1. How can we keep track of all of this “stuff”?
Measuring Matter
How can we keep track of all of this “stuff”?

2. Measuring Matter in Two Ways
Qualitative Measurements: which are usually descriptive like observations.
Now it is important to start making…
Quantitative Measurements: are in the form of numbers and units.

3. The Powers of Ten
Picture a microscopic cell.
Picture the galaxy.

4. A Short History of Standard Units
Humans did not always have standards by which to measure temperature, time, distance, etc.
It wasn’t until 1790 that France established the first metric system.

5. The First Metric System
The French established that one meter was one ten-millionth of the distance of the from the Equator to the North Pole.
One second was equal to 1/86,400 of the average day.

6. Today’s Standards
The techniques used today to establish standards are much more advanced than in 1790…
One meter is equal to the distance traveled by light in a vacuum in 1/299,792,458 of a second.
One second is defined in terms of the number of cycles of radiation given off by a specific isotope of the element cesium.

7. S.I. Units
The International System of Units is used ALMOST exclusively worldwide (the U.S. is one the of the exceptions).
ALL science is done using S.I. units.

8. The United States’ System of Measurement
In 1975, the U.S. government attempted to adopt the metric system with little success.
The U.S. currently uses the English System of Measurement.

9. So why S.I.?

10. So why S.I.? Decimals are more “computationally friendly”
Multiples of ten
Eliminate LARGE numbers by using prefixes
Scientifically based

11. Measurements and SI Units
Quantitative measurements must include a number AND a unit.
Base units are used with prefixes to indicate fractions or multiples of a unit.
Try to fill in your table.

12. Base Units Base Unit Symbol Measures… kilograms kg mass meter m
distance
Liter L
volume
Kelvin K
temperature
degrees Celsius °C
mole
mol
amount of a substance
second
S or sec
time
ampere A
electric current
candela cd
light intensity
Joule J
energy

13. Prefixes
Prefixes combine with base units to indicate fractions or multiples of a unit.

14. Prefixes Prefix Symbol Meaning M K H D d c m µ n p f mega-
times larger
kilo- K
1 000 times larger
Hecto- H
100 times larger
deca- D
10 times larger than base unit
Base Unit
deci- d
10 times smaller than base unit
centi- c
100 times smaller
milli- m
1 000 times smaller
micro-
times smaller
nano- n
times smaller
pico- p
times smaller
femto- f
times smaller

15. SI Prefixes
yotta- zetta- exa- peta- tera- giga- mega- kilo- hecto- deca- deci- centi- milli- micro- nano- pico- femto- atto- zepto- yocto-
Y Z E P T G M k h da d c m µ n p f a z y

16. More Details: Length meters centimeters (for smaller units of length)
millimeters (very small units of length)
kilometers (for large units of length)
These are the most commonly used.

17. More details: Mass
gram
kilogram
Measured using balances.

18. More details: Volume liters milliliters (for small volumes)
microliter (for extremely small volumes)
Measured using a graduated cylinder, pipet or buret (more accurate), volumetric flask or even a syringe.

19. Volume is a derived units…
Some metric units are derived from S.I. units.
Volume is L x W x H = cm x cm x cm = cm3
One cm3 is the same as one mL.
Also, dm x dm x dm = dm3
One dm3 is the same as one L.

20. Conversions Using Factor Label Method
Multiplying any number by an equality does NOT change the value.
An equality is two measurements that are equal in amount but have different units and numbers.
Examples:
one dozen bagels = 12 bagels
10 mm = 1 cm

21. Steps of Factor Label Method
Write down the units you are given.
Write X and a Line.
Write unit you want to cancel on the bottom of the line, the unit you want to keep on the top of the line. Find and plug in your equality. (hint: the larger unit will always get a 1 next to it)
Cancel units and do the math.
Voila!

22. Let’s do one together… 0.600L = _______mL 1000 mL 0.600L X = 600 mL 1
TRY THE REST ON YOUR OWN !!!!

23. Temperature Scales
There are three temperature scales in use in this country that you need to be familiar with.

24. Temperature:
A measure of the average kinetic energy of the particles in a sample.

25. Fahrenheit 18th-century German physicist Daniel Gabriel Fahrenheit
Based his scale on an ice-salt mixture and normal body temperature
Freezing point for water = 32°F
Boiling point for water = 212°F

26. Celsius Scale Swedish guy, Anders Celsius in 174
Freezing point at 0°C.
Boiling for water at 100°C.
Below 0 is negative.

27. Kelvin Scale English guy, William Kelvin Measures molecular movement
Theoretical point of ABSOLUTE ZERO is when all molecular motion stops (no negative numbers)
Divisions (degrees) are the same as in Celsius

28. Absolute Zero
Theoretical point where there is absolutely no movement of molecules in matter and a measure of ZERO ENERGY
This is not something that we ever witness, scientists have only theorized this point

29. Conversion Factors You need to know these conversion factors!
K = °C + 273
°C = K – 273
On Table T
NOT on Table T

30. Practice Conversion Problems
Room temperature is approximately 23°C. What is this temperature in Kelvin?
Ethanol has a boiling point of 351 K. That won’t help us if we have a thermometer reading degrees Celsius, so convert it.

31. Work on the conversion problems in your packet…

32. Uncertainty in Measurement

33. Uncertainty in Measurement
No measurement can be perfect.
Scientists need to account for some degree of uncertainty in measurements.
Refer to terms accuracy and precision.
An ideal measuring device is accurate and precise and does not have a great deal of uncertainty.

34. Accuracy
Accuracy is when you are close to the actual value of what you are trying to measure (For example you throw three darts and they are all close to the bull’s eye).

35. Precision
Precision is a measure of how close each measurement is to the others. For example if you are at the driving range and all of the golf balls head towards the pond, that is precision (but not accuracy).

37. Uncertainty in Measurement
Uncertainty occurs in every measurement made and must be accounted for.
In chemistry, we use different tools, each of which has certain limitations.
We use the ± to indicate uncertainty in measurement.

38. Uncertainty Equation
YOU MUST MEMORIZE THIS EQUATION AS YOU WILL PERFORM IT ALMOST DAILY!
ΔF = (δF / δX1) ΔX1 + (δF / δX2) ΔX2 + …(δF / δXn) ΔXn

39. Just Kidding!
However, there is a system we use in chemistry that helps to minimize uncertainty by only including those values that have certainty and one that is uncertain.
We call them SIGNIFICANT FIGURES!

40. Using Significant Figures
Why are significant figures important?
Have you ever multiplied two numbers and come up with a really LONG decimal?
Well, those numbers are INSIGNIFICANT with respect to scientific calculations.
And now for a short story…

41. Counting Sig Figs The Atlantic-Pacific Rule
If the decimal is Absent (A), start counting on the Atlantic (right) side. Go to the first NON-zero number and count everything after that.
If the decimal is Present (P), start counting on the Pacific (left) side. Go to the first NON-zero number and count everything after that.

42. Quick Self-Assessment
How many sig figs are in each of these numbers:
98,000 m
0.123 L L
8765 cm
40,506 m
20.00 mL

43. More Practice
Round each number to the number of sig figs shown in parentheses.
m (4)
m (2)
8792 m (2)

44. Sig Figs in Calculations
When adding or subtracting measurements, report to the LEAST number of DECIMAL PLACES.
For example: =
You will report this with one decimal place as

45. Sig Figs in Calculations
When multiplying or dividing measurements, you want to report the answer with the same number of sig figs as the measurement with the least number of sig figs.
For example:
7.55 m x 0.34 m = square meters
You will report this with 2 sig figs as 2.6 square meters because 0.34 m contains 2 sig figs.

46. Percent Error in Experimentation
When trying to determine how accurate your experimental value (“what you got in the lab”) is compared with the theoretical (“what it is supposed to be”), we use a simple formula.
experimental – theoretical x 100%
theoretical
Theoretical VS Experimental

47. Example of Percent Error% %
When you calculate the density of chemical X experimentally you get 1.13 g/mL. The actual density according to the literature is 1.16 g.mL. What is your percent error?

48. Common Units = g/mL or g/cm3
Density
Density is a derived unit which is found by dividing a substances mass by its volume.
Density Equation
D=M/V
Common Units = g/mL or g/cm3

49. Density Practice Problems
A student measures the mass of a piece of metal to be 4.0g and it has a volume of 1.5mL what is the density of this metal?
The Density of CO2 gas is 1.8 grams per liter. What is the mass of 0.2L of CO2 gas?