1. PRINT OFF
TALKS

2. This image looks 3-D and appears to waver across the page.

3. Burning Magnesium Metal in an Open Container
strip of
magnesium
metal
white
powder
white
powder
An Investigation of the Process of … “Burning Magnesium Metal in an Open Container”
Description
This slide shows an experiment similar to the work done by Antoine Lavoisier in 1778, which demonstrated that mass is conserved in chemical reactions.
Basic Concepts
Mass is neither created nor destroyed when substances react with one another (the law of conservation of mass)
Gases, such as air, are matter and thus have mass and volume.
Teaching Suggestion
Use this slide to teach the law of conservation of mass. Explain that, at the time Lavoisier was investigating the process of burning, scientists thought that a change in mass could occur during a chemical reaction. They observed that when some materials were burned, the substance produced had a larger mass that the starting material. Lavoisier, however, postulated that the extra mass came from the air. When he excluded outside air from the reaction, the total mass did not change.
Questions
In diagram (A), how does the product of the reaction differ from the original magnesium strip? How are they similar?
List three possible explanations that could account for the fact that the white powder has more mass than the magnesium strip in diagram (A). Then, for each of your explanations, decide whether the experiment shown in diagram (B) supports or refutes the explanation. Explain your reasoning.
Why is there no change in total mass when magnesium is burned in a closed flask?
In diagram (B), if you were to determine the mass of the white powder alone, how would it compare to the mass of the magnesium strip? How would the mass of the white powder in diagram (B) compare with the mass of white powder in (A)? Explain your answers.
What would you expect to happen if the stopper were removed from the closed flask containing the white powder in diagram (B)? Explain your answer.
Suppose that a person dropped 8.4 grams of baking soda into a beaker containing 150 g of vinegar. Bubbling occurred, releasing a gas. When the bubbling stopped, the mass of the material in the beaker was 154 g. What was the mass of the gas produced in this reaction?
A friend tells you that mass is not conserved when a candle burns, because the candle gets smaller and smaller. Explain to your friend what must really happen when a candle burns.
strip of
magnesium
metal
Before burning
After burning

4. Burning Magnesium Metal in a Closed Container
nichrome
wire connected
to a battery
strip of
magnesium
metal
white
powder
Before burning
After burning
mass conserved
2 Mg + O2  2 MgO

5. Lighting a Bunsen Burner
inner,
less
transparent,
brighter,
greenish-blue
cone
gas valve
outer,
transparent,
dim blue
cone
air adjustment
DEMONSTRATION
Show the correct way to light a burner. Most important thing to recall is to light the match before you turn on the gas.
Remind students to hold the match at an angle with their hand below the top of the tube – not above the tube. You do not adjust a burner by turning down the amount of gas from the inlet valve. The gas should be “ON” full-blast at all times.
Use a long, clear glass rod and rubber tubing to make a ‘flame thrower’. Students will observe that the glass tube does not get hot (as it only carries the gas through it). They should also observe the effect of pure gas vs. adding more air to the mixture on the color of the flame. They can then practice adjusting the air input into the burner to get a proper flame.
Discuss that an orange flame lacks oxygen and will not support complete combustion of methane. I run my hand through the orange flame several times and then show the carbon soot that is deposited on my hand.
Introduce a chemical equation and chemical symbols
methane + oxygen  carbon dioxide + water energy
CH O  CO H2O + light & heat
gas
adjustment
Always light match BEFORE turning on gas.

6. Flame Temperature Distribution
Meker
Burner Flame
1640oC
1660oC
1670oC
1720oC
1680oC
1775oC
1540oC
1550oC
1560oC
1470oC
520oC
1450oC
350oC
300oC
Bunsen / Tirrill
Burner Flame
“The German chemist, Robert Bunsen ( ), invented the Bunsen burner in 1855 because he needed a colorless flame. He wanted to use it for flame tests for analyzing salts present in mineral waters. His burner, which mixes air with gas before it burns, has been widely used in chemistry laboratories ever since.”
Eyewitness Science “Chemistry” , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 31
The Bunsen / Tirrill burner and Meker burner differ not only in the higher maximum
temperature produced, but also in heat distribution within the flame.

7. It would be easy to see what is most toxic if it was labeled this clearly!

8. Scientific Method An Overview Question Research Hypothesis Procedure/
First
What does the scientist want
to learn more about?
Then
Research
Gathering of information
Scientific Method
An Overview
Next
Hypothesis
An “Educated” guess of an
answer to the question
Then
Procedure/
Method
Written and carefully
followed step-by-step
experiment designed to test
the hypothesis
Next
Data
Information collected during
the experiment
And
And
Observations
Written description of what
was noticed during the
experiment
Finally
Conclusion
Was the hypothesis correct
or incorrect?

9. Units for Measuring Mass
1 kg
(1000 g)
1 lb
1 lb
0.20 lb
Christopherson Scales
Made in Normal, Illinois USA
1 kg = lb

10. Quantities of Mass Giga- Mega- Kilo- base milli- micro- nano- pico-
Earth’s atmosphere
to 2500 km
1018 g
Quantities of Mass
1015 g
1012 g
Ocean liner
Giga-
Mega-
Kilo-
base
milli-
micro-
nano-
pico-
femto-
atomo-
109 g
Indian elephant
106 g
Average human
103 g
1.0 liter of water
100 g
10-3 g
10-6 g
Grain of table salt
10-9 g
10-12 g
10-15 g
10-18 g
Typical protein
10-21 g
Uranium atom
10-24 g
Water molecule

11. Dissolving of Salt in Water
Na+
ions
Water molecules
Cl-
ions
When sodium chloride crystals are dissolved in water, the polar water molecules exert attracting forces which weaken the ionic bonds. The process of solution occurs the ions of sodium and chloride become hydrated.
NaCl(s) + H2O  Na+(aq) + Cl-(aq)

12. Elements, Compounds, and Mixtures
oxygen atoms
hydrogen
atoms
hydrogen
atoms
“Elements, Compounds, and Mixtures”
Description: This slide shows the molecular composition of an element, a compound, and two mixtures.
Basic Concepts
All samples of a substance have the same molecular composition and intensive properties and are homogeneous.
Elements and compounds are substances; mixtures are not.
The elements making up a compound combine in fixed ratios.
Mixtures can be separated by physical methods.
Mixtures that have a uniform composition throughout are homogeneous; those that have parts with different compositions are heterogeneous.
Teaching Suggestions
Use this transparency to help students visualize the molecular composition of elements, compounds, and mixtures and to review the definitions of these terms. Make sure students understand the difference between the terms matter and substance. Remind students that elements and compounds are always homogeneous, while mixtures can be either homogeneous or heterogeneous.
Questions:
Which of the bottles pictured above contain(s) matter? Which contain(s) a single substance? Explain your answers.
How many elements are present in each molecule of water shown in bottle (b)? What is the relative number of atoms of each element in a water molecule?
As you know, ice is frozen water. In other words, ice and water are the same substance, in different phases. What would you expect the ratio of hydrogen atoms to oxygen atoms to be in a molecule of ice? Explain your reasoning.
Bottle (c) and bottle (d) both contain mixtures. How are these mixtures similar? How are they different?
Suppose you find an unlabeled bottle containing a clear liquid. Can you tell by looking at it whether the material is a compound or a mixture? Explain your answer.
How can you prove that a sample of sea water is a mixture?
Classify the following items as elements, compounds or mixtures; rice pudding, copper, carbon dioxide, air, milk, magnesium chloride, granite, mercury, and maple syrup.
A chocolate-chip cookie with more chips in one part of the cookie than another can be used to demonstrate a heterogeneous mixture. Name two other materials that can be classified as heterogeneous mixtures. Explain your reasoning.
(a)
an element
(hydrogen)
(b)
a compound
(water)
(c)
a mixture
(hydrogen
and oxygen)
(d)
a mixture
(hydrogen
and oxygen)

13. Paper Chromatography of Water-Soluble Dyes
orange red yellow
Filter paper
(stationary phase)
Suggested
red dye
is not
homogeneous
Orange mixture of
red
and
yellow
Initial spots of dyes
Direction of Water
(mobile phase)
movement

14. Conservation of Mass + + High voltage electrodes Before reaction glass
chamber
High
voltage
After reaction
0 g H2
40 g O2
300 g (mass
of chamber) +
385 g total O2
H2O H2
5.0 g H2 O2
“Conservation of Mass” (Lavoisier)
Description: This slide illustrates a reaction between hydrogen and oxygen in a nonstoichiometric mixture of these gases.
Basic Concepts
·         Mass and atoms are conserved in chemical reactions.
·         When non-stoichiometric quantities of substances are mixed, they react in stoichiometric proportions. Any reactants in excess remain unreacted.
Teaching Suggestions
Explain that the first diagram shows the amount of oxygen and hydrogen in a closed chamber. A spark passes between the electrodes, causing the O2 and H2 to react rapidly. The second diagram shows what is in the chamber after the reaction.
Use this slide to illustrate that reactants combine in the stoichiometric proportions. Stress that is is not sufficient to know the amounts of starting materials present. One must also know the amounts of reactants that will take part in the reaction.
Questions
What is the ratio of the mass of O2 to H2 before the reaction?
What is the ratio of the number of moles of O2 to H2 before the reaction?
How do you account for the fact that the mass of the chamber and its contents is the same before and after the reaction.
Why is some oxygen left in the chamber after the reaction?
What are the masses of H2 and O2 that take part in the reaction?
What is the ratio of the mass of O2 to H2 taking part in this reaction? What is the ratio of the number of moles of O2 to H2 taking part in the reaction? Why is this mole ratio different from the mass ratio?
If there were twice as much H2 in the chamber (10 g) but the same amount of O2 (80g), what would you expect to find in the chamber after the reaction? Explain your answer.
80 g O2
45 g H2O
? g H2O
300 g (mass
of chamber) +
385 g total

15. Crookes Tube Crookes tube Cathode (-) Anode (+) (Cathode ray tube)
William Crookes
Glow
Cathode
(-)
Anode
(+)
Mask holder
Mask holder

16. The Effect of an Obstruction on Cathode Rays
shadow
High
voltage
source of
high voltage
cathode
yellow-green
fluorescence

17. - Crooke’s Tube + voltage source magnet vacuum tube metal disks
William Crookes - +
Sir William Crookes ( ) was the British scientist who invented the cathode ray tube. His work paved the way to the discovery of the electron.
The Crookes tube can be thought of as the forerunner of the modern fluorescent light tube - a partially evacuated tube fitted with electrodes and filled with mercury vapor and a little argon gas.
magnet
vacuum tube
metal disks

18. The Effect of an Electric Field on Cathode Rays
High
voltage
cathode
source of
high voltage
positive
plate
negative
anode _ +
Charged particles tend to move away from particles with the same charge and toward particles with the opposite charge. When the cathode rays bent away from the negative pole of the magnet and toward the positive pole, this rule caused Thomson to realize the cathode rays were negatively charged.

19. Television Picture Tube
Blue beam
Green beam
Red beam
Glass window
Shadow mask
Fluorescent
screen with
phosphor dots
Red beam
Green beam
Blue beam
Shadow mask
Fluorescent screen
Electron
gun
Electron
beam
Deflecting
electromagnets
The modern television picture tube is a cathode ray tube (CRT) that is based on the same principles used by Thomson.
A changing electric field is used to direct the path of a beam of charged particles across a phosphorescent surface.

20. Geiger Counter (-) (+) e- + e- + e- + + e- Ionization of fill gas
takes place along
track of radiation
(-)
Speaker gives
“click” for
each particle
(+)
Metal tube
(negatively
charged) e- + e-
Window + e- + + e-
Radiation cannot be seen, heard, felt, or smelled. Thus warning signs and radiation detection instruments must be used to alert people to the presence of radiation and to monitor its level. The Geiger counter is one such instrument that is widely used.
Other devices used to detect and measure ionizing radiation: scintillation counter, film badge
Free e- are attracted to (+) electrode, completing the circuit and generating a current.
The Geiger counter then translates the current reading into a measure of radioactivity.
Ionizing
radiation
path
Free e- are attracted to
(+) electrode, completing
the circuit and generating a current. The Geiger counter then translates the current reading into a measure of radioactivity.
Atoms or molecules
of fill gas
Central wire electrode
(positively charged)

21. Models of the Atom - Dalton’s model (1803) Greek model (400 B.C.)+ + -
Dalton’s model
(1803)
Greek model
(400 B.C.)
Thomson’s plum-pudding
model (1897)
Rutherford’s model
(1909)
Bohr’s model
(1913)
Charge-cloud model
(present)
1803 John Dalton
pictures atoms as
tiny, indestructible
particles, with no
internal structure.
1897 J.J. Thomson, a British
scientist, discovers the electron,
leading to his "plum-pudding"
model. He pictures electrons
embedded in a sphere of
positive electric charge.
1911 New Zealander
Ernest Rutherford states
that an atom has a dense,
positively charged nucleus.
Electrons move randomly in
the space around the nucleus.
1926 Erwin Schrödinger
develops mathematical
equations to describe the
motion of electrons in
atoms. His work leads to
the electron cloud model.
1913 In Niels Bohr's
model, the electrons move
in spherical orbits at fixed
distances from the nucleus.
“Models of the Atom”
Description: This slide shows he evolution of the concept of the atom from John Dalton to the present.
Basic Concepts
·         The model of the atom changed over time as more and more evidence about its structure became available.
·         A scientific model differs from a replica (physical model) because it represents a phenomenon that cannot be observed directly.
Teaching Suggestions
Use this slide as a review of the experiments that led up to the present-day view of the atom. Ask students to describe the characteristics of each atomic model and the discoveries that led to its modification. Make sure that students understand that the present-day model shows the most probable location of an electron at a single instant.
Point out that most scientific models and theories go through an evolution similar to that of the atomic model. Modifications often must be made to account for new observations. Discuss why scientific models, such as the atomic models shown here, are useful in helping scientists interpret heir observations.
Questions
Describe the discovery that led scientists to question John Dalton’s model of the atom ad to favor J.J. Thomson’s model.
What experimental findings are the basis for the 1909 model of the atom?
What shortcomings in the atomic model of Ernest Rutherford led to the development of Niels Bohr’s model?
A friend tells you that an electron travels around an atom’s nucleus in much the same way that a planet revolves around the sun. Is this a good model for the present-day view of the atom? Why or why not?
Another friend tells you that the present-day view of an electron’s location in the atom can be likened to a well-used archery target. The target has many holes close to the bull’s-eye and fewer holes farther from the center. The probability that the next arrow will land at a certain distance from the center corresponds to the number of holes at that distance. Is this a good model for the present-day view of the atom? Why or why not?
Suppose that, it the future, an apparatus were developed that could track and record the path of an electron in an atom without disturbing its movement. How might this affect the present-day model of the atom? Explain your answer.
How does developing a model of an atom differ from making a model of an airplane? How are these two kinds of models the same?
Drawing on what you know in various fields of science, write a general statement about the usefulness of scientific models.
Timeline: Wysession, Frank, Yancopoulos Physical Science Concepts in Action, Prentice Hall/Pearson, 2004 pg 114
1924 Frenchman Louis
de Broglie proposes that
moving particles like electrons
have some properties of waves.
Within a few years evidence is
collected to support his idea.
1932 James
Chadwick, a British
physicist, confirms the
existence of neutrons,
which have no charge.
Atomic nuclei contain
neutrons and positively
charged protons.
1904 Hantaro Nagaoka, a
Japanese physicist, suggests
that an atom has a central
nucleus. Electrons move in
orbits like the rings around Saturn.

22. + + + + + + From the time of Dalton to Schrödinger, our model
Thomson (1904)
(positive and negative charges) + +
Rutherford (1911)
(the nucleus) + + + + .
Bohr (1913)
(energy levels - orbits) .
Schrödinger (1926)
(electron cloud model – orbitals)
From the time of Dalton to Schrödinger, our model
of the atom has undergone many modifications.

23. 6Li 7Li 3 p+ 3 n0 3 p+ 4 n0 2e– 1e– 2e– 1e– + + Lithium-6 Lithium-7
Nucleus
Neutron
Proton
Nucleus
Neutron
Proton
Electrons +
Electrons +
Nucleus
Nucleus
Lithium-6
Lithium-7
Neutrons 3
Protons 3
Electrons 3
Neutrons 4
Protons 3
Electrons 3

24. Electron aligned against
Spin Quantum Number, ms
North
South S N -
Electron aligned with
magnetic field,
ms = + ½
Electron aligned against
magnetic field,
ms = - ½
The electron behaves as if it were spinning about an axis through its center.
This electron spin generates a magnetic field, the direction of which depends
on the direction of the spin.

25. Nitrogen N = 1s22s22p3 H He Li C N Al Ar F Fe La Energy Level Diagram
6s p d f
Bohr Model
5s p d
4s p d
Arbitrary Energy Scale
3s p N
Hund’s Rule “maximum
number of unpaired orbitals”.
2s p 1s
Electron Configuration
NUCLEUS
N = 1s22s22p3
H He Li C N Al Ar F Fe La
CLICK ON ELEMENT TO FILL IN CHARTS

26. METALS Metals and Nonmetals Nonmetals Metalloids H He Li Be B C N O F1 He 2 1 Li 3 Be 4 B 5 C 6
Nonmetals N 7 O 8 F 9 Ne 10 2 Na 11 Mg 12 Al 13 Si 14 P 15 S 16 Cl 17 Ar 18 3 K 19 Ca 20 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Ga 31 Ge 32 As 33 Se 34 Br 35 Kr 36 4
METALS Rb 37 Sr 38 Y 39 Zr 40 Nb 41 Mo 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 Cd 48 In 49 Sn 50 Sb 51 Te 52 I 53 Xe 54 5
Metalloids Cs 55 Ba 56 Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Tl 81 Pb 82 Bi 83 Po 84 At 85 Rn 86 6 * Fr 87 Ra 88 Rf
104 Db
105 Sg
106 Bh
107 Hs
108 Mt
109 7 W La 57 Ce 58 Pr 59 Nd 60 Pm 61 Sm 62 Eu 63 Gd 64 Tb 65 Dy 66 Ho 67 Er 68 Tm 69 Yb 70 Lu 71 Ac 89 Th 90 Pa 91 U 92 Np 93 Pu 94 Am 95 Cm 96 Bk 97 Cf 98 Es 99 Fm
100 Md
101 No
102 Lr
103

27. The Electromagnetic Spectrum
HIGH
ENERGY
Decreasing wavelength
LOW
ENERGY
Increasing frequency
Increasing photon energy
AM radio
Short wave
radio
Television channels FM
Radar
Microwave
Radio Waves V i s b l e L g h t
Gamma Rays UV
Rays
“The Electromagnetic Spectrum”
Description: This slide depicts the electromagnetic spectrum from gamma rays through radio waves.
Basic Concepts
·         All forms of electromagnetic radiation are not identical
·         All forms of electromagnetic radiation travel at the same speed in a vacuum (the speed of light, c = 3.00 x 108 m/sec).
·         Wavelength and frequency are inversely proportional for a wave traveling at a constant speed.
·         Energy and frequency are directly proportional for electromagnetic waves traveling at the speed of light.
Teaching Suggestions
Use this transparency to review the relationship of visible light to other types of radiation. Explain that all of the rays and waves shown are types of electromagnetic radiation. Point out that they differ essentially from each other only in energy level, wavelength, and frequency. Try the analogy of an ocean wave to help students understand electromagnetic waves. Question 6 can be used to assess the students understanding of wave velocity, wavelength, and frequency.
Questions:
List the ways in which visible light is different from the other types of radiation shown in the diagram.
List the ways in which all of the types of radiation shown in the diagram are similar.
You are told that sound waves cannot travel in a vacuum. Are sound waves a types of electromagnetic radiation? Explain your logic.
Radio waves can go around an obstruction if the obstruction is smaller than the radio wave’s wavelength. What would you expect to happen if visible light were beamed at a thin wire 2 x 10-5 centimeter thick? Explain your answer.
For electromagnetic waves traveling at the speed of light, the wavelength is inversely proportional to frequency, as expressed by the equation c = fl, where c = speed of light in vacuum (3.00 x 108 meters/second), f = frequency, and l= wavelength. Using this equation, calculate the frequency of a 3-meter radio wave traveling at the speed of light. Compare your answer with the diagram.
Suppose that at a particular beach the ocean waves are traveling at a speed of 2 meters/second. If you know that the distance between waves is 10 meters, can you calculate how often they hit the shore? Explain your answer.
For electromagnetic waves traveling at the speed of light, the energy of a single photon is expressed by the equation E = hf, where E = energy, f = frequency, and h = Planck’s constant, 6.6 x joules/hertz.
Which has more energy, a photon of visible light or a photon of radar, if both traveling at the speed of light?
Do you think you can calculate the energy of an ocean wave using this energy equation? Explain your answer.
infrared
X- Rays
R O Y G B I V
Red Orange Yellow Green Blue Indigo Violet

28. Electromagnetic Spectrum
Visible spectrum
HIGH
ENERGY
Violet
Blue
Green
Yellow
Orange
Red
LOW
ENERGY
400 nm
500 nm
600 nm
700 nm
White
Light
g rays
X-rays
Ultraviolet
Infrared
Microwave
Radio waves
Long
Wave
Water waves transmit energy through space by the periodic oscillation of matter.
• Energy that is transmitted, or radiated, through space in the form of periodic oscillations of an electric and a magnetic field is called electromagnetic radiation.
Electromagnetic radiation
– Consists of two perpendicular waves, one electric and one magnetic, propagates at the speed of light, abbreviated c, and has a value of x 108 m/s
– Is radiant energy that includes radio waves, microwaves, visible light, X -rays, and gamma rays
– Various kinds of electromagnetic radiation all have the same speed (c) but differ in wavelength and frequency
– Frequency of electromagnetic radiation is inversely proportional to the wavelength
c =  or  = c/
– Energy of electromagnetic radiation is directly proportional to its frequency (E  ) and inversely proportional to its wavelength (E  1/)
Radar TV FM
Short
Wave
10-2nm
10-1nm
100nm
101nm
102nm
103nm
10-3cm
10-2cm
10-1cm
100cm
101cm
1cm
101m
102m
103m
104m
Wavelength, l
1019Hz
1018Hz
1017Hz
1016Hz
1015Hz
1014Hz
1013Hz
1012Hz
1011Hz
1010Hz
109Hz
100 MHz
10 MHz
1 MHz
100 KHz
Frequency, n
Electromagnetic spectrum

29. The Photoelectric effect and the frequency of light
IR, infrared
Red
Orange
Yellow
Green
Blue
Violet
UV, ultraviolet
Lens
Quartz prism
A beam of white light is dispersed into its wavelength components by a quartz prism and falls on a metal sample (potassium, in this case).
Light of the highest frequencies (violet and ultraviolet) produces the most energetic photoelectrons (longest arrows).
Light of lower frequencies (for example, orange) results in less energetic photoelectrons (shorter arrows)
Light with a frequency lower than 4.23E14/sec (710 nm) produces no photoelectric effect on all potassium,
regardless of how bright (intense) the light is.
The kinetic energies depend on the frequency (color) of light. Dim blue light produces photoelectrons with higher energies than does bright red light.
Moreover, if the frequency of the light falls below a certain minimum value, called the threshold frequency, we don't observe the photoelectric effect at all.
Slit
Light source
UV v b g y o r IR
Potassium

30. Discovering the Periodic TableNe Ar Kr Xe Po Rn Ra Eu Lu Pa Ac C S Fe Cu Ag Sn Au Hg Pb
Ancient Times Tc Hf Re At Fr Pm Np Pu Am Cm Bk Cf Es Fm Md No Lr He Sc Ga Ge Rb Ru In Cs Tl Pr Nd Sm Gd Dy Ho Tm Yb La Cr Mn Li K N O F Na B Be H Al Si Cl Ca Ti V Co Ni Se Br Sr Y Zr Nb Mo Rh Pd Cd Te I Ba Ta W Os Ir Mg Ce Tb Er Th U P Zn As Sb Pt Bi
Midd Rf Db Sg Bh Hs Mt
1965-
The Noble Gases
At the start of the 1890s, no one had any idea that there was a separate group of gases in the periodic table, the noble gases. Noble gases are familiar to us from their use in neon signs and helium balloons. By 1900 this whole new group had been identified and isolated. While trying to determine an accurate atomic mass for nitrogen, British physicist Lord Raleigh ( ) discovered that nitrogen prepared from ammonia was noticeably lighter than nitrogen that came from the atmosphere. He and William Ramsay ( ) both studied “atmospheric” nitrogen. By removing the nitrogen from it, they produced a tiny quantity of another gas. Since it did not react with anything they called it argon, from the Greek word for lazy. The discovery of helium followed a year later in Ramsay and his assistant Morris Travers ( ) then started to search for additional elements in this new group. They attempted this by fractional distillation of large quantities of liquid air and argon. In 1898, their efforts were rewarded; they had prepared krypton, neon, and xenon.
Eyewitness Science “Chemistry” , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 32

31. Binary Compounds NaCl sodium chlor ine ide (Na1+ Cl1-) CaSHe 2 C 6 N 7 O 8 F 9 Ne 10 B 5 H 1 Al 13 Si 14 P 15 S 16 Cl 17 Ar 18 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30 Ga 31 Ge 32 As 33 Se 34 Br 35 Kr 36 Y 39 Zr 40 Nb 41 Mo 42 Tc 43 Ru 44 Rh 45 Pd 46 Ag 47 Cd 48 In 49 Sn 50 Sb 51 Te 52 I 53 Xe 54 Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Tl 81 Pb 82 Bi 83 Po 84 At 85 Rn 86 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Rf
104 Db
105 Sg
106 Bh
107 Hs
108 Mt
109 Be 4 Ca 20 Sr 38 Ba 56 Ra 88 Mg 12 * 2+ 3+ 1+ 2+
Binary compounds that contain a metal of fixed oxidation number
(group 1, group 2, Al, Zn, Ag, etc.), and a non-metal.
To name these compounds, give the name of metal followed by the
name of the non-metal, with the ending replaced by the suffix –ide.
Examples:
NaCl
sodium chlor
ine
ide
(Na1+ Cl1-)
CaS
calcium sulf ur
ide
(Ca2+ S2-)
AlI3
aluminum iod
ide
ine
(Al I1-) 3

32. Molecular Shapes AB2 Linear AB3 Trigonal planar AB2E Angular or Bent
Tetrahedral
AB3E
Trigonal
pyramidal
AB2E2
Angular
or Bent
AB5
Trigonal bipyramidal
A hyperlink to additional information is found on the central atom of the diagrams.
AB4E
Irregular tetrahedral
(see saw)
AB3E2
T-shaped
AB2E3
Linear
AB6
Octahedral
AB6E
Square pyramidal
AB5E2
Square planar

33. sp2 hybrid orbitals shown together
Ground-state B atom 2s 2p 2s 2p
B atom with one electron “promoted”
sp2
hybrid orbitals
Energy
sp2 2p px py pz s
B atom of BH3 orbital diagram
p orbitals H B
three sps hybrid orbitals
sp2 hybrid orbitals shown together
(large lobes only)
hybridize
s orbital

34. Nitrogen is in excess – or hydrogen is limiting reagent.
Chemical Equations
Chemical Equations
N2 (g) +
3 H2 (g)
2 NH3 (g) +
“Microscopic recipe”
1 molecule N2 +
3 molecules H2
2 molecules NH3
“Macroscopic recipe”
1 mol N2 +
3 mol H2
2 mol NH3
Experimental Conditions Reactants Products
Before reaction
1 mol N2 +
3 mol H2
2 mol NH3
2 molecules N2
3 molecules H2
0 molecules NH3
After reaction
1 molecules N2
0 molecules H2
2 molecules NH3
Nitrogen is in excess – or hydrogen is limiting reagent.

35. TABLE OF SOLUBILITIES IN WATER
aluminum ss s n i d
ammonium
barium
calcium
copper (II)
iron (II)
iron (III)
lead
magnesium
mercury (I)
mercury (II)
potassium
silver
sodium
zinc
acetate
bromide
carbonate
chloride
chromate
hydroxide
iodide
nitrate
phosphate
sulfate
sulfide
Legend
SOLID
AQUEOUS
i = insoluble
ss = slightly soluble
s = soluble
d = decomposes
n = not isolated

36. Limiting Reactants + CB + 4 T CT4 plus 16 tires excess 8 car bodies
8 car bodies
48 tires
8 cars
CB T CT4

37. The Greenhouse Effect HIGH ENERGY, SHORT l LIGHT PASSES
EASILY THROUGH ATMOSPHERE
CO2 MOLECULES
ENERGY RELEASED
AS HEAT
The atmospheric greenhouse effect
– CO2 is a greenhouse gas that absorbs heat before it can be radiated from the earth into space
– CO2 in the atmosphere can result in increased surface temperatures (the greenhouse effect)
– Temperature increases caused by increased CO2 levels are superimposed upon much larger variations in the earth’s temperature
LOWER ENERGY, LONGER l
LIGHT IS BLOCKED BY CO2 AND CH4; ENERGY DOESN’T
ESCAPE INTO SPACE;
ATMOSPHERE HEATS UP

38. Elastic vs. Inelastic Collisions
POW 8 v1 v2
elastic collision v3 v4 8
inelastic collision

39. Molecular Velocities the Maxwell speed distribution
molecules sorted by speed
many different molecular speeds
Fractions of particles
the Maxwell speed distribution
speed

40. Kinetic Theory and the Gas Laws10 10 10
“Kinetic Theory and the Gas Laws”
Description: This slide shows three cylinders filled with equal masses of the same gas under different conditions. In cylinder (b), the pressure and temperature both have doubled from their values in cylinder (a). In cylinder (c), only the temperature has been doubled. The velocities of the gas particles are represented graphically to help students interpret the macroscopic changes in terms of the kinetic theory.
Basic Concepts
The kinetic theory can explain the combined gas law relationships of an ideal gas.
The pressure of a confined gas increases when the particles of the gas collide with the container walls more frequently or with greater force.
Teaching Suggestions
Use this slide to review the kinetic theory as it applies to the temperature, volume, and pressure of gases. Ask students whether the velocities of the gas particles would be the same for all gases at the same temperature. Point out that the average kinetic energy would be the same for all gases, but more massive gas particles would have slower velocities because mass and velocity are inversely proportional in the formula for kinetic energy: ½ mv2.
Questions
The three cylinders in the diagram contain equal masses of the same gas under different conditions. Compare and contrast the conditions of the gas in each of the cylinders with respect to the pressure, volume, and temperature of the gas.
Cylinder (b) illustrates what happens when the gas in cylinder (a) is heated at constant volume. What happens to the average kinetic energy of the gas particles? How does this affect the velocity of the particles?
Use the kinetic theory to explain why the internal gas pressure in cylinder (b) is greater than the internal gas pressure in cylinder (a). What has been done to cylinder (b) to keep the gas volume the same as in cylinder (a)? Explain in terms of the equilibrium between internal and external pressure.
Suppose you are given the information that the temperature in both cylinder (b) and cylinder c) is double the temperature in cylinder (a). Compared with cylinder (a), how much greater is the volume in cylinder (c)? Explain how your answers illustrate the combined gas law (PV/T = constant).
Predict what would happen if you doubled the mass on top of the piston in cylinder (a), with the temperature constant. Use the kinetic theory to describe how this change would affect each of the following: volume, density, temperature, average velocity of the particles, rate of collision of the particles with the container wall, energy of these collisions, and internal pressure.
Suppose you take a basketball outside on a cold winter day. At first the basketball bounces normally, but then you notice that it starts to lose some of its bounce. Use the kinetic theory of gases to explain what is happening. Think of a few explanations. How could you determine which explanation is the best one?
(a)
(b)
(c)
original temperature
original pressure
original volume
increased temperature
increased pressure
original volume
increased temperature
original pressure
increased volume

41. Molar Volume 1 mol of a gas @ STP has a volume of 22.4 L nO = 1 mole
273 K
nO = 1 mole
(32.0 g)
VO = 22.4 L
P = 1 atm 2
273 K
nHe = 1 mole
(4.0 g)
VHe = 22.4 L
P = 1 atm 2
273 K
nN = 1 mole
(28.0 g)
VN = 22.4 L
P = 1 atm 2
MOLAR VOLUME
One mole of any gas occupies 22.4 liters at standard temperature and pressure (STP).
Timberlake, Chemistry 7th Edition, page 268

42. Volume and Number of Moles
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 413

43. 95 mm Hg
105.9 kPa
X atm
1.51 atm
324 mm Hg
X kPa
251.8 kPa
844 mm Hg
X mm Hg
183 mm Hg
X kPa
0.44 atm
218 mm Hg
X atm
72.4 kPa
125mm Hg
85.3 kPa
X mm Hg
X mm Hg
712 mm Hg
145.9 kPa
118.2 kPa
783 mm Hg
X mm Hg
528 mm Hg
X mm Hg
106.0 kPa

44. BIG BIG = small + height 760 mm Hg 112.8 kPa = 846 mm Hg
height = BIG - small
101.3 kPa
X mm Hg =
846 mm Hg -
593 mm Hg
X mm Hg =
253 mm Hg
253 mm Hg
STEP 1) Decide which pressure is BIGGER
STEP 2) Convert ALL numbers to the unit
of unknown
STEP 3) Use formula Big = small + height
small
0.78 atm
height
X mm Hg
760 mm Hg
0.78 atm =
593 mm Hg
1 atm

45. Evaporation
H2O(g)
molecules
(water vapor)
H2O(l)
molecules

46. Evaporation
H2O(g)
molecules
(water vapor)
H2O(l)
molecules

47. How Vapor Pressure is Measured
760 mm mm = 880 mm Hg
atmospheric
pressure
1 atm = 760 mm Hg
dropper
containing
a liquid
atm. pressure
120 mm
vapor from
the liquid
mercury
Animation by Raymond Chang
All rights reserved

48. Nonvolatile Solvent solute molecules

49. Pinhole
Gas
Vacuum

50. Solvents at the hardware store
…………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………………
Glass
Cleaner
Glass
Cleaner
Glass
Cleaner
specs
paint thinner
sunnyside
naptha
mineral
spirits
odorless
t.r.p.s.
turpentine
xylol
lacquer
thinner
acetone
methyl ethyl
ketone
boiled
linseed oil
carbo-sol
brush cleaner
adhesive
remover
specs
paint thinner
sunnyside
naptha
sunnyside
naptha
sunnyside
mineral
spirits
sunnyside
mineral
spirits
sunnyside
odorless
paint thinner
sunnyside
odorless
paint thinner
sunnyside
t.r.p.s.
sunnyside
turpentine
sunnyside
xylol
sunnyside
lacquer
thinner
sunnyside
lacquer
thinner
sunnyside
acetone
sunnyside
methyl ethyl
ketone
sunnyside
boiled
linseed oil
sunnyside
carbo-sol
sunnyside
carbo-sol
sunnyside
brush cleaner
sunnyside
adhesive
remover
sunnyside
made-up
stuff
sunnyside
specs
paint thinner
sunnyside
naptha
sunnyside
odorless
paint thinner
sunnyside
mineral
spirits
sunnyside
xylol
sunnyside
turpentine
sunnyside
t.r.p.s
sunnyside
lacquer
thinner
sunnyside
acetone
sunnyside
denatured
alcohol solvent
sunnyside
methyl ethyl
ketone
sunnyside
boiled
linseed oil
sunnyside
adhesive
remover
sunnyside
brush
cleaner
sunnyside
t.r.p.s
sunnyside
lacquer
thinner
acetone
denatured
alcohol solvent
methyl ethyl
ketone
boiled
linseed oil
adhesive
remover
brush
cleaner
turpentine
xylol
mineral
spirits
naptha
odorless
paint thinner
specs
paint thinner
sunnyside
1 GAL
paint thinner
sunnyside
1 GAL
paint thinner
sunnyside
1 GAL
paint thinner
sunnyside
2GAL
LATEX
sunnyside
1 GAL
kerosene
sunnyside
1 GAL
kerosene
sunnyside
2 GAL
kerosene
sunnyside
2 GAL
Cool
Stuff
Juice
sunnyside
muriatic acid
sunnyside
muriatic acid
sunnyside
sunnyside
muriatic acid
sunnyside
muriatic acid
…………………………………………………………………………………………
Mr. C’s sweat!
Solvents at the hardware store
Solvents at the hardware store.

51. Concentration = 1 “fishar”
# of fish
volume (L)
1 fish
Concentration =
1 (L)
Concentration = 1 “fishar”
V = 1000 mL
V = 1000 mL
V = 5000 mL
n = 2 fish
n = 4 fish
n = 20 fish
Concentration = 2 “fishar”
[ ] = 4 “fishar”
[ ] = 4 “fishar”

52. Concentration = # of moles volume (L) V = 1000 mL V = 1000 mL
n = 8 moles
[ ] = 32 molar
V = 1000 mL
V = 1000 mL
V = 5000 mL
n = 2 moles
n = 4 moles
n = 20 moles
Concentration = 2 molar
[ ] = 4 molar
[ ] = 4 molar

53. Spec-20 Instrument Wavelength Light control Amplifier Sample holder
control knob
Sample holder
cover
Amplifier
control knob
Light control
knob

54. Calcium (Ca2+) and Magnesium (Mg2+) Ions in Untreated Water
Sodium (Na1+)
Ions
on beads
Calcium (Ca2+)
and
Magnesium (Mg2+)
Ions
on beads
The water softening system consists of a mineral tank and a brine tank. The water supply pipe is connected to the mineral tank so that water coming into the house must pass through the tank before it can be used.
The mineral tank holds small beads (also known as resin) that carry a negative electrical charge. The positively charged calcium and magnesium (called ions) are attracted to the negatively charged beads. This attraction makes the minerals stick to the beads as the hard water passes through the mineral tank. (See figure 3.)  Eventually the surfaces of the beads in the mineral tank become coated with the calcium and magnesium minerals. To clean the beads, a strong sodium (salt) solution held in the brine tank is flushed through the mineral tank. Sodium ions also have a positive electrical charge, just not quite as strong as that of calcium and magnesium. This large volume of sodium ions overpowers the calcium and magnesium ions and drives them off of the beads and into the solution. The sodium solution carrying the minerals is then drained out of the unit. Some sodium ions remain in the tank attached to the surfaces of the beads. 
Sodium (Na1+)
Ions in
Treated
Water

55. p p Applied pressure needed to prevent volume increase Osmotic
Pure
solvent
Solution
Net
movement
of solvent
Semipermeable
membrane
Solute
molecules
Solvent
The pressure required to allow for no transport of solvent across the membrane is called the OSMOTIC pressure
and obeys the relation:
                                                                                                                     
Osmotic pressure has a great effect on living CELLS, because their walls are a semipermeable membrane.

56. Titration M1V1 = M2V2 M V = M V M V n = M V n HCl must be ~ __x
muriatic acid
sunnyside
HCl must be ~ __x
more concentrated
than the NaOH. 6
? M of HCl
30.0 mL of 2.0 M of NaOH
10.5 mL
If it requires 10.5 mL of ? M HCl to titrate 30.0 mL of 2.0 M NaOH to its endpoint:
what is the concentration of the HCl?
(x M)(10.5 mL) = (2.0 M)(30.0 mL)
M1V1 = M2V2
X = 5.7 M
10.5 mL of HCl
M V = M V
HCl(aq)  H+(aq) + Cl-(aq) H+ H+
OH-
OH-
0.1 M
0.1 M
0.1 M
M V n = M V n H+ H+
OH-
OH-
H2SO4(aq)  2 H+(aq) + SO42-(aq)
30.0 mL of NaOH with bromthymol blue indicator
0.1 M
“0.2 M”
0.1 M
Endpoint of titration is reached…color change.
proper term is Normality (N)
Al(OH)3(aq)  Al3+(aq) OH-(aq)
0.1 molar H2SO4 is 0.2 normal

57. Commercial vinegar is sold
Acid
Base
Calibration Curve
vinegar ammonia
Vinegar Ammonia
1 mL
3 mL
5 mL
10 mL
15 mL
ammonia
vinegar
Using 3 mL vinegar…
titrate with M NaOH solution.
Calculate % acetic acid in vinegar. % = part/whole x100 A)
moles HC2H3O2 =
moles NaOH
NaOH M
mol L
molNaOH = M x L
mol = (0.130 M)( L)
mol = mol NaOH
therefore, you have ...
mol HC2H3O2
TEACHER NOTE: I mix 26 g of NaOH in 5 L of distilled water to yield a M NaOH solution.
When using household vinegar – it requires ~20 mL of M NaOH / 3.0 mL of vinegar.
You may choose to add 52 g of base to reduce volume of base to reach endpoint.
I use phenolpthalein indicator – its cheap. It is not the proper indicator for use with a weak acid (vinegar) and weak base (ammonia).
A more accurate result would be achieved using the indicator bromthymol blue. B)
x g HC2H3O2 = mol HC2H3O2
0.153 g HC2H3O2 C)
% =
% =
Commercial vinegar is sold
as % acetic acid
% = % acetic acid

58. Photoelectric Effect Light photons Electrons ejected from the surface
cathode
anode
Symbolic representation
of a photoelectric cell
evacuated glass
envelope
Light photons
Photoelectric Cell
Electrons ejected
from the surface
Sodium metal

59. When light strikes a metal surface, electrons are ejected.
Photoelectric Effect
Light
Electron
Nucleus
Metal
When light strikes a metal surface, electrons are ejected.

60. Animation by Raymond Chang
Alpha, Beta, Gamma Rays
Lead block
(+)
(-)
b rays
(negative charge)
Aligning
slot
(no charge)
g rays
a rays
Radioactive
substance
(positive charge)
Three kinds of radiation – α particles, β particles and γ rays
1. Distinguished by the way they are deflected by an electric field and by the degree to which they penetrate matter
2. α particles and β particles are deflected in opposite directions; α particles are deflected to a much lesser extent because of their higher mass-to-charge ratio.
3. γ rays have no charge and are not deflected by electric or magnetic fields.
4. α particles have the least penetrating power, and γ rays are able to penetrate matter readily.
Photographic
plate
Electrically charged
plates
(detecting screen)
Animation by Raymond Chang
All rights reserved

61. Uranium Radioactive Decay
238
4.5 x 109 y
24 d
1.2 m
2.5 x 105 y
8.0 x 104 y
1600 y
3.8 d
3.0 m
27 m
160 ms
5.0 d
138 d
stable
U-238
Th-234 a
234
Pa-234 b
U-234 b
Th-230 a
230
Ra-226 a
226
Rn-222 a
222
Mass number
Po-218 a
218
Pb-214 a
Impossible for any nuclide with Z > 85 to decay to a stable daughter nuclide in a single step, except via nuclear fission
• Radioactive isotopes with Z > 85 usually decay to a daughter nucleus that is radioactive, which in turn decays to a second radioactive daughter nucleus, and so forth, until a stable nucleus finally results
• This series of sequential alpha- and beta-decay reactions is called a radioactive decay series
Some nuclei spontaneously transform into nuclei with a different number of protons, producing a different element
These naturally occurring radioactive isotopes decay by emitting subatomic particles
Should be possible to carry out the reverse reaction, converting a stable nucleus to another more massive nucleus by bombarding it with subatomic particles in a nuclear transmutation reaction
Po-214 b
214
Bi-214 b
Pb-210 a
Bi-210 b
Po-210 b
210
Pb-206 a
206 81 82 83 84 85 86 87 88 89 90 91 92
Atomic number

62. positron emission and/or
140
209 83 Bi
a decay
Nuclear Stability
130
120
184 74 W
110
100 90
Decay will occur in such a way as to return a nucleus to the band (line) of stability.
b decay 80
107 47 Ag
Neutrons (N) 70 60 50 56 26 Fe 40 30
positron emission and/or
electron capture 20 20 10 Ne 10 10 20 30 40 50 60 70 80 90
Protons (Z)