2. What is Chemistry? Chemistry is: the study of matter & the changes it undergoes Composition Structure Properties Energy changes

3. Matter Pure Substance - definite composition -cannot be physically separated Element -one type of atom Compound -two or more types of atoms chemically bonded Mixture -two or more varying parts -can be physically separated Homogeneous -evenly distributed -uniform properties Solution -very small particles -light passes through Heterogeneous -unevenly distributed -varying composition Suspension -large visible particles -particles settle out -blocks light Colloid -unevenly suspended particles -shows a visible beam of light (positive Tyndall Effect)

4. States of Matter PhaseVolumeShapeStructure SOLID LIQUID GAS PLASMA Draw an arrow next to the table above in the order of increasing energy. Definite Indefinite

5. Three States of Matter Plasma

6. Taking Measurements

7. The SI System Around 1793, scientists all over the world began to agree upon a single measurement system called Le Systeme International d’ Unites or SI System 7 base units The idea was to create a unifying system of weights and measurements

8. QuantityUnitSymbol Masskilogramkg Lengthmeterm Timeseconds Amount of Substance molemol TemperatureKelvinK Electric current AmpereA Luminous intensity candelacd Note: This is the SI standard unit but the BASE unit is the GRAM What’s missing?

9. Combinations of base units Volume: amount of space taken up by an object Most common unit: cm 3 = mL Density: ratio of mass to volume Common units: g/cm 3 of g/mL or g/L Does not change for a given substance Derived Units

10. PrefixSymbolMeaningNumerical Value Giga-G10 9 1,000,000,000 Mega-M10 6 1,000,000 Kilo-k10 3 1,000 Hecto-h10 2 100 Dekada10 1 10 BASEg, l, m10 0 1 Deci-d10 -1.1 Centi-c10 -2.01 Milli-m10 -3.001 Micro-m10 -6.000001 Nano-n10 -9.000000001

12. Practice Problems 1. 5.6 cm to m 2. 56 mg to g 3. 340 mm to cm 4. 1.2 ML to L 0.056 m 0.056 g 34 cm 1,200,000 L

13. Scientific Notation 602,000,000,000,000,000,000,000 6.02 x 10 23 0.0000000000000000000000199 1.99 x 10 -23 Some numbers are very large or very small, so a short hand notation is needed! Too large: Too small:

14. N x 10 n N is a number between 1 and 10 n is a positive or negative integer if n is a negative number, the full number is a small decimal if n is a positive number, the full number is a large number General Notation:

15. Practice 3.69 x 10 -4 4.382 x 10 -2 8.37 x 10 -7 1.245 x 10 5 8.7900 x 10 8 2.6091 x 10 2 0.00000568 0.00436 0.00000000002 2460000000 3456965 3400450

16. Factor-Label Method (Dimensional Analysis)

17. Using Factor-Label Method Sample Problems: Converting 9.8 g to kg 9.8 g x 1 kg = 0.0098 kg 1000. g Converting 9.8 kg to g 9.8 kg x 1000. g = 9800 g 1 kg “1” goes in front of larger unit!

18. Density Practice Density Formula m D V D = m V What is the density of carbon dioxide gas if 0.196 g occupies a volume of 100 mL?

19. Making Measurements

20. 24.62what is certain? what is uncertain? Measuring always involves some estimation Certain Digits: A digit that represents a mark on a scale or a non-blinking number on a display. Uncertain (Estimated) Digits: A digit that represents the space between the marks on a scale or a blinking number on a display.

21. Sig Figs: Using the Pacific/Atlantic Rule Step 1: Ask yourself: is the decimal point Present or Absent ? Step 2: Determine which way to start counting If the decimal point is Present, start counting from the LEFT If the decimal point is Absent, start counting from the RIGHT PACIFICPACIFIC ATLANTICATLANTIC resent bsent

22. Pacific/Atlantic Rule Step 3: Start counting on Pacific or Atlantic side from the first NON-ZERO number. Count all numbers after the first non-zero number including zeros. Examples: a)1234 = ________ sig figs b)1204 = ________ sig figs c)0.00234 = _______ sig figs d)1230 = ______ sig figs e)1234.0 = ______ sig figs Absent  Present 

23. Using Sig. Figs. In Calculations Addition/Subtraction Rule Answer should contain least # of decimal places Multiplication/Division Rule Answer should contain least sig figs.

24. Do Now: Precision of Lab Instruments 1.Record the following quantities to the correct number of decimal places. ________ L ________ mL _______ o C 2.Convert your answer in A to milliliters: ________ mL 3.Add your answer from A & B. Record using correct sig. figs. ________ mL

25. Analyzing Measurements

26. Accuracy & Precision in Measurements Accuracy: closeness of measurements to correct value Precision: closeness of a set of measurements to each other (assuming they’re made in the same way) When recording a measurement, an instrument that provides the most digits past the decimal is most precise.

28. Accuracy vs. Precision Example: A student measures the density of a sample of nickel. The density of nickel is 8.9 g. mL -1 Is this accurate or precise? Density Result (g.mL -1 ) Trial 17.8 Trial 27.7 Trial 37.8

29. Percentage Error Accuracy of an individual value (or average) can be compared to the correct/accepted value % Error = Experimental – Accepted x 100 Accepted

30. Percentage Error What is the percentage error for a mass measurement of 17.7 g, given that the correct value is 21.2 g? A volume is measured experimentally as 4.26 mL. What is the percentage error, given that the accepted value is 4.15 mL?

31. Scientific Method

32. SCIENTIFIC METHOD logical approach to solving problems Observation Problem Hypothesis Experiment Data Analysis Conclusion

33. You have 15 seconds to count how many letter “F”s you see in the following statement.

34. FEATURE FILMS ARE THE RESULT OF YEARS OF SCIENTIFIC STUDY COMBINED WITH THE EXPERIENCES OF YEARS.

35. Qualitative (think “quality”): observations using words Example: Quantitative (think “quantity”): observations using numbers and units. Example: Types of Observations

36. How observant are you?!?

37. PHYSICAL AND CHEMICAL CHANGES

38. physical property: characteristic that can be observed or measured without changing the identity of the substance. melting point, boiling point, density physical change: change in a substance that does not involve a change in the identity of the substance. dissolving, cutting, melting, and boiling Physical Properties and Changes

39. chemical property: a substance’s ability to undergo changes that transform it into different substances Example: combustibility, reactivity chemical change: change in which one or more substances are converted into different substance Example: rusting, cooking food Chemical Properties and Changes

40. Evidence of a Chemical Change 1.Color change 2.Temperature change 3.Production of a gas 4.Change in odor 5.Formation of a precipitate Precipitate: insoluble solid that separates out of solution

41. NOTE: Solubility and Phase changes are PHYSICAL!!!!

42. In an experiment… System: specific portion of the experiment that has been selected for study Constant: experimental conditions that do not change Control: experimental condition that is used as a standard for comparison Variable: experimental condition that does change

43. SpongeBob loves to garden and wants to grow lots of pink flowers for his pal Sandy. He bought a special Flower Power fertilizer to see if it will help plants produce more flowers. He plants two plants of the same size in separate containers with the same amount of potting soil. He places one plant in a sunny window and waters it every day with fertilized water. He places the other plant on a shelf in a closet and waters it with plain water every other day. 1. What are Spongebob’s constants in his experiment? 2. What are Spongebob’s variables in his experiment? 3. What did Spongebob do wrong? 4. What should SpongeBob do to test the effectiveness of Flower Power fertilizer? Describe an experiment.

44. Amount of Fertilizer (g)Plant Growth (cm) 65 99 1517 2322 Fertilizer  Growth  Direct Relationship Independent Variable Dependent Variable Title Appropriate scale Axis labeled “Best fit” line

45. Direct Relationshi ps When 2 quantities divided by each other gives a constant value K (constant value) = Y/X Ex: Density

46. Inverse Relationshi ps When 2 quantities multiplied by each other gives a constant value K = X Y Ex: Boyle’s Law K = PV