Atomic Theory. Essential Questions What are we made of? How are scientific models developed? Do...

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 PerfumeGunpowderGuns Hot air balloons Telegraph Electric motor Internal combustion engine

Car BatteryRechargeable batteryPhotographyX-ray photographyBunsen burnerGas lights Incandescent light bulbElectric lights Glass blowingSewing machine

Cathode ray tube ( TV) Submarine Vacuum technology Asphalt Tin cans CandleElectricity Conduction of electricity Graphite pencil

Before 1800Glass (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

3. Law of Multiple Proportions

1. Law of Conservation of Mass

Antoine Lavoisier• born 1740• turned to science

in his 20’s• “father of modern

chemistry”• invented a

balance that read to 0.0005 g

1. Law of Conservation of Massmass 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

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% Sn56.74 g

and6.74 g O x 100% = 11.88% O56.74 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

and6.74 g O x 1 mol O = 0.4212 mol O

16.00 g O

2. Law of Definite Proportions

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

Last time

• The 3 laws• Law of conservation of mass, • Law of definite proportions• Law of multiple proportions (Dalton)

Dalton’s theory of the atom

Dalton’s Atomic Theory1. 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 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 Theory1. 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 Model of the Atom

Subatomic Particle Discovery

Electron was the 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

Nucleons

1. 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)

Atomic Theory Video Questions #1The Earliest Models - Mists of Prehistory

1. How do Democritus's theories compare with today’s theories?

2. Who did not believe in atoms?

3. Why did the church oppose atomism?

4. What did alchemists try to do?

5. What did Ben Franklin believe lightning to be?

6. What did Lavoisier show?

Atomic Theory Video Questions #1The Earliest Models - Mists of Prehistory

1. How do Democritus's theories compare with today’s theories? pretty well 3/5

2. Who did not believe in atoms? Plato, Aristotle

3. Why did the church oppose atomism? Human spirit claimed to be explained by physical means - atoms

4. What did alchemists try to do? Introduced the idea of observation and experimentation, while trying to change common metals into silver and gold

5. What did Ben Franklin believe lightning to be? A fluid – excess = + charge, deficit = - charge

6. What did Lavoisier show? Mass of reactants = mass of products (Law of Conservation of Mass)

Atomic Theory Video Questions #1Smaller than the Smallest

1. What was Dalton trying to figure out when he discovered atomic theory?

2. What did Faraday think of electricity?

3. How did Millikan establish an electron’s charge?

4. What is the raisin bun model?

Atomic Theory Video Questions #1 Smaller than the Smallest

1. What was Dalton trying to figure out when he discovered atomic theory? Why water absorbs more of one type of gas than another

2. What did Faraday think of electricity? = force of affinity that holds atoms together

3. How did Millikan establish an electron’s charge? Microscopic oil droplets which carry a charge – pass X-rays through plates free electrons

4. What is the raisin bun model? Positively-charged dough with negative charges sprinkled throughout

Atomic Theory Video Questions #1The Rutherford Model

1. What do electrostatic forces affect?

2. What are the two components of the interaction of charged particles?

3. What does his structure of the atom mimic?

4. What was wrong with Rutherford’s model?

Atomic Theory Video Questions #1The Rutherford Model

1. What do electrostatic forces affect? The path/direction of alpha particles as well as their speed

2. What are the two components of the interaction of charged particles? Perpendicular – change in direction, parallel – change in speed

3. What does his structure of the atom mimic? The solar/planetary system

4. What was wrong with Rutherford’s model? Electrons should lose energy as they accelerate/move around the nucleus, showing a frequency shift but there is no frequency shift

Nucleons2. 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

Particle Location ChargeRelative

Mass Symbol

Electron

Proton

Neutron

Outside Nucleus

Inside Nucleus

Inside Nucleus

Subatomic Particles

Particle Location ChargeRelative

Mass Symbol

Electron

Proton

Neutron

Outside Nucleus

Inside Nucleus

Inside Nucleus

-1

+1

0 1

1

1/1840 e-

p+

no

See diagram on board of relative sizes of the parts of the atom

record in your notes

Atomic Theory Video Questions #2Introducing the Players

1. What particle gives the atom its positive charge?

2. What are the three building blocks of atoms?

3. What positively identifies the atom?

4. What defines an isotope?

5. Do isotopes have the same chemical and physical properties?

6. Does the nucleus change in a chemical reaction?

7. Why would the electrons in Rutherford’s model eventually crash into the nucleus?

Atomic Theory Video Questions #2Introducing the Players

1. What particle gives the nucleus its positive charge? proton

2. What are the three building blocks of atoms? Electrons, protons, neutrons

3. What positively identifies the atom? # protons

4. What defines an isotope? # neutrons

5. Do isotopes have the same chemical and physical properties? Basically yes – only mass differs

6. Does the nucleus change in a chemical reaction? no

7. Why would the electrons in Rutherford’s model eventually crash into the nucleus? An accelerating mass emits EMR; loses energy and crashes into the nucleus

Atomic Mass

• Atomic number = # protons (also = # electrons)

• Mass number = # protons + # neutrons

• # neutrons = mass number – atomic number

Isotopes

Isotope#

Protons#

Neutrons#

Electrons

Hydrogen-1 protium

11H 1 0 1

Hydrogen-2 deuterium

21H 1 1 1

Hydrogen-3 tritium

31H 1 2 1

Isotopes of Hydrogen and Carbon

Isotopes

Isotopes WS1. What do the following symbols represent?

a. e- _________________b. n0 _________________c. p+ _________________

2. Which subatomic particles are found in an atom’s nucleus?

3. Which subatomic particle identifies an atom as that of a particular element?

4. Explain why atoms are neutral even though they contain charged particles.

5. What do the numbers, 39, 40, and 41 after the element name potassium refer to?

6. Write the symbolic notation for each of the following isotopes:a. potassium-39 ________________b. potassium-40 ________________c. potassium-41 ________________

7. Write an equation showing the relationship between an atom’s atomic number and its mass number.

Isotopes WS1. What do the following symbols represent?

a. e- ___electron_____b. n0 ___neutron_______c. p+ ___proton_______

2. Which subatomic particles are found in an atom’s nucleus? protons, neutrons

3. Which subatomic particle identifies an atom as that of a particular element? protons

4. Explain why atoms are neutral even though they contain charged particles. # protons = # electrons

5. What do the numbers, 39, 40, and 41 after the element name potassium refer to? mass number

6. Write the symbolic notation for each of the following isotopes:a. potassium-39 _____39

19K______b. potassium-40 _____40

19K______c. potassium-41 _____41

19K______

7. Write an equation showing the relationship between an atom’s atomic number and its mass number. Mass # = atomic # + # of neutrons

Warm UpCompare 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 PracticeMarbles

• 3 Marbles– Mass = 1.59 g, 1.51 g, 1.76 g– Average Mass = (1.59 g + 1.51 g + 1.76 g)

3

= 1.62 g

Average Atomic Mass PracticeMarbles

The same calculation can be done as

4.86 g x 0.3333… = 1.62 g.

Average Atomic Mass PracticeMarbles

• 100 Marbles– 79 are 1.59 g, 10 are 1.51 g, and 11 are 1.76 g– 79 x 1.59 g =+126 g

10 x 1.51 g =+ 15.1 g

11 x 1.76 g = 19.4 g

160.5 g divided by 100

= 1.60 g

Average Atomic Mass PracticeMarbles

or, using relative abundances:

0.79 x 1.59 g =+1.26 g

0.10 x 1.51 g =+0.151 g

0.11 x 1.76 g = 0.194 g

1.60 g

Average Atomic Mass PracticeAtoms

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 PracticeAtoms

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 V-51 (50.94 amu, 99.750%).

Use your periodic table to check your answer.

Atomic #, Mass # and Average Atomic Mass

Atomic # Mass # Average Atomic Mass

# p+ #p+ + #no Weighted average of all

isotopeswhole # whole # Decimal, limited

by sfs

Found on PT NOT on PT Found on PT

Warmup (from Atomic Theory WS #1)

2. Give the number of protons, electrons, and neutrons in each of the following atoms.

a. 10847Ag b. 40

20Ca 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. Give the number of protons, electrons, and neutrons in each of the following atoms.

a. 10847Ag 47 p+, 47 e-, 61 n0

b. 4020Ca 20 p+, 20 e-, 20 n0

c. 2311Na 11 p+, 11 e-, 12 n0

3. Name each isotope, and write it in symbolic notation.a. Atomic number 26; mass number 56 iron-56, 56

26Fe

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. 7533As3- b. 180

74W6+

14. 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. 7533As3- 33 p+, 36 e-, 42 n0 b. 180

74W6+ 74 p+, 68 e-, 106 n0

14. 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

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 (1852-

1908) - 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, 4

2He, a) with 2+ charge

e.g. 22688Ra 222

86Rn + 42He (+ 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

10n 1

1p + 0-1b

e.g. 146C 14

7N + 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 (0

0g)

usually accompany alpha and beta radiation

e.g. 23892U 234

90Th + 42He + 2 0

0g

Types of RadiationNuclei with lower 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 neutronPositron = particle with same mass as an e-, but opposite charge

Proton neutron + positron 1

1p 10n + 0

1b

e.g. 116C 11

5B + 01b

Types of RadiationNuclei with lower 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

11p + 0

-1e 10n

e.g. 0-1e + 81

37Rb 8136Kr + X-ray photon

Radioactive Particles WS

Positron same mass as e-’s 0+1b 1/1840 1+

Electron capture electrons 0-1e-1 1/1840 1-

(Added to the reactants side)

Radioactive Particles WS1. 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 WS1. 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:

23592U + 1

0n 23692U 92

36Kr + 14156Ba + 3 1

0n + 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 11H 4

2He + energy

Mass is not conserved.

Mass is converted into energy via E = mc2

Creation of Natural Elements

Other fusion reactions occur in the sun:

42He + 4

2He 84Be + g (gamma ray)

42He + 8

4Be 126C + g

2. Synthetic Elements

a. Nuclear bullets

i. Bombard nuclei of elements with small particles such as p+, n ,

42He (a particles) & e- (0

-1b particles)

ii. Elements # 93-100

2. Synthetic Elements

iii. 1919 first experiment:

147N + 4

2He 178O + 1

1H

(Rutherford)

2. Synthetic Elementsb. Crashing nuclei

i. Accelerators hurl nuclei into each other at very high speeds.e.g. 12

6C + 24496Cm 254

102No + 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

24494Pu + 48

20Ca 289114Fl + 3 1

0n (1999, Russia)

Nuclear equations

Complete the following equations :

21483Bi 4

2He + _____

23993Np 239

94Pu + ______

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.

2210Ne + 244

95Am 263105Db + 3 1

0n

2210Ne + 244

95Am

2210Ne + 244

95Am

5. Each of the radioisotopes in the table decays within 20 seconds to 10 hours. Write a nuclear equation for each decay.

26695Am 4

2He +263

105Db 0-1b +

262103Lr + 0

-1e- 266

106Sg 42He +

6. 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; 266

105Db

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.

2210Ne + 244

95Am 263105Db + 3 1

0n

2210Ne + 244

95Am 262103Lr + 4

2He

2210Ne + 244

95Am 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.

24495Am 4

2He + 24093Np

263105Db 0

-1b + 263106Sg

262103Lr + 0

-1e- 262102No

266106Sg 4

2He + 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 :

21483Bi 4

2He + _21081Tl_

23993Np 239

94Pu + __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.

24195Am 4

2He + 23793Np

Write a balanced nuclear equation for the beta decay of bromine-84.

8435Br 0

-1b + 8436Kr

Nuclear equations

Complete the following equations:

21483Bi 4

2He +

23993Np 239

94Pu +

24195 Am 4

2He +

8435Br _______ + 0

-1b

Nuclear equations (KEY)

Complete the following equations:

21483Bi 4

2He + 21081Tl

23993Np 239

94Pu + 0-1b

24195 Am 4

2He + 23793Np

8435Br 84

36Kr + 0-1b

Warmup – Nuclear Equations

1. 94Be + 1

1H 42He + _______

2. 23892U + 4

2He 2 10n + _______

3. 156C + 1

0n _______

4. 13755Cs ________ + 0

-1b

Warmup – Nuclear Equations

1. 94Be + 1

1H 42He + 6

3Li

2. 23892U + 4

2He 2 10n + 240

94Pu

3. 156C + 1

0n 166C

4. 13755Cs 137

56Ba + 0-1b

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

– Examples include: microwaves, TV, Radio, X-rays

←Memorize

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 nNote 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 Spectrum1. 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 Spectrum1. 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 Spectrum6. 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 = lx n

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)

EMR Practice Problemsc = lx n

1. 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? (2.63 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)

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 ProblemsE = 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 ProblemsE = 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.

EMR Practice Problems1. 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.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.

8. 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?

EMR Practice Problems1. 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)

EMR Practice Problems5. 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) An X-ray source, emitting X-rays with a wavelength of 175.4 pm. (1.71 x 1018 s-1)

8. 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?(Note: v is not required for the calculation.)(2.67 x 102s)

Warmup - EMR1. 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 of EMR with a wavelength of 5.26 mm? What type of EMR is it? (A: 5.70 x 1010 s-1)

5. Determine the energy, in joules, of a photon whose frequency is 3.55 x 1017 Hz. (h = 6.626 x 10-34 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 ModelElectrons…

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 Model1 = lowest energy4 = highest energy

As e- goes furtherAs e- goes further away from the nucleus, it increases in potential energy

E3 - E1 = a quantum of energy in the form of EMR++ 1 2 3 4

Bohr ModelElectrons…

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 levels.

Bohr’s model did not explain the line spectra of atoms with >1 electron.

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

ground state

excited state

ENERGY IN PHOTON OUT

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

VisibleBalmer Series

Infrared Paschen Series

UV Lyman Series

7 6 5 4 3 2 1

E2

E1

The Hydrogen Spectrum

Atomic Theory Video Questions #3The Bohr Atom

1. What does a change in color indicate?

2. What did Max Planck propose about how matter emits energy?

3. Which has more energy: an electron near the nucleus or an electron far from the nucleus?Why?

4. What element did Bohr use in his experiments?

5. How much energy must be carried by a free electron in order to knock another electron out of orbit?

6. What happens if an electron of more than 13.6 eV hits another electron?

7. What did Bohr propose happens when an electron goes down an energy level?

Atomic Theory Video Questions #3The Bohr Atom KEY

1. What does a change in color indicate? Change in frequency

2. What did Max Planck propose about how matter emits energy? In packets called quanta

3. Which has more energy: an electron near the nucleus or an electron far from the nucleus?Why? Far – more potential energy

4. What element did Bohr use in his experiments? hydrogen

5. How much energy must be carried by a free electron in order to knock another electron out of orbit? > 10.2 eV

6. What happens if an electron of more than 13.6 eV hits another electron? Liberates it from the nucleus ionized

7. What did Bohr propose happens when an electron goes down an energy level? Emits photon equivalent to amount of energy

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 dida. 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? _______________

Atomic Theory Video Questions #3Spectra KEY

1. What is the orbit of an electron closest to the nucleus called? ground state

2. What region of the spectrum of excited hydrogen gas dida. Balmer predict? visibleb. Paschen predict? infraredc. Lyman predict? UV

3. 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? 2b. the Paschen series? 3c. 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

s

pd f

Atomic Orbitals: predict 90% probability of location of electrons (electron cloud)

Each orbital can contain a maximum of 2 electrons, spinning in opposite directions.

Orbitals

Orbitals WS1. 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 WS1. 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 EnergyLevel (n)

(or “Shell”)distance from

nucleus

Sublevel (l)(or “Subshell”)

shape of probability

cloud

3-D (m)(position in

space)x, y, z

Spin (s)+ ½- ½

Total # e-

2n2

1 s only 1 orientation 1 e- + ½1 e- - ½

2(1)2 = 2

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 2(4)2 = 32

5 s, p, d, f

6 s, p, d, (f)

7 s, p, d, (f)

s: holds 2 e- x 1 orbital = 2 e- totalp: holds 2 e- x 3 orbitals = 6 e- totald: holds 2 e- x 5 orbitals = 10 e- totalf: holds 2 e- x 7 orbitals = 14 e- total

• n = # of sublevels per level• n2 = # of orbitals per level• Sublevel sets: 1 s, 3 p, 5 d, 7 f

px py pz

1. The p orbitals

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 WS1. 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 WSKEY

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? 2p

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%

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 Diagramss ps p ds p d fs p d fs p d fs p d f

• Oxygen (8 e-)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 electron configurations:

2. Write electron configurations (only) for:a) helium b) carbona) neon b) aluminuma) argon b) brominea) krypton b) palladiuma) xenon b) leada) radon b) uranium

c) Write all b)’s in noble gas (condensed) configurationd) Show how to jump into an Aufbau diagrame) s,p,d,f blocks in the periodic table

Atomic Theory Video Questions # 4Electron Arrangement

1. What 2 parts of Bohr’s model did Schrödinger and Heisenberg keep?

2. How did they change Bohr’s model?

3. 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 # 4Electron 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

3. 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 p orbitals? 3 d orbitals? 5

6. How many electrons can be in any one orbital? 2, of opposite spin

7. Which orbitals are filled with electrons first? Lowest energy

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? 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

© 1998 by Harcourt Brace & Company

s p

d (n-1)

f (n-2)

1234567

67

Atomic Theory and the Periodic Table

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

s-block

1st Period

1s11st column of s-block

1

2

3

4 5

6

7

Periodic Patterns

• Example - Hydrogen

1

2

3

4

5

6

7

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.

[Ar]

1

2

3

4 5

6

7

4s2 3d10 4p2

Periodic Patterns

• Example - Germanium

1

2

3

4 5

6

7

Stability

• Ion Formation– Atoms gain or lose electrons to become more stable.– Isoelectronic with the Noble Gases.