Atomic Theory

ATOMIC THEORYYou too can be as smart as

Einstein (almost)

History - Greeks

The elementsEarth – dry, heavyWater – wet, heavyAir – cool, lightFire – warm, light

The composition of a substance could be estimated from its properties.

History - Greeks

These ideas were based on observation, logic and reason, but not experimentation.

 Democritus (460 B.C. - 370 B.C.)

History - Greeks

Matter is made of small, hard indivisible particles called atoms, which exist in the void.

These atoms differ in size and shape, but not in any other way.Quantitative differences (how much) vs. Qualitative differences

(what kind)

History - French

Antoine Lavoisier (1743-1794)

Discoverer of Oxygen (disputed)

His work refuted the phlogiston theory

Responsible for the law of conservation of matter.

History – French 1800

Claude Louis Berthollet Joseph Louis Proust

History - French

Berthollet – “compounds do not have a fixed composition”.

 Cu + S CuxSy

  Every time he tried the experiment

he got a different result.

History - French

Proust - compounds have a fixed composition.

2H2 + O2 2H2O 

He always got the same result. Proust’s argument is called The Law

of Definite Proportions. He was proved to be right.

History - English

John Dalton (ca. 1804)

The father of modern atomic theory

Schoolteacher Colorblind –

studied colorblindness

Dalton’s Atomic Theory

The points of Dalton’s theory All matter is made of atoms Atoms are indivisible and indestructible All atoms of one element are exactly

alike, and atoms of different elements are different.

Atoms combine in small whole number ratios to form compounds.

Dalton’s Atomic Theory

The Law of Multiple Proportions:If two elements combine to make two

different compounds, the ratios of the elements involved are small

whole numbers.  Examples: CO and CO2

CuS and Cu2S

H2O and H2O2

Inside the Atom

J. J. Thomson and the Electron (1897)

The Electron

Thomson discovered the electron - he called it a “corpuscle”.

He used an instrument called a Crookes tube.

Cathode (-) Evacuated tube Anode (+)

The Electron

He noticed a stream of charged particles coming from the cathode, called cathode rays.

Thomson proposed the "plum pudding" atomic model - negatively charged corpuscles swarm inside a cloud of massless positive charge.

Ernest Rutherford and the Nucleus

The Nucleus

The gold foil experiment (1909)

Gold Foil Experiment

Most of the alpha particles went straight through, and a few were bounced straight back.

 Rutherford’s interpretation: The atom has a small, hard, dense and positively charged nucleus. The electrons are outside the nucleus.

The Proton and the Neutron

Discovery of the proton: Henry Moseley (1913) Moseley bombarded metals with x-rays Each successive element had one more

positive charge – called “atomic number”

Rutherford proved that the nucleus of nitrogen contains hydrogen nuclei – a “proton” (1918-19)

Discovery of the neutron – James Chadwick (1932)

Parts of the Atom

Name Charge Mass (amu)

Location Discoverer

Electron -1 1/2000 outside nucleus

Thomson

Proton +1 1 nucleus Moseley/Rutherford

Neutron 0 1 nucleus Chadwick

Isotopes

Atomic number = number of protons in the nucleus

Atomic number determines the identity of the element

Mass number = protons + neutrons Number of electrons = number of protons Isotopes: two atoms of the same

element with different numbers of neutrons

C-12 and C-13 are isotopes of carbon

Nomenclature and symbols

Nuclear symbols

13C Write the nuclear symbol for lead-

206.

206Pb

6

82

Periodic table

20

CaCalcium

40.078

Atomic number

Symbol

Name

Average atomic mass

Average atomic mass

Average mass of all the isotopes of an element

Average is weighted Example: Boron has two isotopes, B-10

and B-11

B-10: 19.9%B-11: 80.1%

Average atomic mass of boron:10x0.199 = 1.9911x0.801 = 8.811

Average atomic mass = 1.99 + 8.811 = 10.8amu

Outside the Nucleus

Niels Bohr and the stepwise atom (ca. 1918)

Rutherford-Bohr Model of the Atom (1911-1913)

Rutherford suggested that electrons orbit around the nucleus like planets around the sun.

This did not explain emission spectra, which gave sharp lines.

He theorized that electrons could only travel in certain sized orbits, and not anywhere in between.

Rutherford-Bohr Model of the Atom (1911-1913)

Bohr Model of the Atom

The orbits were called energy levels. Each orbit has a specific energy.

 Electrons can jump from one level to another; as they do, they absorb or emit energy.

Quantum Mechanics

Erwin Schrödinger and probable cause (ca. 1935)

Quantum Mechanics

Schrödinger’s work showed that electrons do not move in actual “orbits”.

Electrons move randomly and form “probability clouds”. The shape of these clouds is similar to the shape of Bohr’s orbits.

The position and momentum of an electron cannot be determined simultaneously (Heisenberg Uncertainty Principle)

Quantum Mechanics

Schrödinger’s “electron cloud”

Electron Energy Level Populations

Bohr suggested that electrons inhabit energy levels around the nucleus.

Each level has a specific energy associated with it.

The outermost (highest energy) level is called the “valence shell”.

The electrons in the valence shell are called the “valence electrons”.

The valence electrons are the most important electrons in the chemistry of the atom.

Electron Energy Level Populations

Electron Energy Level Populations

The number of levels depends on the number of electrons.

The first level (K) holds two electrons. The second level holds eight electrons. The third level holds 18, and the fourth

32. No atom can have more than eight

electrons in its valence shell. When the valence shell reaches eight

electrons, the next two electrons are put in a higher level. Then the lower level can be filled.

Lewis Electron Dot Structures

Lewis dot structures show how many electrons are in the valence shell of an atom.

Lewis dot structure for sodium The first electron always goes to the right

of the symbol. The second is paired with the first.

Lewis Dot Structures

Lewis dot structure of magnesium The third goes on top.

Lewis dot structure of aluminum

Lewis Dot Structures

The fourth goes on the left, and is not paired. The fifth goes on the bottom, and successive electrons are paired until a total of eight is reached.

Lewis dot structure of silicon

Lewis Dot Structures

Lewis dot structure of oxygen

Atomic Spectra

Bohr’s model based on atomic spectra

Obtaining emission atomic spectra Energy is applied to a gas or liquid sample.

Flame test (for samples in solution) Gas discharge tube

The energy makes an electron or two jump to a higher energy level.

The electrons fall back down to a lower level, and give off energy in the form of light – bright lines against a dark background.

Atomic Spectra

Absorption spectra – light is passed through a sample and analyzed – looks like a rainbow with dark lines

Interpreting atomic spectra The light given off is viewed through a

spectroscope. The spectroscope has either a prism or

a grating, which splits the light into its component colors.

Atomic Spectra

Atomic Spectra

Only a few sharp lines appear in the spectrum.

Each line corresponds to a specific electron transition.

Transition = jump from one energy level to another

Light Energy

Light energy travels in the form of waves.

Light Energy

Color depends on frequency. High frequency = violet end of

spectrum Low frequency = red end of spectrum

Energy also depends on frequency, so each color has its own energy. Blue or violet is higher energy than red or green.

When a specific color line is seen in a spectrum, the energy of the electron transition responsible can be calculated.

Electromagnetic spectrum

Light Energy and Bohr’s Model

Bohr reasoned that since only certain lines are seen in atomic spectra, only certain energies must be allowed in electron orbits.