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 element compound
represented by chemical symbols on the periodic table of elements
Examples: H (hydrogen), He (helium)
Smallest unit of an element is an atom
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
L need to know first 20 elements and symbols
K need to know first 36 elements and symbols

4. Mixtures A physical blend of 2 or more pure substances
POGIL activity: classification of matter

5. Describing matter (3:03) Chemical properties Physical properties
Link to Mr. Edmonds song

6. 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.

7. Physical Properties Describe bulk quantities, not single atoms
Observed or measured without changing the composition
Review of properties (2:37)
Link to prezi

8. 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)

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

10. Mass a measure of the amount of matter in an object
measured using an electronic balance
Record all numbers on the screen, even zeroes
Measurement practical at lab tables

11. 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

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

13. 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

14. 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)

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

16. 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

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

18. Systems
During a chemical reaction (change), 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.

19. Law of Conservation of Mass
Matter can be changed from one form into another, mixtures can be separated or made, and pure substances can be decomposed, but the total amount of mass remains constant.
The total mass of the universe is constant within measurable limits; whenever matter undergoes a change, the total mass of the products of the change is, within measurable limits, the same as the total mass of the reactants. 

20. Law of Conservation of Mass

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

22. Methods for Determining Density
Observe (density column, float or sink)
Particle diagram (spacing of particles)
Calculation (d=m/v)
Graph of volume vs. mass (calculate slope)

23. Density Which container is more dense? Which container has more mass?B
Which container is more dense?
Which container has more mass?
Which container has more volume?

24. 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?

25. 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

26. Density calculations (2:33)
Link to Mr. Edmonds song

27. 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

28. 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

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

30. Density graph Slope = (y2 – y1) / (x2 – x1) Density of water
Calculate the slope of the line.
Include: formula, substitution, and answer with units
Slope = (y2 – y1) / (x2 – x1)
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 plotted data points) from best-fit line for calculating slope
Density of water

31. 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)

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

33. 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
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

35. 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.

36. 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.

37. Scientific notation

38. 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.

39. STUDENT HANDOUT STOPS HERE. YOU SHOULD KEEP GOING!

40. Good measurements Read graduated cylinder at eye level
Record volume at the bottom of the meniscus (3:04)
Link to Mr. Edmonds song

41. Sig Figs and Song (2:26)
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
Link to Mark Rosengarten song

42. Significance in Measurement
Record the volume of the purple liquid in the graduated cylinder, in milliliters.
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. 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)

44. Significance in Measurement
All of the digits up to and including the estimated digit are called significant digits (or figures).
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

45. How many sig figs in 23.01? 4

46. Graduated cylinder reading by expert student is 23.01 ml.
How far apart were the scale divisions on the cylinder, in mL?
Every 10 mL
Every 1 mL
Every 0.1 mL
Every 0.01 mL
0.1 ml

47. Which digit is uncertain in 23.01 mL?
The 1

48. How many sig figs? a. 0.000341 kg = 0.341 g = 341 mg
b. 12 µg = g = kg
c Mg = kg = g
3, 2, 4

49. 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
These are on the bottom of the reference materials

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

51. 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

52. 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

53. Rounding
(1:35)
song

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

55. Dimensional Analysis An organized method for converting between units.
Uses “conversion factors”
Example: m = 1 km
Can be written as a fraction:
1000m/1km or 1 km/1000m

56. Dimensional Analysis Example: Convert 25 m into km.
25 𝑚 1 × 1 𝑘𝑚 1000𝑚 = ?
The number and unit that you want to convert goes first. (It can be written as a fraction by placing a 1 as the denominator.)
Find a conversion factor that relates the unit you started with and the unit you want in your answer. (Sometimes, you will need more than one conversion factor.)
1 km = 1000m
Write the conversion factor as a fraction, with the wanted unit in numerator and the given unit in denominator.
Solve.
(25x1)/(1x1000) = 0.025
All units, except the desired unit, should cancel.
25 m = km

57. Conversion practice 1256 mg = ________ g 14 kg = __________ g
1.89 L = __________ mL

58. 109 g = __________ kg
cL = __________ mL
L = __________ kL

59. In chemistry, the term used to describe how much matter we have is the…..
Mole
In chemistry, the most important term used to describe more than one of something and/or how much matter we have is the mole.

60. 1 mole = 6.02 X 10 23 particles What’s a Mole?
Atoms
6.02 x 1023 particles
Smallest particle of an element that maintains the identity (#protons) of that element.
The mole is to chemistry as the dozen is to eggs.
The mole is defined as 6.02 x 1023 particles. (click)
In chemistry we have a variety of different particles that we keep track of.
For this unit we will focus on atoms. Recall that atoms are the smallest particles of an element.
Don’t worry about memorizing Avogadro’s, you can locate this number in the constant and conversions section on the Chemistry EOC Reference Materials.
© 2007 Microsoft Corporation. All rights reserved. Microsoft, Windows, Windows Vista and other product names are or may be registered trademarks and/or trademarks in the U.S. and/or other countries.
The information herein is for informational purposes only and represents the current view of Microsoft Corporation as of the date of this presentation. Because Microsoft must respond to changing market conditions, it should not be interpreted to be a commitment on the part of Microsoft, and Microsoft cannot guarantee the accuracy of any information provided after the date of this presentation. MICROSOFT MAKES NO WARRANTIES, EXPRESS, IMPLIED OR STATUTORY, AS TO THE INFORMATION IN THIS PRESENTATION.

61. How big is a Mole?
(4:32)

62. 4/23/2017 4:13 AM
How can we determine how many moles of matter we have in a sample of matter?
Particles and atoms are too small to be counted by hand….but, we can
measure mass (and volume).
We need to know the relationship between moles and mass.
6.02 x 1023 particles
So if use our 1 dozen equals 12 eggs analogy. How can we determine how many moles of matter we have in a sample of matter?
We can relate the # particles to # moles using Avogadros number.
We also know that particles are so small, there is no way to physically count out the exact number we need by hand (click)
We can measure mass and volume as a way to describe how much matter we have, but that does not tell us how many moles.(click)
We could use the same strategy as we did with the popcorn kernels.
© 2007 Microsoft Corporation. All rights reserved. Microsoft, Windows, Windows Vista and other product names are or may be registered trademarks and/or trademarks in the U.S. and/or other countries.
The information herein is for informational purposes only and represents the current view of Microsoft Corporation as of the date of this presentation. Because Microsoft must respond to changing market conditions, it should not be interpreted to be a commitment on the part of Microsoft, and Microsoft cannot guarantee the accuracy of any information provided after the date of this presentation. MICROSOFT MAKES NO WARRANTIES, EXPRESS, IMPLIED OR STATUTORY, AS TO THE INFORMATION IN THIS PRESENTATION.

63. Molar Mass: mass of one mole or 6.02 x 10 23 particles.
Equivalent amounts of Silicon
28.1 g Si = 1 mole Si = x atoms Si
The relationship between # moles and mass can be determined using the molar mass of that substance.
In this case the matter we are dealing with is Silicon. You can determine the mass of 1 mole of silicon by going to the periodic table and finding the molar mass. We usually round this number off to the tenths place when doing calculations. According to the periodic table the mass of 1 mole of silicon is 28.1 g. This also means that 28.1 g of silicon would contain 6.02 x atoms of silicon.
28.1 g of silicon equals 1 mole of silicon equals 6.02x1023 atoms of silicon are equivalent quantities. In other words, they all describe the exact same amount of matter. Notice that since silicon is an element the particles present are called atoms. To help us remember these relationships we can use a small graphic for moles, particles. and molar mass.
6.02 x 1023 particles
Molar Mass

64. Compare these quantities of Silicon (<, > or +)
Equivalent amounts of Silicon
28.1 g Si = 1 mole Si = x atoms Si
80 g Si
2 mole Si
0.015 mole Si
3.200 x 1023 atoms Si
56.2 g Si
1.2 x 1024 atoms Si
7.8 x 1013 atoms Si
1.11 x 1023 atoms Si
Correct answers, greater, less, equal, less
Then ask students to determine the correct number of significant figures in each quantity.

65. More Practice….. 10 g Na 10 mol Na 55.8 g Fe 6.022 x 1023 atoms Fe
8.00 g He mol He
2.00 mol B g B
14.0 g Al x 1023 atoms Al
2.000 mol Li x 1023 atoms Li
4.00 mol Ca g Ca
66 g Ag mol Ag
2.00 mol Cu x 1024 atoms Cu
32.0 g P x 1024 atoms P
10 g Na < 10 mol Na
55.8 g Fe = x 1023 atoms Fe
8.00 g He > 1.00 mol He
2.00 mol B < 32 g B
14.0 g Al < x 1023 atoms Al
2.000 mol Li > x 1023 atoms Li
4.00 mol Ca > g Ca
66 g Ag < mol Ag
2.00 mol Cu = x 1024 atoms Cu
32.0 g P < x 1024 atoms P

66. mass-mole-particle problems
Use dimensional analysis.
Conversion factors:
1 mole = 6.02 x 1023 particles (atoms or molecules)
1 mole = ___________ grams (find the molar mass on the periodic table)

67. Mass-mole-particle problems
If you have a 5.93 gram sample of calcium, how many moles of calcium are in your sample?
Jerry has a 7.58 mole sample of iron (Fe), how many atoms of iron are in Jerry’s sample?