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Introduction to Chemistry and Measurement

Introduction to Chemistry and Measurement. What is Chemistry?. The study of all substances (matter) and the changes they undergo. EX: Burning Paper (chemical change) Melting Ice (physical change)

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Introduction to Chemistry and Measurement

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  1. Introduction to Chemistry and Measurement

  2. What is Chemistry? The study of all substances (matter) and the changes they undergo. EX: Burning Paper (chemical change) Melting Ice (physical change) Chemistry is the central science because it overlaps many other sciences, like biology and Earth science.

  3. Scientific Method It is a systematic way to answer question about the natural world. Steps: • 1. A scientist makes an observation. • 2. The observation leads to a question. • 3. Experiment - a test of the hypothesis • 4. Conclusion – a possible explanation of the results • 5. Natural Law – describes how nature behaves. It can be duplicated many times.

  4. Scientific Theory It explains why nature behaves in a certain way. It is based on empirical evidence.

  5. Measurement It is a major part of science. • 1. Every measurement needs a number. (value) • 2. All measurements need to include units. All sciences use the metric system • 1. They’re called the International System of Units • 2. All scientists use it as a common unit.

  6. International System of Units (SI) 1. Built upon the base 10 method of counting - Length is measured in meters (m) EX: 1m = 3.3 ft - Mass is measured in grams (g) - Time is measured in seconds (s) • 2.  Some units are derived units - Area = l x w - Volume = l x w x h

  7. Exceptions to Base Units Volume and temperature units are not based on log base 10 - Liter – based on cubed meters - Celsius – unit of temperature

  8. Metric Prefixes Used to make units larger or smaller than the base unit EX: 1kg = 1000g Ex) If you have 10 kg, how many grams do you have?

  9. Common Metric Prefixes mega = M 1000000 kilo = k 1000 base unit =meters, grams, liters 1 deci = d 0.1 centi = c 0.01 milli = m 0.001 micro = u 0.0001 nano = n 0.000000001 pico = p 0.000000000001

  10. Accuracy in Measurements Precision is achieved when you obtain the same answer over and over Accuracy is achieved when you obtain a value close to the accepted value Reasons for uncertain measurements 1. Instruments can have flaws or are not calibrated 2. Human error when estimating

  11. True Value 1. Achieved when you take estimated value to the farthest guess 2. Actual value is plus (+) or minus (-) 0.1 units depending on guess measurement EX: 31.7 ± 0.1 is 31 and 7 tenths plus or minus 0.1 4. Accepted Value - the correct value ( measurement)

  12. Significant Figures These are the certain digits and estimated digits of a measurement. EX: 31.7 , the 3 and 1 are certain digits and 7 is the estimated digit YOU CANNOT REPORT DATA THAT IS MORE PRECISE THAN YOUR LEAST PRECISE MEASURMENT

  13. “Rules” of Significant Figures Zero - As a “place keeper” it tells where the decimal point goes and is NOT significant. If it is after the decimal than it is significant. http://dbhs.wvusd.k12.ca.us/webdocs/SigFigs/SigFigRules.html EX: If a balance measures to the nearest 10 grams and measures 1060 grams then the number has 3 significant figures, 1, 0, 6. The last 0 only shows where the decimal goes. Try this: 7,006,500

  14. Atlantic Pacific Rule a. measurements with a decimal point uses the Pacific rule, count from the left using the first non zero number and going to the end. b. Measurements without a decimal point uses the Atlantic rule, count form the right using the first non zero number and going to the end. c. Start all counting with first non zero number then count all the way to the end.

  15. Calculations An exact number does NOT affect number of significant figures in answer EX: 1000m = 1km ALWAYS In multiplication and division the measurement with the least significant figures tells the number of significant figures allowed in reported answer In addition or subtraction the significant figures depends on the number with the least significant figures (least accurate) ONLY THE FINAL ANSWER is put into significant figures Rounding – 5 or greater round up, 4 or below rounds down

  16. Scientific Notation Makes numbers easier to work with, especially really large numbers and decimals. Step 1: Move the decimal to the right or the left form a whole number between 1 and 10. Leave the other digits after the decimal. Step 2: Set up the scientific notation using the number from step 1. You will not have an exponent yet. Step3: Count the number of spaces the decimal has moved. Step 4: Determine if the decimal moved to the right or the left. Movement to the right creates a negative exponent. Movement to the left creates a positive exponent. Step 5: Place the exponent in the scientific notation.

  17. Try It! Complete practice problems.

  18. Percent Error When calculating percent error, we compare the measured value to the accepted value. Formula from Table T: Measured Value – Accepted Value X 100 Accepted Value

  19. Try it! You measured 0.26 g of product produced from your experiment. The accepted value is 0.28g . What is the percent error in your experiment?

  20. Ratios Ratios are found by comparing two quantities. An example of ratios is speed (m/s). The more common example is density. (Table C) Density is calculated by Density = mass volume Density is expressed in units of gram(g) per mL or cubic centimeter.

  21. Problem Solving Step 1: Determine what the question is asking and what needs to be solved. Step 2: Make a plan to solve the problem. Step 3: Solve it! Step 4: Check to see if the answer makes sense. If not, make another plan.

  22. Dimensional Analysis Dimensional – Using dimensions with units Analysis – Analyzing information Dimensional Analysis: Conversions between different units.

  23. Steps to Dimensional Analysis 1. Write what you need to convert with the units you want on bottom 2. Choose a unit equality and write the conversion factor(s) for the corresponding unit equality 3. Set up problem to get desired units and so other units will cancel 4. Multiply and cancel units (remember significant figures!) 5. Check answer to see if it makes sense.

  24. Step 1 The units to convert are given in the problem. Ex: If there is 756 L in the container, how many gallons are there? Write: 756L X

  25. Step 2 Unit Equalities show how different units are related. Unit equality: 3.785 L = 1 gal Write: 1 gal 3.785L

  26. Step 3 Set up problem to cancel units: Write: 756 L x 1 gal = X 3.785L

  27. Step 4 Multiply and cancel units: 756 L x 1 gal = 199.73579gal X 3.785L With Significant Figures ~ 2.00 x 102 gal

  28. Step 5 Check that the answer is reasonable. Since we had more than 3.785L in our sample then we should have more than 1gal in our answer. Therefore, our answer is reasonable.

  29. Things to Watch Out For • Make sure conversion factors are set up to give the desired units for your answer. • If the undesired units do not cancel, check to see if the problem is set up correctly. • Always make sure to include units with the final answer. • Impossible conversions like seconds to kilograms and liters to degrees.

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