1. GENERAL CHEMISTRY 1

2. Chapter 1

3. Chemistry Definition – Study of structure and interaction of matter, including energy changes. Will discuss energy in a later chapter. Matter is anything that has mass and occupies space.

4. Scientific Method: Systematic approach to scientific work Observation Hypothesis – Attempt to explain why Law – Summary of observation without trying to explain why Experimentation – Test hypotheses Theory – Hypothesis that stands the test of time

5. Matter: 3 Types: Elements – Basic substances – Can’t be decomposed to a simpler substance. Only 115 known elements. Compounds – 2 or more elements chemically combined together in a fixed, definite proportion by weight. Can only be decomposed into its component elements by a chemical reaction. Elements and compounds are called Pure Substances.

6. Mixtures – 2 or more pure substances physically mixed together in any proportions whatsoever. Can be decomposed, frequently fairly easily, using physical methods, into its components. Each component retains its own properties. 2 Types: Homogeneous – All parts identical Heterogeneous – Non-uniform composition

7. One other quick definition is for an ion – a charged particle. Positively charged particles are specifically called cations, while negatively charged ones are specifically called anions. We will discuss these in much more detail later on.

8. 3 States of Matter: Solid – Keeps its shape and size Liquid – Keeps its size but takes the shape of its container Gas – Takes the size and shape of its container. Any substance can be a solid, liquid or gas, depending on the conditions.

9. All substances can be described by their characteristics or properties. Physical Property – Can be studied without changing the substance into a new substance. Chemical Property- Can only be studied by changing (or attempting to change) the substance into a new substance.

10. We study these properties by observing physical & chemical changes. Physical Changes are changes in the appearance of a substance, but not in its identity. Chemical changes are actual changes in the identity of the substance. A new substance or substances is formed.

11. Physical Sciences require frequent measurements. The official measurement system is called the International System or SI, which defines basic units for measuring various quantities and derived units from them for other quantities. You can check in your book for these on page 16. The SI system replaced the older Metric System, which is very similar and only differs in minor areas of definition. I will explain the metric system.

14. Basic units & abbreviations:

15. The scientific temperature scale is called the Celcius (  C) scale as opposed to the American scale of Fahrenheit (  F). A third scale, also commonly used in science is called the Absolute or Kelvin scale (F). They all measure the same thing, intensity of heat, just using different units.

16. They can be converted into each other. The most common formulas are:  C = (  F – 32) / 1.8  F =  C x 1.8 + 32 K =  C + 273.15

17. Another method:  C = (  F + 40) / 1.8 – 40  F = (  C + 40)1.8 – 40

18. Accuracy vs. Precision Accuracy describes how close a measurement is to the correct value. We usually don’t know how accurate we are (otherwise we wouldn’t be taking the measurement) Precision describes how close measurements are to each other. This depends on the quality of the measuring instrument. This gives a reasonable idea as to how accurate our measurement is. Good precision usually, but not always, indicates good accuracy. If we keep getting the same measurement, then it most likely is accurate.

19. Significant Figures: Help indicate the precision of the measuring instrument. Any measurement cannot be more precise than the measuring instrument allows.

20. Assume the above picture is a ruler measuring in cm. What is the measurement at the arrow? More than 13 but less than 14. Not accurate to say 13 cm or 14 cm. We can do better. Mentally, we can divide the space between the smallest marks into 10 parts. In our case, these mental marks are tenths of cm. We then estimate how far along these mental marks our arrow is. What do you think here?

21. Our final measurement will be reported with one digit left of the decimal point (in the tenths position). We know we are making a guess, but it is a reasonable guess, but there still is some uncertainty, probably + or – 0.1cm. It is not considered reasonable to go any further ( our minds are not capable of dividing a small space into 100 parts, only into 10 parts). This last digit is called the last significant digit (digit means numeral).

22. We always measure until we have to make a guess, then we stop. This indicates to anybody reading our measurement without having the instrument in front of them, to know how precise it is. Or, in other words, what the smallest division on the device is.

23. In many calculations, need to know how many significant digits (or significant figures) are present in a measurement. All non-zero digits are always significant. The possible confusion lies with zeroes.

24. Zeroes in Sig Figs Leading zeroes - To the left of first non-zero digit. All are not significant. Middle zeroes - In between 2 significant digits. All are significant. Trailing zeroes - To the right of the last non- zero digit. Trailing zeroes to the right of the decimal point are significant. Trailing zeroes in whole numbers may or may not be significant. Only the measurer knows. The ambiguity can be removed by using scientific notation. We will discuss this shortly.

25. When calculations are done involving measurements, there are specific rules to follow for handling significant figures (From now on we will use S.F. to refer to significant figures)

26. S.F. in Calculations Multiplication and/or Division - The final answer will have the same number of significant digits (figures) as the measurement with the least number of significant digits. Exact values and counting units are considered to have an infinite number of significant digits.

27. Addition and/or Subtraction - The final answer will have the same number of decimal positions as the measurement with the least number of decimal positions. Again, exact values and counting units are considered to have an infinite number of significant digits. Mixed Calculations - Round off when you switch from one type of calculation to the other.

28. In all cases the final answer is rounded off to proper significant figures.

29. Scientific Notation. Makes use of the exponential powers of 10: for example: 2000 = 2 x 1000 = 2 x 10 3 We need to be able to write a number in normal decimal notation (1045.23) or in scientific notation ( 1.04523 x 10 3 ). We need to be able to go back and forth. My approach is to realize that for every unit of exponent we have to move the decimal point one place. If the exponent is decreasing, then the decimal point is moved to the right, while it is moved to the left if the exponent is increasing. For this approach remember that a normal decimal number, such as 156.2, is the same as 156.2 x 10 0

30. There is a difference between exponential notation and scientific notation. Exponential notation (frequently used in engineering applications, allows any # of digits to the left of the decimal point, while scientific notation allows only one non-zero digit to the left of the decimal point.

31. Scientific notation avoids S.F. ambiguity, because the non-exponential part of the number will contain only the S.F. of that number. (Remember: a number in scientific notation is still only one number, even though it looks like a multiplication of 2 numbers. Modern scientific calculators all can easily make use of scientific notation. Get familiar with your particular calculator.

32. NOTE: When doing calculations involving measurements, remember that almost all measurements in science have units, such as g, mL, km etc. Units are treated like numbers in multiplication and division, while in addition and subtraction, only like units can be added or subtracted. All answers must have correct units or it is not a correct answer.

33. Two important physical properties that you will be studying in the third experiment this semester are density and specific heat. Density measures the mass of a specific volume of the substance. It is essentially constant for the solid state and liquid state (slight variation in liquid state) and can be used to help identify a substance.

35. Specific Heat (SH) measures the amount of heat needed to raise the temperature of 1 gram of a substance by 1  C. Therefore the total heat involved in a temperature change can be calculated if we know the specific heat of the substance. Determining the SH is also useful in helping to identify a substance. The amount of heat is usually symbolized by the letter q.

36. q = (SH)(mass)(t final – t initial )

37. Dimensional Analysis Conversion factors Ratio of 2 measurements = 1.00000 Choose the conversion factor that cancels the appropriate unit or units. Also, if possible, make sure the conversion factors using book values has enough significant figures so that the number of sig. figs. from measured values is not decreased. Permanent & temporary

38. Do examples 1.6, 1.7 and 1.8 in text (pages 29-30). In most cases we will solve problems using this method, although at times I will use algebra instead. Either method is acceptable.