1. Welcome to AP Chemistry !

2. DOR: Introductory Information
What was your favorite summer event?
Was the AP Chemistry summer assignment difficult for you?
As a review, what should I know about you so I can help you succeed in AP Chemistry? (ie. What learning methods are best for you?)

4. Scientific Measurements
SI Units
Significant Figures

5. Measurement system universal to scientists.
SI Units
Measurement system universal to scientists.
Measurement standards (base vs. derived)
Base unit—
Quantity we can MEASURE
Derived unit--
Quantity provided by CALCULATION
Quantities are represented by measurements.
We will use the gram (g) unit a lot !!!
Mass vs. weight
Weight = gravitational pull on matter
Derived SI Units = combine SI base units through multiplication or division.
100ml =1L, 1000 cm3 = 1 L so cm3= ml, we measure volumes a LOT in chemistry

6. SI Units (cont.) Base Mass = kilogram (kg) Length = Meter (m)
Time = Second (s)
Temperature = Kelvin (°K)
Derived
Volume = m3, we will use cm3/ml
ml is NOT an SI unit (1ml = 1cm3)
Density = kg/m3, we will use g/cm3 or g/ml

7. Scientific Measurements
SI Units
Significant Figures
ALWAYS WRITE YOUR UNITS ON YOUR ANSWER !!!!----units provide meaning to numbers.

8. Significant Figures
The number of digits in a scientific measurement known with certainty and one estimated or uncertain digit.
Method of reporting scientific measurements and calculations.

9. Significant Figure Basic Rules
1. All numbers that are NOT zeroes are significant
Ex. 211, 345 =3
2. All zeroes BETWEEN 2 significant figures are significant.
Ex. 2003, 1203 = 4
3. All zeros that are FINAL and PAST a decimal are significant.
Ex =2
4. Final zeros are only significant if a DECIMAL POINT is added to the end of the number.
Ex. 100 = 1
Ex = 3
Ex = 6
Rule #1: easiest to remember.
Rule #3: people have the hardest time with, have to have BOTH conditions.
**Put a check mark on significant figures.
#1 Example 3 sig figs
#2 Example 4, 4 sig figs
#3 Example 2 sig. figs.

10. Examples
1) 567 2) 3)
4) 200
**Remember to put a check mark on significant figures. 2 5 1

11. Significant Figures: Adding/Subtracting
# of digits to RIGHT of decimal = # of digits RIGHT of decimal from number with the LEAST digits.
Ex kg kg =
75.3 kg
Ex g – g =
10.62 g

12. Significant Figures: Multiplying/Dividing
# of significant figures = # of significant figures in number with LEAST significant figures.
Ex / 3.1 =
~ 3.9
Ex * =
~ 56.2

13. Significant Figure Practice---If Necessary
1) ) 804.5
2) )
3) ) 1002
4) )
5) ) g / 4.13 cm3 =
11) = ) 2.0 * 4.35 =
13) ) 31400
15) 3500
) 5
) 6
) 4
) 6
) 4.35 g/cm3
11) ) 8.7
13) ) 3
15) 2

14. Dimensional Analysis

15. Dimensional Analysis Method of converting one unit to another
Conversion factor—ratio relationship between 2 units
Write units, BE SURE THEY CANCEL ! ! !
“Where are we starting? Where are we trying to go?”

16. KNOW
2.54 cm = 1 inch

17. Example 1:
Science fiction often uses nautical analogies to describe space travel. If the starship U.S.S. Enterprise is traveling at warp factor 1.71, what is its speed in knots?
1.71 Warp = 5.00 times the speed of light
Speed of light = 3.00x108 m/s
1 knot = 2000 yd/h exactly
2.95 x 109 knots

18. Example 2:
Apothecaries use the following set of measures in the English system:
20 grains ap = 1 scruple
2 scruples = 1 dram ap
8 dram ap = 1 oz. ap
1 dram ap = 3.888g
What is the mass of 1 scruple in grams?

19. Example 3:
BMI is a measurement to determine if a person is obese. A BMI > 25 is considered overweight. Calculate the BMI for a 100 lb person with a height of 52 inches. Is this person overweight?
Hint: Do conversion factors FIRST, then cancel units

20. Conversion Factors: In Detail
= 1 value
12 inches = 1 ft.
2.54 cm = 1 inch
5280 ft. = 1 mile SO
(12 inches = 1 ft.) 3 = 13 = 1 same conversion factor

21. In Detail Cont.
Use same conversion factor but raise the factor to the correct power for conversion.
Cubic inches to ft3
(12 inches = 1 ft)3
When converting between squared and cubic units (m2 to cm2 /m3 to cm3)

22. Example 4:
Volume of 2.56 cm3 convert to m3

23. Temperature Conversions
°F = 1.8 (T°C) + 32
°K = T°C

24. Accuracy
How close measurements are to the true/accepted value.

25. Precision How close measurements are to each other.
These measurements are not necessarily accurate.

26. Percent Error % Error =(True – Experimental/True)x 100
How much error you have in your lab measurements.
Comparison between your experimental value/measurement to the true/accepted value.
% Error =(True – Experimental/True)x 100

27. Percent Error Problems
1) True value = 21.2g, Measured value = 17.7g
2) True value = 4.15ml, Measured value = 4.26 ml.
17% 2) 2.7%

28. Density and SI Units

29. Density Physical property of matter/substances
Used for substance identification
Provides information on how solids/liquids interact
Ratio of a substance’s mass and volume
Density = Mass/Volume
Units = SI Unit (kg/m3), we will use g/cm3 or g/ml

30. Example 1: A student records V1= 2. 7ml and V2= 3
Example 1: A student records V1= 2.7ml and V2= 3.4 ml after placing an object in a graduated cylinder. The mass of an empty beaker is 1.13g and the mass of both the beaker and substance is 4.13g. What is the object’s density?
4.3 g/ml

31. Example 2: It has been estimated that there are 4 x 10-6 mg of gold per liter of sea water. At a price of $22.30 per gram of gold, what would be the value of gold in 1.00 cubic kilometers of the ocean?  

32. Example 3: A metallic sphere has a diameter of 0
Example 3: A metallic sphere has a diameter of inches and a mass of ounces. What is the density of this object, in units of g/cm3?
2.72 g/cm3

33. Specific Gravity
Removes gravitational influence on a weighed object, density should not vary with weight !
Ratio of density of substance/density of reference
= Dsubstance/Dreference
Liquids/solids, Dwater = g/ml at 3.98°C (decreases with increase/decrease in temperature)
Gases, Dair = 1.29 g/L
At STP (1 atm, 0°C), 1 mole of gas = 22.4 L SO density can be calculated with molecular mass of gas.

34. Ex. 1 A solid object has a mass of 3. 04 kilograms and a volume of 3
Ex. 1 A solid object has a mass of 3.04 kilograms and a volume of 3.0 x 104 cm3. What is the density and specific gravity of the solid object?
Density = g/ml
Specific gravity = 0.101

35. Ex. 2 Phosphorus trifluoride is present at STP conditions
Ex. 2 Phosphorus trifluoride is present at STP conditions. What is its density and specific gravity?
3.93 g/L = density
Specific gravity = 3.05

36. Homework
Nitrogen dioxide density = 2.05 g/L
Specific gravity = 1.59