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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.