1. AP Chemistry Olympic High School Mr. Daniel

2. My Website I post all my assignments and give access to notes and homework sheets via my website. To visit the site go to the school website, click on ‘staff’ and scroll to my name http://olhs.cksd.wednet.edu/

3. The Scientific Method Hypothesis: a tentative explanation or prediction based on experimental observations Law: a description of observed natural phenomena Theory: a unifying principle that explains a body of facts and the laws based on them. A good theory can be modified to explain new information.

4. What is chemistry? “Science that deals with the properties, composition, and structure of substances (elements and compounds), the reactions and transformations they undergo, and the energy released or absorbed during those processes... Often called the "central science," chemistry is concerned with atoms as building blocks..., with everything in the material world, and with all living things.” Encyclopaedia Britannica, 2002. “The study of the properties of materials and the changes that materials undergo.” Chemistry: The Central Science, Brown, LeMay, Bursten 7th ed.

5. Why should we study chemistry / How does chemistry affect us? Food science Forensics chemistry Drugs Environmental chemistry: geo- and atmospheric Materials (construction materials, clothing, paints, computer components, etc) New energy sources Et cetera...

6. Classification & Properties of Matter Matter – Anything that occupies space and has mass. Substance – Pure matter with unique properties; it cannot be separated by physical processes. Element – Cannot be subdivided by chemical or physical processes into simpler substances. Atom – The smallest particle of an element that retains the characteristic chemical properties of that element.

8. Compounds Compounds – Contains more than one type of atom, but all molecules (or repeat units) are the same, e.g., water (H 2 O), sodium chloride (NaCl), and “fool’s gold” pyrite (FeS 2 ). Chemical formula – gives the ratio of atoms of each element in a compound. Molecule – smallest discrete units that retain the composition and chemical characteristics of the compound. These groups of atoms are held together with covalent bonding

9. Mixtures Mixture – Have variable composition and can be separated into component parts by physical methods. Mixtures contain more than one kind of molecule, and their properties depend on the relative amount of each component present in the mixture. Homogeneous Mixture (solution) – Uniform composition. Air: principle components include O 2, N 2 & CO 2 Brass: solid solution of Cu and Zn Heterogeneous Mixture – Non-uniform composition. Chocolate Chip Cookie: Chocolate, flour, water, etc Blood– red & white blood cells, plasma, etc.

10. NeN2N2 SO 3 Homogeneous Mixture

11. Physical Properties Physical Properties: Properties that can be observed and measured without changing the composition of a substance. color, state of matter, melting and boiling points density, solubility, electric conductivity, malleability and ductility, viscosity, etc...  Extensive Properties: Depends on the amount of a substance present, e.g., mass and volume.  Intensive Properties: Does not depend on the amount of a substance present, e.g., boiling point and density.

12. States of Matter Gas: volume varies with temperature & pressure; no definite shape; gas molecules are far apart. Solid: fixed volume; rigid shape; particles are packed closely together. Liquid: fixed volume; no definite shape; atoms or molecules are arranged randomly

13. Chemical Properties Chemical Properties: Describe the way a substance may change or react to form other substances. Examples include: Ethanol burns in air (reacts with oxygen in the air) Sodium reacts vigorously with water Corrosion of metal parts (rust) Etc.

14. Chemical Change 2H 2 (g) + O 2 (g)  2H 2 O(g)

15. SI (Metric) Units Measured Property Name of Unit Abbreviation Mass kilogram (not gram!) kg Lengthmeterm Timeseconds TemperaturekelvinK Amount of substance molemol Electric current ampereA Table 1.2 Some SI Base Units

16. Metric Prefixes (Table 1.3) PrefixAbbreviationMeaningExample mega-M 10 6 1 megaton = 1×10 6 tons kilo-k 10 3 1 kilogram (kg)= 1×10 3 grams deci-d 10 -1 1 decimeter (dm) = 1×10 -1 m centi-c 10 -2 1 centimeter (cm) = 1×10 -2 m milli-m 10 -3 1 millimeter (mm) = 1×10 -3 m micro- 10 -6 1 micrometer (  m) = 1×10 -6 m nano-n 10 -9 1 nanometer (nm) = 1×10 -9 m pico-p 10 -12 1 picometer (pm) = 1×10 -12 m femto-f 10 -15 1 femtometer (fm) = 1×10 -15 m

17. Temperature Fahrenheit scaleCelsius scale water freezes32 °F0 °C body temperature98.6 °F 37 °C water boils212 °F 100 °C T (°C) = 5 °C / 9 °F [ T(°F) – 32] Kelvin scale = same size unit as Celsius, but the lowest temperature that can be achieved is absolute zero K = °C +273 °C = K - 273

18. What about Volume? 1 m 3 = 1000 dm 3 (L) = 1,000,000 cm 3 1,000,000 cm 3 = 1  10 6 cm 3 Liter (L) is not a base SI unit. It is a derived unit.

19. Density Density = Mass Volume Units used are :g/mL (liquids), g/cm 3 (solids), g/L (gases)

20. Precision and Accuracy

21. Significant Figures!!

22. Significant Figure Significant figures are a scientist’s way of reporting on how accurately a measurement was made.Significant figures are a scientist’s way of reporting on how accurately a measurement was made. Accuracy depend on two things:Accuracy depend on two things: The device used to measureThe device used to measure The care taken by the individualThe care taken by the individual

23. Significant Figures A Significant figure is defined as including all certain digits plus one estimated digit.A Significant figure is defined as including all certain digits plus one estimated digit. What would be the measurement of the liquid in the graduated cylinder?

24. Significant Figures Which measurement is the best? What is the measured value?

25. Rules for Counting Significant Figures Non-zeros always count as significant figures : 3456 km has 4 significant figures

26. Rules for Counting Significant Figures Zeros Leading zeros never count as significant figures: 0.0486 mL has 3 significant figures

27. Rules for Counting Significant Figures Zeros Captive zeros always count as significant figures: 16.07 cm has 4 significant figures

28. Rules for Counting Significant Figures Zeros Trailing zeros are significant only if the number contains a written decimal point: 9.300 g has 9300 g has 9300. g has 4 significant figures 2 significant figures

29. Rules for Counting Significant Figures Two special situations have an unlimited number of significant figures: 1.Counted items a)23 people, or 425 thumbtacks 2.Exactly defined quantities b)60 minutes = 1 hour

30. Significant Figures 1.Read the number from left to right and count all the digits, starting with the first digit that is NOT zero if a decimal is present. 1200 1200. 1200.0 0.00120 0.01002 2 sig figs 4 sig figs 5 sig figs 3 sig figs 4 sig figs

31. Significant Figures 1) In multiplication or division, the number of significant figures in the answer should be the same as that in the quantity with the fewest significant figures. 0.080 × 125 =? 0.080 × 125 = 10 → 10. or 1.0 × 10 2 (2)(3)(answer should have 2 sig figs) 2) When the number is rounded off, the last digit to be retained is increased by one only if the following digit is 5 or greater.

32. Significant Figures 3) When adding or subtracting numbers, the number of decimal places in the answer is equal to the number of decimal places in the number with the fewest digits after the decimal. 10.1 + 100.001 + 1010 =? 10.1(10 -1 place) 100.001(10 -3 place) 1010.(10 0 place) 1120.101 → 1120. or 1.120×10 3

33. Dimensional Analysis (factor label method) Use units to guide you through calculations. Conversion factors are fractions where the numerator and denominator are the same quantity with different units. Many conversion factors are EXACT. 1000 mm 2.54 cm 1 ft 1 lb 1 m 1 in 12 in 453.6 g