Unit 3 – History of Atomic Theory

Democritus of Abdera (c. 460 – c. 370 BCE) A greek philosopher that believed all matter was composed of discrete, indivisible

particles called (atomos). The more accomplished philosophers (Aristotle and στομωσPlato) did not accept this theory. They believed all matter was made of four elements: earth, air, water and fire.

Antoine Lavoisier (1743-1794)

Following extensive studying of combustion and the discovery of carbon, dioxide, oxygen, nitrogen and hydrogen gases, Lavoisier regarded measurement as the essential operation of chemistry. Through careful experiments, he weighed the reactants and products of various reactions and suggested the LAW OF CONSERVATION OF MASS

Joseph Proust (1754-1826)

Proust furthered Lavoisier’s experiments by weighing the ratios of reactants in specific chemical compounds. He found that for specific compounds the ratio of their combining was constant. This led to the LAW OF DEFINITE PROPORTION.

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John Dalton (1766-1826)

Dalton developed the first atomic model of the atom that offered explanations for the simple laws that govern chemistry. solid sphere model His theory has three main postulates.

1. An element is composed of tiny particles called atoms. All atoms of a given element show the same chemical properties. *This is not true today because of the discovery of neutrons and isotopes.

2. Atoms of different elements have different properties. In an ordinary chemical reaction, no atom of any element disappears or is changed into an atom of another element.

3. Compounds are formed when atoms of two or more elements combine. In a given compound, the relative numbers of atoms of each kind are definite and constant. In general, these relative numbers can be expressed as integers or simple fractions. This led to the LAW OF CONSTANT COMPOSITION.

Henri Becquerel (1788-1878)

Becqueral accidentally found that a piece of a mineral containing uranium could produce an image on a piece of photographic paper in the absence of light. He attributed this phenomenon to a spontaneous emission of radiation by the uranium, which he called radioactivity. Studies demonstrated three types of radioactive emission: gamma ( ) rays, γbeta ( ) particles and alpha particles ( ) particles. A gamma ray is essentially high energy β α“light” (no mass and no charge); a beta particle is a high speed electron (essentially no mass and a negative charge); and an alpha particle is essentially a helium nucleus (heavy mass = 4 and 2+ charge). More models of radioactivity have been discovered but the alpha particles were used in early crucial experiments.

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J J Thomson (1856 – 1940)

Thomson did the first important experiments that led to the understanding of the composition of the atom. He studied electrical discharges in partially evacuated tubes called cathode ray tubes.

Thomson found that when high voltage was applied to the tube a “ray” (he called it the cathode ray) was produced. Due to the fact that the ray emanated from the negative pole of the electric field the ray was composed of negatively charged particles. This led to the discovery of electrons. Also from this discovery he calculated the charge-to-mass ratio of the electron. He reasoned that all atoms contain electrons and that to be electrically neutral, atoms must also contain positive charges. Thomson postulated that an atom contained a diffuse cloud of positive charge with negative electrons embedded in it. This is often called the plum pudding or raisin bun model of the atom because the electrons are like raisins dispersed in the dough (the positive charge).

plum pudding model

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spherical cloud of positive charge

electrons

Max Planck (1858-1947)

Planck studied the radiation profiles emitted by solid bodies that were heated to incandescence. From these experiments, he determined that the results could not be explained by the typical (Newtonian) physics of the time. The loss or gain of energy could only be obtained by whole number multiples of the quantity, h , where h is Planck’s νconstant (6.626 x 10-34 J·s/particle) and (nu) is the frequency of the electromagnetic νradiation. A small “packet” of energy was called a quantum (plural, quanta).

Robert Millikan (1868-1953)

Millikan performed experiments with charged oil drops. These experiments allowed him to determine the magnitude of the electron charge. Using this value and the charge-to-mass ratio determined by Thomson, Millikan was able to determine the mass of an

electron. Oil drop experiment

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Ernest Rutherford (1871-1937)

Rutherford, who also performed many experiments exploring radioactivity, carried out an experiment to test Thomson’s model of the atom. The experiment directed alpha particles at a thin sheet of gold foil. He reasoned that if Thomson’s model were accurate, the massive alpha particles would crash through the foil with little deflection. Although most of the particles passed straight through, many of the particles were deflected at large angles and a very few were reflected back at him. This led to three conclusions:

1. The atom is mostly empty space. Most alpha particles went straight through.

2. The atom has a positively charged center or nucleus. Due to the fact that some particles came close to the center and the deflection path was more similar to an electric field of like charges.

3. The nucleus is the concentration of the atom’s mass. The very few that were deflected were on a collision course with a much more massive center of charge.

Gold Foil Experiment

If Thomson was correct with his model, then a completely different outcome would have occurred. This Gold-Foil Experiment created the nuclear model of the atom, with the positively charged nucleus at the center.

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James Chadwick (1891-1974)

Capitalizing on the theories proposed by Rutherford and others, Chadwick produced experimental evidence of the existence of the atomic particle that had a mass nearly equal to that of a proton but no charge. This was a result of alpha particle bombardments of beryllium, lithium and boron. He named this particle the neutron according to the equation: 9Be + → α 12C + 1n

Louis de Broglie (1892-1987)

Utilizing Albert Einstein’s special theory of relativity (energy has mass) and the incorporations of Planck’s equation, de Broglie supported the view that light could not only be quantized but that it also could possess wave properties as all matter can. This is called the wave-particle duality of light. It depends upon the nature of the quantity you are investigating. Light, which was previously thought to be purely wavelike, was found to have particle (masslike properties) and the opposite is also true. The relationship is called de Broglie’s equation = h/mv, where is the wavelength (in meters), h is Planck’s λ λconstant, m is the mass of the particle (in kilograms), and v is the velocity of the particle (in meters/second).

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Niels Bohr (1885-1962)

As Rutherford’s graduate assistant, Bohr held the planetary model of the atom and incorporated the work of Planck and de Broglie to propose that the electrons must be balanced by the attraction for the nucleus to resist flying off the atom. However a constantly accelerating particle (like the electron) should lose energy and then eventually fall into the nucleus. Due to this apparent contradiction, Bohr applied a theory to the hydrogen atom that claimed that energy could only occur in specified increments. These increments were called energy levels and could also have only whole number values. Electrons in a normal state (ground state) could gain energy and be promoted to a higher level (excited state) but eventually that energy could not be sustained and would be emitted as electromagnetic radiation (visible light in some cases). The specific lines of radiation emitted were consistent with observed emission spectra. This lead to the quantized planetary model of the (hydrogen) atom and Bohr’s equation: E = -B/n2, where E is energy of radiation, B is Bohr’s constant, and N is the energy level. Also energy changes can be calculated for the hydrogen atom: E = EΔ hi - Elo = -1312 kJ/mol (1/nhi

2 - 1/nlo2)

where “hi” represents the excited level and “lo” represents the ground state. Bohr’s findings were in agreement with 0.01% of observed results. However, when investigating helium or higher “multi-electron” atoms the percent error grew exponentially.

By the mid – 1920’s it became obvious that Bohr’s model could not be made to work for these higher atoms. Three physicists, Werner Heisenberg, Louis de Broglie and Erwin

Schrödinger (1887 - 1961) proposed a new branch of physics to help explain the contradictions. This new branch of physics is called quantum mechanics. Schrödinger gave greater emphasis to the wave properties of the electron and the probability of finding an electron in a certain volume of the atom. This led to Schrödinger’s equation that derived certain values for describing the likelihood of locating an electron at certain radii from the nucleus and describing the electron’s properties in each of these volumes in 3-D space.

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Erwin SchrödingerSchrodinger’s equation for the x-axis:

d2Ψd x2 + 8π 2

h2 (E-V) = 0Ψ

Solving for all three axes leads to three quantum numbers to explain the location and properties in multi-electron atoms.

Werner Heisenberg (1901-1976)

Heisenberg stated that due to the physical methods of detection of the electron and small particles it is impossible to be certain of both its location and momentum according to the equation: x Δ · (mv) Δ ≥ h/4 . This is Heisenberg’s Uncertainty Principle. π

Current model of the atom shows the “electron cloud” as a probability distribution.

Noyau = nucleusProbabilité de présence de l’électron = probability of the presence of the electron

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scientist discovery model name model picture

John Dalton

JJ Thomson

Ernest Rutherford

Niels Bohr

Current model

(deBroglie, Heisenberg,

Schrödinger)

Development of the Atomic Models:

Development of Atomic Theory Studyguide:

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Democritus: “____________” (indivisible particles) “thought” about it (philosopher – no proof)

Dalton: Atomic theory (had experimental evidence)

1.

2.

3.

___________ model

Thomson: Cathode ray tube Discovered the __________ _________________ model

Millikan: _____________ experiment What two things did he determine?

Rutherford: _____________ experiment What did he think would happen?

What really happened?

Discovered the __________

Quantum Theory – extremely small particles traveling at velocities close to the speed of light

Planck: loss or gain of energy in small "packets" called quanta

Bohr: Quantized energy levels Definite orbits ____________ model

deBroglie: wave-particle duality

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duality of nature light can act like a particle (photon)/electron can act like energy (wave) λ = h/mv (matter has wavelength)

Einstein: photoelectric effect called light particles photons

Current Model: probability distribution

Heisenberg: Heisenberg uncertainty principle:

Schrödinger Wave Equation:

Date of Atomic Theory Quiz: _______________________

Warm-up:

Write the name of each model underneath the picture that represents it:

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solid sphere (Dalton), plum pudding (Thomson), nuclear (Rutherford), planetary (Bohr), quantum mechanical (current)

Atomic Theory Matching:

1. Millikan

2. Thomson

3. Dalton

4. Rutherford

5. Bohr

6. Schrödinger

7. Heisenberg

8. deBroglie

a. atomic theory

b. discovered electron

c. uncertainty principle

d. wave equation

e. oil drop experiment

f. gold foil experiment

g. solid sphere model

h. planetary model

i. wave-particle duality

j. plum pudding model

k. discovered the nucleus

l. quantized energy levels

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ee e

e

e e

ee

e

The Atom – Modern Atomic Theory

Subatomic Particles:

symbol location Relative charge

Relative mass (amu)

Actual mass (g)

Proton

Neutron

Electron

In a neutral atom, the number of ________ is always the same as the number of _________.

Why does the atom not fall apart?

How does the mass of the electron compare to the mass of the proton?

Heisenberg Uncertainty Principle:

The nucleus makes up most of the __________ of an atom.

The electron cloud makes up most of the ___________ of an atom.

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Atomic Numbers:

Atomic Number – number of ________ in an atom – determines the atom’s identity (located above element on periodic table.)

______________ found that each element has a different number of protons in its nucleus.

T or F: The atomic number of an element can change. Why or why not?

An individual atom is electrically neutral, which means the number of protons equals the number of electrons.

So, an element’s atomic number also indicates the number of electrons in a neutral atom.

Ions:

When an atom gains or loses electrons, it acquires a net electrical charge and is called an ion.

Anion:

Cation:

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Examples of Ions:

1. A magnesium ion loses 2 electrons, what is the charge on a magnesium ion?

2. A fluorine atom gains 1 electron, what is the charge on the fluorine ion?

3. An iron atom loses 3 electrons, what is the charge of the iron atom?

Isotopes:

All atoms of a given element have to have the same number of __________ in their nucleus, but not necessarily the same number of ___________.

One isotope of each element is more common than another.

In nature, elements are found as a mixture of isotopes.

For example, in a drop of water, 99.9844 % of the hydrogen isotopes are hydrogen – 1.

isotope name # of protons # of neutrons # of electrons* hydrogen

(hydrogen - 1)1 0 1

deuterium(hydrogen – 2)

1 1 1

tritium (hydrogen – 3)

1 2 1

* indicates the most common isotope of H (look at its average atomic mass)

The major difference between isotopes of the same element is their mass. The more neutrons an element has, the higher the mass of that element is.

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Mass Number – sum of the _________ and ________ in an atom

Atomic Mass – average mass of all the ________ of a particular element

Isotope – atom of an element that has a different number of ________ in the nucleus

Isotopes are named by their mass numbers.

For example: All chlorine nuclei have 17 protons in their nuclei, but some have

18 neutrons Mass # = ____

and some have

20 neutrons Mass # = ____

What is the most common isotope of chlorine?

Examples:

1. The chlorine isotope with 17 protons and 18 neutrons is called:

2. The chlorine isotope with 17 protons and 20 neutrons is called:

3. The symbol for chlorine – 35 is:

4. The symbol for chlorine – 37 is:

* mass # on top and atomic # on bottom

Because Atomic # = # of protons

and Mass # = # of protons + # of neutrons

How do you find the # of neutrons in an isotope?

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WS 3.2: Subatomic Particles

Complete the following table.

Element SymbolAtomic Number

Mass Number

Protons Neutrons Electrons

1. fluorine ______ ______ 19 ______ ______ ______

2. iron ______ ______ ______ ______ 30 ______

3. ________ ______ ______ 27 ______ ______ 13

4. ________ K ______ ______ 19 20 ______

5. ________ Br ______ 80 ______ ______ ______

6. ________ ______ ______ 197 ______ ______ 79

7. ________ ______ 47 ______ ______ 61 ______

8. ________ ______ ______ ______ 14 15 ______

9. ________ ______ ______ 47 ______ ______ 22

10. ________ ______ ______ 55 25 ______ ______

11. ________ ______ 1 2 ______ ______ ______

12. ________ ______ ______ 50 24 ______ ______

13. ________ ______ ______ 118 ______ 68 ______

14. ________ ______ ______ 58 ______ ______ 28

15. ________ ______ ______ ______ 53 74 ______

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3.3 Atomic Particles Worksheet

1. Complete the following table:symbol number of

protons in nucleus

number of neutrons in

nucleus

number of electrons

net charge

33 42 3+

128Te2- 54

16 16 16

81 123 1+

195Pt

2. Give the protons and neutrons in the nucleus of each of the following neutral elements:a. plutonium-238 g. lead-207b. copper-65 h. europium-151 (Eu)c. nitrogen-15 i. cobalt-60d. hydrogen-3 (also called tritium) j. chromium-54e. chromium -52 k. silver-107f. helium-4 l. silver-109

3. Complete the following table:symbol atomic number mass number number of

neutrons in nucleus

number of electrons

238U

20 40 18

23 28 20

89Y3+

Br1- 44

15 31 18

8 6

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WS 3.4: IsotopesPart I: Complete the following table.

(You need to complete the nuclear symbol as well.)

Nuclear Symbol Hyphen notation Protons Neutrons Electrons

1. 40 Ca calcium - 40 _______ _______ _______

2. ___________ ___________ 20 22 _______

3. 43 Ca ___________ _______ _______ _______

4. 44 Ca ___________ _______ _______ _______

5. ___________ ___________ _______ 26 20

6. 16 O ___________ _______ _______ _______

7. ___________ oxygen - 18 _______ _______ _______

8. 108Ag ___________ _______ _______ _______

9. ___________ ___________ 35 45 _______

10. ___________ ___________ _______ 125 82

11. 27Al ___________ _______ _______ _______

12. ___________ ___________ _______ 16 16

13. 56 Fe ___________ _______ _______ _______

14. ___________ selenium - 78 _______ _______ _______

15. ___________ iron - 55 _______ _______ _______

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Part II: The isotopes of sulfur are sulfur-32, sulfur-33, sulfur-34, and sulfur-36. Write the nuclear symbol of each.

Part III: Answer each question. Show your work.

1. What is the atomic mass of silicon if 92.21% of its atoms have mass 27.977 amu, 4.70% have mass 28.976 amu, and 3.09% have mass 29.974 amu?

2. Determine the average atomic mass for magnesium using the following information.

24 Mg 78.99 % 23.985 amu

25 Mg 10.00 % 24.986 amu

26 Mg 11.01 % 25.983 amu

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3. Neon has two isotopes. Neon-20 has a mass of 19.992 amu and neon-22 has a mass of 21.991 amu. In an average sample of 100 neon atoms, 90.0% of them are neon-20 and 10.0% of them are neon-22. Calculate the average atomic mass of neon.

4. What is the atomic mass of hafnium if, out of every 100 atoms, 5 atoms have mass 176.0 amu, 19 atoms have mass 177.0 amu, 27 have mass 178.0 amu, 14 have mass 179.0 amu, and 35 have mass 180.0 amu?

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Radioactivity

As atomic number increases (protons increase), stable nuclei will have more neutrons than protons. For the larger elements, a point is reached beyond which there are no stable nuclei. Nuclei with more than 83 protons (after bismuth) are all unstable and will eventually break up into smaller pieces.

Definition: Radioactvity =

There are three types of radioactive decay:1. Alpha (α) nucleus emits an ____________ particle

An alpha particle is a ___________ nucleus (two protons and two neutrons)

2. Beta (β)

a. Beta-minus decay: a neutron turns into a proton by emitting an _________ and an antineutrino (changes to the next element on the periodic table)

b. Beta-plus decay: a proton turns into a neutron by emitting an __________ and a neutrino (changes to the previous element on the periodic table)

3. Gamma (γ) = nucleus changes from a higher energy state to a lower energy state

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The EMR emitted is a _____________.

Half-life: (t1/2) is the amount of time required for the amount of something to fall to half its initial value (radioactive half-life)

1. What is the half-life of the material in the graphic below?

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2. Can you predict how much material will be left after 6 half-lives?

Unit 3 TEST DATE: _____________________________

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