Models of the Atom

Dalton’s model (1803)

Thomson’s plum-pudding model (1897)

Rutherford’s model (1909)

Bohr’s model (1913)

Charge-cloud model (present)

Greek model(400 B.C.)

1800 1805 ..................... 1895 1900 1905 1910 1915 1920 1925 1930 1935 1940 1945

1803 John Dalton pictures atoms astiny, indestructible particles, with no internal structure.

1897 J.J. Thomson, a Britishscientist, discovers the electron,leading to his "plum-pudding" model. He pictures electronsembedded in a sphere ofpositive electric charge.

1904 Hantaro Nagaoka, aJapanese physicist, suggests that an atom has a centralnucleus. Electrons move in orbits like the rings around Saturn.

1911 New Zealander Ernest Rutherford statesthat an atom has a dense,positively charged nucleus. Electrons move randomly in the space around the nucleus.

1913 In Niels Bohr'smodel, the electrons move in spherical orbits at fixed distances from the nucleus.

1924 Frenchman Louis de Broglie proposes thatmoving particles like electronshave some properties of waves. Within a few years evidence is collected to support his idea.

1926 Erwin Schrödinger develops mathematicalequations to describe the motion of electrons in atoms. His work leads to the electron cloud model.

1932 James Chadwick, a British physicist, confirms the existence of neutrons, which have no charge. Atomic nuclei contain neutrons and positively charged protons.

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Foundations of Atomic Theory

Law of Definite Proportions

The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.

Law of Multiple Proportions

If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first elements is always a ratio of small whole numbers.

Law of Conservation of Mass

Mass is neither destroyed nor created during ordinary chemical reactions.

Conservation of Atoms

John Dalton

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 204

2 H2 + O2 2 H2O

4 atoms hydrogen2 atoms oxygen

4 atoms hydrogen2 atoms oxygen

H

H

O

O

O

O

H

H

H

H

H

H

H2

H2

O2

H2O

H2O

+

Legos are Similar to Atoms

Lego's can be taken apart and built into many different things.

H

H

O

O

O

O

H

H

H

H

H

HH2

H2

O2

H2O

H2O

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Atoms can be rearranged into different substances.

45 g H2O? g H2O

Conservation of Mass

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 204

Highvoltage

Before reaction

electrodes

glasschamber

5.0 g H2

80 g O2

300 g (mass of chamber)+385 g total

H2O2

Highvoltage

After reaction

0 g H2

40 g O2

300 g (mass of chamber)+385 g total

O2

H2O

Daltons’ Models of Atoms

Carbon dioxide, CO2

Water, H2O

Methane, CH4

Radioactivity (1896)1. rays or particles produced by

unstable nuclei

a. Alpha Rays – helium nucleus

b. Beta Part. – high speed electron

c. Gamma ray – high energy x-ray

2. Discovered by Becquerel –

exposed photographic film

3. Further work by CuriesAntoine-Henri Becquerel

(1852 - 1908)

Background Information

Cathode Rays• Form when high voltage is applied across

electrodes in a partially evacuated tube.• Originate at the cathode (negative electrode)

and move to the anode (positive electrode)• Carry energy and can do work• Travel in straight lines in the absence of an

external field

Cathode Ray Experiment

1897 Experimentation

• Using a cathode ray tube, Thomson was able to deflect cathode rays with an electrical field.

• The rays bent towards the positive pole, indicating that they are negatively charged.

J.J. Thomson

• He proved that atoms of any element can be made to emit tiny negative particles.

• From this he concluded that ALL atoms must contain these negative particles.

• He knew that atoms did not have a net negative charge and so there must be balancing the negative charge.

J.J. Thomson

William Thomson (Lord Kelvin)

• In 1910 proposed the Plum Pudding model– Negative electrons

were embedded into a positively charged spherical cloud.

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56

Spherical cloud ofPositive charge

Electrons

Thomson Model of the Atom

• J.J. Thomson discovered the electron and knew that electrons could be emitted from matter (1897).

• William Thomson proposed that atoms consist of small, negative electrons embedded in a massive, positive sphere.

• The electrons were like currants in a plum pudding.

• This is called the ‘plum pudding’ model of the atom.

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Ernest Rutherford (1871-1937)

• Learned physics in J.J. Thomson’ lab.

• Noticed that ‘alpha’ particles were sometime deflected by something in the air.

• Gold-foil experiment

Rutherford

PAPER

Rutherford

PAPER

Animation by Raymond Chang – All rights reserved.

Rutherford ‘Scattering’

• In 1909 Rutherford undertook a series of experiments• He fired (alpha) particles at a very thin sample of gold foil• According to the Thomson model the particles would only

be slightly deflected• Rutherford discovered that they were deflected through large

angles and could even be reflected straight back to the source

particlesource

Lead collimator Gold foil

Rutherford’s Apparatus

beam of alpha particles

radioactive substance

gold foil

circular ZnS - coated

fluorescent screen

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120

Rutherford received the 1908 Nobel Prize in Chemistry for his pioneering work in nuclear chemistry.

Rutherford’s Apparatus

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120

beam of alpha particles

radioactive substance

fluorescent screencircular - ZnS coated

gold foil

What he expected…

What he got…richochetingalpha particles

The Predicted Result:

expected path

expected marks on screen

mark onscreen

likely alphaparticle path

Observed Result:

Interpreting the Observed Deflections

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120

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gold foil

deflected particle

undeflected particles

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Rutherford Scattering (cont.)

Rutherford interpreted this result by suggesting that the particles interacted with very small and heavy particles

Particle bounces off of atom?

Particle attracts to atom?

Particle goes through atom?

Particle path is alteredas it passes through atom?

.

Case A

Case B

Case C

Case D

Table: hypothetical description of alpha particles

alpha rays don’t diffract

alpha rays deflect towards a negatively charged plate and away from a positively charged plate

alpha rays are deflected only slightly by an electric field; a cathode ray passing through the same field is deflected strongly

... alpha radiation is a stream of particles

... alpha particles have a positive charge

... alpha particles either have much lower charge or much greater mass than electrons

observation hypothesis

(based on properties of alpha radiation)

Copyright © 1997-2005 by Fred Senese

Explanation of Alpha-Scattering Results

Plum-pudding atom

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Alpha particles

Nuclear atom

Nucleus

Thomson’s model Rutherford’s model

Interpreting the Observed Deflections

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120

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gold foil

deflected particle

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Rutherford’sGold-Leaf Experiment

Conclusions:

Atom is mostly empty space

Nucleus has (+) charge

Electrons float around nucleus

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120

Bohr’s Model

Nucleus

Electron

Orbit

Energy Levels

Bohr Model of Atom

The Bohr model of the atom, like many ideas in the history of science, was at first prompted by and later partially disproved by experimentation.

http://en.wikipedia.org/wiki/Category:Chemistry

Increasing energyof orbits

n = 1

n = 2

n = 3

A photon is emittedwith energy E = hf

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An unsatisfactory model for the hydrogen atom

According to classical physics, lightshould be emitted as the electron circles the nucleus. A loss of energywould cause the electron to be drawncloser to the nucleus and eventuallyspiral into it.

Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 294

• Bohr’s contributions to the understanding of atomic structure:

1. Electrons can occupy only certain regions of space,

called orbits.

2. Orbits closer to the nucleus are more stable —

they are at lower energy levels.

3. Electrons can move from one orbit to another by absorbing or emitting energy, giving rise to characteristic spectra.

• Bohr’s model could not explain

the spectra of atoms heavier

than hydrogen.

Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.