Chapter

Chapter

Electrons in the Atom

Timeline of the Models*Timeline of the Models*

460 BCDemocritus

“Atoms”

1800John Dalton

“Billiard Ball”

1897JJ

Thompson

“Plum Pudding”

1911Rutherford

“Planetary”

1912

Niels Bohr

“Bohr Model”

Today

Many scientists

“Modern”

?

*There are many more models. These are the ones we’ll cover in class.

John Dalton (c. 1800 AD)John Dalton (c. 1800 AD)

English chemist, John Dalton, performed experiments with various chemicals and showed matter seemed to consist of “indivisible” particles (atoms).

Though he didn’t know about atoms’ structure, he did know about the Law of Conservation of Matter and based his theory on this.

Dalton was an avid weather watcher and discoverer of color blindness among other things.

Dalton’s IdeasDalton’s Ideas

““Billiard Ball” or “Marbles”Billiard Ball” or “Marbles”Dalton's model says atoms are tiny,

indivisible, indestructible particlesBelieved each atom had a certain

mass, size, and chemical behavior determined by what kind of elements they make up.

See next slide for the details….

John Dalton Theory --1800(a.k.a “the marble guy”)

Atoms are the smallest particles of nature-indivisible and indestructible

All atoms of the same element are identical

Atoms of different kinds can combine to form compounds

Chemical reactions are atoms recombining to form new substances

J.J. Thomson (c. 1897)J.J. Thomson (c. 1897)

In 1897, English physicist J.J. Thompson discovered the electron and proposed a model for the structure of the atom.

Using a CATHODE RAY TUBE, Thomson discovered electrons have a negative charge and thought that the rest of matter must have a positive charge to offset the negative electron.

His ExperimentHis Experiment

JJ Thompson’s ModelJJ Thompson’s Model

“Plum Pudding”“Plum Pudding” Because the beam of light traveled

from to the positive end of the tube he concluded that the light had a negative charge

Because the beam could push a paddle wheel he concluded that the particle had mass.

Thompson's model says atoms are positively charged spheres with negatively charged electrons randomly located throughout.

Side Trip…Alpha Particles!Side Trip…Alpha Particles!

Around this time scientists also discovered alpha rays (particles), which had a positive charge.

Some physicists thought these alpha particles were made up of the positive parts of JJ Thompson’s atom.

Ernest Rutherford (1911)Ernest Rutherford (1911)

Rutherford as a student worked under J.J. Thompson supervision at the famous Cavendish Laboratories.

In 1911 Ernest Rutherford bombarded atoms with alpha rays to investigate the inside of the atom.

The results were, to say the least, unexpected!

Gold Foil ExperimentGold Foil Experiment

Rutherford used Radium as the source of the alpha particles and shot them at a thin gold foil like aluminum foil but made of gold

A fluorescent screen sat behind the gold foil on which he could observe the alpha particles’ impact.

When the particles bounced back or were deflected Rutherford reasoned that it hit something massive and positive. This mass became know as the nucleus.

When the alpha particles went straight through it hit nothing.

This happened most often so the atom is mostly empty space.

Rutherford’s Rutherford’s “Planetary“Planetary Model Model

Rutherford’s model said the negative electrons orbited a positive center (NUCLEUS)like our planets orbit the sun.

The nucleus contained most of the mass of the atom

And the distance between the positive center (nucleus) and the electrons was huge-like a marble in the center of a football field.

The atom was mostly empty space!!!!

One little problem…One little problem…

The theory of electricity and magnetism predicted that opposite charges attract each other and the electrons should gradually lose energy and spiral inward toward the nucleus.(BOOM! No more atom.)

Niels Bohr (1912)

In 1912 a Danish physicist, Niels Bohr came up with a theory that said the electrons do not spiral into the nucleus and came up with some rules for what does happen.

This was a pretty radical approach, because for the first time rules had to fit the observation regardless of how they conflicted with the theories of the time!(Aristotle would have been furious).

Previous experiments-White light gives off all wavelengths of energy- all colors.

The explanation-The explanation-

An electron absorbs energy it jumps farther away form the nucleus

As the electron falls back closer to the nucleus it gives off the energy as colored light.

How Light Relates to Electron LocationHow Light Relates to Electron Location

Bohr observed that only certain colors were given off

Therefore the electron could only orbit at certain distances from the nucleus

Bohr’s Rules

RULE 1: Electrons can orbit only at certain allowed distances from the nucleus (energy-levels).

RULE 2: An atom absorbs energy when an electron gets boosted from a low-energy orbit to a high-energy orbit.

Rule 3: Atoms radiate energy when an electron jumps down from a higher-energy orbit to a lower-energy orbit.

Bohr’s Model of the atom. Light can excite electrons around atoms and this gives rise to “quantum levels”.

Are We Done Yet?Are We Done Yet?

Almost… Cliff Notes version is that Niels Bohr came really close, and when you add the works of Arnold Sommerfeld, Wolfgang Pauli, Louis de Broglie, Erwin Schrödinger, Max Born, and Werner Heisenberg, we arrive at today’s model…

Today’s Model!-Electron CloudToday’s Model!-Electron Cloud

Today's model says electrons are not confined to fixed orbits.

They occupy volumes of space outside the nucleus.

Models of the Atom

Dalton – indivisible sphere Thomson Model – a ball of positive charge

containing a number of electrons “plum pudding”

Rutherford Model – atom has a positively charged nucleus

Bohr – electrons travel around the nucleus in definite orbits (energy levels)

Quantum Mechanical Model – no definite shape and no definite electron orbitals

Models of the Atom

Energy level – the region around the nucleus where the electron is likely to be moving (discrete levels, like stair-steps)

Quantum – the amount of energy required to move an electron from its present energy level to the next higher one

Atomic orbital – cloud shaped regions where electrons are thought to be located

S – orbital = Spherical P – orbital = Peanut D – orbital = Double peanut F – orbital = Far too complex (Flower) Nodes – regions where the probability of

finding an electron is very low (eg: No electrons)

Models of the Atom

Principle Energy Level

Number of Sublevels

Types of Orbitals

n = 1 1 1s (1 orbital)

n = 2 2 2s (1 orbital), 2p (3 orbitals)

n = 3 3 3s (1 orbital), 3p (3 orbitals), 3d (5 orbitals)

n = 4 4 4s (1 orbital),

4p (3 orbitals), 4d (5 orbitals) 4f (7 orbitals)

Table 5.1 Summary of Principle Energy Levels, Sublevels, and Orbitals

Energy Level n 1 2 3 4

Maximum number of electrons allowed

2 8 18 32

Increasing energy

(increasing distance from nucleus)

Electron Arrangement in Atoms

Electron Configurations – the way electrons are arranged around the nucleus of atoms

Aufbau Principle – electrons enter orbitals of lowest energy first

Pauli exclusion Principle – An atomic orbital may describe at most two electrons


Hund’s rule – When electrons occupy orbitals of equal energy, one electron enters each orbital until all orbitals contain one electron with parallel spins

Once all orbitals of equal energy have one electron with parallel spin, the next electron to enter the orbital has the opposite spin.


Cr =

Cu =

SEE OVERHEAD

Physics and the Quantum Mechanical Model

Electromagnetic radiation – includes radio waves, microwaves, visible light, infrared light, ultraviolet light, X-rays, and gamma rays


Amplitude – the height of the wave from the origin to the crest

Wavelength – the distance between the crests Frequency – the number of wave cycles to

pass a given point per unit of time

CrestCrest


Hertz – the metric unit of cycles per second

Spectrum – when sunlight passes through a prism, the light separates into a spectrum of colors

Atomic emission spectrum – passing light emitted by an element through a prism gives a spectrum of the element


c = ג x vc = speed of light

v = frequency

ג = wavelength


Photons – light particlesPhotoelectric effect – electrons called

photoelectrons are ejected by metals when light shines on them.Li, Na, K, Cs, and Rb are the best metals to

demonstrate the effect.


Ground State – the lowest energy level (n = 1)

Energy levels above the ground state are denoted N = 2,3,4,5,6,ect…

Electrons absorb energy to move up an energy level and give off energy to move down an energy level

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