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Atomic Structure

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This electron microscope

high-resolution image

shows magnification of the

thin edges of a piece of

mica. The white dots are

"empty tunnels" between

layers of silicon-oxygentetrahedrons, and the black

dots are potassium atoms

that bond the tetrahedronstogether. Note the 10

Angstrom width, which is

0.000001 mm.

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First Definition of the Atom

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The Early Greeks thought of everything as beingmade up of four basic elements.

Earth.

Air.Fire.

Water.

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Democritus and Leucippus thought that matter wasdiscontinuous or made up of individual particles.

Democritus called these fundamental particles atoms.

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Atomic Structure Discovered

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IntroductionIn 1661 Robert Boyle defined an element as a simple

substance which could not be broken down into simpler

substances.

We now define an element as a pure substance that can

not be broken down into simpler things by either

chemical or physical methods.

Since elements always combine in fixed ratios, this lendssupport to the idea of elements being made of discrete

particles.

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(A) Oxygen andlead combine toform yellow leadoxide in a ratio of1:13. (B) If 1 atomof oxygencombines with 1atom of lead, the

fixed ratio inwhich oxygen andlead combinemust mean that 1

atom of lead is 13times moremassive than 1atom of oxygen.

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Reasoningthe existence of

atoms from theway elementscombine in fixed-weight rations.

(A) If matter werea continuous,infinitelydivisible material,there would be noreason for one amount to go with another amount. (B) Ifmatter is made up of discontinuous, discrete units (atoms),then the units would combine in a fixed-weight ratio. Sincediscrete units combine in a fixed-piece ratio, they must

also combine in a fixed-weight-based ratio.

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Discovery of the ElectronA cathode ray is a beam of electrons that moves between

metal plates in an evacuated tube from a negative to a

positive terminal. The electron beam is seen as a green beam.

These rays can be deflected by a magnet.

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A vacuum tube with metal plates attached to a high

voltage source produces a greenish beam called

cathode rays. These rays move from the cathode

(negative charge) to the anode (positive charge).

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In 1897 JJ Thompson place a positively charges plate

on one side of the tube and a negatively charged plate on

the other side of the tube.

The beam was deflected away from the negative plate toward

the positive plate.

Thompson realized that the particles that made up the beam

must be negatively charged, since like charges repel and

opposite charges attract.

By balancing the deflections made by the magnet with that

made by the electrical field, Thompson was able to calculate the

ratio of the charge to mass of an electron as 1.7584 X 1011coulomb/kilogram

These particles were later named electrons.

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What appears to be visible light coming through the slit in thisvacuum tube is produced by cathode ray particles striking a

detecting screen. You know it is not light, however, since the beam

can be pulled or pushed away by a magnet and since it is attracted

to a positively charged metal plate. These are not the properties of

light, so cathode rays must be something other than light.

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A cathode ray passed between two charged plates is

deflected toward the positively charged plate.

The ray is also deflected by a magnetic field.

By measuring the deflection by both, J.J. Thomson was

able to calculate the ratio of charge to mass.

He was able to measure the deflection because thedetecting screen was coated with zinc sulfide, a substance

that produces a visible light when struck by a charged

particle.

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In 1906 Robert Millikan passed mineral oil through a

vaporized into an apparatus where he could observe the

drops with a magnifier and make measurement on themas they drifted downward.

he found that the least charge on any of the droplets was 1.60 X

10-19 coulombs and that larger droplets always had a charge that

was some multiple of this value.

Knowing Thompsons work of charge to mass ratio and the

charge on an individual electron, it was possible to calculate the

mass of the electron as 9.11 X 10-31 kg.

Thompson proposed that an atom was a blob of positively

charged matter in which electrons were stuck like raisins in

plum pudding.

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Millikan measured the charge of an electron by

balancing the pull of gravity on oil droplets with an

upward electrical force.

Knowing the charge-to-mass ratio that Thomson had

calculated, Millikan was able to calculate the charge on

each droplet.He found that all droplets had a charge of 1.60 x 10-19

coulombs or multiples of that charge.

The conclusion was that this had to be the charge of an

electron

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The Nucleus

Ernst Rutherford determined that there was a positively

charge nucleus associated with the atom, that was

surrounded by electrons.

Rutherford calculated that the radius of the nucleus to be about10-13 cm and the radius of the atom to be about 10-8 cm.

Electrons therefore took up about 100,000 times the radius of

the nucleus.

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Rutherford and his co-workers studied alpha particle

scattering from a thin metal foil.The alpha particles struck the detecting screen, producing

a flash of visible light.

Measurements of the angles between the flashes, the

metal foil, and the source of the alpha particles showed

that the particles were scattered in all directions,

including straight back toward the source

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In 1917 Rutherford broke up the nucleus of the nitrogen

atom by bombarding it with alpha particles and was able

to identify a particle with a positive charge called aproton.

He also thought that there were neutral particles in the nucleus

called neutrons.

The atom has a tiny, massive nucleus made up of protons and

neutrons.

Negatively charged electrons, whose charge balances the

charge on the protons, move around the nucleus at a distance of

about 100,000 times the radius of the nucleus.

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Rutherfords's nuclear model of the atom explained

the alpha scattering results as positive alphaparticles experiencing a repulsive forced from the

positive nucleus

Measurements of the percent of alpha particlespassing straight through and of the various angles

of scattering of those coming close to the nuclei

gave Rutherford a means of estimating the size ofthe nucleus.

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From measurements of alpha particle scattering,Rutherford estimated the radius of an atom to be100,000 times greater than the radius of the nucleus.This ratio is comparable to that of the (A) thickness

of a dime to the (B) length of football field.

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The Bohr Model

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The Quantum Concept.

In 1900 Max Plankintroduced the idea that matter emits

and absorbs energy in discrete units called quanta.

In 1905 Albert Einstein extended the quantum concept

to include light and that light consist of discrete units

called photons.The energy of a photon is directly proportional to the

frequency of vibration.

E=hf

where E = energy

h = Planks constant = 6.63 X 10-34 Js

f = frequency

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(A) Light from incandescent solids, liquids, or dense

gases, produces a continuous spectrum as atoms

interact to emit all frequencies of visible light (B)Light from an incandescent gas produces a line

spectrum as atom emit certain frequencies that are

characteristic of each element.

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Atomic hydrogen produces a series of characteristic line

spectra in the ultraviolet, visible, and infrared parts of the

total spectrum. The visible light spectra always consist of

two violet lines, a blue-green line, and a bright red one.

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Bohrs Theory

Allowed Orbitals

An electron can only orbit around an atom in specific orbits

Radiationless Orbits

An electron in an allowed orbit does not emit radiant energy as

long as it remains in the orbit.

Quantum Leaps

An electron gains or loses energy only by moving from one

allowed orbit to another.

The lowest energy state is known as the ground state

Higher states are known as excited states

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Each time an electron males a "quantum leap," moving froma higher energy orbit to a lower energy orbit, it emits a

photon of a specific frequency and energy value.

An energy level

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An energy leveldiagram for ahydrogen atom, notdrawn to scale. The

energy levels (n)are listed on the leftside, followed bythe energies of

each level in J andeV. The color andfrequency of thevisible light

photons emitted arelisted on the rightside, with the arrowshowing the orbit

moved from and to.

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These fluorescent lights emit light as electrons of mercuryatoms inside the tube gain energy from the electric current.As soon as they can, the electrons drop back to their lower-energy orbit, emitting photons with ultraviolet frequencies.Ultraviolet radiation strikes the fluorescent chemical coating

inside the tube, stimulating the emission of visible light.

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Quantum Mechanics

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Quantum mechanics states that light and matter,

including electrons, have a dual nature of both

particles and waves.

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Matter Waves.

Louis de Broglie reasoned that particles must also have a

dual nature.

He reasoned that the electron should have a certain

wavelength that would fit into its orbit around the

nucleus. =h/mv

where is the wavelength

h is Planks constant

m is the mass

v is the velocity

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(A) A schematic of de Broglie wave, where the standingwave pattern will just fit in the circumference of an orbit.

This is an allowed orbit. (B) This orbit does not have a

circumference that will match a whole number of

wavelengths; it is not an allowed orbit.

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Wave Mechanics

Electrons do emit light in certain wavelengths based on

their energy levels (orbital radius)

Since waves spread out from the electron, the wave

mechanic model predicts an area where an electronwould be found, and not a specific place where it would

be found.

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The Quantum Mechanics Model

Quantum mechanics describes the energy levels of an

electron wave with four quantum numbers.

distance from nucleus

energy sublevel

orientation in space. direction of spin

Principal quantum number (n)

Describes main energy level of the electron in terms of its

distance from the nucleus.

n = 1, 2, 3, 4, 5, 6, 7

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Angular momentum quantum number

Defines energy sublevels within the main energy

levels

s, p, d, or f designating the type of orbital and also the

orbital shape.

The Heisenberg Uncertainty Principle states thatyou cannot measure the momentum and exact position

of an electron at the same time.

What you can measure is the probability that an

electron will be found in a certain area, called an

orbital.

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(A)An electron distribution sketch representing probabilityregions where an electron is most likely to be found. (B) A

boundary surface, or contour, that encloses about 90 percentof the electron distribution shown in (A). This three-dimensional space around the nucleus, where there is thegreatest probability of finding an electron, is called an

orbital.

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Magnetic quantum number

Defines the orientation in space of the orbitals relative

to an electrical field.

The s orbital has one orientation

The p sublevel can have 3 orientations

The d sublevel can have 5 orientations

The f sublevel can have 7 orientations.

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(A) A contour representation of an s orbital. (B) A

contour representation of a p orbital.

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Spin quantum number Describes the direction of spin of an electron in its

orbit.

Electrons occur in pairs and each of the orientations

for a sublevel can have one electron pair.

Experimental

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Experimental

evidence

supports the

concept thatelectrons can

be considered

to spin oneway or the

other as they

move about anorbital under

an external

magnetic field.

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Pauli Exclusion Principle

No two electrons can have the same set of quantum numbers.

At least one of the quantum numbers must differ.

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Electron Configuration

This is a shorthand designation for electron orientation.

The lowest possible energy level is n=1.

If one electron already occupies this energy level, a second can

only occupy it if it has a different spin quantum number.

Electron configurations tells you the quantum numbers ofthe electron.

Energy sublevel

Principle quantum number1s2 two electrons

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There are threepossible

orientations of thep orbital, andthese are called

px, py, and pz.

Each orbital canhold twoelectrons, so atotal of sixelectrons are

possible in thethree orientations;thus the notation

p6.

A matrix showing

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gthe order in whichthe orbitals are

filled. Start at thetop left, then movefrom the head ofeach arrow to the

tail of the oneimmediately belowit. This sequencemoves from thelowest-energy levelto the next higherlevel for each