1. Properties of and Changes in Matter
Unit 1
Properties of and Changes in Matter

2. Chemistry Study of matter and its transformations
Everything is or is made from 118+ different types of atoms, in many combinations

3. Pure substances The simplest type of matter can be an element
represented by chemical symbols on the periodic table of elements
Examples: H (hydrogen), He (helium)
Smallest unit of an element is an atom
or a compound
chemical combinations of elements, represented by a chemical formula
Examples: H2O (water), NaCl (sodium chloride)
Smallest unit of a compound is a molecule or formula unit

4. Describing matter
Chemical properties
Physical properties

5. Chemical Properties Describe the type of changes matter can undergo
Become evident during a chemical reaction
Examples:
A chemical property of metals is the ability to react with acids.
A chemical property of carbon dioxide gas is that no combustion reaction can take place in its presence.

6. Physical Properties Describe bulk quantities, not single atoms
Observed or measured without changing the composition
Review of properties

7. Physical Properties Examples: Color
State of matter (solid, liquid, gas)
Luster
Texture (granular, powdery)
boiling point and melting point
Solubility
Density
Size (mass, volume, length)

8. Measurement
A quantitative observation that includes a number and a unit

9. Mass a measure of the amount of matter in an object
measured using an electronic balance
Record all numbers on the screen, even zeroes

10. Volume a measure of the amount of space occupied
Use a graduated cylinder for liquids
Use v = l x w x h for regular solids
Use water displacement for irregular solids

11. Physical Properties Intensive (also called intrinsic) OR
INdependent of sample size
Examples: color, state of matter, luster, texture, boiling point, melting point, solubility, density OR
Extensive (also called extrinsic)
dependent on sample size
Examples: mass, volume, length

12. Changes Physical changes
Usually involve small amounts of energy (compared to chemical changes)
Involve the particles moving closer together or farther apart
Don’t change the identity of the substance
Only change physical properties of substance

13. Physical change examples
all phase changes (freezing, melting, boiling/evaporation, condensation, depostion, sublimation)
changing size or shape (rolling out a ball of play-do or cutting a piece of paper)
warming or cooling (like heating water from 50°C to 60°C)
Dissolving (sometimes…chemists are still arguing about this)

14. Changes Chemical changes involve more energy than physical changes
Result in a different substance than you started with

15. Chemical change indicators (only need one, but more than one may be seen)
Formation of a gas (will see bubbles)
Formation of a precipitate (an insoluble substance)
Formation of a new odor
Release of light (energy)
Internal temperature change(hotter or colder)
Unexpected color change

16. Review Brain pop movie (measuring matter and property changes)
Tutorial and online practice

17. Systems
During a chemical reaction, the SYSTEM is everything you are considering part of the reaction. Everything else is part of the SURROUNDINGS.
An open system interacts with its surroundings and allows energy or matter to enter or leave.

18. Density Ratio of mass to volume Describes the spacing of the particles
As the spacing between particles decreases, the density increases

19. Density Which container is more dense? Which container has more mass?
Which container has more volume?

20. Density Which container is more dense? Which container has more mass?
Which container has more volume?
If you cut sample B in half, what happens to the:
Mass of B?
Volume of B?
Density of B?

21. Density TRUE OR FALSE Mass of A is greater than Mass of B.
Mass of A is greater than Mass of C.
Mass of B is greater than Mass of C.
Volume of A is less than Volume of B.
Volume of A is less than Volume of C.
Volume of B is less than Volume of C.
Density of A is greater than Density of B.
Density of B is greater than Density of C. c A B

22. Density calculations

23. Density of water Density of water is 1.00 g/ml or 1.00 g/cm3
not drawn to scale!
1 cm3 = 1 mL = 1 cc

24. Density calculations
Find the density of a material that has a mass of 100 grams and takes up 25 cubic centimeters of space. Include: formula, substitution, and answer with units

25. Density calculations
What is the density of the solution in the container? Include: formula, substitution, and answer with units More practice

26. Density graph Density of water Keys to a good line graph: -title
Calculate the slope of the line.
Include: formula, substitution, and answer with units
Keys to a good line graph:
-title
-labelled axes
-even increments
-Use all available space
-plot data points
-circle each data point
-draw a best-fit line through circles
-choose 2 points (NOT data points) from best-fit line for calculating slope
Density of water

27. Metric prefixes King Henry Died By Drinking Chocolate Milk
King - Kilo (K + base unit)
Henry – Hecta (H + base unit)
Died- Deka (Da + base unit)
By – Base – meter, liter, gram
Drinking – Deci (d + base unit)
Chocolate – Centi ( c + base unit)
Milk – Milli (m + base unit)

28. Tera-
1012
Giga-
109
Mega-
106
Micro-
10-6
Nano-
10-9

29. How many jumps does it take?
Ladder Method 1 2 3
KILO 1000 Units
HECTO 100 Units
DEKA 10 Units
DECI 0.1 Unit
Meters Liters Grams
CENTI 0.01 Unit
MILLI Unit
How do you use the “ladder” method?
1st – Determine your starting point.
2nd – Count the “jumps” to your ending point.
3rd – Move the decimal the same number of jumps in the same direction.
4 km = _________ m
Starting Point
Ending Point
How many jumps does it take? 4. 1
__. 2
__. 3
__.
= 4000 m

30. Conversion practice 1000 mg = ________ g 14 kg = __________ g
1 L = __________ mL
109 g = __________ kg
160 cL = __________ mL
250 L = __________ kL

31. accuracy
precision
Correctness
Agreement with the true/accepted value
Check by using a different method
Poor accuracy results from procedural or equipment flaws
Reproducibility
Degree of agreement among several measurements of the same quantity
Check by repeating measurements
Poor precision results from poor technique

32. Accuracy & Precision Group 1 Group 2 Group 3 Accuracy ______
Three different groups of students measure the mass of a medal, with a known value of grams. Evaluate each group’s data for its accuracy and precision (low or high):
Group 1
Group 2
Group 3
Trial 1
5.003 g
Trial 2
5.002 g
Trial 3
5.001 g
Trial 1
5.400 g
Trial 2
5.202 g
Trial 3
5.905 g
Trial 1
5.503 g
Trial 2
5.499 g
Trial 3
5.501 g
Accuracy ______
Precision ______
high
Accuracy ______
Precision ______
low
Accuracy ______
Precision ______
high
low

34. Accuracy and Precision
Sometimes there is a difference between the accepted value and the experimental value.
This difference is known as error.
Error = accepted value – experimental value
Error can be positive or negative depending on whether the experimental value is greater than or less than the accepted value.

35. Accuracy and Precision
Often it is useful to calculate relative error, or percent error.
Percent error = error x 100%
accepted value
The percent error will always be a positive value.

36. Good measurements Read graduated cylinder at eye level
Record volume at the bottom of the meniscus

37. Sig Figs and Song Any measurement has some degree of uncertainty.
When recording measurements, always include one estimated (uncertain) digit (place value position).
Example: This graduated cylinder has markings every 1 mL. You estimate the tenths place.
43.0 mL

38. Significance in Measurement
Measurements always involve a comparison.
When you say that a table is 6 feet long, you're really saying that the table is six times longer than an object that is 1 foot long.
The foot is a unit; you measure the length of the table by comparing it with an object like a yardstick or a tape measure that is a known number of feet long.

39. Significance in Measurement
The comparison always involves some uncertainty.
If the tape measure has marks every foot, and the table falls between the sixth and seventh marks, you can be certain that the table is longer than six feet and less than seven feet.
To get a better idea of how long the table actually is , though, you will have to read between the scale division marks.
Measurements are often written as a single number rather than a range.
When you write the measurement as a single number, it's understood that the last figure had to be estimated. Consider measuring the length of the same object with two different rulers.

40. Significance in Measurement
For each of the rulers, give the correct length measurement for the steel pellet as a single number rather than a range
Ruler on left…1.4 in.
Ruler on right…1.47 in.

41. Significance in Measurement
A zero will occur in the last place of a measurement if the measured value fell exactly on a scale division.
For example, the temperature on the thermometer should be recorded as 30.0°C.
Reporting the temperature as 30°C would imply that the measurement had been taken on a thermometer with scale marks 100°C apart!

42. Significance in Measurement
The graduated cylinder on the right has scale marks 0.1 mL apart, so it can be read to the nearest 0.01 mL.
Reading across the bottom of the meniscus, a reading of 5.72 mL is reasonable (5.73 mL or 5.71 mL are acceptable, too).

43. Significance in Measurement
Determine the volume readings for the two cylinders to the right, assuming each scale is in mL.
Cylinder on left 3.0 mL
Cylinder on left 0.35 mL

44. Signficance in Measurement
Numbers obtained by counting (exact numbers) have no uncertainty unless the count is very large.
For example, the word 'sesquipedalian' has 14 letters.
"14 letters" is not a measurement, since that would imply that we were uncertain about the count in the ones place.
14 is an exact number here.
Conversion factors are exact numbers and have no uncertainty.
Conversions (3 feet = 1 yard)

45. Significance in Measurement
Very large counts often do have some uncertainty in them, because of inherent flaws in the counting process or because the count fluctuates.
For example, the number of human beings in Arizona would be considered a measurement because it can not be determined exactly at the present time.

46. Significance in Measurement
All of the digits up to and including the estimated digit are called significant digits.
Consider the following measurements. The estimated digit is in purple and is underlined:
Measurement Number of Distance Between Markings
Significant Digits on Measuring Device
142.7 g g
103 nm nm
x 108 m x 108 m

47. Significance in Measurement
A sample of liquid has a measured volume of mL. Assume that the measurement was recorded properly.
How many significant digits does the measurement have? 1, 2, 3, or 4?
The correct answer is . . . 4

48. Significance in Measurement
Suppose the volume measurement (23.01 mL) was made with a graduated cylinder. How far apart were the scale divisions on the cylinder, in mL? 10 mL, 1 mL, 0.1 mL, or 0.01 mL?
The correct answer is . . .
0.1 mL

49. Significance in Measurement
Which of the digits in the measurement (23.01 mL) is uncertain? The “2,” “3,” “0,” or “1?”
The correct answer is . . . 1

50. Significance in Measurement
Usually one can count significant digits simply by counting all of the digits up to and including the estimated digit.
It's important to realize, however, that the position of the decimal point has nothing to do with the number of significant digits in a measurement.
For example, you can write a mass measured as g as kg.
Moving the decimal place doesn't change the fact that this measurement has FOUR significant figures.

51. Significance in Measurement
Suppose a mass is given as 127 ng.
That's µg, or mg, or g.
These are all just different ways of writing the same measurement, and all have the same number of significant digits: THREE.

52. Significance in Measurement
If significant digits are all digits up to and including the first estimated digit, why don't those zeros count?
If they did, you could change the amount of uncertainty in a measurement that significant figures imply simply by changing the units.
Say you measured 15 mL
The 10’s place is certain and the 1’s place is estimated so you have 2 significant digits
If the units to are changed to L, the number is written as L
If the 0’s were significant, then just by rewriting the number with different units, suddenly you’d have 4 significant digits
By not counting those leading zeros, you ensure that the measurement has the same number of figures (and the same relative amount of uncertainty) whether you write it as L or 15 mL

53. Significance in Measurement
Determine the number of significant digits in the following series of numbers:
kg = g = 341 mg
12 µg = g = kg
Mg = kg = g
3, 2, 4

54. Significance in Measurement
Any zeros that vanish when you convert a measurement to scientific notation were not really significant figures. Consider the following examples: kg
1.234 x 10-2 kg
4 sig figs
Leading zeros ( kg) just locate the decimal point. They're never significant.

55. Significance in Measurement
kg x 10-2 kg 5 sig figs Notice that you didn't have to move the decimal point past the trailing zero ( kg) so it doesn't vanish and so is considered significant m x 10-5 m Again, the leading zeros vanish but the trailing zero doesn't.

56. Significance in Measurement
84,000 mg
8.4 x 104 mg
2 sig figs (at least)
The decimal point moves past the zeros (84,000 mg) in the conversion. They should not be counted as significant.
32.00 mL
3.200 x 101 mL
4 sig figs
The decimal point didn't move past those last two zeros.

57. Significance in Measurement
When are zeros significant?
From the previous frame, you know that whether a zero is significant or not depends on just where it appears.
Any zero that serves merely to locate the decimal point is not significant.

58. Significance in Measurement
All of the possibilities are covered by the following rules:
1. Zeros sandwiched between two significant digits are always significant.
km kg Mg
5, 4, 6

59. Significance in Measurement
2. Trailing zeros to the right of the decimal point are always significant.
3.0 m µm µm
2, 5, 5
3. Leading zeros are never significant.
m µm µm
1, 3, 5

60. Significance in Measurement
4. Trailing zeros with no decimal are not significant.
3000 m µm 92,900,000 miles 1, 3, 3

61. Significance in Measurement
How many significant figures are there in each of the following measurements? g g g g g
100 g
7, 6, 6, 4, 3, 1

62. Significance in Measurement
Rounding Off
Often a recorded measurement that contains more than one uncertain digit must be rounded off to the correct number of significant digits.

63. Significance in Measurement
Rules for rounding off measurements:
1. All digits to the right of the first uncertain digit have to be eliminated. Look at the first digit that must be eliminated.
2. If the digit is greater than or equal to 5, round up.
g rounded to 2 figures is 1.4 g.
1090 g rounded to 2 figures is 1.1 x 103 g.
g rounded to 3 figures is 2.35 g.
3. If the digit is less than 5, round down.
g rounded to 4 figures is g.
1090 g rounded to 1 figures is 1 x 103 g.
g rounded to 5 figures is g.

64. Significance in Measurement
Try these:
rounded to 3 figures 1,756,243 rounded to 4 figures rounded to 2 figures
x x 101

65. Significant figures
When using or evaluating someone else’s measurements, you need to understand the following significant figure rules:

66. Sig fig rules
Non-zero digits and zeros between non-zero digits are always significant.
Example:
145.6 grams has 4 sig figs
has 6 sig figs

67. Leading zeros are not significant.
Example:
has 4 sig figs
has 3 sig figs

68. Zeros to the right of all non-zero digits are only significant if a decimal point is shown.
Example:
has 6 sig figs
993,000 has 3 sig figs

69. For values written in scientific notation, the digits in the coefficient are significant.
Example:
4.65 x 108 has 3 sig figs
4.650 x 108 has 4 sig figs

70. Scientific notation

71. Scientific notation
Expresses numbers as a coefficient multiplied by a base raised to a power. Example: 456 = 4.56 x 102 Positive exponents indicate that a value is greater than = 4.56 x 10-3 Negative exponents indicate that a value is less than 1.