Kenneth E. Schnobrich

ATOMIC STRUCTURE

?

A Brief History

• About 460 B.C. - a Greek Philosopher, Democritus, developed the idea of atoms ( atomos) particles as small indivisible • About 400 B.C. - a number of Greek philosophers said matter consisted of FIRE, EARTH, WATER, and AIR.

A Brief History

• NEXT 2000 YEARS - Alchemy (a pseudoscience) dominated - they were concerned with turning base metals into gold. During this time Hg, S, and Sb were discovered. Alchemists also discovered how to make mineral acids.

• 1754-1826 Joseph Proust showed that a given compound always had the same proportions by mass. Law of Definite Proportions.

A Brief History

• 1766-1844 - John Dalton (following the work of Robert Boyle) discovered that atoms can combine in more than one way. He proposed the Law of Multiple Proportions. He theorized that the basic unit was the atom . • 1808 - Michael Faraday worked on the electrolysis of molten salts and coined the word ion ( Greek meaning wanderers).

A Brief History

• 1808 - John Dalton published “A New System of Chemical Philosophy” which proposed his theory of atoms – All elements are composed of tiny, discrete, indivisible and indestructible particles called “atoms”.

– All atoms of a given element are identical – Atoms of of different elements are different; they have different masses and properties.

– Chemical combinations of these “atoms” compose all matter - different atoms combine differently to form compounds.

DALTON’S ATOM

JOHN DALTON - he envisioned the atom as a hard spherical unit of matter (the ultimate unit) Lithium Chlorine LiCl Sulfur Li 2 S

DALTON’S ATOM

JOHN DALTON - he envisioned the atom as a hard spherical unit of matter Oxygen Hydrogen H 2 O H 2 O 2

HISTORY (cont.)

• 1875 - Eugen Goldstein discovered the existence of a charged stream from the cathode using a “Crookes Tube” and called them “Cathode Rays”

HISTORY (cont.)

• 1886 - Eugen Goldstein discovered the existence of positively charged particles he called “Canal Rays”

HISTORY (cont.)

• 1897 - J.J. Thomson using a modified “Crookes Tube” determined that the “Cathode Rays” behaved like charged particles and measured the charge.

HISTORY (cont.)

• 1907 - J.J. Thomson proposed his “Raisin Pudding Model” of the atom.

Negatively charged Electrons Positive Matrix

*Atoms are neutral

HISTORY (cont.)

1911 - Ernest Rutherford suggested the atom was “nuclear” based on a famous experiment - “The Scattering Experiment”. He also suggested that the proton was the fundamental unit of positive charge

HISTORY (cont.)

Metal Foil(Au) Alpha Particles

ASSUMPTIONS RUTHERFORD’S WORK

• Most of the atoms mass is concentrated in the nucleus.

• All of the positive charge is concentrated in the nucleus • Neutral atoms have equal numbers of protons and electrons.

• The protons and neutrons are located in the nucleus of the atom.

HISTORY (cont.)

THOMSON MODEL RUTHERFORD MODEL

HISTORY (cont.)

• 1932 - Chadwick discovered and determined the properties of the neutron. Proton Neutron Electrons Nucleus

SUBATOMIC PARTICLES

PARTICLE CHARGE MASS LOCATION SYMBOL PROTON NEUTRON ELECTRON +1 0 -1 1 AMU NUCLEUS 1 H 1 or 1 p 1 1 AMU NUCLEUS 1/1836 AMU OUTSIDE 0 -1 n 1 e 0

LOOKING AT THE ATOMS STRUCTURE

• Atomic Number = # protons and electrons in a neutral atom • Atomic Mass Number* = sum of the protons and neutrons • #Neutrons = Mass# - Atomic# *ATOMIC MASS MAY VARY (ISOTOPES)

ISOTOPES AVERAGE ATOMIC MASS

Most of the elements on the periodic table have several Isotopes. The Mass that you see is the weighted average of known isotopes.

Example: Carbon has two stable isotopes C 12 = 98.89% and C 13 = 1.108% 12(0.9889) + 13(0.01108) = 12.01

Average Atomic Mass

FORMING AN ION

• METALS - usually like to lose electrons to form positive ions called CATIONS .

• NONMETALS - usually like to gain electrons to form negative ions called ANIONS .

FORMING IONS

Na Cl + electron Na +1 + electron Cl -1

Atomic Mass Atomic Number

54.94

Mn

25 2-8-13-2

Electron Arrangement

+2 +3 +4 +7

Oxidation States

THE KERNEL AND VALENCE ELECTRONS

19

K

39

2-8-8-1

Valence Electrons level – those in the outer principal energy Kernel – the nucleus and all of the electrons except those in the valence level

THE KERNEL AND VALENCE ELECTRONS

19

K

39

2-8-8-1

Valence Electrons = 1 Kernel – has a charge of +1

*Now Lewis Dot Structures

LEWIS DOT STRUCTURES

19

K

39 16

S

-2

2-8-8-1 2-8-8

K [ ]

x x

-2

PUTTING IT TOGETHER

PARTICLE N N -3 Sn Sc +3 Na PROTONS ELECTRONS NEUTRONS

CONTINUOUS SPECTRUM

VISIBLE REGION OF THE SPECTRUM

HYDROGEN AND HELIUM LINE SPECTRUM

HYDROGEN HELIUM

More spectra

BRIGHT-LINE SPECTRA

BRIGHT-LINE SPECTRA ARE LIKE FINGER PRINTS THE . EACH ELEMENT HAS ITS OWN CHARACTERISTIC SET OF BRIGHT LINES IN VISIBLE REGION OF THE SPECTRUM.

Hydrogen Helium Carbon

THE BOHR MODEL

Bohr’s Model was based on the simplest atom, Hydrogen. Bohr based his model on the following: (1) Electrons do not follow the rules of large macroscopic bodies.

(2) Electrons in atoms have only specific energies .

(3) Electrons are only in specific orbits outside the nucleus ( ground state ).

(4) When an electron moves from one orbit to another it absorbs or releases energy of a specific frequency.

(5) When electrons absorb energy they move to an excited state (higher energy orbit).

2 1

THE BOHR MODEL Hydrogen

Excited State Ground State 2 1 2 1 Energy Absorbed Energy Released

THE QUANTUM MODEL

As the science of spectroscopy grew and the resolution of the bright-line spectra of an element improved and the atom.

dual nature of the electron was explored scientists formulated a new picture of the This new model of the atom retains some of the original features but changes the concept of electron location. The electron, instead of occupying a specific orbit now is thought to occupy a region of 3-D space called the orbital .

THE QUANTUM MODEL

Dual Nature of the Electron – the electron to this point, had been described as being “ particle-like ” in nature, but it also exhibits “ wave-like behavior .” DeBroglie – was the first to suggest that, based on its extremely small size, the electron does have a measureable wavelength.

Double-click on the You Tube video

THE QUANTUM MODEL

•

After viewing the video we see that Erwin Schrodinger allows us to describe the electrons in an atom with a set of 4 Quantum Numbers .

The quantum numbers help us to describe the relative energies and probable locations of the electrons.

•

The P rincipal Q uantum Bohr Model. The PQN N umber (n) – corresponds very closely with the energy levels described in the can only have small whole number values (n = 1, 2, 3, 4, 5, 6 etc) . The greater the value of “n” the greater the energy and distance from the nucleus for the electron.

THE QUANTUM MODEL

•

The

•

S ublevel Energy Level. The SQN value of the Q uantum PQN .

N umber ( l ) – describes the sublevels the electrons can occupy within a Principal – has values that are determined by the

• •

It can have values from 0 … n-1 So, if n = 0, l = 0

•

If n = 2, l = 0, 1 (which means, in the second P rincipal E nergy L evel, there are two available sublevels the electron can occupy.

•

There are also corresponding letter values for the sublevels 0(s); 1(p); 2(d); 3(f)

THE QUANTUM MODEL

•

In the 2 nd PEL there were two sublevels, 0, 1 or s and p.

•

Within a PEL , as the value of l increases the energy and distance from the nucleus increases.

•

In the 3 rd PEL, there are three sublevels, 0, 1, and 2 or s, p, and d sublevels.

•

The O rbital Q uantum within a sublevel.

N umber (m) (also sometimes called the Magnetic Quantum number) - describes the number of orbitals (3-D orientations in space)

THE QUANTUM MODEL

•

The OQN’s are determined by the values for l

•

m can have values from 0… +/ l

•

So, if l sublevel.

= 0, m = 0, which means that there is only one possible 3-D description (or orbital) in that

•

If l = 1, m = 0, +1, -1, which means in the “p” sublevel there are three, 3-D descriptions (or orbitals), in that sublevel.

Along the X axis Along the Y axis Along the Z axis

THE QUANTUM MODEL

•

If an orbital is located in an s sublevel it is referred to as an s-orbital and has a spherical distribution along the X, Y, and Z axes.

•

If an orbital is located in an p sublevel it is referred to as an p-orbital and has a “dumbell” distribution” along the X, Y, and Z axes.

THE QUANTUM MODEL

•

Of course there are other orbital shapes but they are complicated and for our purposes, our concerns will be limited to the s and p orbital shapes.

•

The fourth quantum number is the S pin Q uantum N umber – based on the Stern/Gerlach experiment it is thought that an electron can have one of two possible spins, +1/2 and -1/2 (it spins on its axis).

•

Since no two electrons can have exactly the same set of four quantum numbers , only two electrons can occupy an orbital , provided they have opposite spins .

THE QUANTUM MODEL

•

Based on the work of many scientists, including deBroglie, Shrodinger, and Heisenberg, we now know that –

•

we can only speak in terms of the probable location of the electrons

•

the bright line spectra available for the elements gives us additional information on the energy associated with the electrons

THE QUANTUM MODEL OF HYDROGEN

Note: the electron is pictured as a cloud or region of space where you will most probably find the electron.

Nucleus

Quantum Atom Relationships

Increasing Energy ORBIT Principal Energy Level SUBLEVELS s p d f 1 orbital 3 orbitals 5-orbitals 7-orbitals

Electron Filling

• When we fill the energy sublevels that are several rules we must follow – The Aufbau Principle – you must always fill from lowest energy to highest energy • Hund’s Rule – you must completely half-fill an energy sublevel before you start pairing electrons • Pauli Exclusion Principle – no two electrons can have the same set of four quantum numbers in a given orbital, they must have opposite spins to exist in the same orbital.

Filling the Sublevels & Orbitals

• When filling the sublevels and orbitals remember the rules • It is also important to remember that for multi-electron atoms some of the sublevels do overlap from an energy standpoint.

• there is a simplified relationship to help us with this overlap – Sublevel Energy = n +

l (

n is the PQN and

l

is the SQN • ). It is why the 4s sublevel fills before the 3d sublevel (see the diagram on the next slide).

4s = 4 + 0 = 4 and 3d = 3 + 2 = 5 • the sublevel energy of 4s is lower than that of 3d, therefore, the 4s sublevel fills before the 3d sublevel.

General Sublevel Arrangement

3d 4s 3p 3s 2p 2s 1s

General Sublevel Arrangement 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f

* The idea of sublevel overlap can be much more complicated for larger atoms

General Sublevel Arrangement

3d 4s 3p 3s 2p 2s 1s

For 19 K 39 the sublevel filling would look like this

General Sublevel Arrangement

3d 4s 3p 3s 2p 2s 1s

For 7 N 14 filling would look like this the sublevel Note: the sublevel is half-filled, the electrons have parallel spins (the same)

Electron Arrangement

Let’s take a sample and show you how the electron arrangement can be written in three formats.

19

K

39 19

K

39 19

K

39

2-8-8-1 1s

2

2s

2

[Ar]4s 2p

1 6

3s

2

3p

6

4s

1 Principal Energy Level

4s

1 # of electrons Energy sublevel