CHEM 101

1. CHEM 101 Chapter 1 Matter and Measurement Virginia State University Summer 2008 Dr. Victor Vilchiz 5/19/08

2. Physical Properties • There are two types of physical properties: • Extensive Properties • Intensive Properties • Extensive Properties are those that depend on the amount of substance present (mass, volume). • Intensive Properties are independent of substance amount (color, taste, density).

3. Density • It has been mentioned several times… but what is density? • Density is the distribution of something over a given amount of space. • Mass density: given amount of mass in a given volume. • Charge density: amount of charge in a given space. • Particle density: amount of ………….

4. Observations and Measurement • Last time we said that if you analyze something w/o any measurements all you will get is qualitative information. • Qualitative: tells you the quality (what is present) but it does not tell you % • If you start making measurements then the information obtained is quantitative. • Quantitative: tells you the quantity of components in the sample.

5. Measurements • Before we continue talking about measurements we must agree on several things… • Units • Standards • References • In reality all three are the same thing. • There are several unit systems in use but we will concentrate in one and only one (SI)

6. SI Units • First things first… it is not the SI system… that’s redundant. • SI = Systeme Internationale • Established in the 1960’s by the world’s scientific community. • Based on the metric system • Until the late 1990’s two countries did not use this system • US and Liberia • We are now the only ones, but we are slowly changing.

7. The Fundamental SI Units

8. Common Derived Units

9. Temperature Scales • There are many temperature scales but 3 of them are the most important. • Fahrenheit • Celcius/Centigrade • Kelvin • We will discuss their origins and points of reference.

10. Fahrenheit • Proposed in the early 1700’s by Gabriel Fahrenheit • It is believed that he used 4 reference points to build his scale • 0°F Lowest temperature attained by a mixture of salt/ice/water • 30°F freezing point of water • 100°F human body temperature • 240°F boiling temperature of water

11. Fahrenheit • Since his proposal all these temperatures have changed. • 0° was discarded since it did not specify how to get the mixture (bad science); some sources claim (1:1:1 ammonium chloride, ice, water) • 30°32° we now know water freezes at 32°F (not bad)

12. Fahrenheit • 100°  98.6°F (it is believed that he did not really reported a human’s temperature but the temperature of a pig’s rectum.) • 240°212°F (big difference) however, considering that this had to be an extrapolation since during his time thermometers were made using alcohol which boils before water it was a good estimate.

13. Centigrade/Celsius • First proposed by Anders Celsius • Developed by using a rudimentary thermometer (column filled with mercury) • The column is brought into contact with a mixture of water and ice and the height of the liquid mercury is marked. • Column brought into contact with boiling water and height of mercury column is marked.

14. Celsius • Lower mark is set to be 0 and higher mark 100. • The column is divided in 100 degrees (grades) • Hence the scale was originally referred to as CentiGrade (100 degrees) • Named changed to avoid confusion with other centigrade (100th of a degree).

15. Kelvin • Also known as the absolute temperature scaled • Proposed by Lord Kelvin. • The scale has “absolute zero” as its zero point. • Absolute zero: in this scale there can’t be negative temperatures. • The scale was derived from P, T work done in gases. • Kelvin observed that if he extrapolated his results to P=0 he always got the same temperature. This is now defined as absolute zero and/or -273.15°C.

16. Comparison of temperature scales.

17. Temperature Conversions • The conversion from Celsius to Kelvin is simple since the two scales are simply offset by 273.15° • The conversion of Fahrenheit to Celsius, and vice versa, can be accomplished with the following formulas.

18. Temperature vs Heat • Temperature is a relative measurement, remember how the Centigrade scale was developed. An object is hotter/colder than a reference point. • Heat is energy available to flow from system to surroundings or vice versa

19. Metric System • The system is based on factors of 10 • The basic unit (besides the kilogram) has a “factor” of 100 =1 • You need to become familiar with the following prefixes: • Giga, Mega, kilo, deci, centi, milli, micro, nano, and pico.

20. Metric Prefixes

21. Common Non-SI Units • There are three common non-SI units used in scientific measurements • Mass: while the SI Unit is the kg… we usually report mass in grams. • Volume: The derived unit of Volume is the m3 yet in measurements we vary the unit depending on the substance… liquids are usually measured in mL, gases in L • 1000L=1m3 and 1mL=1cm3 • Temperature: we report temperatures in °C

22. Making Measurements • Now that reference points (units) have been agreed on we can proceed to actual measurements. • So you are making a measurement… what should you worry about... • There are two things that you should pay close attention to. • Reliability • Reproducibility

23. Reliability • Reliability • Can we trust you to make good measurements? • Are your values close to the true values? • Can you hit the Bull’s Eye? • AKA Accuracy

24. Reproducibility • Reproducibility • Can we trust you to get the same result every time? • Did you get lucky? • Are your values close to the average? • Can you hit the same spot over and over? • AKA Precision • Note: some books will tell you that you cannot be accurate if you are not precise… that is false!!! It is possible to hit the 4 corners of a square and you average will be the center (accurate)

25. Measurement and Significant Figures • The first panel has no precision but may be accurate • The second panel shows precision but no accuracy • The third panel shows both precision and accuracy

26. Errors • Since you are not 100% precise there will always be an error. % error = (mv - tv)/tv *100 mv = measured value tv = true value • No matter how careful you are there is always some error. • There are two types of errors • Systematic • Random

27. Uncertainty • Uncertainty: this is the degree of “unknown” or unreliability of your measurement. • To minimize this unreliability you need two things… • Practice in making measurements • Precise equipment… the more precise the more expensive the instrument will be.

28. Systematic Error • When something goes wrong we tend to blame the instrument we are using. • At times it is the machine that is wrong • The error can be fixed by recalibrating the instrument. • Example: You use a mass balance and when you check the mass of your standards no matter which one you use you seem to be off by 10grams on the heavy side. • You recalibrate by adjusting your machine by 10 grams or subtracting 10 from the measured value.

29. Random Error • This is just part of life and no matter what you do you cannot remove it from your experiment. • The only thing you can do is to perform many trials and average out the error. • While in systematic error the measure value is always either too high or too low… on random error at times it is low and at times it is high and the value will vary.

30. Reporting Values • Ok so you are careful, reliable, and can reproduce measurements… NOW WHAT? • You need to know how to report numbers. • At times it is best to report numbers in scientific notation… a good rule of thumb is.. If its bigger than 16 use Sci Notation… if it is smaller than 0.1 use Sci notation.

31. Measurement and Significant Figures • To indicate the precision of a measured number (or result of calculations on measured numbers), we often use the concept of significant figures. • Significant figures are those digits in a measured number (or result of the calculation with a measured number) that include all certain digits plus a final one having some uncertainty.

32. Measurement and Significant Figures • All measurements, no matter how careful they are taken, have an error, uncertainty, associated with them.

33. Measurement and Significant Figures • Number of significant figures refers to the number of digits reported for the value of a measured or calculated quantity, indicating the precision of the value. • To count the number of significant figures in a measurement, observe the following rules:

34. Measurement and Significant Figures • All nonzero digits are significant. • Zeros between significant figures are significant. • Zeros preceding the first nonzero digit are not significant. • Zeros to the right of the decimal after a nonzero digit are significant. • Zeros at the end of a nondecimal number may or may not be significant. (Use scientific notation.)

35. Significant Figures 0.0486 has 3 sig figs. • 3456has • 4sig figs. • 16.07has • 4sig figs. • 1.300has • 4sig figs. • 1310has • 3sig figs. • 1310.has • 4sig figs.

36. Measurement and Significant Figures • When multiplying and dividing measured quantities, give as many significant figures as the least found in the measurements used. • When adding or subtracting measured quantities, give the same number of decimals as the least found in the measurements used.

37. Rounding • It is good practice to carry 2 extra significant figures through the calculation and only round at the end • 1st non-significant figure >5 round up • 1st non-significant figure <5 round down • 1st non-significant figure =5 then: • If last significant figure is odd round up • If last significant figure is even round down

38. Measurement and Significant Figures • 14.0 g /102.4 mL = 0.136718 g/mL

39. Measurement and Significant Figures • 14.0 g /102.4 mL = 0.136718 g/mL only three significant figures

40. Measurement and Significant Figures • 14.0 g /102.4 mL = 0.137 g/mL only three significant figures

41. Measurement and Significant Figures • Anexact numberis a number that arises when you count items or when you define a unit. • For example, when you say you have nine coins in a bottle, you mean exactly nine. • When you say there are twelve inches in a foot, you mean exactly twelve. • Note that exact numbers have no effect on significant figures in a calculation. They have infinite number of sig figs if you will.

42. Units: Dimensional Analysis • In performing numerical calculations, it is good practice to associate units with each quantity. • The advantage of this approach is that the units for the answer will come out of the calculation. • And, if you make an error in arranging factors in the calculation, it will be apparent because the final units will be nonsense.

43. Units: Dimensional Analysis • Dimensional analysis (or the factor-label method) is the method of calculation in which one carries along the units for quantities. • Suppose you simply wish to convert 131m to millimeters. • Note that the units have cancelled properly to give the final unit of mm.

44. Units: Dimensional Analysis • The ratio (1000 mm/1 m) is called aconversion factor. • The conversion-factor method may be used to convert any unit to another, provided a conversion equation exists. • Relationships between certain U.S. units and metric units are given in Table 2.2

45. General Chemistry IVirginia State University Chapter 2 Dr. Vilchiz Summer 2008

46. Alchemy • Is in essence the ancestor of modern chemistry • Alchemy was more than a science • It was a philosophy • A way of life • Alchemists strived to reach pureness and perfection. • Alchemists venerated gold as the symbol of perfection

47. Alchemy • It was believed that to posses gold will make you rich and pure • To drink gold meant to live forever • Alchemy became the movement to find a way to transform (transmute) matter into gold. • While these believes may seem silly and/or far-fetched they were widely accepted.

48. The search for the elixir of life • It was known that it was possible to take iron and make steel and that if you mixed copper and zinc you will get brass. • Why wouldn’t be possible to make gold? • Needless to say the search for the elixir was futile and eventually alchemy gave way to new scientific questions and approaches.

49. Alchemy’s Legacy • While it might be true that alchemy failed to produce an answer to its driving force it was not by any means a waste. • Many process we now use were discovered or developed during the alchemists years. • Distillation • Fermentation • Putrefaction • Many elements were also discovered • Bi, Zn, As, Co, and P

50. Alchemy to Chemistry • Where does Chemistry come from? • We are not 100% sure where the name comes from but there are several possibilities • It could had come from Egypt Khem = turn black • It could had come from GreeceCheo=to cast • It could had come from ChinaChin-I=gold making juice.