1. Pre-AP Chemistry Chapter 2 “Measurements and Calculations”

2. Scientific Method Observe State the Problem Formulate Hypotheses Test Hypotheses with controlled experiments Record and analyze data Report Conclusions

3. Units of Measurement SI(International System) Base Units QuantityUnitSymbol lengthmeterm masskilogramKg timeseconds temperatureKelvinK amount of matter molemol

4. SI Prefixes Prefix Symbol Factor teraT10 12 gigaG10 9 megaM10 6 kiloK10 3 centic10 -2 millim10 -3 micro  10 -6 nanon10 -9 picop10 -12

5. Derived Units derived unit - a unit obtained by combinations of fundamental units Example - volume (cm 3 ) –V = l X w X h –V(cm 3 ) = (cm) X (cm) X (cm) 1cm 3 = 1mL Volume is measured using the base unit the Liter (L)

6. Density density - the mass per unit volume of a material density = mass/volume D(g/cm 3 ) = m(g) / V(cm 3 ) D(g/mL) = m(g) / V(mL) For gases –D(g/L) = m(g) / V(L)

7. Sample Density Problem Problem: Calculate the density of 10g of a material occupying a volume of 2.5mL. (10 pts) D = m/V, m = 10g, V = 2.5mL 5 pts D = 10g / 2.5mL 2 pts D = 4 g/mL – (1pt) (2pts) 3pts

8. Heat and Temperature Heat is thermal energy, the energy of particles in motion. Heat can be measured in joules. –joule - the SI unit of measuring energy Temperature is the measure of the effects of heat. Temperature can be measured in degrees Celsius.

9. Temperature Scales The boiling and freezing points of water as well as room temperature will be used to compare temperature scales. –Fahrenheit Celsius Kelvin b.p. 212 0 100 0 373 0 r.t. 72 0 24 0 297 0 f.p. 32 0 0 0 273 0 a.z. -454 0 -273 0 0 0

10. Temperature Conversions K= 0 C + 273 0 C = 5/9 (F-32) 0 F = 9/5 C + 32 How would you convert Fahrenheit to Kelvin?

11. Absolute Zero Absolute zero is a theoretical temperature. It is the temperature at which there is no molecular motion. Why is it a theoretical temperature? Particles of matter are constantly moving. This motion creates friction which gives all matter a temperature.

12. Accuracy and Precision accuracy - the nearness of a measurement to an accepted value precision - the agreement between a set of measurements

13. How do we measure accuracy? Percentage Error O - A %E = ------------ X 100% A O – observed or experimental value A - accepted value

14. Sample Problem Suppose you performed an experiment and found the volume of a 10g sample of aluminum to be 3.5mL. –Calculate the density of the sample using your data. –The accepted density of aluminum is 2.7g/mL. Calculate the percentage error of your measurements.

15. Significant Digits The following rules are used to determine the number of significant digits. 1. Nonzero digits are always significant. 2. All final zeros after the decimal point are significant. 3. Zeros between two other significant digits are always significant. 4. Zeros used solely for spacing the decimal point are not significant.

16. Operations With Significant Digits Addition and Subtraction –The answer should be rounded off so that the final digit is in the same place as the leftmost uncertain digit. Example 34.9 +4.54 39.44.......the correct answer is 39.4

17. Operations With Significant Digits Multiplication and Division The answer should be rounded off to the same number of significant digits as the measurement with the least number of significant digits. Example –2.34 X 6.5 = 15.21 –the correct answer is 15

18. Scientific Notation scientific notation - an expression of numbers as powers of 10 Express 93,000,000 in scientific notation 93,000,000

19. Scientific Notation Express 0.000 000 000 189 in scientific notation. 0.000 000 000 189

20. Operations with Scientific Notation Adding or subtracting: The numbers must be adjusted so that the exponents of the numbers are the same. The answer is written in scientific notation. Example: 4.2 x 10 4 + 7.9 x 10 3 becomes 4.2 x 10 4 g The correct answer: + 0.79 x 10 4 g 5.0 x 10 4 4.99 x 10 4 g correct significant figures

21. Multiplication In multiplication, the first values are multiplied and the exponents are added. The answer is written in scientific notation with correct significant figures. Example: (5.23 x 10 6 mm) (7.1 x 10 -2 mm) = 37.133 x 10 4 mm 2 = 3.7 x 10 5 mm 2

22. Division with scientific notation For division operations, the first numbers are divided, and the exponent of the denominator is subtracted from the numerator. Example: 5.44 x 10 7 m 8.1 x 10 4 m = 0.671604983 x 10 3 m = 6.7 x 10 2 m

23. Direct Proportions Two quantities are directly proportional if dividing one by the other gives a constant value. The graph of the two quantities is a straight line

24. Inverse Proportions Two quantities are inversely proportional to each other if their product is constant. The graph is a hyperbola

25. Solving Quantitative Problems Step 1: Analyze - Read the problem at least twice. Identify the known and unknown. Step 2: Plan - Using the known and unknown information, determine the route or formula(s) to be used for solving the problem. Step 3: Compute - Substitute the values into the formula, perform the appropriate operations, round the answer off to the appropriate significant digits, and don’t forget the units! Step 4: Evaluate - Ask yourself, does the answer make sense?