1. Chapter 1

2. 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

3. 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.

4. Scientific Method:(summarizes how scientists work) 1. Observations 2.Hypotheses 3. Experimental Testing 4. Theories 2a. Laws

5. 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).

6. 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.

7. 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

8. 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.

9. 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?

10. 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.

11. 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

12. 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.

13. 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

14. 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.

15. Compounds also have symbols based on the symbols of the elements composing the compound. Examples: NaCl, NH 4 NO 3, Mg(OH) 2. These are called molecular formulas. Later in the semester we will learn about other ways to represent the formulas of compounds.

16. 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.

17. 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.

18. 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.

19. 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)

20. 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

21. 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.

22. 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

23. 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.

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

27. 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.

28. Practice 1. Convert 5.36 m into cm

29. 2. Convert 3468 nm into m

30. 3. Convert 5.92km to cm An important relationship that you need to know is that: 1mL = 1 cm 3 = 1 cc = 1 cubic centimeter

31. 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 10 14 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).

32. You will need to know how to;  Convert from decimal to scientific notation  Convert from scientific to decimal notation Let’s practice some.