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Chemistry: Chapter 1

Chemistry: Chapter 1. Introduction to Chemistry. Scientific Law Observation of a natural event Summary of what occurs Does not try to explain why something occurs only tells what occurs. Scientific Theory Explanation of events Tries to explain why something occurs

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Chemistry: Chapter 1

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  1. Chemistry: Chapter 1 Introduction to Chemistry

  2. Scientific Law Observation of a natural event Summary of what occurs Does not try to explain why something occurs only tells what occurs Scientific Theory Explanation of events Tries to explain why something occurs Supported by several experiments Must be able to predict what happens by using the theory Distinguish between a scientific law and a scientific theory.

  3. Observation Seeing something that makes you ask a question Formulate a question What exactly do you want to learn? Research the question Find out what others have said about your question Develop a hypothesis Use the information you found in research to develop an educated guess in answer to your question Experiment/Collect Data Test only one variable at a time Use a control Control is a version of the experiment where nothing is changed Draw conclusions Analyze the data Did the data support your hypothesis? Explain and apply the steps of the scientific method.

  4. Describe the relationship between pure science and technology. • Pure science studies things that may never be useful • Pure science seeks only to know • Technology is useful • Technology is applied science • Studying far off galaxies is pure science • Creating a vaccine for a disease is technology

  5. Independent variable The variable that you change in the experiment If you place one plant in the window and one in the closet the variable you are changing is the amount of light Amount of light would be the independent variable Dependent variable The variable that changes because of the change in the independent variable The plant in the closet is only 6 cm tall. The plant in the window is 12 cm tall. Height of the plant would be the dependent variable. Distinguish between an independent and dependent variable.

  6. Chemistry: Chapter 2 Data Analysis

  7. SI Units • For a measurement to make sense, it requires both a number and a unit. • Many of the units you are familiar with, such as inches, feet, and degrees Fahrenheit, are not units that are used in science.   • Scientists use a set of measuring units called SI, or the International System of Units. The abbreviation stands for the French name Système International d'Unités. • SI is a revised version of the metric system. • There are seven primary base units you need to learn.

  8. Base Units

  9. Metric Prefixes • The metric unit for a given quantity is not always a convenient one to use. • A metric prefix indicates how many times a unit should be multiplied or divided by 10. • Learn these prefixes…

  10. Identify the SI base units and compare the base units to derived units • Derived units are made up of more than one base unit • Examples of derived units : g/cm3, g/mol, m3 • If it is not on the list of base units, it is a derived unit.

  11. Derived Units

  12. Solve density problems. • Density is the Mass divided by the Volume • D = M/V • The equation can be rearranged to solve for any of the three variables. • M = D x V • V = M/D • Example • A block of aluminum occupies a volume of 15.0 mL and has a mass of 40.5 g. What is its density? • M = 40.5 g • V = 15.0 mL • D = M/V D = 40.5g/15.0 mL D = 2.7 g/mL

  13. Density Problems cont. • What is the mass of the ethyl alcohol that exactly fills a 200.0 mL container? The density of ethyl alcohol is 0.789 g/mL. • D = 0.789 g/ml (3 sig figs) • V = 200.0 mL (4 sig figs) • M = D x V M = 0.789g/ml x 200.0 mL • M = 158 g (3 sig figs)

  14. Density Problems cont. • What volume of silver metal will have a mass of exactly 2500.0 g. The density of silver is 10.5 g/cm3. • D = 10.5 g/cm3 • M = 2500.0 g • V = M / D V = 2500.0 g/ 10.5 g/cm3 • V = 238 cm3 (3 sig figs)

  15. Density Problems cont. • A rectangular block of copper metal has a mass of 1896 g. The dimensions of the block are 8.4 cm by 5.5 cm by 4.6 cm. From this data, what is the density of copper? • First calculate the volume. (V = l x w x h) • V = 8.4 cm x 5.5 cm x 4.6 cm V = 212.52 cm3 • There are only 2 sig figs in the problem so you must round to 210 cm3 • Now calculate the density. • M = 1896 g • V = 210 cm3 • D = M/V D = 1896 g/ 210 cm3 D = 9.03 g/cm3 • Can only have 2 sig figs therefore the answer would be 9.0 g/cm3

  16. Density Practice • Mercury metal is poured into a graduated cylinder that holds exactly 22.5 mL. The mercury used to fill the cylinder has a mass of 306.0 g. From this information, calculate the density of mercury. • A flask that has a mass of 345.8 g is filled with 225 mL of carbon tetrachloride. The mass of the flask and carbon tetrachloride is found to be 703.55 g. From this information, calculate the density of carbon tetrachloride. • Calculate the density of sulfuric acid if 35.4 mL of the acid has a mass of 65.14 g. • Find the mass of 250.0 mL of benzene. The density of benzene is 0.8765 g/mL. • A block of lead has dimensions of 4.50 cm by 5.20 cm by 6.00 cm. The block ‘s mass is 1587 g. From this information, calculate the density of lead. • What is the volume of a substance with a mass of 0.35 g and a density of 0.9 g/ml?

  17. Measuring Temperature • A thermometer is an instrument that measures temperature, or how hot an object is. • The two temperature scales that you are probably most familiar with are the Fahrenheit scale and the Celsius scale • You can convert from one scale to the other by using one of the following formulas.

  18. The Kelvin Scale • The Kelvin is the SI unit of temperature. • On the Kelvin scale, water freezes at about 273 K and boils at 373 K. • It is easy to convert from Celsius to Kelvin. Kelvin = °C + 273 • No degree sign is needed with Kelvin.

  19. Scientific Notation • Scientific Notation is based on powers of the base number 10. • The number 123,000,000,000 in scientific notation is written as : 1.23 X 1011 • The first number 1.23 is called the coefficient. • It must be greater than or equal to 1 and less than 10. • The second number is called the base. • It must always be 10 in scientific notation. The base number 10 is always written in exponent form. • In the number 1.23 x 1011 the number 11 is referred to as the exponent or power of ten.

  20. 300 000 000 Only 1 number allowed in front of the decimal Count the number of times the decimal must be moved this number becomes the exponent If the original number is larger than one the exponent is positive 3.0 x 108 0.000 000 03 If the number is smaller than one the exponent is negative 3.0 x 10-8 Use scientific notation to represent very large and small numbers

  21. Write the following numbers in scientific notation • 800 000 000 m • 0.0015 kg • 60 200 L • 0.00095 m • 8 002 000 km • 0.000 000 000 06 kg • 602 000 000 000 000 000 000 000 atoms

  22. Write the following numbers in standard form • 4.5 x 105 • 7.009 x 109 • 4.6 x 104 • 3.2 x 1015 • 3.115 x 10-8 • 6.05 x 10-3 • 1.99 x 10-10 • 3.01 x 10-6

  23. Adding and Subtracting • Exponents must be the same or the operation cannot be performed. If the exponents do not agree, change the decimal and the power of ten notation of either number so as to agree with the other. • Add or subtract the number. • Keep the same power of ten.

  24. Multiplying and Dividing • MULTIPLICATION: • Multiply the numbers. • Add the exponents for the power of ten. • DIVISION: • Divide the numbers as in any division problem. • Subtract the denominator power of ten exponent from the numerator power of ten exponent.

  25. Dimensional Analysis • Suppose you want to convert the height of Mount Everest, 8848 meters, into kilometers. Based on the prefix kilo-, you know that 1 kilometer is 1000 meters. This ratio gives you two possible conversion factors. • Since you are converting from meters to kilometers, the number should get smaller. Multiplying by the conversion factor on the left yields a smaller number. • Notice that the meter units cancel, leaving you with kilometers (the larger unit).

  26. 300 mm = _______m 2 km = ________hm 0.90 cm = _____mm 5.67 dm = _____km 3.6 Gm = _______m 4.5 mg= _________g 34 kg = ________mg 45 ms = ________s 6.7 nm = ________m 37 kg = ________cg 23 ml = ________ kl 9.7 dam = ______m 0.0054 cg = _____mg 0.5 cm = _______mm 0.68 Mg = _______cg 1Gm= _________nm Convert the following measurements

  27. Limits of Measurement • Precision--Precision is a gauge of how exact a measurement is. • The precision of a measurement depends on the number of digits in the answer. • Significant figures are all the digits that are known in a measurement, plus the last digit that is estimated. • The fewer the significant figures, the less precise the measurement is.

  28. Uncertainty • When you make calculations with measurements, the uncertainty of the separate measurements must be correctly reflected in the final result.   • The precision of a calculated answer is limited by the least precise measurement used in the calculation. • So if the least precise measurement in your calculation has two significant figures, then your calculated answer can have at most two significant figures.

  29. Accuracy • Accuracy is the closeness of a measurement to the actual value of what is being measured. • Although an instrument is precise, it does not have to be accurate.

  30. Precision or Accuracy?

  31. Define significant figures and know when to use them. • When numbers are measured, measurements are always taken to the first number that is estimated (guessed). • This is the last significant figure. • If the measurement is 100, the actual number could be anywhere from 50-149. • If the measurement is 100.0 then it can only vary from 99.5 and 100.4. • Significant figures indicate precision of measurement.

  32. Use significant figures in problem solving. • The answer to a problem cannot have more significant figures than the number in the problem with the fewest significant figures. • Once the correct number of significant figures has been reached, zeroes are used as place holders • 76543210 written in 3 significant figures would be 76500000 • Cannot drop the zeros. You wouldn’t want me to pay you 10 dollars if I owed you 1000. Dropping zeroes changes the value of the number. Round then hold the place with zeroes to keep the value the same. • If the number following the last significant figure is 5 or greater the last significant figure goes up one. • If the number following the last significant figure is 4 or less the last significant figure stays the same.

  33. Rules for Significant Figures • All non-zero numbers are always significant. • Zeros between non-zero numbers are always significant. • All final zeros to the right of the decimal place are significant. • Zeros that act as placeholders are not significant. Convert quantities to scientific notation to remove the placeholder zeros. • Counting numbers and defined constants have an infinite number of significant figures.

  34. 1.560 1560 0.01560 300000 290100000 0.000002390 0.000000000002 14.9800 100 100.0 20000.0 3009000 Determine the number of significant figures in each of the following:

  35. Rounding • A calculated value with eight significant figures is not appropriate when you only need four significant figures.

  36. Rules for Rounding • If the digit to the immediate right the last significant figure is less than five, do not change the last significant figure. • If the digit to the immediate right of the last significant figure is greater than five, round up the last significant figure. • If the digit to the immediate right of the last significant figure is equal to five and is followed by a nonzero digit, round up the last significant figure. • If the digit to the immediate right of the last significant figure is equal to five and is not followed by a nonzero digit, look at the last significant figure. If it is an odd digit, round it up. If it is an even digit, do not round up.

  37. Round the following to 3 significant figures: • 4.900 • 4.905 • 20087 • 653456 • 928227 • 5.596 • 300.0 (try scientific notation)

  38. Rules for Significant Figures with Operations • When adding and subtracting numbers, your answer should have the same number of digits to the right of the decimal point.

  39. Examples • 43.2 cm+ 51.0cm+ 48.7 cm • 258.3 kg + 257.11 kg + 253 kg • 0.0487 mg + 0.05834 mg + 0.00483 mg • 93.26 cm – 81.14 cm • 5.236 cm – 3.14 cm • 4.32 x 103 cm – 1.6 x 103 cm

  40. Multiplying and Dividing • When you multiply or divide numbers, your answer must have the same number of significant figures as the measurement with the fewest number of significant figures.

  41. Examples • 24 m x 3.26 m • 120 m x 0.10 m • 1.23 m x 2.0 m • 53.0 m x 1.53 m • 4.84 m/2.4 s • 60.2 m/20.1 s • 102.4 m/51.2 s • 168 m/58 s

  42. Percent Error • Percent error is used to calculate the accuracy of experimental data values and the accepted value. • Percent error is the ratio of an error to its accepted value and can be found using the following formula:

  43. Organizing Data • Scientists accumulate vast amounts of data by observing events and making measurements. • Interpreting these data can be a difficult task if they are not organized.   • Scientists can organize their data by using data tables and graphs. • These tools make it easier to spot patterns or trends in the data that can support or disprove a hypothesis

  44. Data Tables • The simplest way to organize data is to present them in a table. • The table relates two variables—an independent variable and a dependent variable.

  45. Line Graphs Typical Line Graph A line graph is useful for showing changes that occur in related variables. In a line graph, the independent variable is generally plotted on the horizontal axis, or x-axis. The dependent variable is plotted on the vertical axis, or y-axis, of the graph. A direct proportionis a relationship in which the ratio of two variables is constant. An inverse proportion, a relationship in which the product of two variables is a constant.

  46. Bar Graphs Typical Bar Graph • A bar graph is often used to compare a set of measurements, amounts, or changes. • The bar graph makes it easy to see how the data for one thing compares with the data for another.

  47. Circle Graphs • A circle graph is a divided circle that shows how a part or share of something relates to the whole.

  48. Communicating Data • A crucial part of any scientific investigation is reporting the results.   • Scientists can communicate results by writing in scientific journals or speaking at conferences. • Scientists also exchange information through conversations, e-mails, and Web sites. • Young scientists often present their research at science fairs

  49. Peer Review • Different scientists may interpret the same data differently. This important notion is the basis for peer review, a process in which scientists examine other scientists' work. • Peer review encourages comments, suggestions, questions, and criticism from other scientists. • Peer review can also help determine if data were reported accurately and honestly.

  50. Lab Equipment Erlenmeyer Flask Graduated Cylinder Triple Beam Balance Watch glass Beakers Volumetric Flask

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