1. QUALITATIVE VS. QUANTITATIVE QUALITATIVE = DESCRIPTIVE, NON- NUMERICAL. EX. : THE MORNINGS ARE GETTING VERY COLD. QUANTITATIVE = NUMERICAL, DEFINITE MEASUREMENT EX. : THE TEMPERATURE WAS 65°F THIS MORNING.

2.  Know your metric relationships!  kilo -hecta - deca -(base) - deci - centi – milli ( 10 3 - 10 2 -10 1 - (1) – 10 -1 - 10 -2 - 10 -3 )  K ing H enry D ied B y D rinking C hocolate M ilk

3. In science, we deal with some very LARGE numbers: 1 mole = 602000000000000000000000 In science, we deal with some very SMALL numbers: Mass of an electron = 0.000000000000000000000000000000091 kg Scientific Notation

4. Imagine the difficulty of calculating the mass of 1 mole of electrons! 0.000000000000000000000000000000091 kg x 602000000000000000000000 x 602000000000000000000000 ???????????????????????????????????

5. Scientific Notation: A method of representing very large or very small numbers in the form: M x 10 n M x 10 n  M is a number between 1 and 10  n is an integer

6. 2 500 000 000 Step #1: Insert an understood decimal point. Step #2: Decide where the decimal must end up so that one number is to its left up so that one number is to its left Step #3: Count how many places you bounce the decimal point the decimal point 1234567 8 9 Step #4: Re-write in the form M x 10 n

7. 2.5 x 10 9 The exponent is the number of places we moved the decimal.

8. 0.0000579 Step #2: Decide where the decimal must end up so that one number is to its left up so that one number is to its left Step #3: Count how many places you bounce the decimal point the decimal point Step #4: Re-write in the form M x 10 n 12345

9. 5.79 x 10 -5 The exponent is negative because the number we started with was less than 1.

10.  Rules for using scientific notation  Numbers are represented as a value between 1 <10  Exponents (+),(-) describe the decimal movement  (+) number grows larger  (-) number gets smaller Ex: 1.2 x 10 3 = 1200 Math function and scientific notation:  Be sure to use the [2 nd ][EE] function on your calculator! Ex. 5.0 3 x 10 14 / 2.9 x 10 7 = ?

11.  Standard Unit of measure  Metric based measurements Length = meter (m) Mass = kilograms (kg) Time = seconds (s) Temperature = Kelvin (K) Amount of substance = mole (mol) Electric Current = ampere (A)

12. QuantitySymbolAbbrev.Definition Length l m Length light path in 3.0 x 10 8 s (speed of light m/s) MassmkgBased on standard mass Timets Atomic clock: Based on Ce-133 isotope (clock-based on frequency of energy as electrons transition between orbitals) TemperatureTK kelvin: 0 K is absolute zero (no molecular movement. Triple pt. water = 273.15K AmountNmol Mole: amt of substance equivalent to 12g C-12. (Avogadro’s # 6.023 x 10 23 ) Electric CurrentlA Measures the amt. of electric charge passing a pt. in a electric circuit/time (Coulomb/s )

13. Atomic clock

14. Temperature:

15. Measurement produced by multiplying or dividing SI units. quantitysymb ol Abbrev.Derivation AreaAm2L x W VolumeVm3L X W X H DensityD g/cm3 (mL) Mass/volume Molar massM g/mole Mass/ amt. of substance Molar volume VmVm m3/mole Volume. /amt. of a substance EnergyEJoulesForce X Length

16.  The amount of space occupied by an object  Some useful relationships: 1Liter = 1dm 3 = 10 3 cm 3 = 1000mL 1cm 3 = 1mL

17. Volume:

18.  Describes the ratio of mass to volume  Density is a physical property that is independent to amount ( intensive property ) D = Mass/Volume= (g/mL or g/cm 3 )

19. Measuring density:  Solids -  Liquids  Displacement

20. Calculations using density:  Conversions using density: Ex. Aluminum has a density of 2.699 g/cm3 What is mass of a 11.25 cm3 block of Al? 11.25 cm 3 2.699 g= 30.36 g Al 1cm 3 What is the volume of a 3. 75g sample of Al? 3.75 g Al 1cm 3 = 1.38 cm 3 2.699g

21.  A relationship showing the equality between two different units that allows for conversion between these units.  Examples:  1 week = 7 days  1 dozen eggs = 12 eggs  ???

22.  Accuracy - close to the true value  Precision - close within measurements  % Error- describes how close your data is to the accepted value Accepted – Experimental x 100 Accepted Value

23. Accuracy & Precision

24.  Used to express the accuracy of a number.  Used for “measured” numbers.  Rules for determining Sig. Figs. 1. All non-zero numbers are significant 2. Captive zeros are significant 3. Leading zeros are not significant 4. Trailing zeros are significant if a decimal is present

25. Significant figures:  Determine the number of sig figs 1. 3060 2..000789 3. 1.40 x 10 6 4. 1000 5. 2030. 6. 45,0060

26. Answers:  Determine the number of sig figs 1. 3060 (3) 2..000789 (3) 3. 1.40 x 10 6 (3) 4. 1000 (1) 5. 2030. (4) 6. 45,0060 (5)

27.  Addition/Subtraction  Limited to the placement of the least significant figure. ex. 12.00 + 8.3 + 14 =  Multiplication/Division  Limited to the lowest number of significant figures. Ex. 12.0 x 8.3 X 14 =

28. Let’s set up a problem: What is the mass of a block of gold that has a volume of 2.0 cubic inches? 98 grams.. Worth $4,520.00!

29. Problem Solving  Steps to Problem Solving Approach: A. Analyze: determine starting point and plan steps required to get a solution (unknown). B. Plan: set up a strategy to solve the problem. C. Solve: conduct appropriate calculations based on plan. This may require multiple steps. D. Evaluate: Review answer to see if it seems reasonable.

30. Problem Solving  Conversion Factors - a ratio of equivalent measurements ( 1 inch = 2.54 cm)  Dimensional Analysis- the technique for solving problems using unit conversions based on conversion factors Ex. 6. 42 inches = ? cm

31. Problem Solving  Multi-step Problems – use more than one conversion factor: ex. 5 days = ? minutes  Complex Problems- involves ratios of two units: ex..45 Km/hrs = ? m/s Golden Rule for conversions…always show your work!