1. Scientific Measurement
Chemistry
Unit 2

2. Measurements
Qualitative Measurement – give results in a descriptive nonnumerical form
For example
It is hot today
It’s over yonder
I’ll be ready in a bit
That is too heavy to pick up
The new store is huge

3. Measurements
Qualitative measurements are relative to the person giving them, and can be confusing to other people
For example
If you live in Death Valley, it may not be hot
What is a yonder?
Should I start watching television?
I can lift that with one hand
Not compared to the mall

4. Measurements
Quantitative measurements – give results in a definite form. They contain both numbers and units
For example
The temperature is 99◦F
It’s 5 miles south
I will be ready in 2 minutes
That weighs 200 pounds
The new store is ½ a block

5. Measurements
For the purpose of science it is very important that we express measurements in a quantitative form
It is more precise
It is more accurate
It is reproducible
It removes any bias

6. Accuracy
Accuracy – is a measure of how close a measurement is to the actual or true value of that measurement
If we put a thermometer in a glass of ice water it should read 0◦C
This is the accepted value
Right away we can tell if our thermometer is accurate or not
Calibrating thermometers, football passes, strikes in baseball, standard weights and scales, soccer goal is big

7. Precision
Precision – a measure of how close a series of measurements are to each other
Requires more than one measurement
The closer the measurements are the greater the precision
If we put three thermometers in the same cup of ice water, and all three measure 1◦C, we can say our measurement is precise.
Is our measurement accurate?
Leading somebody in a football pass, strike zone or ball, soccer goal

8. Accuracy vs. Precision Accuracy is about the true or accepted value
Precision has nothing to do with the true value
Accuracy can be achieved with one measurement
Precision takes a series of measurements. Then are evaluated to see how close they are to each other

9. DOH!!!! Error Everybody makes a mistake!!!!!!!!
Simpson’s get a good Homer impression

10. Error Error = Experimental value – accepted value
Experimental value – the value obtained during a lab
Accepted value – the correct value based on reliable sources

11. Error Error can be positive or negative
If the experimental value is more than the accepted value the error is positive
If the experimental value is less than the accepted value the error is negative

12. Error Error is expressed as a percent
% error = (exp. Value – accep. Value) x 100%
accep. Value
% error = (980 – 1000) x 100% = -2%
1000
Error is a %, show equation, give ex. Of a light bulb w/1000 hrs only burning for 980.

13. Scientific Notation
Numbers in Chemistry can be extremely large or small, and cumbersome to express
For example Avagadro’s number:
602,200,000,000,000,000,000,000
Or the mass of one atom of carbon: g

14. Scientific Notation
We need an easier way to deal with these quantities
That is why we use scientific notation
Scientific notation is the product of two numbers
Avagadro’s number = x 1023
1 atom of carbon = x 10-23g

15. Scientific Notation Scientific notation is the product of two numbers
The first number is called the coefficient
The coefficient must be a number greater than or equal to 1 and less than 10

16. Scientific Notation The second number is 10 to an exponent
If the exponent is positive it indicates how many times you should multiply the coefficient by 10
If the exponent is negative it indicates how many times you should divide the coefficient by 10

17. Scientific Notation
Since our number system is in base 10 we can just move the decimal the number of times the exponent indicates
For a positive exponent we will move the decimal point to the right
3.12 x 103 = 3120
For a negative exponent we will move the decimal point to the left
1.56 x 10-4 =

18. Scientific Notation Multiplication of numbers in scientific notation
Multiply the coefficients
Add the exponents
(4.00 x 106) x (2.00 x 103) = x 109
(3.20 x 104) x (2.30 x 104) = x 108

19. Scientific Notation Division of numbers in scientific notation
Divide the coefficients
Subtract the exponents
9.00 x 108 / x 104 = x 104
8.17 x 107 / x 103 = x 104

20. Scientific Notation
Addition and subtraction of numbers in scientific notation
Remember the exponent means how many places to move the decimal point
The decimal points must be aligned to add
In order to line up the decimal points we need to make the exponents the same

21. Scientific Notation Take the following example
3.66 x x 103
If they were not in scientific notation they would look like this
Therefore we need to change one of the exponents to equal the other
0.366 x x 103

22. Scientific Notation 0.366 x 103 + 4.12 x 103
Now we can add the coefficients and the exponent will be 103
0.366 x x 103 = x 103

23. Scientific Notation
For subtraction the same rule applies, the exponents must be the same
Subtract the coefficients from each other
For example
3.66 x x 103
0.366 x x 103
0.366 x x 103 = x 103

24. Significant Figures
Significant figures – include all digits that are known plus a last digit that is estimated
We estimate all the time
We can only estimate 1 digit beyond what is known
In a measurement we must be able to determine the significant figures

25. Significant Figures
There are six rules for determining how many significant figures there are
Rule #1 – in a measurement every non zero digit is assumed to be significant
For example
All of these measurements have four significant figures

26. Significant Figures
Rule #2 – All zeros between non zero digits are significant
For example
10002
98.023
8.0604
580.06
All of these numbers have five significant figures

27. Significant Figures
Rule # 3 – All zeros to the left of the first non zero digit are not significant
For example
0.52
0.021
All of these numbers only have two significant figures

28. Significant Figures Rule # 3 cont.
Writing in scientific notation helps prevent confusion with the zeros
5.2 x 10-1
8.4 x 10-4
2.1 x 10-2
3.7 x 10-5

29. Significant Figures
Rule # 4 – Zeros at the end of a number, and to the right of a decimal point are significant
For example
58,001.10
All of these numbers have seven significant figures

30. Significant Figures
Rule # 5 – Zeros at the end of a number, and to the left of the decimal point are not significant
For example
1230.
107,000.
45,600
1,540,000
All these numbers have three significant figures

31. Significant Figures Rule # 5 – continued
Again this is why scientific notation helps prevent confusion
You know what is next
1.23 x 103
1.07 x 105
4.56 x 104
1.54 x 106

32. Significant Figures
Rule # 6 – Unlimited significant figures occur in two situations
The first situation is when something is counted
For example
There are ______ people in class today
There are 50 stars on the U.S. flag
There are six lab stations in the room
There are three doors in the room

33. Significant Figures Rule # 6 continued
The second situation with unlimited significant figures is when the number is an exactly defined quantity
For example
There are 100 yards between football end zones
There are four quarters in a dollar
There are 12 months in a year
There are 1,000 milliliters in a liter

34. Significant Figures
When we are using significant figures in calculations it is important to understand that the answer cannot be more precise than the measurements used in the calculation.
This is why we need to be able to determine how many significant figures are in a number

35. Significant Figures For example: You measure a room to be 6.6m by 5.9m
Both of these measurements have two significant figures
When you multiply them to find the area
6.6m x 5.9m = m2
The area has four significant figures
Is this a valid answer?

36. Significant Figures
No, because the calculation cannot be more precise than the measurements
To get the correct amount of significant figures, we must round
38.94m2 would be rounded to 39m2

37. Significant Figures Rules for rounding
If the digit after the correct amount of significant figures is between 0 – 4 just drop the rest of the digits
If the digit after the correct amount of significant figures is between 5 – 9 increase the last significant figure by 1 and drop the rest of the digits.

38. Significant Figures Multiplication and Division
Round the answer to the same amount of significant figures as the measurement with the least amount of significant figures

39. Significant Figures Addition and Subtraction
Round the answer to the number of decimal places (not digits) as the measurement with the least amount of decimal places

40. SI Units All measurements need a unit to be clearly understood
The metric system, established in France in 1790 uses base 10
The advantage of the metric system is that conversions are very easy

41. SI Units
In 1960 the International System of Units (abbreviated SI) was adopted by international agreement
It is a revision of the metric system
It consists of seven base units
All other units of measurement can be derived from these seven

42. SI Units The seven base units in the SI system are: Length – meter (m)
Mass – kilogram (k)
Temperature – kelvin (K)
Time – second (s)
Amount of substance – mole (mol)
Luminous intensity – candela (cd)
Electric current – ampere (A)

43. SI Units
From these base units we can derive other units with the use of a prefix
mega (M) – 1,000,000 times larger (106)
kilo (k) – 1,000 times larger (103)
deci (d) – 10 times smaller (10-1)
centi (c) – 100 times smaller (10-2)

44. SI Units Cont. milli (m) – 1,000 times smaller (10-3)
micro (μ) – 1,000,000 times smaller (10-6)
nano (n) – 1,000 million times smaller (10-9)
pico (p) – 1 trillion times smaller (10-12)

45. Density Density = mass / volume Density can be affected by temperature
Specific gravity = density of substance / density of water
There are no units for specific gravity
Specific gravity is measured by a hydrometer

46. Temperature Temperature determines the direction of heat transfer
Celsius (◦C) – named for Anders Celsius sets the freezing point of water at 0◦C and the boiling point at 100◦C

47. Temperature Recall that the SI unit for temperature is Kelvin
This temperature is named after Lord Kelvin
The freezing point of water is 273K and the boiling point is 373K
Notice the (◦) symbol is not used

48. Temperature 0K is known as absolute zero
It is the temperature where all molecular vibration ceases
Converting between Celsius and Kelvin
K = ◦C + 273
◦C = K - 273

49. Converting Units
Conversion factors are a way to change a measurements units into different units
They are expressed as a ratio and are equal to one
For example
1000mL 1L