Transcript File

ATOMIC THEORY
ESSENTIAL QUESTIONS
• What are we made of?
• How are scientific models developed?
• Do atoms exist or are they just concepts invented by scientists? What
evidence is there in your everyday life for the existence of atoms?
• How did the understanding of the atom affect historical events?
• How have historical events affected the model of the atom?
ESSENTIAL QUESTIONS
• What do we think the atom “looks like” now?
• If the atom is mostly empty space, why doesn’t my butt fall through the chair?
• How are light and electrons related?
• How do we “see” where electrons are located in the atom?
• Why is the location of electrons so important?
HISTORICAL BACKGROUND
• Greek Philosophers
• Democritus (460-370 BCE)
•
“atomism”
• Aristotle (384-322 BCE)
•
•
•
•
Earth
Air
Fire
Water
• No Experiments
HISTORICAL BACKGROUND
• Alchemy (Up to the Middle Ages)
• Transmutation of other metals into gold
• Phlogiston
•
•
Imaginary element
Believed to separate from combustible bodies when burned
HISTORICAL BACKGROUND
• Early Experimental Chemists
• Henry Cavendish (1731-1810)
(hydrogen – “inflammable air”)
•
•
•
•
Joseph Priestley (1733-1804) (oxygen)
Antoine Lavoisier (1743-1794) (oxygen)
Karl Wilhelm Scheele (1742-1786) (oxygen)
Count Amedeo Avogadro (1776-1856)
(gases  mole)
INVENTIONS
2. Which of the following were invented
a) before 1800?
b) between 1800 and 1900?
c) after 1900?
BEFORE 1800,
BETWEEN 1800 AND 1900,
OR AFTER 1900??
Glass
Mercury thermometer
Barometer
Perfume
Gunpowder
Guns
Hot air balloons
Telegraph
Electric motor
Internal combustion
engine
Car
Battery
Rechargeable battery
Photography
X-ray photography
Bunsen burner
Gas lights
Incandescent light
bulb
Electric lights
Glass blowing
Sewing machine
Cathode ray tube (
TV)
Submarine
Vacuum technology
Asphalt
Tin cans
Candle
Electricity
Conduction of
electricity
Graphite pencil
BEFORE 1800
Glass (3000 BC)
Perfume (Egypt – BC)
Electricity (static
electricity 600 BC)
Candle (~ 200 BC,
China, whale fat)
Glassblowing (50 BC)
Gunpowder (800s,
China)
Handguns (1400s)
Graphite pencil (1564,
England)
Muskets (1600s)
Barometer (1608)
Vacuum technology
(1650, Germany)
Mercury thermometer
(1714)
Conduction of
electricity (1729, Ben
Franklin 1747)
Hot air balloons
(1783)
Gas lights (1792)
Battery (1799, Volta)
BETWEEN 1800 AND 1900
Sewing machine
(~1800)
Tin cans (1810)
Asphalt (1824, Paris)
Photography (1826)
Electric motor (1831)
Incandescent light
bulb (1835, Scotland)
Telegraph (1838)
Bunsen burner (1855)
Rechargeable battery
(1859, Germany)
Electric lights (1870s)
Cathode ray tube (
TV) (1878)
Internal combustion
engine (1886;
Daimler, Maybach and
Benz )
X-ray photography
(1895)
Car (1897)
Submarine (late
1800s, some earlier)
THREE LAWS (LATE 1700S, ~1800)
1.
Law of Conservation of Mass
2.
Law of Definite Proportions
2. Law of Multiple Proportions
1. LAW OF CONSERVATION OF MASS
Antoine Lavoisier
•
•
born 1740
•
“father of modern
chemistry”
•
invented a balance
that read to 0.0005
g
turned to science in
his 20’s
1.
LAW OF CONSERVATION OF
MASS
MASS OF PRODUCTS = MASS OF
REACTANTS
Lavoisier heated tin in air  tin oxide
• e.g.
2 Sn(s)
+
O2(g)  2 SnO(s)
50.00 g + 6.74 g  56.74 g
• first experimental evidence for law of conservation of mass
2. Law of Definite Proportions
Joseph Proust (1799)
Compounds always
contain elements in the
same proportion by mass,
no matter how they are
made or where they are
found.
2. Law of Definite Proportions
SnO will always be
50.00 g Sn x 100% = 88.12% Sn
56.74 g
and
6.74 g O x 100% = 11.88% O
56.74 g
no matter where SnO is found or how it is made
2. Law of Definite Proportions
In modern, molar terms:
50.00 g Sn x 1 mol Sn = 0.4212 mol Sn
118.71 g Sn
and
6.74 g O x 1 mol O = 0.4212 mol O
16.00 g O
3. LAW OF MULTIPLE PROPORTIONS
John Dalton (1803)
Different compounds made
from the same elements:
The ratio of mass of an
element in the first
compound relative to the
mass of the same element in
a second compound is a
fixed whole number.
3. LAW OF MULTIPLE PROPORTIONS
e.g. SnO2 vs SnO (two oxides of tin)
Begin with 50.0 g Sn in both cases:
56.74 g SnO contains 6.74 g O
63.48 g SnO2 contains 13.48 g O
Ratio of the masses of O in these two compounds:
13.48 g O in SnO2 = 2.00
6.74 g O in SnO
3. LAW OF MULTIPLE PROPORTIONS
In modern, molar terms:
6.74 g O x 1 mol O in SnO = 0.4212 mol O
16.00 g O
and
13.48 g O x 1 mol O in SnO2 = 0.8414 mol O
16.00 g O
Mole ratio of O’s: 0.8414 mol in SnO2= 2.000
0.4212 mol in SnO
DALTON’S ATOMIC THEORY
1.
All matter is made of extremely small particles called atoms.
2.
All atoms of a given element are identical (mass, physical and
chemical properties).
3.
Atoms of different elements have different masses, and different
physical and chemical properties.
4.
Different atoms combine in simple whole number ratios to form
compounds.
5.
In a chemical reaction, atoms are combined, separated, or
rearranged.
6.
Atoms cannot be created, divided into smaller particles, or
destroyed.
DALTON’S ATOMIC THEORY
1.
2.
3.
4.
5.
6.
All matter is made of extremely small particles called
atoms.
All atoms of a given element are identical (mass,
physical and chemical properties).
Atoms of different elements have different masses,
and different physical and chemical properties.
Different atoms combine in simple whole number ratios
to form compounds.
In a chemical reaction, atoms are combined,
separated, or rearranged.
Atoms cannot be created, divided into smaller
particles, or
destroyed.
DALTON’S MODEL OF THE ATOM
• The 3 laws
• Law of conservation of mass,
• Law of definite proportions
• Law of multiple proportions (Dalton)
 Dalton’s theory of the atom
SUBATOMIC PARTICLE DISCOVERY
Electron was first
discovered!
J.J. Thomson (1897)
Received the Nobel Prize in
Physics in 1906 for his work
on the electron
1.
Experiment: Cathode Rays
• Passed electricity through partially evacuated tube
of gas
• Observed a ray of light passing from one electrode
to the other
• Ray moved a paddle wheel inside the tube
2.
Conclusions
•
Ray must be (-) because it moved toward (+)
electrode
•
Ray must be made of particles (moved the
paddle wheel)
•
Particles must be in all atoms (same results with
different gases)
J. J. THOMSON’S MODEL OF THE
ATOM
positively-charged “dough” embedded with small
negatively-charged particles called electrons
1.
NUCLEONS
Protons: 1909/1910
Ernest Rutherford
Received the Nobel
Prize in Chemistry in
1908 for his work in
radioactivity
GOLD FOIL EXPERIMENT
Alpha particles (He nucleus, + charge) shot through gold foil were deflected in
peculiar ways, inconsistent with the plum pudding model of the atom  Rutherford’s
idea that there is an area of concentration of positive charge
GOLD FOIL EXPERIMENT
GOLD FOIL EXPERIMENT
Hans Geiger and Marsden performed the experiments;
Rutherford interpreted them
COMPARISON BETWEEN THOMSON’S AND
RUTHERFORD’S MODELS OF THE ATOM
WARM UP – THREE LAWS AND
DALTON’S THEORY
1.
Name the three laws and briefly describe what they say.
Give examples for each one if you can.
2.
What are the five parts of Dalton’s theory?
a) Make a table with a column describing his theory, then next to this
column, indicate which parts of his theory are still considered to be
true, and which ones are no longer considered to be correct.
b) Analyze your table – is there a common theme that unites the parts
of Dalton’s theory that are no longer considered to be correct?
If so, what is the common theme? (describe it)
NUCLEONS
2.
Neutrons: 1932
James Chadwick
received the Nobel
Prize in Physics in 1935
for the discovery of the
neutron
(worked with Hans
Geiger, then Ernest
Rutherford)
NUCLEONS
2.
Uncharged particles in the nucleus with mass were
pushed out of beryllium when bombarded with alpha
particles. These particles accounted for the “missing
mass” in the nucleus.
SUBATOMIC PARTICLES
Relative
Particle Location Charge
Mass
Symbol
Electron
Outside
Nucleus
Proton
Inside
Nucleus
Neutron
Inside
Nucleus
SUBATOMIC PARTICLES
Relative
Particle Location Charge
Mass
Symbol
Electron
Outside
Nucleus
-1
1/1840
e-
Proton
Inside
Nucleus
+1
1
p+
Neutron
Inside
Nucleus
0
1
no
See diagram on board of
relative sizes of the parts of
the atom
 record in your notes
ATOMIC MASS
• Atomic number = # protons (also = # electrons)
• Mass number = # protons + # neutrons
• # neutrons = mass number – atomic number
ISOTOPES
#
#
#
Protons Neutrons Electrons
Isotope
Hydrogen-1
protium
1 H
1
1
0
1
Hydrogen-2
deuterium
2
1H
1
1
1
Hydrogen-3
tritium
3
1H
1
2
1
WARM UP
Compare the atomic models of Dalton, Thomson, and Rutherford:
For each model,
1.
draw a diagram,
2.
use words to describe the model,
3.
use words to describe the experiment that provided the
evidence to change the preceding model into the next one,
4.
explicitly state the connection that prompted Thomson and
Rutherford made to propose the changes from the previous
model to their new model
LAST TIME
Mass number
Atomic number
Number of protons
Number of neutrons
Number of electrons
WHOLE NUMBERS
AVERAGE ATOMIC MASS
• 1 atomic mass unit (amu)=1/12 mass of a C-12 atom
• average Atomic Mass = weighted average (by
abundance)
• shown on Periodic Table
AVERAGE ATOMIC MASS PRACTICE
POGIL
AVERAGE ATOMIC MASS PRACTICE
ATOMS
What is the average atomic mass of antimony?
The isotopes of antimony and their percent
abundances are Sb-121 (120.90 amu, 57.21%)
and Sb-123 (122.90 amu, 42.79%)
Use your periodic table to check your answer.
AVERAGE ATOMIC MASS PRACTICE
ATOMS
What is the average atomic mass of vanadium?
The isotopes of vanadium and their percent
abundances are V-50 (49.95 amu, 0.250%) and V51 (50.94 amu, 99.750%).
Use your periodic table to check your answer.
WARMUP
(FROM ATOMIC THEORY WS #1)
2.
Give the number of protons, electrons, and neutrons in each
of the following atoms.
a. 10847Ag b. 4020Ca
c. 2311Na
3.
Name each isotope, and write it in symbolic notation.
a. Atomic number 26; mass number 56
b. Atomic number 29; mass number 64
c. Atomic number 17; mass number 37
WARMUP KEY
(FROM ATOMIC THEORY WS #1)
2.
3.
Give the number of protons, electrons, and neutrons in each of the following
atoms.
a. 10847Ag 47 p+, 47 e-, 61 n0
b.
40 Ca
20
20 p+, 20 e-, 20 n0
c.
23 Na
11
11 p+, 11 e-, 12 n0
Name each isotope, and write it in symbolic notation.
a. Atomic number 26; mass number 56 iron-56,
56 Fe
26
b. Atomic number 29; mass number 64 copper-64, 6429Cu
c. Atomic number 17; mass number 37 chlorine-37, 3717Cl
WARMUP
(FROM ATOMIC THEORY WS #1)
4.
How many protons, electrons and neutrons are in each of the following isotopes?
a. Uranium-235
b. Hydrogen-3
c. Silicon-29
Additional problem:
How many protons, electrons and neutrons are in the following ion isotopes?
a.
14.
75
333As
b.
180
6+
74W
An element has three naturally occurring isotopes
Isotope 1 has a mass of 19.992 amu, 90.48% abundance
Isotope 2 has a mass of 20.994 amu, 0.27% abundance
Isotope 3 has a mass of 21.991 amu, 9.25% abundance
a. Calculate the (average) atomic mass of the element.
b. Identify the element, using the periodic table.
WARMUP
(FROM ATOMIC THEORY WS #1)
4.
How many protons, electrons and neutrons are in each of the following isotopes?
a. Uranium-235
92 p+, 92 e-, 143 n0
b. Hydrogen-3
1 p+, 1 e-, 2 n0
c. Silicon-29
14 p+, 14 e-, 15 n0
Additional problem:
How many protons, electrons and neutrons are in the following ion isotopes?
a.
14.
75
333As
33 p+, 36 e-, 42 n0
b.
180
6+
74W
74 p+, 68 e-, 106 n0
An element has three naturally occurring isotopes
Isotope 1 has a mass of 19.992 amu, 90.48% abundance
Isotope 2 has a mass of 20.994 amu, 0.27% abundance
Isotope 3 has a mass of 21.991 amu, 9.25% abundance
a. Calculate the (average) atomic mass of the element. 20.18 amu
b. Identify the element, using the periodic table. neon
ATOMIC #, MASS # AND AVERAGE
ATOMIC MASS
Atomic #
Mass #
Average
Atomic Mass
# p+
#p+ + #no
whole #
whole #
Weighted
average of all
isotopes
Decimal, limited
by sfs
Found on PT
NOT on PT
Found on PT
RADIOACTIVITY
• Nuclear Reaction  change in the identity of the elements
• Radioactivity = radiation emitted by atoms with an unstable
n0:p+ ratio
RADIOACTIVITY
• Smaller
Elements (atomic # < 20)
• Stable ratio = 1 n0:1 p+
i.e. Mass # = 2 x atomic #
• Larger Elements
- Stable ratio = 1.5 n0:1p+
- All elements with atomic # > 83 are radioactive
RADIOACTIVITY
• Unstable nuclei  emit radiation and change their identities
• This is called radioactive decay
HISTORICAL
FIGURES
• Wilhelm
Roentgen (1845•
•
1923) - discovered X-rays –
1895
X-ray of his wife’s hand
Nobel Prize in Physics, 1901
HISTORICAL
FIGURES
• Henri
Becquerel (18521908) - discovered
radioactivity in U - 1896
• Nobel Prize in Physics,
1903, shared with the
Curies, for his discovery of
spontaneous radioactivity
HISTORICAL
FIGURES
• Ernest Rutherford –
identified different
types of radiation,
and explored their
properties (beg.
1898)
• Nobel Prize in
Chemistry in 1908
HISTORICAL
FIGURES
• Pierre
(1859-1906) and Marie
Curie (1867-1934) - discovered
radium and polonium – 1898; first
used the term “radioactivity”
• Nobel Prize in Physics 1903, Pierre
and Marie, with Henri Becquerel
• Nobel Prize in Chemistry 1911
(Marie only) for discoveries of
radium and polonium
TYPES OF RADIATION
1.
Alpha radiation
(most common in elements with atomic # > 83,
 increase the number of neutrons)
Alpha particles = 2 p+ + 2 n0
(He nucleus, 42He, a) with 2+ charge
e.g.
226 Ra
88

222 Rn
86
+
4
2He
(+ energy)
TYPES OF RADIATION
2.
Beta radiation
(most common in elements with high n0:p+ ratio
 decrease the number of neutrons)
Beta particles = 1 e- (0-1b) with 1- charge
Neutron  proton + beta particle
1 n
0
e.g.

14 C
6
1 p
1

-1b
+
0
14 N
7
+
0
-1b
(+ energy)
TYPES OF RADIATION
Note:
The sum of the mass #s and atomic #s on both sides of the
equation are the same
TYPES OF RADIATION
3. Gamma radiation
Gamma rays = high-energy radiation with no mass and
no charge (00g)
usually accompany alpha and beta radiation
e.g.
238 U
92

234 Th
90
+
4 He
2
+ 2 00g
TYPES
Nuclei withOF
lowerRADIATION
neutron:proton ratios than optimal:
4.
Positron Emission
(most common in lighter elements with low n0:p+ ratio)
 more neutrons by converting a proton into a neutron
Positron = particle with same mass as an e-, but opposite
charge
Proton  neutron + positron
1 p  1 n + 0 b
1
0
1
e.g.
11 C
6

11 B
5
+
0
1b
TYPES
Nuclei withOF
lowerRADIATION
neutron:proton ratios than optimal:
5.
Electron Capture
(most common in elements with a high n0:p+ ratio)
 more neutrons by pulling in an e- which combines with a
proton to form a neutron
Proton + electron  neutron
1 p + 0 e  1 n
1
-1
0
e.g.
0 e
-1
+ 8137Rb 
81 Kr
36
+ X-ray photon
RADIOACTIVE PARTICLES WS
Positron
same mass as e-’s0+1b
Electron capture electrons
(Added to the reactants side)
0 e-1
-1
1/1840
1/1840
1+
1-
RADIOACTIVE PARTICLES WS
1.
Which radioactive emission has the greatest mass? Least mass?
2.
Why do you think gamma rays are drawn as wavy lines?
3.
Which charged plate are the alpha particles attracted to? Explain.
4.
Which charged plate are the beta particles attracted to? Why do
the beta particles have a greater curvature than the alpha particles?
5.
Explain why the gamma rays do not bend toward one of the
electrically charged plates.
RADIOACTIVE PARTICLES WS
1.
Which radioactive emission has the greatest mass? Least mass?
alpha – greatest; gamma – no mass
2.
Why do you think gamma rays are drawn as wavy lines?
Gamma rays have no mass and are EMR, which is often drawn as wavy lines.
3.
Which charged plate are the alpha particles attracted to? Explain.
To the negatively-charged plate, as the alpha particles are positively-charged.
4.
Which charged plate are the beta particles attracted to? Why do the beta particles have a greater
curvature than the alpha particles?
to the positively-charged plate, as the beta particles are negatively-charged. They have a smaller mass, so
are more greatly influenced by the electric field.
5.
Explain why the gamma rays do not bend toward one of the electrically charged plates.
Gamma rays have no charge, therefore, they are not attracted to either plate.
NUCLEAR FISSION
• Nuclear fission = the splitting of a nucleus into smaller,
more stable fragments, accompanied by a large release
of energy
e.g. Uranium-235:
235 U
92
+ 10n  23692U  9236Kr + 14156Ba + 3 10n + energy
(unstable)
The new neutrons (10n)  fission of more U-235 (= chain
reaction, a self-sustaining process)
NUCLEAR FISSION
• Chain reaction requires a critical mass (= minimum
amount of starting material to maintain a chain
reaction)
• supercritical mass may  violent nuclear explosion
• results in radioactive waste
• Practical examples = nuclear power plant, atomic
bomb
NUCLEAR FUSION
• Nuclear Fusion = the process of binding smaller
atomic nuclei into a single larger and more stable
nucleus, requiring a huge amount of energy to
initiate, followed by a large release of energy
1. CREATION OF NATURAL
ELEMENTS
Elements are created by nuclear reactions
a. Hydrogen, other light elements
- from the Big Bang
CREATION OF NATURAL ELEMENTS
b. Elements #2-92 (except Fr, Pr, Te, At)
Nuclear fusion occurs in stars (naturally)
Occurs in hydrogen bomb (artificially) > 2 x 107oC
The sun converts 3 x 1014 g of H into He every second.
4
1 H
1

4 He
2
+ energy
Mass is not conserved.
Mass is converted into energy via E = mc2
CREATION OF NATURAL ELEMENTS
Other fusion reactions occur in the sun:
+ g
4 He
2
+
4 He
2

4 He
2
+
8 Be
4
 126C + g
8 Be
4
(gamma ray)
2. SYNTHETIC ELEMENTS
a. Nuclear bullets
i. Bombard nuclei of elements with
small particles such as p+, n ,
4 He (a particles) & e- (0 b particles)
2
-1
ii. Elements # 93-100
2. SYNTHETIC ELEMENTS
iii. 1919 first experiment:
14 N
7
+
4 He
2
(Rutherford)

17 O
8
+
1 H
1
2. SYNTHETIC ELEMENTS
b. Crashing nuclei
i. Accelerators hurl nuclei into each other at very high speeds.
e.g. 126C + 24496Cm  254102No + 2 10n
carbon
curium
nobelium
neutron
ii. Elements beyond #100
iii. These elements are very unstable:
e.g. Element 109 existed for only 3.4 x 10-3 sec (3 atoms)
2. SYNTHETIC ELEMENTS
c.
Superheavy elements
(“transuranium” elements)
Stability of nucleus of atom depends on filling "shells" within nucleus with
alternating p+ and n.
The more filled shells, the more stable it would be.
e.g. Element 114
244
94Pu
+
48
20Ca

289
114Fl
+ 3 10n (1999, Russia)
NUCLEAR EQUATIONS
Complete the following equations :
214 Bi
83
 42He + _____
239 Np
93
 23994Pu + ______
PRODUCTION OF TRANSURANIUM
ELEMENTS
PRODUCTION OF TRANSURANIUM
ELEMENTS WS
1. Does the diagram illustrate a natural transmutation reaction or an induced transmutation reaction?
2. What is the name and nuclear symbol of the isotope produced in the reaction?
3. Write a nuclear equation to show how dubnium-263, lawrencium-262, and seaborgium-266 can be produced from a nuclear reaction of neon-22 and americium-244.
5.
22 Ne
10
+
244 Am
95
 263105Db + 3 10n
22 Ne
10
+
244 Am
95

22 Ne
10
+
244 Am
95

Each of the radioisotopes in the table decays within 20 seconds to 10 hours. Write a nuclear equation for each decay.
 42He +
266 Am
95
263
105Db
262
103Lr
266
106Sg
6.
 0-1b +
+ 0-1e- 
 42He +
Which, if any, of the four isotopes listed in the table would you expect to find at Earth’s surface? Why?
PRODUCTION OF TRANSURANIUM
ELEMENTS WS
1. Does the diagram illustrate a natural transmutation reaction or an induced
transmutation reaction? Induced transmutation
2. What is the name
and nuclear symbol of the isotope produced in the reaction?
Dubnium-266; 266105Db
3. Write a nuclear equation to show how dubnium-263, lawrencium-262, and
seaborgium-266 can be produced from a nuclear reaction of neon-22 and americium244.
22 Ne
10
+
244 Am
95
 263105Db + 3 10n
22 Ne
10
+
244 Am
95
 262103Lr + 42He
22 Ne
10
+
244 Am
95
 266106Sg + 0-1b
PRODUCTION OF TRANSURANIUM
ELEMENTS WS
5. Each of the radioisotopes in the table decays within 20 seconds to 10 hours. Write a
nuclear equation for each decay.
244 Am
95
 42He + 24093Np
263
105Db
 0-1b + 263106Sg
262
103Lr
+ 0-1e-  262102No
266
106Sg
 42He + 262104Rf
6. Which, if any, of the four isotopes listed in the table would you expect to find at
Earth’s surface? Why? None – they all have very short half-lives.
NUCLEAR EQUATIONS
Complete the following equations :
214 Bi
83
 42He + ______
239 Np
93
 23994Pu + ______
NUCLEAR EQUATIONS
Complete the following equations :
214 Bi
83
 42He + _21081Tl_
239 Np
93
 23994Pu + __0-1b____
NUCLEAR EQUATIONS
Write a balanced nuclear equation for the alpha decay of
americium-241.
Write a balanced nuclear equation for the beta decay of
bromine-84.
NUCLEAR EQUATIONS
Write a balanced nuclear equation for the alpha decay of americium-241.
241 Am
95
 42He +
237 Np
93
Write a balanced nuclear equation for the beta decay of bromine-84.
84 Br
35

0 b
-1
+
84 Kr
36
NUCLEAR EQUATIONS
Complete the following equations:
214 Bi
83
 42He +
239 Np
93
241
95
 23994Pu +
Am  42He +
84 Br
35

_______
+ 0-1b
NUCLEAR EQUATIONS (KEY)
Complete the following equations:
214 Bi
83
 42He +
239 Np
93
241
95
210 Tl
81
 23994Pu + 0-1b
Am  42He + 23793Np
84 Br
35

84 Kr
36
+ 0-1b
WARMUP – NUCLEAR EQUATIONS,
INCLUDING NUCLEAR FUSION
1.
9 Be
4
2.
238 U
92
3.
15 C
6
4.
137 Cs
55
+
+
1 H
1
+

4 He
2
1 n
0
4 He
2
+ _______
 2 10n + _______
 _______
 ________ +
0 b
-1
WARMUP – NUCLEAR EQUATIONS
1.
9 Be
4
2.
238 U
92
3.
15 C
6
4.
137 Cs
55
+
+
1 H
1
+

4 He
2
1 n
0


4 He
2
+
6 Li
3
 2 10n +
16 C
6
137 Ba
56
+
0 b
-1
240 Pu
94
NEXT STEPS - PROPERTIES OF ELECTRONS
•
Wave nature of light – EMR
(James Maxwell, 1864)
•
Particle nature of light – quantum
(Max Planck, late 1800s)
•
Emission of light and other EMR from heated elements 
emission spectra
ELECTROMAGNETIC RADIATION
EMR
= energy that exhibits wave-like
behavior as it travels through
space
James Maxwell (1864) –
unified electric and
magnetic forces into
electromagnetic force
ELECTROMAGNETIC RADIATION
Unified the electric and
magnetic forces

electromagnetic force (emf)
James Maxwell (1864) –
unified electric and
magnetic forces into
electromagnetic force
ELECTROMAGNETIC RADIATION
• Speed of EMR always the same
c = 3.00 x 108 m/s
←Memorize
• Examples include: microwaves, TV, Radio, X-rays
ELECTROMAGNETIC RADIATION
Wavelength = λ (lambda)
•
usually in nm
Frequency = n (nu) or f
•
Waves per second = Hz (Hertz) = cycles/s or s-1
Speed = c, measured in m/s
c =λ n
Note the inverse relationship between λ and n
ELECTROMAGNETIC SPECTRUM
DEMO
EMR speed is the same, while frequency and wavelength
change – red vs. blue light
EMR SPECTRUM
EMR SPECTRUM
1.
What kinds of waves have the longest wavelength? What kinds of waves have the shortest wavelength?
2.
Which waves have the lowest frequency?
3.
Which has a higher frequency: microwaves or X rays?
4.
Which waves can be seen by the eye?
5.
Sequence the different segments of the visible spectrum in order from shortest wavelength to longest wavelength.
6.
Sequence the following types of waves from lowest frequency to highest frequency: ultraviolet rays, infrared rays, gamma rays,
radio waves, and green light.
7.
Compare the wavelengths and frequencies of each kind of wave. What is the relationship between frequency and wavelength?
8.
What is the wavelength of a radio station emitting its signal at 95.5 MHz? Estimate your answer to the nearest power of ten.
EMR SPECTRUM
1.
What kinds of waves have the longest wavelength? Radio waves
2.
What kinds of waves have the shortest wavelength? Gamma rays
3.
Which waves have the lowest frequency? Radio waves
4.
Which has a higher frequency: microwaves or X rays? X-rays
5.
Which waves can be seen by the eye? Visible portion of the
spectrum
EMR SPECTRUM
6.
Sequence the different segments of the visible spectrum in order from shortest
wavelength to longest wavelength.
Violet, Indigo, Blue, Green, Yellow, Orange, Red
7.
Sequence the following types of waves from lowest frequency to highest
frequency:
radio waves, infrared waves, green light, ultraviolet waves, gamma rays
8.
Compare the wavelengths and frequencies of each kind of wave. What is the
relationship between frequency and wavelength? Inversely proportional
9.
What is the wavelength of a radio station emitting its signal at 95.5 MHz?
Estimate your answer to the nearest power of ten. About 3 m, or 3 x 100 m
EMR PRACTICE PROBLEMS FROM
TEXTBOOK (PP. 121, 124)
c=lxn
Microwaves are used to transmit information. What is the
wavelength of a microwave having a frequency of 3.44 x
109 Hz? (8.72 x 10-2 m)
1.
EMR PRACTICE PROBLEMS
C=LXN
What is the frequency of green light which has a wavelength of 4.90 x 10-7
m? (6.12 x 1014 s-1)
2.
An X-ray has a wavelength of 1.14 x 10-10 m. What is its frequency?
x 1018 s-1)
3.
What is the speed of an electromagnetic wave that has a frequency of 7.8 x
106 Hz? (3.00 x 108 m/s)
4.
A popular radio station broadcasts with a frequency of 94.7 MHz.
What is the wavelength of the broadcast? (1 MHz = 106 Hz) (3.17 m)
(2.63
QUANTUM
• Max Planck (1858-1947)–
Nobel Prize in Physics, 1918,
for his discovery of energy
quanta
• Revolutionary concept in
physics
THE IDEA OF THE QUANTUM
• Quantum = the smallest discrete amount of energy that can exist
independently, esp. as EMR
• 1 quantum = 1 photon
• E = hn, where h = a constant
• The amount of energy in EMR increases with increasing frequency
DEMO
Photons – glow in the dark
EMR PRACTICE PROBLEMS
E = HN
Planck’s constant = h = 6.626 x 10-34 J·s
Example: Tiny water drops in the air disperse the
white light of the sun into a rainbow. What is the
energy of a photon from the violet portion of the
rainbow if it has a frequency of 7.23 x 1014 s-1?
(4.79 x 10-19 J)
EMR PRACTICE PROBLEMS
E = HN
Planck’s constant = h = 6.626 x 10-34 J·s
5.
What is the energy of each of the following types of EMR?
a. 6.32 x 1020 s-1
(4.19 x 10-13 J)
b. 9.50 x 1013 Hz
(6.29 x 10-20 J)
c. 1.05 x 1016 s-1
(6.96 x 10-18 J)
6.
Name the types of radiation in each part of #5.
1.
EMR PRACTICE PROBLEMS
What is the frequency of EMR with a wavelength of 235 pm? What type of EMR is this?
2.
What is the frequency of EMR with a wavelength of 0.614 cm? What type of EMR is this?
3.
What is the wavelength of EMR with a frequency of 8,512 Hz? What type of EMR is this?
4.
What is the wavelength of EMR with a frequency of 625 x 1017 Hz? What type of EMR is this?
5.
If the speed of light is 3.00 x 108 m/s, calculate the wavelength of the electromagnetic
radiation whose frequency is 7.500 x 1012 Hz.
6.
Determine the frequency of light with a wavelength of 4.257 x 10-7 cm.
7.
For the following sources, calculate the missing member of the wavelength/frequency pair.
8.
a)
FM radio waves with a frequency of 94.7 Hz.
b)
A laser with a wavelength of 1064 nm.
c)
An X-ray source, emitting X-rays with a wavelength of 175.4 pm.
How long would it take a radiowave with a frequency of 7.25 x 105 Hz to travel from
Mars to Earth if this distance between the two planets is approximately 8.00 x 107 km?
1.
EMR
PRACTICE PROBLEMS
What is the frequency of EMR with a wavelength of 235 pm? What type of
EMR is this?
(1.28 x 1018 s-1; X-rays)
2.
What is the frequency of EMR with a wavelength of 0.614 cm? What type
of EMR is this?
(4.89 x 1010 s-1; microwaves)
3.
What is the wavelength of EMR with a frequency of 8,512 Hz? What type
of EMR is this?
(3.52 x 104 m; radio waves)
4.
What is the wavelength of EMR with a frequency of 625 x 1017 Hz? What
type of EMR is this?
(4.80 x 10-12 m; X-rays, gamma rays)
5.
EMR PRACTICE PROBLEMS
If the speed of light is 3.00 x 108 m/s, calculate the wavelength of the electromagnetic
radiation whose frequency is 7.500 x 1012 Hz.
(4.00 x 10-5 m)
6.
Determine the frequency of light with a wavelength of 4.257 x 10 -7 cm. (7.05 x 1016 s-1)
7.
For the following sources, calculate the missing member of the wavelength/frequency pair.
a) FM radio waves with a frequency of 94.7 Hz. (3.17 x 106 m)
b) A laser with a wavelength of 1064 nm. (2.82 x 1014 s-1)
c)
8.
An X-ray source, emitting X-rays with a wavelength of 175.4 pm. (1.71 x 1018 s-1)
How long would it take a radiowave with a frequency of 7.25 x 105 Hz to travel
from Mars to Earth if this distance between the two planets is approximately 8.00 x 10 7 km?
(Note: v is not required for the calculation.)
(2.67 x 102s)
WARMUP - EMR
1.
Which color has the shorter wavelength (l) – blue or red?
2.
Which color has the higher frequency (n) – blue or red?
3.
What is the wavelength of light with a frequency of
4.90 x 1016 s-1)? (A: 6.12 x 10-9 m)
4.
What is the frequency of10EMR
with a wavelength of 5.26 mm? What type of
EMR is it? (A: 5.70 x 10 s-1)
5.
Determine the energy,-34
in joules, of a photon whose frequency is 3.55 x 1017
Hz. (h = 6.626 x 10 J s)
(A: 2.35 x 10-16 J)
WARMUP - HONORS
5.
When sodium is heated, a yellow spectral line whose
energy is 3.37 x 10-19 J per each photon is produced.
a. What is the frequency of this light?
(A = 5.09 x 1014 s-1)
b. What is its wavelength?
(A = 5.89 x 10-7 m)
PLANETARY MODEL – NEILS BOHR
• Neils Bohr
(1885-1962)
(Danish physicist)
• Studied with Thomson and Rutherford
• Refined Rutherford’s model in 1913
• Received the Nobel Prize in 1922 for
his work on the structure of atoms
PLANETARY MODEL – NEILS BOHR
• Neils Bohr
(1885-1962)
(Danish physicist)
• Incorporated Planck’s idea of quanta
of energy
• Provided an explanation for the
spectral lines of hydrogen
BOHR MODEL
Electrons…
1.
are arranged in circular paths around
nucleus, “orbits”.
2. have fixed energy levels to prevent them from
falling into nucleus.
Electrons closest to the nucleus have
lowest Etotal = KE + PE (most stable).
BOHR’S MODEL
1 = lowest energy
4 = highest energy
E3 - E1
+
+ 1
2
34
As e- goes further
= a quantum of energy in
the form of EMR
As e- goes further away
from the nucleus, it
increases in potential
energy
BOHR MODEL
Electrons…
3. must gain or lose energy to change energy
levels.
EMR is emitted from the atom when electrons fall
down to a lower energy level.
4. in different energy levels are not the same
distance apart.
A “quantum” = amount of energy needed to
make the leap between energy level.
BOHR MODEL
Bohr’s model did not explain the line spectra of
atoms with >1 electron.
DEFINITIONS RELATED TO SPECTRA
Spectrum
= whole range of related qualities
[Latin: appearance, from specere – to view]
Electromagnetic spectrum = all EMR arranged according to l
DEFINITIONS RELATED TO SPECTRA
Emission = any radiation of energy by EM waves
[Latin: emitto – to send out, to utter, to hurl, to set free]
Emission spectrum = the spectrum into which light or other EMR from
any source can be separated
Continuous spectrum = a spectrum whose source emits light of every
l in a continuous band
Bright-line spectrum = pattern of bright lines on a dark background.
Source = glowing gas that radiates in special l’s characteristic of the
chemical composition of the gas
CONTINUOUS WHITE LIGHT
SPECTRUM
LINE-EMISSION SPECTRUM
excited state
ENERGY IN
PHOTON OUT
ground state
COMPARISON OF SPECTRA
COMPARISON OF CONTINUOUS, LINE
AND ABSORPTION SPECTRA
THE HYDROGEN SPECTRUM
Each photon emitted has a
characteristic λ which
contributes a line to the
spectrum
E2
E1
Visible
Balmer Series
UV
Lyman Series
Infrared Paschen
Series
765 43 21
THE HYDROGEN SPECTRUM
ATOMIC THEORY VIDEO QUESTIONS
#3
SPECTRA
1.
What is the orbit of an electron closest to the nucleus called?
2.
What region of the spectrum of excited hydrogen gas did
a. Balmer predict?
________________
b. Paschen predict? ________________
c. Lyman predict?
________________
3.
What was wrong with Balmer’s formula?
4.
What made Bohr’s mathematical model so special?
5.
Does the electron emit radiation when it is bumped up an energy level or when it falls back down?
6.
What level (n) does the electron fall to in order to produce:
a. the Balmer series?
_______________
b. the Paschen series? _______________
c. the Lyman series?
_______________
1.
ATOMIC THEORY VIDEO QUESTIONS
#3
SPECTRA KEY
What is the orbit of an electron closest to the nucleus called? ground state
2. What region of the spectrum of excited hydrogen gas did
a. Balmer predict?
visible
b. Paschen predict? infrared
c. Lyman predict?
3.
UV
What was wrong with Balmer’s formula? Nobody knew why it worked
4. What made Bohr’s mathematical model so special? Based on a possible structure of the atom
5. Does the electron emit radiation when it is bumped up an energy level or when it falls back down?
when it falls
6. What level (n) does the electron fall to in order to produce:
a. the Balmer series?
2
b. the Paschen series? 3
c. the Lyman series?
1
MODERN (QUANTUM) THEORY
• Wave nature of the electron
• Louis de Broglie received the
Nobel Prize for Physics in 1929
for his discovery of the wave
nature of electrons
MODERN (QUANTUM) THEORY
• Quantum Mechanics
• Erwin Schrödinger
received the Nobel Prize
for Physics in 1933 for his
work in atomic theory
• Wave equation: electrons
as waves (1926)
• Foundation of the quantum
theory of the atom
MODERN (QUANTUM) THEORY
Orbital
MODERN (QUANTUM) THEORY
• Quantum Mechanics
• Werner Heisenberg,
• received Nobel Prize for
Physics in 1932 for quantum
mechanics
• Heisenberg’s Uncertainty
Principle
• We cannot simultaneously
measure an electron’s position
and its velocity
ORBITALS AND QUANTUM #’S
• Orbital in the shape of an “electron cloud” contains all (90%)
locations of an electron
• Each e- in an atom has its own set of four
quantum #’s — n, l, m, and s
ORBITALS AND QUANTUM #’S
• Principal energy levels = Bohr’s orbits
• Total # of e- in one principal energy level = 2n2
• Sublevels (l), magnetic position (m), and
spin (s)—additions to classify e- energies
Atomic Orbitals: predict 90% probability of location of
electrons (electron cloud)
Each orbital can contain a maximum of 2 electrons, spinning
in opposite directions.
p
s
d
f
ORBITALS
ORBITALS WS
1.
What is the shape of an s orbital?
2.
What is the relationship between the size of an s orbital and the principal energy level in which it is found?
3.
What is the shape of a p orbital? How many p orbitals are there in a sublevel?
4.
How many electrons can each orbital hold?
5.
Look at the diagrams of the p orbitals. What do x, y and z refer to?
6.
How many d orbitals are there in a given sublevel? How many total electrons can the d orbitals in a sublevel hold?
7.
Which d orbitals have the same shape?
8.
What point in each diagram represents an atom’s nucleus?
9.
How likely is it that an electron occupying a p or a d orbital would be found very near an atom’s nucleus? What part of the diagram supports your conclusion?
ORBITALS WS
1.
What is the shape of an s orbital? spherical
2.
What is the relationship between the size of an s orbital and the principal energy level in which it is found? Size increases with increasing principal energy level
3.
What is the shape of a p orbital? Dumbbell How many p orbitals are there in a sublevel? 3
4.
How many electrons can each orbital hold? 2
5.
Look at the diagrams of the p orbitals. What do x, y and z refer to? 3 perpendicular axes
6.
How many d orbitals are there in a given sublevel? 5 How many total electrons can the d orbitals in a sublevel hold? 10
7.
Which d orbitals have the same shape? 4 out of the 5
8.
What point in each diagram represents an atom’s nucleus? The origin – where the x, y and z axes intersect
9.
How likely is it that an electron occupying a p or a d orbital would be found very near an atom’s nucleus? Very unlikely What part of the diagram supports your conclusion? The shapes
of the orbitals come to a point at the intersection of the three axes, making the possibility of an electron being found there very unlikely.
ORIGIN OF ORBITAL NAMES
• s – sharp (or use sphere)
• p – principal (peanut)
• d – diffuse (daffodil or daisy)
• f – fundamental (funky)
•
Names come from the spectrum analysis,
e.g. the hyperfine splitting of the d-line of the sodium spectrum
Principal Energy
Level (n)
(or “Shell”)
distance from
nucleus
s:
p:
d:
f:
Sublevel (l)
(or “Subshell”)
shape of
probability
cloud
3-D (m)
(position in
space)
x, y, z
Spin (s)
+ ½
- ½
1 e- + ½
1 e- - ½
Total # e2n2
2(1)2 = 2
1
s
only 1 orientation
2
s, p
3 orientations for
p:
px, py, pz
2(2)2 = 8
3
s, p, d
etc.
2(3)2 = 18
4
s, p, d, f
5
s, p, d, f
6
s, p, d, (f)
7
s, p, d, (f)
holds 2 eholds 2 eholds 2 eholds 2 e-
x
x
x
x
1 orbital
3 orbitals
5 orbitals
7 orbitals
2(4)2 = 32
= 2 e- total
= 6 e- total
= 10 e- total
= 14 e- total
• n = # of sublevels per level
• n2 = # of orbitals per level
• Sublevel sets: 1 s, 3 p, 5 d, 7 f
1. THE P ORBITALS
px
py
pz
2. The d orbitals
Electron spin
• An orbital can hold 2 electrons that spin in opposite directions.
THREE RULES FOR FILLING ORBITALS
• Aufbau Principle
• Fill in order of increasing energy levels
• Hund’s Rule
• Fill all orbitals at same energy level with at
least 1 e-, before adding the second e-
• Pauli Exclusion Principle
• only 2 e- per orbital, of opposite spin
ELECTRON CONFIGURATION
ELECTRON CONFIGURATION WS
1.
What does each small box in the diagram represent?
2.
How many electrons can each orbital hold?
3.
How many electrons can the d sublevel hold?
4.
Which is associated with more energy: a 2s or a 2p orbital?
5.
Which is associated with more energy: a 2s or a 3s orbital?
6.
According to the Aufbau Principle, which orbital should fill first, a 4s or a 3d orbital?
7.
Which orbital has the least amount of energy?
8.
What is the likelihood that an atom contains a 1s orbital?
9. Sequence the following orbitals in the order that they should fill up according to the Aufbau Principle: 4d, 4p, 4f, 5s, 6s, 3d, 4s.
ELECTRON CONFIGURATION WS
KEY
1.
What does each small box in the diagram represent? An orbital
2.
How many electrons can each orbital hold? 2
3.
How many electrons can the d sublevel hold? 10
4.
Which is associated with more energy: a 2s or a 2p orbital?
5.
Which is associated with more energy: a 2s or a 3s orbital? 3s
6.
According to the Aufbau Principle, which orbital should fill first, a 4s or a 3d orbital? 4s
7.
Which orbital has the least amount of energy? 1s
8.
What is the likelihood that an atom contains a 1s orbital? 100%
2p
9. Sequence the following orbitals in the order that they should fill up according to the Aufbau Principle: 4d, 4p, 4f, 52, 62, 3d, 4s:
4s, 3d, 4p, 5s, 4d, 6s, 4f
ELECTRON CONFIGURATION
• Order of Filling—Aufbau Diagram
s
s
s
s
s
s
s
p
p
p
p
p
p
• Oxygen (8 e-)
d
d
d
d
d
f
f
f
f
Orbital Diagram
Electron Configuration
1s2 2s2
2p4
PRACTICE WITH ORBITAL DIAGRAMS AND
ELECTRON CONFIGURATIONS:
1.
Draw your own Aufbau diagram, then write orbital
diagrams and electron configurations for
oxygen
calcium
gallium
PRACTICE WITH COMPLETE AND
CONDENSED (NOBLE GAS) ELECTRON
CONFIGURATIONS:
2.
Write electron configurations
a) helium
b)
a) neon
b)
a) argon
b)
a) krypton
b)
a) xenon
b)
a) radon
b)
(only) for:
carbon
aluminum
bromine
palladium
lead
uranium
c) Write all b)’s in noble gas (condensed) configuration
d) Show how to jump into an Aufbau diagram
e) s,p,d,f blocks in the periodic table
ATOMIC THEORY VIDEO QUESTIONS
#4
ELECTRON ARRANGEMENT
1.
What 2 parts of Bohr’s model did Schrödinger and Heisenberg keep?
2.
How did they change Bohr’s model?
2.
What is the probability distribution of the first energy level called?
4.
What is the second type of probability distribution shaped like?
What is it called?
5.
How many orientations do s orbitals have?
p orbitals?
d orbitals?
6.
How many electrons can be in any one orbital?
7.
Which orbitals are filled with electrons first?
8.
What is the second “rule” about filling orbitals?
9.
Which electron orbitals have the most bearing on the chemical properties of a particular atom?
10.
What practical application do the periodic table groupings provide to working chemists?
ATOMIC THEORY VIDEO QUESTIONS # 4
ELECTRON ARRANGEMENT KEY
1.
What 2 parts of Bohr’s model did Schrödinger and Heisenberg keep?
Positively-charged nucleus, energy levels for electrons
2.
How did they change Bohr’s model? No definite orbits, but rather probability electron clouds called orbitals
2.
What is the probability distribution of the first energy level called? s
4.
What is the second type of probability distribution shaped like? Dumb-bells What is it called? p
5.
How many orientations do s orbitals have? 1
6.
How many electrons can be in any one orbital? 2, of opposite spin
7.
Which orbitals are filled with electrons first?
8.
What is the second “rule” about filling orbitals? 1 e- per orbital of same energy, then add second e-
9.
Which electron orbitals have the most bearing on the chemical properties of a particular atom?
p orbitals?
3 d orbitals? 5
Lowest energy
Last ones to be filled
10.
What practical application do the periodic table groupings provide to working chemists?
Helps predict chemical reactions
ATOMIC THEORY SCIENTISTS
ATOMIC THEORY AND THE
PERIODIC TABLE
s
1
2
3
4
5
6
7
f (n-2)
p
d (n-1)
6
7
© 1998 by Harcourt Brace & Company
E. PERIODIC PATTERNS
• Period #
• energy level (subtract for d & f)
• A/B Group #
• total # of valence e-
• Column within sublevel block
• # of e- in sublevel
PERIODIC PATTERNS
• Example - Hydrogen
1
2
3
4
5
6
7
1
1s
1st Period
1st column
of s-block
s-block
PERIODIC PATTERNS
• Shorthand Configuration
• Core e-: Go up one row and over to the Noble Gas.
• Valence e-: On the next row, fill in the # of e- in each sublevel.
1
2
3
4
5
6
7
PERIODIC PATTERNS
• Example - Germanium
1
2
3
4
5
6
7
[Ar]
2
4s
10
3d
2
4p
STABILITY
• Ion Formation
• Atoms gain or lose electrons to become more stable.
• Isoelectronic with the Noble Gases.
1
2
3
4
5
6
7