1. SOL Review 1

2.  Accuracy: measure of how close a measurement comes to the actual true value  Precision: measure how close a series of measurements are to one another

3.  Accepted value: correct value based on reliable references.  Experimental value: value measured in the lab.  Error = experimental value – accepted value

4. Absolute error = experimental value – accepted value experimental value- your measured value accepted value- correct value from a reliable reference Percent Error = │experimental value – accepted value│ x 100% accepted value Example: What is your percent error if your thermometer reads 101.3 ºC in boiling water? Answer: │101.3 o C – 100 o C │ x 100% = 1.3% 100 o C

5.  A measurement is only as accurate and precise as the instrument that measured it.  The length is read to the level of the precision of the instrument PLUS one estimated digit.

6.  A number expressed as the product of two factors:  A number from 1 to 10, raised to a power of 10  If the number is greater than 1, the exponent is POSITIVE  If the number is less than 1, the exponent is NEGATIVE  463,000 = 4.63 x 10 5  0.0777 = 7.77 x 10 -2  879,000,000 meters = 8.79 x 10 8 meters  0.000147grams = 1.47 x10 -4 grams  1054 miles = 1.054 x 10 3 miles

7.  You must be able to move from one set of units to another, while maintaining the correct number of significant digits in the measurement.  DIMENSIONAL ANALYSIS is a technique to solve conversion problems using units.  A conversion factor is any ratio of equivalent measurements: 1 dollar or 4 quarters 4 quarters 1 dollar

8.  There are 1000 mL in 1 L.  Therefore, the conversion factors are: 1000 mL or 1 L. 1 L 1000 mL Answer: 2.75 L x 1000mL = 2750 mL 1 L (using this conversion factor, the liters cancel out, leaving you with mL)

9. Because numbers used in chemistry are measurements, the calculated answers are only as good as the measurements they come from. 1. Any number other than 0 is significant. 96.2 g -3 sig figs 2. Zeroes between significant figures are significant. 5.029 m -4 sig figs 3. One or more final zeroes after a decimal are significant. (After a significant figure AND a decimal point) 82.0 m-3 sig figs

10. 4. Initial zeros are not significant. 0.0071 m - 2 sig figs 0.0000998 cm - 3 sig figs 5. Final zeros without a decimal point are not significant. They are place holders. 300 m - 1 sig fig 7200 cm - 2 sig figs

11. Round the answer to the same number of decimal places as the measurement with the fewest decimal places. 9.88sec + 1.2sec 11.08  report 11.1 sec(1 decimal place) 90 sec - 43.1 sec 46.9 sec  report 50 sec ( no decimal places)

12. Round the answer to the same number of significant digits as the measurement with the smallest number of significant digits. 36.5 cm x 2.1 cm 76.65 cm 2  report 77 cm 2 (2 sig figs) 3 cm x 4 cm = 12cm 2  report 10 cm 2 ( 1 sig fig)

14. Density = __Mass__ Volume Usually expressed in grams per ml or cubic centimeter (g/ml or g/cm 3 ) Volume is determined by: Length x width x height (cm 3 ) Density is an Intensive property (does not depend on the amount of matter you have) Mass is an Extensive property (depends on the amount of matter)

15. What is the density of a block of metal that is 2.00 cm x 3.00 cm x 3.50 cm and has a mass of 35.5 g? Answer: Volume of block = 2 cm x 3 cm x 3.5 cm = 21 cm 3 Density = mass = 35.5 g = 1.69 g volume 21 cm 3 cm 3

16. We can also use density as a conversion factor: What is the mass of an object that has a volume of 23.2 ml and a density of 2.52 g/ml? Answer: 23.2 ml x 2.52 g = 58.5 g ml