Download presentation

Presentation is loading. Please wait.

1
Chemistry 1A Mr. Kimball dkimball@bakersfieldcollege.edu http://www2.bakersfieldcollege.edu/dkimball

2
Welcome to Chemistry 2A Podcasts A little about myself A little about you –New? Major? ESL? International? –Learning Disorders Sign roll sheet Get phone numbers of others in class

3
The Class Syllabus http://www2.bakersfieldcollege.edu/dkimball

4
Some students prefer to skim through a course. If you really want to succeed you need to go deep!

5
1.Take Responsibility (it’s your education) 2. Have Confidence (you can do it!) 3. Don‘t Procrastinate (study for final now) 4. Read/Listen Precisely (ignore things that just aren’t there) 5. Practice (use practice tests) 6. Persistence (get up one more time than you fall down) 7. Recognize Patterns (most things are done the same way) 8. Use Pictures (outline problem) 9. Think Sequentially (one step at a time) 10. Do Neat work (so you can check it) 11. Group Study (explain things to each other) 12. Try Something New (don’t keep repeating failures) 13. Get Help Learning Skills Learning Skills Power Point RCPPPPPPSNGNHRCPPPPPPSNGNH

6
Using This Book Concept check problems are found within the chapter with the solutions right there with the problem. Exercise problems are found within the chapter with select answers in the back of the book. Homework is assigned from the Internet. You should check the Eduspace link from my Main Web site under Chemistry 1a for instructions.

7
Using This Book A CD comes with the book that has tutorials, practice quizzes and other aids. Those not doing well in the class should consider using some of these aids. There is also a web site. It has practice tests, flash cards, animations, etc. (http://college.hmco.com/chemistry/general/ebbing/general_chem/8e/students/index.html )http://college.hmco.com/chemistry/general/ebbing/general_chem/8e/students/index.html

8
Chemistry and Measurement 1.1 Modern Chemistry 1.2 Experiment and Explanation 1.3 Law of Conservation of Mass 1.4 Matter

9
Aristotle Earth Air Fire Water

10
Democritus (460-370 BC) Greek Philosopher who first coined the word “atomos”.

11
What Is Chemistry? Chemistry is the study of the composition, structure, and properties of matter and energy and changes that matter undergoes. –Matter is anything that occupies space and has mass. –Energy is the “ability to do work.”

12
Archimedes Archimedes lived in Syracuse on the island of Sicily.

13
Archimedes A comparison of Archimedes’ Pulleys and Study! Big Study! Little Study!

14
Galileo Galilei Father of the Scientific Method Heavy things fall faster than light things???? Aristotle

15
Experiment and Explanation Experiment and explanation are the heart of chemical research. –An experiment is an observation of natural phenomena carried out in a controlled manner so that the results can be duplicated and rational conclusions obtained. –After a series of experiments, a researcher may See some relationship or regularity in the results.

16
Experiment and Explanation If the regularity or relationship is fundamental and we can state it simply, we call it a law. –A law is a concise statement or mathematical equation about a fundamental relationship or regularity of nature. –An example is the law of conservation of mass, which says that mass, or quantity of matter, remains constant during any chemical change.

17
Experiment and Explanation Explanations help us organize knowledge and predict future events. –A hypothesis is a tentative explanation of some regularity of nature. –If a hypothesis successfully passes many tests, it becomes known as a theory. –A theory is a tested explanation of basic natural phenomena.

18
Experiment and Explanation The general process of advancing scientific knowledge through observation, laws, hypotheses, or theories is called the scientific method.

19
The Scientific Method Your Assignment: 1.Formulate a Problem. 2.Observe and collect Data. 3.Interpret Data. 4.Test your Interpretation. Examples: 1.Pons and Fleishman, Univ. of Utah. 2.Horoscope 3.Weather.

20
Matter: Physical State and Chemical Constitution There are two principal ways of classifying matter: –By its physical state as a solid, liquid, or gas. –By its chemical constitution as an element, compound, or mixture.

21
Solids, Liquids, and Gases Solid: the form of matter characterized by rigidity; a solid is relatively incompressible and has a fixed shape and volume. Liquid: the form of matter that is a relatively incompressible fluid; liquid has a fixed volume but no fixed shape. Gas: the form of matter that is an easily compressible fluid; a given quantity of gas will fit into a container of almost any size in shape.

22
Elements, Compounds, and Mixtures To understand how matter is classified by its chemical constitution we must first look at physical and chemical changes. –A physical change is a change in the form of matter but not in its chemical identity. –Physical changes are usually reversible. –No new compounds are formed during a physical change. –Melting ice is an example of a physical change.

23
Elements, Compounds, and Mixtures (cont’d) A chemical change, or chemical reaction, is a change in which one or more kinds of matter are transformed into a new kind of matter or several new kinds of matter. –Chemical changes are usually irreversible. –New compounds are formed during a chemical change. –The rusting of iron is an example of a chemical change.

24
A physical property is a characteristic that can be observed for material without changing its chemical identity. Examples are physical state (solid, liquid,or gas), melting point, and color. A chemical property is a characteristic of a material involving its chemical change. –A chemical property of iron is its ability to react with oxygen to produce rust. Elements, Compounds, and Mixtures (cont’d)

25
Matter Pure SubstancesMixtures Compounds (molecules) Homogeneous (solutions) Heterogeneous (most things) Elements (atoms) Hydrogen Oxygen Copper Zinc Granite Sand Wood Orange Juice Air Sodas Ocean Water Alcoholic drinks Water Alcohol Sugar Salt Separate by Physical Processes -filtering -distillation -centrifuging Separate by Chemical Processes -burning -fermentation -rusting

26
Separation by distillation.

27
Elements: sulfur, arsenic, iodine, magnesium, bismuth, mercury. Photo courtesy of American Color.

28
A mixture of potassium dichromate and iron fillings. Photo courtesy of James Scherer. Return to slide 15.

29
A magnet separates the iron filling from the mixture. Photo courtesy of James Scherer. Return to slide 15.

31
1.5 Measurement and Significant Figures 1.6 SI Units 1.7 Derived Units 1.8 Units and Dimensional Analysis Chemistry and Measurement

32
Measurement and Significant Figures Measurement is the comparison of a physical quantity to be measured with a unit of measurement -- that is, with a fixed standard of measurement. –The term precision refers to the closeness of the set of values obtained from identical measurements of a quantity. –Accuracy is a related term; it refers to the closeness of a single measurements to its true value.

33
Precision vs. Accuracy

34
Measurement and Significant Figures (cont’d) 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.

35
Scientific Notation Useful with very large and very small numbers. Decimal always after first digit. Use x 10 n where n is the number of decimal places you must move the decimal to get it just after the first digit. Positive exponents represent large numbers. 2,340,000,000,000,000 = 2.34 x 10 15 Negative exponents represent small numbers. 0.00000000000000234 = 2.34 x 10 -15

36
Measurement Accuracy How long is this steel rod? There is no such thing as a totally accurate measurement!

37
Significant Figures Numbers that measure or contribute to our accuracy. The more significant figures we have the more accurate our measurement. Significant figures are determined by our measurement device or technique.

38
Rules of Determining the Number of Significant Figures 1. All non-zero digits are significant. 203 = 3 sig figs 1.003 = 4 sig figs 1,030.2 = 5 sig figs 2. All zeros between non-zero digits are significant. 234 = 3 sig figs 1.333 = 4 sig figs 1,234.2 = 5 sig figs

39
Rules of Determining the Number of Significant Figures 3. All zeros to the right of the decimal and to the right of the last non-zero digit are significant. 0.0200 = 3 sig figs 0.1220 = 4 sig figs 0.000000012210 = 5 sig figs 4. All zeros to the left of the first non-zero digit are NOT significant. 2.30 = 3 sig figs 1.000 = 4 sig figs 3.4500 = 5 sig figs

40
Rules of Determining the Number of Significant Figures 5.Zeros to the right of the first non-zero digit and to the left of the decimal may or may not be significant. They must be written in scientific notation. 2300 = 2.3 x 10 3 or 2.30 x 10 3 or 2.300 x 10 3 2 sig figs 3 sig figs 4 sig figs

41
Rules of Determining the Number of Significant Figures 6. Some numbers have infinite significant figures or are exact numbers. 233 people 14 cats (unless in biology lab) 7 cars on the highway 36 schools in town

42
How many significant figures are in each of the following? 1) 23.34 2) 21.003 4 significant figures 3).0003030 4) 210 5) 200 students 6) 3000 5 significant figures 2 or 3 significant figures infinite significant figures 1, 2, 3, or 4 significant figures

43
Using Significant Figures in Calculations Addition and Subtraction 1.Line up the decimals. 2.Add or subtract. 3.Round of to first full column. 23.345 +14.5 + 0.523 = ? 23.345 14.5 + 0.523 38.368 = 38.4 or three significant fingures

44
Using Significant Figures in Calculations Multiplication and Division 1.Do the multiplication or division. 2.Round answer off to the same number of significant figures as the least number in the data. (23.345)(14.5)(0.523) = ?177.0368075 = 177 or three significant figures

45
Length MassVolumeTime metergramLitersecond SI System Km=1000m Kg=1000gKL=1000L1min=60sec 100cm=1m1000mg=1 g1000mL=1L60min=1hr 1000mm=1m British 12in=1ft16oz=1 lb4qt=1gal(same) 3ft=1yd2000 lb=1 ton2pts=1qt 5280ft=1mile Footpoundgallonsecond

46
Table 1.5 Relationships of Some U.S. and Metric Units LengthMassVolume 1 in = 2.54 cm1 lb = 0.4536 kg1 qt = 0.9464 L 1 yd = 0.9144 m1 lb = 16 oz4 qt = 1 gal 1 mi = 1.609 km1 oz = 28.35 g 1 mi = 5280 ft 1 lb = 454 g 1 L = 1.06 qt

47
Table 1.3 SI Prefixes MultiplePrefixSymbol 10 6 megaM 10 3 kilok 10 -1 deciD 10 -2 centiC 10 -3 millim 10 -6 micro 10 -9 nanon 10 -12 picop

48
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.

49
Unit Conversion Sodium hydrogen carbonate (baking soda) reacts with acidic materials such as vinegar to release carbon dioxide gas. Given an experiment calling for 0.348 kg of sodium hydrogen carbonate, express this mass in milligrams. x 0.348 kg x 10 3 g 1 kg 10 3 mg 1 g = 3.48 x 10 5 mg

50
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 20 yards to feet. –Note that the units have cancelled properly to give the final unit of feet.

51
Units: Dimensional Analysis The ratio (3 feet/1 yard) is called a conversion 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 1.5.

52
Unit Conversion Suppose you wish to convert 0.547 lb to grams. –From Table 1.5, note that 1 lb = 453.6 g, so the conversion factor from pounds to grams is 453.6 g/1 lb. Therefore,

53
Temperature The Celsius scale (formerly the Centigrade scale) is the temperature scale in general scientific use. –However, the SI base unit of temperature is the kelvin (K), a unit based on the absolute temperature scale. –The conversion from Celsius to Kelvin is simple since the two scales are simply offset by 273.15 o.

54
Figure 1.23: Comparison of Temperature Scales

55
Temperature The Fahrenheit scale is at present the common temperature scale in the United States. –The conversion of Fahrenheit to Celsius, and vice versa, can be accomplished with the following formulas

56
Derived Units The SI unit for speed is meters per second, or m/s. –This is an example of an SI derived unit, created by combining SI base units. –Volume is defined as length cubed and has an SI unit of cubic meters (m 3 ). –Traditionally, chemists have used the liter (L), which is a unit of volume equal to one cubic decimeter.

57
where d is the density, m is the mass, and V is the volume. Generally the unit of mass is the gram. The unit of volume is the mL for liquids; cm 3 for solids; and L for gases. Derived Units The density of an object is its mass per unit volume,

58
A Density Example A sample of the mineral galena (lead sulfide) weighs 12.4 g and has a volume of 1.64 cm 3. What is the density of galena? Density = mass volume = 12.4 g 1.64 cm 3 = 7.5609 = 7.56 g/cm 3

Similar presentations

© 2017 SlidePlayer.com Inc.

All rights reserved.

Ads by Google