If the energy is absorbed, it is called an endothermic reaction. Energy is the ability to do work (move objects).

• 3 Types for now:1. Potential – stored energy for later use (from position or chemical make-up)

2. Kinetic – Energy of motion.

3. Heat

• Energy is frequently converted from one type to another. Examples: Burning oil to produce heat (chemical potential heat energy) or turning on a light switch (electrical light energy).

• Temperature is the measure of intensity (how hot or cold) of heat.

• Besides studying the identity and behavior of matter in a

descriptive sense (qualitative analyses), chemistry is

also very much of a quantitative science, involving many

measurements and calculations based on these

measurements.

• For measurements all scientists use the Metric System

rather than the English system. (actually the International

System (SI)) rather than the English system). For our

purposes we will consider the classical Metric System.

• There are many advantages. A few of these are:

1. Simpler - Many fewer words and meanings to learn

2. More logical – Like our money system, the Metric system is based on powers of ten

3. Easier calculations – Frequently only need to move a decimal point.

BASIC UNITS

• Name Abbrev Type of ApproximateMeasurement Eng Equivalent

• meter m length 39.36 inches (about 1 yard)

• liter L volume 1.06 quarts

• gram g mass 0.035 oz (about 1/30 oz)

METRIC SYSTEM

• Prefix Abbreviation Value

• giga G 1 billion times (1 x 109 )

• mega M 1 million times (1 x 106 )

• kilo k 1 thousand times (1 x 103 )

• deci d one-tenth (1/10)

centi c one-hundredth ( 1/100 )

milli m one-thousandth ( 1/ 1000 )

micro one-millionth ( 1 / 10 6 )

nano n one-billionth ( 1 / 10 9 )

pico p one-trillionth ( 1 / 10 12 )

Finally, time is measured the same way in all systems. The basic unit is the second (abbreviated as s)

Let’s practice some conversions:

1. If converting from a large unit to a smaller unit, move the decimal point to the right. Fill in any empty spaces with zeroes. The number of decimal places to move equals the number of zeroes in the prefix.

2. If converting from a smaller unit to a larger unit, move the decimal point to the left and continue as above.

Practice

1. Convert 5.36 m into cm

2. Convert 3468 nm into m

3. Convert 5.92km to cm

An important relationship that you need to know is that:

1mL = 1 cm3 = 1 cc = 1 cubic centimeter

• Frequently scientist have very large or very small

numbers to deal with.

• To make it easier they use scientific notation.

123000000000000. g becomes 1.23 x 1014 g, while

0.0000000000538 mm becomes 5.38 X 10-11 mm.

• For large #’s, the positive exponent indicates the # of

decimal positions after the first non-zero digit. For very

small #’s, the negative exponent indicates the number of

decimal positions from the decimal point to the right of

the first non-zero number (going from left to right).

You will need to know how to;

Convert from decimal to scientific notation

Convert from scientific to decimal notation

Let’s practice some.