1. Introduction and Measurements Mrs. Daniels Chemistry.2 Aug. 28-29, 2002 Rev. summer 2006

2. Major Units of Measurement Length Mass Volume Time Temperature Amount of substance Density Meter - m Kilogram - kg Liter - L Second - s Kelvin – K Mole - mol g/mL

3. Major Units of Measurement Mass – the amount of matter in an object Volume - the amount of space occupied by an object (3 dimensional) Temperature – a measure of the average amount of energy being produced by the movement of molecules/atoms within a substance Mole – a unit of measurement of the amount of a substance Density- mass per unit of volume

4. Prefixes for Units Kilo-k1,000 Deci-d0.1 (1/10) Centi-c0.01 (1/100) Milli-m0.001 (1/1000) Micro-  0.000001 (1 millionth) Nano-n0.000000001 (1 billionth)

5. Glassware Graduated Cylinder- measures volume of a liquid in mL Beaker – larger graduations, volume of liquid in mL Funnel – for getting one solution into another type of container without spilling and also for filtering Cover glass- to place over beakers to eliminate splashing and to minimize evaporation

6. Glassware Flasks: –Erlenmeyer flask – measures volume of a liquid in mL –Round bottom flask –chemical reactions –Volumetric flask - contains a set volume of liquid when filled exactly to the line

7. Other equipment Test tubes – for small scale reactions Bunsen burner – used for heating and burning Heating plate – for heating without open flame Ring stand – for holding any equipment that needs to be elevated Thermometers – always measure in Celsius Mortar and Pestle – for grinding Crucible – for heating to extremely high temps Buret – for acid/base titrations

9. Scientific Method Let’s review the basic steps to the scientific method: Observations Problem (or question) Hypothesis Experiment & data collection Analysis and conclusion Repeat and/or revise

10. Making a Graph A Line Graph MUST Always Include: –Title –Labeled axes (horizontal and vertical) –Dependent (x axis) vs. independent (y axis) variables located properly –Units –Data –A best-fit line (not connect the dots) drawn with a straight edge

12. Practice with Conversions ? m = 10 km 1 L = ? mL 400cm = ? mm ? °C = 273 K ? dm = 1 m 7.8 mm = ? m ? s = 1 hr 1 mL = ? cm 3

14. Scientific Notation One digit to the left of the decimal Keep all other significant digits after the decimal point (we’ll discuss shortly) Expressed by writing this number x 10 raised to a power That exponent will be: –Positive if the decimal needs to be moved to the right to express the original number –Negative if the decimal needs to be moved to the left to express the original number

15. Scientific Notation Ex. 3,122 would be written 3.122 x 10 3 The decimal in 3.122 x 10 3 needs to be moved to the RIGHT “3” places to convert it back into 3,122 Ex. 0.0141 would be written 1.41 x 10 -2 The decimal in 1.41 x 10 -2 would have to be moved to the LEFT “2” places in order to convert it back into 0.0141

16. Significant Figures Rules for Significant Figures: 1. Every nonzero digit is significant 2. Zeros between nonzero digits are significant 3. Zeros at the end of a number to the right of a decimal are significant 4. Leftmost zeros appearing in front of nonzero digits are NOT significant 5. Zeros to the rightmost end of a number with no decimal place are NOT significant 6. Exact numbers have infinite # of sig figs.

18. Quantitative v. Qualitative QUANTITATIVE Definite form Numbers Need instruments to determine QUALITATIVE No numbers Descriptive only

19. Precision v. Accuracy Precision is a measure of how close or consistent a series of measurements are to one another –Ex. In archery, this would be a tight group Accuracy is a measure of how close a measurement comes to the true value –Ex. In archery, this would be hitting the bullseye