1. Chapter 1
Matter & Measurement

2. Chemistry is…
…the study of the composition, structure, and properties of matter and the changes it undergoes
C2H5OH O2  2 CO2 + 3 H2O + Energy
Reactants  Products

3. Matter Mass A measure of the amount of matter
Anything that has mass and occupies
space
Mass
A measure of the amount of matter

4. Atom Element The smallest unit of an element that maintains
the properties of that element
Element
A pure substance made of only one kind of atom

5. Properties of Matter Extensive properties
depend on the amount of matter that is present.
Extensive properties
Volume
Mass
Energy Content (think Calories!)
do not depend on the amount of matter present.
Intensive properties
Melting point
Boiling point
Density

6. Physical Change
A change in a substance that does not involve a change in the identity of the substance.
Example:
Phase Changes

7. Phase Differences
Solid – definite volume and shape; particles packed in fixed positions.
Liquid – definite volume but indefinite shape; particles close together but not in fixed positions
Gas – neither definite volume nor definite shape; particles are at great distances from one another
Plasma – high temperature, ionized phase of matter as found on the sun.

8. Three Phases

9. Copper Phases - Solid

10. Copper Phases - Liquid

11. Copper Phases – Vapor (gas)

12. Chemical Change
A change in which one or more substances are converted into different substances.
Heat and light are often evidence of a chemical change.

15. Separation of a Mixture
The constituents of the mixture retain their identity and may be separated by physical means.

16. Separation of a Mixture
The components of dyes such as ink may be separated by paper chromatography.

17. Filtration:

18. Separation of a Mixture
Distillation

19. Separation of a Compound The Electrolysis of water
Compounds must be separated by chemical means.
With the application of electricity, water can be separated into its elements
Reactant  Products
Water  Hydrogen + Oxygen
H2O  H O2

20. Measurement

21. Nature of Measurement Part 1 - number Part 2 - scale (unit) Examples:
Measurement - quantitative observation
consisting of 2 parts
Part 1 - number
Part 2 - scale (unit)
Examples:
20 grams
6.63 x Joule seconds

22. The Fundamental SI Units (le Système International, SI)

23. SI Prefixes Common to Chemistry
Unit Abbr.
Exponent
Kilo k
103
Deci d
10-1
Centi c
10-2
Milli m
10-3
Micro 
10-6

24. Temperature Scales

25. The Thermometer
Determine the temperature by reading the scale on the thermometer at eye level.
Read the temperature by using all certain digits and one uncertain digit.
Certain digits are determined from the calibration marks on the thermometer.
The uncertain digit (the last digit of the reading) is estimated.
On most thermometers encountered in a general chemistry lab, the tenths place is the uncertain digit.

26. Do not allow the tip to touch the walls or the bottom of the flask.
If the thermometer bulb touches the flask, the temperature of the glass will be measured instead of the temperature of the solution. Readings may be incorrect, particularly if the flask is on a hotplate or in an ice bath.

27. Reading the Thermometer
Determine the readings as shown below on Celsius thermometers:
_ _ . _ C 8 7 4
_ _ . _ C 3 5

28. Volume Instruments

29. Reading the Meniscus
Always read volume from the bottom of the meniscus. The meniscus is the curved surface of a liquid in a narrow cylindrical container.

30. Try to avoid parallax errors.
Parallax errors arise when a meniscus or needle is viewed from an angle rather than from straight-on at eye level.
Correct: Viewing the meniscus at eye level
Incorrect: viewing the meniscus from an angle

31. Measuring Volume
Determine the volume contained in a graduated cylinder by reading the bottom of the meniscus at eye level.
Read the volume using all certain digits and one uncertain digit.
Certain digits are determined from the calibration marks on the cylinder.
The uncertain digit (the last digit of the reading) is estimated.

32. Use the graduations to find all certain digits
There are two unlabeled graduations below the meniscus, and each graduation represents 1 mL, so the certain digits of the reading are…
52 mL.

33. Estimate the uncertain digit and take a reading
The meniscus is about eight tenths of the way to the next graduation, so the final digit in the reading is
0.8 mL
The volume in the graduated cylinder is
52.8 mL.

34. 10 mL Graduate
What is the volume of liquid in the graduate? 6
_ . _ _ mL 6 2

35. Uncertainty in Measurement
A digit that must be estimated is called uncertain. A measurement always has some degree of uncertainty.

36. Why Is there Uncertainty?
Measurements are performed with instruments
No instrument can read to an infinite number of decimal places
Which of these balances has the greatest uncertainty in measurement?

37. Precision and Accuracy
Accuracy refers to the agreement of a particular value with the true value.
Precision refers to the degree of agreement among several measurements made in the same manner.
Neither accurate nor precise
Precise but not accurate
Precise AND accurate

39. Significant Digits: Atlantic/ Pacific Rule
When the decimal is ABSENT, go to the Atlantic side of the number, start counting digits when you reach a non-zero number. Record

40. Practice
45,000 hrs
78,700 kilometers
3,000 liters
two
three
one

41. Pacific - when the decimal is PRESENT, go to the Pacific side of the number, start counting digits when you reach a non-zero number

42. Practice
grams mL
100.0 meters
.00506 4 5
none

43. Rules for Counting Significant Figures - Details
Nonzero integers always count as significant figures.
3456 has
4 sig figs.

44. Rules for Counting Significant Figures - Details
Zeros
- Captive zeros always count as
significant figures.
16.07 has
4 sig figs.

45. More practice 56,000,000 seconds 33,000 candles 60900 milligrams
centimeters
meters 2
0 * 3 6 4

46. Sig Fig Practice #1 1.0070 m  5 sig figs 17.10 kg  4 sig figs
How many significant figures in each of the following?
m 
5 sig figs
17.10 kg 
4 sig figs
100,890 L 
5 sig figs
3.29 x 103 s 
3 sig figs
cm 
2 sig figs
3,200,000 
2 sig figs

47. Density – a derived measure value is found through mathematical computation D=Mass/ Volume

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

49. . 9 8 7 6 5 4 3 2 1
Step #1: Insert an understood decimal point
Step #2: Decide where the decimal must end
up so that one number is to its left
Step #3: Count how many places you bounce
the decimal point
Step #4: Re-write in the form M x 10n

50. 2.5 x 109
The exponent is the number of places we moved the decimal.

51. 0.0000579 1 2 3 4 5 Step #2: Decide where the decimal must end
up so that one number is to its left
Step #3: Count how many places you bounce
the decimal point
Step #4: Re-write in the form M x 10n

52. 5.79 x 10-5
The exponent is negative because the number we started with was less than 1.