1. CHEMISTRY
Mrs. Adams
Room 601

2. Today’s TO DO: 9/7/10 Introduction Agenda Book Syllabus & Expectations
Notebooks
Hand out
Set up
Journal #1
Homework 

3. Notebook
My gift to you…I expect you to BRING it and USE it E-V-E-R-Y DAY!
Set Up
Leave 1st 4 sheets blank for Table of Contents
Date Title Format pg #
What is Chemistry? Journal 1
if same as above, leave blank Notes
Starting with 5th sheet as pg.1, number each page in lower outer corner

4. Journal #1 – What is Chemistry?
Go to pg.1 of your notebook
Write title Journal #1 – What is Chemistry? on top w/date on right,
Write the following prompt:
What do you think of when you hear the word chemistry?
List 10 things that you think are part of chemistry.
Leave a blank line
Answer

5. Homework: Real Life Application of Chemistry
Find a current article showing an application of chemistry to our lives
Write a summary in your own words that includes
main ideas of the article
your thoughts about the article
1-2 questions you still have that were NOT answered in the article
If done on computer, me the article as well as your summary (GREEN )
If hand written, also provide a printout of the article
DUE: Friday, Sept. 10th

6. Journal #2 – Self Reflection
How do you think your friends and family would describe you?
What are your goals & expectations in this class?
What do you hope to accomplish in the next 5 years?
What do you hope to accomplish in the next 15 years?

7. Lab Safety Read the Flinn Lab Safety Directions
To reinforce the concepts in your assigned section, your table will create one of the following to present to the class:
Skit
Poem
Cartoon
Song
Poster or Ad
Your presentation should try to convince us that your rules are THE MOST IMPORTANT rules to follow in this class.

8. Lab To do a lab really well, have your report done well in advance.
In Other Words…
Know WHAT you’re supposed to do BEFORE you do it.
Know WHY you’re doing something BEFORE you do it.

9. Lab BEFORE every lab, you will address: Background information
Materials & procedures
Safety concerns
Hypothesis
Creation of data tables

10. Lab AFTER every lab you will Organize your data into charts and graphs
Analyze your data
Accept or reject your hypothesis
Discuss your conclusion & possible source(s) of error

11. Scientific Method
A way to answer questions about the world based on observations and experiments.
INQUIRE
Ask questions
OBSERVE
Not always visual
EXPERIMENT
Changing a variable to determine
UNDERSTAND
Where does your Hypothesis fit?
Where does your Conclusion fit?

12. Scientific Method in Action
CHEMOTHERAPY
OIL SPILLS
FOOD INDUSTRY
COSMETIC INDUSTRY
For each:
Explain the general idea of why scientific method applies
Write a SPECIFIC
observation or question AND
hypothesis you could test

13. Measurements Always uncertain Instruments never flawless
Some estimation always required
Example: A ruler

14. Numbered lines = centimeters Smaller lines = 0.1 cm = 1 mm
Reliable
Smaller lines = 0.1 cm = 1 mm
Any point between each line
Must be estimated
Not reliable
Ex: Line above the ruler = _____ cm
Written as _______________ cm
2.45
2.45  +/- 0.01

15. Reliability 2 ways to check numbers Precision Accuracy
repeat measurement
test against standard
Precision
how close repeated measurements are to each other
Accuracy
how close measurements are to standard or accepted value

16. Precision vs. Accuracy

17. Sample Problem
In 3 separate trials, Sara calculates the density of water to be
0.88g/ml
0.87g/ml
Is she precise ?
Yes, all close together
Is she accurate?
No, accepted value is 1 g/ml
Density = mass/volume

18. Significant Digits Sig Figs
Number of digits within a value that are considered significant with respect measurement validity
Follows Pacific-Atlantic Rule

19. Decimal PRESENT - PACIFIC
Start at far left of number (like pacific ocean is on far left of US) 
Start counting first Non Zero number
End at rightmost digit (including zeros)

20. Examples 34.067g 0.0007458ml 0.009070g ____ sig figs ____ sig figs

21. Decimal ABSENT – ATLANTIC
far right of number (like the Atlantic Ocean is at far right of US)
Start counting first Non Zero number
End at leftmost digit

22. Examples
2030cm
____ sig figs
2007dm
19,000,000,000g 3 4 2

23. Practice Problems 0.0026701m 19.0550kg 3500V 1,809,000L ____ sig figs4

24. Sig Fig's in Calculations
Exact numbers or conversions do not count as sig figs
Ex: Speed of light ~ 300,000,000 m/s
Can have infinite # of sig figs and must be specified
Sample –
Speed of light expressed to 3 significant digits = 3.00 x 108

25. Multiplication & Division
Answer must have same # of sig figs as lowest sig figs found within problem
Ex: x = = 58.0
(3) (5) (3)

26. Example: Volume = length x width x height
Find the volume an object with sides m x 1.34m x 13.22m
(5 sig figs) (3) (4)
m (10)
m3 (5)
192.7 m (4)
193 m (3)

27. Addition & Subtraction
Largest uncertainty determines number of sig figs
Answer will have lowest sig figs to the right of the decimal from numbers in problem
Ex: =

28. More Examples 34.50g + 3.2345g + 671.1g + 25.345g 2092 ml – 147.54 ml

29. Practice Problems 6.15m x 4.026m = 12.7km / 3.0 =
150ml ml + 209ml ml =
(35.6L + 2.4L) / 4.803  =
2.542m x (16.408m m) =
24.8 m2
4.2 km
~ 440 ml
7.91 L
31.85 m

30. Journal # 3 – Sig Figs What are the rules for significant figures?
Be sure to include those for addition/subtraction and multiplication/division

31. Scientific Notation One digit to the left of the decimal
# of digits to right of decimal is determined by sig fig rules
Example 19,000,000 ml
2 sig figs
1.9 x 107 ml
Example g
4 sig figs
4.569 x 10-4 g 

32. Scientific Notation Practice
3.27 x 104  
32,700 =
1,024,000 = = =
1.024 x 106   
x 10-3   
3.901 x 10-9   

33. Knowing Equations: Density
Density = mass /volume
D=m/v
Know formula & manipulate w/algebra OR
know graphic below M D V

34. Dimensional Analysis Step by step conversion between units
Convert 10.0µm to inches
Conversion factors 
1m=1,000,000µm  
1m = 39.37inches
Start with the given unit, then use you conversion factors to cancel units until to arrive at the unit you want to convert to.
10um x  1m   x  39.37inches = in
      1,000,000um   1m   

35. Practice Problems 250.0 cm to inches ? gal in 39L ? cm in 16in
? seconds in 5 days
? ft in 86cm
? cm3 in 2.3gal
? m in 3.5mi

36. Percent Error % Error = measured – accepted x 100 accepted Ex:
Accepted value for density of water = 1 g/ml
Measured value for density of water in lab = 0.9 g/ml
% Error =
(0.9 – 1)/1 * 100 = 10% Error

37. Journal #4 – Dimensional Analysis
Are there REALLY seconds in a day?
Show your dimensional analysis to defend your answer.

38. Energy & Matter i
Chapter 2

39. Energy Potential to do work or produce heat 3 Main Types – Radiant
Ex: sunlight
Kinetic
Energy of motion
Ex: Mechanical – energy of moving parts
Ex: Thermal – energy from internal particle motion in matter
Potential
Ex: Gravitational – falling water
Ex: Electrical – opposite charges
Ex: Chemical - battery

40. Energy Units calories Joules Measuring calories
Amount of heat needed to raise 1g of water 1oC
1 Calorie in food = 1000 calories
Joules
SI unit of energy
1 cal = J
Measuring calories
Calorimeter

41. Law of Conservation
Energy is neither created nor destroyed in any process
Energy can be transformed from one form to another
Ex: kinetic energy of bat transferred to baseball (kinetic, sound)
Ex: Chemical energy of striking match transformed into heat and light

42. Temperature Celsius Kelvin 0oC = freezing pt 100oC = boiling pt
21oC = room temp
37oC = body temp
Kelvin
SI Unit of temperature
oC + 273

43. Matter can be neither created nor destroyed
Matter & Conservation
Has mass & volume
States
Solid
Liquid
Gas
Changes
Physical
Chemical
Just like Energy,
Matter can be neither created nor destroyed

44. Elements Meet the Elements
Substances that cannot be separated into simpler substances by chemical change
Organized in Periodic Table
Combine chemically to form COMPOUNDS
Meet the Elements

45. Mixtures Blend of 2 or more pure substances (elements or compounds)
Heterogeneous
Visible differences in combined substances
Ex: chocolate chip cookies
Homogeneous
No visible differences in substances
Ex: salt water

46. Separation of Homogeneous Mixtures
Distillation
Tap water = homogeneous mixture
Crystallization
Evaporation of liquids from solids
Chromatography
Separates mixtures by
Solubility
size
charge

47. The Atom
Chapter 3

48. Early Models of the Atom

49. Atoms The Greek Philosopher Democritus
Proposed all matter made up of small, indivisible particles
Called these “atomos” = atoms
Today’s Definition- smallest particles of an element that retain properties of element

50. Democritus Ideas REJECTED
Because he didn’t know what held these particles together
They remained rejected until the 17th century when better technology = closer observations

51. 1700’s Lavoisier Joseph Louis Proust Law of Conservation of matter
Law of constant composition
compounds always contains same elements in same proportions by mass

52. John Dalton Atomic Theory of Matter:
Each element is composed of extremely small particles called atoms
All atoms of an element are identical, but differ from those of other elements
Atoms are neither created or destroyed in a chemical reaction
A compound always has the same relative numbers of atoms.

53. Discovering Atomic Structure

54. Michael Faraday structure of atoms is related to electricity
atoms contain particles that have electrical charge.

55. Static Electricity Benjamin Franklin famous electricity experiment
Conclusions from his kite & key experiment
lightning is a static discharge from clouds
electricity has two kinds of charges.
Positive (+)
Negative(-)

56. Cathode Rays & Electrons
Running electricity through a partially evacuated glass tube occurs in a cathode tube.
Negative end = cathode
Positive end = anode
With a fluorescent lining & addition of
electricity, particles are visible

57. Where's the rest of the mass??
J.J. Thompson
Deflected particles w/magnet
Discovered particles were negatively charged
Named them electrons
Along w/Milikin discovered their mass to be only 1/2000 of full atomic mass …
Where's the rest of the mass??

58. Ernest Rutherford Discovers alpha particles Discovers beta particles
Deflects towards negative plate
Charge =
Discovers beta particles
Deflect towards positive plate
Discovers gamma particles
Not affected by electric plates

59. The Nuclear Atom If electrons are negative, why are atoms neutral?
Must contain positive parts equal to the negative parts.
Where are they?

60. The Gold Foil Experiment (figure 3-14)

61. What does this mean?
The experiment determined that most of the atoms positive charge, as well as the mass, is in the middle, called the nucleus.
Most of the particles pass through the empty space but occasionally one gets close enough to the positive nucleus to deflect

62. Modern Atomic Theory

63. Atoms Smallest units of matter
Composed of protons (+) and neutrons in nucleus and electrons (-) in orbitals

64. 1 proton has the mass of about 2000 electrons

65. Periodic Table Information
Atomic Number
Represents # of protons
Also # of electrons in a stable atom of an element
Discovered by Moseley
Atomic Mass
Sum of protons & neutrons
Electron mass is small and
almost negligible

66. Electrons Electrons move in space around the nucleus
Rutherford visualized it as a mini solar system.

67. Atomic Mass Measured by Atomic Mass Units (AMU)
atomic mass approximately = protons + neutrons
Atomic Mass (AM) = average mass of element’s atoms, including isotopes

68. Unstable Atoms Ions Isotopes Different # of electrons
Atoms with a charge
Isotopes
Different # of neutrons
Often radioactive
Used as diagnostic tracers

69. Ions When an atom gains or loses electrons it acquires a charge
Fewer electrons means positive charge
More electrons means negative charge
Charge of ion = # protons - # electrons

70. Sample
Write the chemical symbol for the ion with 9 protons and 10 electrons
Answer F-
What is the symbol of the ion with 13 protons and 10 electrons?
Answer Al3+
7 Protons and 10 electrons?
N 3-

71. Isotopes Dalton said all atoms of an element are the same.
Not quite true, ISOTOPES have a different number of neutrons
In nature, elements are almost always found as a mixture of isotopes

72. Identifying Isotopes To identify isotopes more specifically
Use the Mass Number
Mass Number = (# protons) + (# neutrons)

73. Fundamental Subatomic Particles
Location
Charge (C)
Mass (g)
Mass (AMU)
Proton
Inside nucleus
x 10-19
1.673 x 10-24 1
Neutron
1.675 x 10-24
Electron
Outside nucleus
x 10-19
9.109 x 10 –28

74. Changes in the Nucleus

75. Radioactive Elements Discovered by Becquerel in late 1800s
Uranium
Pierre & Marie Curie
Radium & Polonium

76. Nuclear Reactions Changes in nucleus Unstable nucleus  radioactivity
Changes composition of nucleus
Alpha & beta radiation comes from nucleus
Unstable nucleus  radioactivity
Not many elements radioactive
Why not?? Seems like all those + protons would cause a lot of repelling…..

77. Composition of Stable Nuclei
As the number of protons increases, it takes more and more neutrons to remain stable.
All atoms above 83 are unstable

78. Radioactive Decay RULE:
sums of mass numbers & atomic numbers are same before & after reaction

79. Types of Radioactive Decay
Alpha Radiation
Stream of high energy alpha particles
Consists of 2 protons and 2 neutrons
Identical to a helium-4 nucleus
Symbol 42He2+ or 42He or 42
Do not cause a health risk
Do not travel far

80. When an atom emits one of these, it is said to be undergoing radioactive decay
Which brings us to the nuclear equation, or a way to keep track of the components

81. Alpha Decay

83. stream of high speed electrons
Beta Radiation
stream of high speed electrons
neutron changes into 1 proton & electron
proton stays in nucleus
electron is propelled out at high speed
Symbol of 0-1e- or 0-1e or 0-1
damaging to skin

84. Beta Decay

86. Extremely energetic form of light energy we cannot see
Gamma radiation
Extremely energetic form of light energy we cannot see
Symbolized by 00
Does not consist of particles
Able to penetrate deeply into substances

88. Types of Radiation Name Identity Charge Alpha () Helium-4 nuclei 2+
Penetrating Ability
Alpha ()
Helium-4 nuclei 2+
Low
Beta ()
Electron 1-
Medium
Gamma ()
High energy particle
None
high

89. Practice Problems
Write the nuclear equation for the alpha decay of uranium 238.
Write a nuclear equation for the beta decay of sodium 24

90. Other Nuclear Reactions
Nuclear Fusion
Atoms collide and join together releasing great amounts of energy
Like in the Sun
Nuclear Fission
Splitting the nuclei of large atoms
Like in Nuclear reactors

91. Bonding
Atoms bond in order to fill their valence shell (outer energy level)
Octet Rule
The idea that most atoms want 8 electrons in their outer shell and will share, steal, or give away electrons in order to fill the valence shell
Exceptions are those who have less than 6 total electrons

92. Covalent Bonding Valence electrons (outer shell) are shared
Form molecules
Single, double or triple bonds are possible

93. Ionic Bonding
Electrons are lost or gained from the outer shell in order to fulfill the octet rule

94. Hydrogen Bonding
Weak bonds formed between molecules that contain polar covalent bonds
Bonding animation: