Unit 2:

Atomic Theory

Video links

overview of atomic historyhttp://www.youtube.com/watch?feature=player_detailpage&v=k1RHY8QcN1s

I. Atomic History

A. The GreeksDemocritus

[Philosopher]

All matter is made of tiny, indivisible parts called ‘atoms’Developed word ‘atomos’ meaning not divisible

John Dalton (1803-1808)

Used experiments with gases to develop the “Atomic Theory”Determined atoms looked like ‘cannonballs’ or solid masses

                                

Dalton’s Atomic Theory

1) All elements are made of atoms2) Atoms of each element are all the same, or have the same masses3) Atoms of different elements are different, or have different masses4) Atoms cannot be created or destroyed5) Atoms combine in small, whole number ratios

J.J. Thomson (1897)

Developed ‘Cathode Ray’ experiment Said atoms consisted of particles smaller than an entire atomDiscovered that the smaller particles within an atom had a negative chargeDiscovered 1st subatomic particle: Electron Founded “Plum Pudding Model”: Electrons were embedded within a positively charged mass

Cathode Ray Tube Experiment

Thomson manipulated cathode rays with a magnet to discover that subatomic particles existed and that they had negative charges

Ernest Rutherford (1898)

Discovered alpha and beta radiation emitted from certain radioactive substancesDeveloped and used Gold Foil ExperimentFirst to separate the smaller parts of the atom

Discovered the nucleus Placed electrons outside the nucleus Stated that atoms are composed of lots of empty space

Rutherford’s Gold Foil Experiment

micro.magnet.fsu.edu/electromag/java/rutherford/

Niels Bohr (1922)

Bohr analyzed work of others and studied atomic spectra, or light, given off by the elementsDescribed the “Atomic Spectra” of elementsDeveloped ‘Solar System’ modelMoved electrons from single, giant pathway into discrete energy levels around the nucleusEach energy level contained 2, 8, 18, 32, etc. electrons total

Bohr Model of the Atom

Stated that electrons moved around the nucleus in ‘orbits’ or energy levelsAs electrons gain energy, they jump up energy levels, then release this energy to generate spectra

Bohr’s model and the atomic spectrum

http://jersey.uoregon.edu/vlab/elements/Elements.html

•The spectral lines in the visible region of the atomic emission spectrum of barium are shown below.

•

                                                                                                                                                                                                                        

•Spectral lines exist in series in the different regions (infra-red, visible and ultra-violet) of the spectrum of electromagnetic radiation. •The spectral lines in a series get closer together with increasing frequency. •Each element has its own unique atomic emission spectrum.

Erwin Schrodinger (1930)

Developed mathematical equations representing electronsElectrons had wave and particle behaviorsCreated “Wave-Mechanical” or “Modern” modelMost scientists use this model todayPlaced electrons in orbitals

“Electron Cloud” ModelCreated paths for electrons within Bohr’s energy levelsOnly 2 electrons per pathElectron paths, or ORBITALSORBITALS, are mathematical equations describing probability densities for electrons Developed sublevels with discrete paths within each energy level

II. Subatomic Particles

A. Particles1) ProtonsProtons

found in the nucleus of an atom charge of +1, mass of

1.0073a.m.u.

2) NeutronsNeutrons found in the nucleus of an atom no charge, mass of 1.0087a.m.u.

A. Subatomic Particles

3) Electrons Found outside the nucleus in regions of

probability [orbitals] Charge of –1, mass of 5.46 x 10-4 a.m.u.,

or 1/1836 a.m.u. Have particle and wave properties

B. Atomic Number

Atomic Atomic numbernumber = the number of protons in the nucleus

All atoms of the same element have the same atomic number

Atoms arranged on PT by increasing atomic numbers

In neutral atoms: Atomic number

equals number of electrons

C. Isotopes

IsotopesIsotopes = atoms of the same element that have differing numbers of neutrons in their nucleus, different mass number, but same atomic numberSame number of protons!!!Changing number of neutrons affects properties [radioactivity…]

D. Atomic Mass

Atomic Mass Atomic Mass Number Number = number of protons plus the number of neutrons in the nucleus

Whole number!!

Mass Number changes when using different isotopes

Written in isotopic notations, just subtract the top from bottom values:

E. Ions

IonsIons = atoms of the same element that have lost or gained electronsHave overall (+) or (-) chargeSame numbers of protons, number of neutrons irrelevantPositive ions: have LOST electronsNegative ions: have GAINED electrons

F. Atomic Mass (average)

Atomic MassAtomic Mass = weighted average of the natural isotopes times their percent abundanceDecimal value on PTAccounts for the natural existence of various isotopes Ex] calculate the atomic mass of carbon given that 98.92% is carbon-12 and 1.108% is carbon-13

Virtual textbook

http://www.chem1.com/acad/webtext/intro/int-1.html#SEC1

III. Electronic StructureA. EMS [Electromagnetic Spectrum]

A. EMS [Electromagnetic Spectrum]

EMS EMS = continuous series of various types of energy, separated by their wavelengths and frequenciesVisible lightVisible light = small portion; only part we can see without instrumentsContinuous spectrumContinuous spectrum = picture of all colors of visible light as they pass through a prism

EMS continued

Wavelength = distance between 2 peaks or troughs of 2 consecutive waves

Symbol = λ [Greek letter “lambda”]

Units are usually in ‘m’ or ‘nm’

Frequency = the number of peaks or troughs that pass a single point in one second

Symbol = ʋ [Greek letter “nu”]

Units are usually in ‘1/s’ or ‘s-1’ or ‘Hz’

Calculations using lambda and nu

c = λν

C = speed of lightC = 3.0 x 10+8 m/s

E = hν

E = energy of photonh = Planck’s constanth = 6.63 x 10-34 Js

All electromagnetic radiation travels at the speed of light

Can calculate the energy of the radiation/electron given the wavelength

Planck’s ConstantPlanck observed hot, glowing matter

Concluded: different substances glow different colors at different temperaturesDetermined: matter releases energy in tiny, discrete packets called ‘quanta’Developed constant to relate energy and temperature, Planck’s constant, “h”

h = 6.63 x 10-34 J*s

Light traveling as waves

All colors of light energy travel at the same speed, just different wavelengths!

Particle vs. Wave Behavior of Light

Wave behavior of light

B. Photoelectric Effect

Einstein used Planck’s idea of quanta and photons to describe the photoelectric effectLight of a certain wavelength shines on clean metal, causing the metal to eject electrons

                                                                                  

C. Bohr’s Model [conclusions made]

Bohr used the idea of ‘quanta’ to explain the bright-line emission spectra

Stated that each element’s atomic spectrum is unique

Electrons exist in ground state energy levels, as listed via the periodic table

Bohr Model of the Atom

Stated that electrons moved around the nucleus in energy levels

Electrons will gain and lose energy at will

This generated the element’s atomic spectrum

Bohr’s model

Useful Websites and References

//www.avogadro.co.uk/light/bohr/spectra.htm shows formation of spectral lines for hydrogen idea of ground vs. excited state

//jersey.uoregon.edu/vlab/elements/Elements.html Periodic table showing the absorption and emission

spectra for each element

Also check out Wikipedia under Bohr atom and Atomic spectra!Also check out Wikipedia under Bohr atom and Atomic spectra!

Creation of an emission spectrum

If electrons absorbabsorb packets of energy, quanta, they temporarily move to into a higher energy level, called the excited stateThe electrons then releaserelease this quanta of energy and fall back down to ground stateThe release of energy generates the bright-line emission spectrum

Examples of Bohr Diagrams

IV. Electron Configurations

A. Energy Levels

These are areas with a high possibility of finding electrons with similar potential energies7 energy levels total

Bohr Diagrams and Energy Levels

Bohr Diagrams show the numbers of protons and neutrons in the nucleus Shows electrons in their respective energy levels

Energy levels hold: 1st holds 2 electrons 2nd holds 8 electrons 3rd holds 18

electrons 4th holds 32

electrons Etc…..

B. Sublevels

Sublevels are divisions within each energy levelRepresent the shapes and orientation in 3D spaceToo many electrons within the energy levels & they lose momentum and will crash into the nucleus--- not good!not good!

1st energy level has 1 sublevel: “s”2nd has 2 sublevels: “s” and “p”3rd has 3 sublevels: “s, p, and d” 4th has 4 sublevels: “s, p, d, and f”

Sublevels and Shapes

“s” is spherical and has a max of 2 electrons“p” is dumbbell shaped and has a max of 6 electrons“d” is cloverleaf shaped and holds up to 10 electrons“f” is a split cloverleaf with a max of 14 electronshttp://micro.magnet.fsu.edu/electromag/java/atomicorbitals/index.html

Order of Sublevel FillingIt does not go in order…

1s2

2s2 2p6

3s2 3p6 3d10

4s2 4p6 4d10 4f14

5s2 5p6 5d10 5f14

6S2 6P6 6d10

7s2 7p6

Orbitals within Sublevels

Each sublevel consists of 1 to 7 orbitals [areas of probability for finding an electron]

Each path or orbital only holds 2 electrons

The 2 electrons within in each orbital each have a different spinThis allows the electrons to exist in the same area without conflicting

C. Extended and Abbreviated Configurations

Electron Configurations = way to describe how the electrons are distributed around an atom and within the energy levels and sublevelsGround state configurations are same order as electrons on PTExcited state configurations have one electron shifted to a higher energy level

Writing Electron Configurations

Electrons add in the same order as the atomic numbers of the PT

Aufbau Principle = adding electrons in the exact order of the PT

                                   

Writing Configurations

Add in order of arrows for Neutral, Ground Neutral, Ground State atomsState atoms

Examples:Examples:

Abbreviated Configurations

Abbreviated Abbreviated configurationsconfigurations show only the placement of electrons added after the last ‘noble gas’Ex]

Bracket the configuration of the last noble gas [group 18] and add remaining electronsEx]

D. Orbital Notations and Rules

Orbital notationsOrbital notations are specialized versions of a full electron configuration showing the spin of each electron within an orbitalDraw the orbitals present for each sublevel and fill with ‘spin-paired’ electrons

Rules for Configurations

1. Hund’s RuleHund’s Rule = electrons in the p, d, and f sublevels must be added to each orbital first, before one flips to spin-pair and fill the orbital

2. Pauli Exclusion Pauli Exclusion PrinciplePrinciple = no 2 electrons may be in the same orbital and have the same spin; no 2 electrons will have the same 4 quantum numbers

Rules cont’

3.3. Heisenberg’s Uncertainty Heisenberg’s Uncertainty PrinciplePrinciple = states that the electron’s momentum and position cannot be accurately determined at the same time

Example…

Excited State vs. Ground State

Excited StateExcited State configurations show one electron has moved into a higher energy level, leaving an unfilled space below

Ground state configurations are written in order of the periodic table

**Total # of **Total # of electrons = electrons = Atomic # Atomic #

for for BOTH !!

E. Lewis Dot Structures

Lewis Dot Structures are pictures showing the placement and number of valence electrons for an element

Structure:

s1s2

p6 p1

p3 p4

p5p2

Valence electrons are s and p outer shell electronsMaximum of 8!Ex]

F. Quantum Numbers

Each electron in an atom is assigned a set of 4 quantum numbersThese numbers tell the exact “address” of an electron, regardless of the elementNo 2 electrons have the same 4 quantum numbers!

Quantum Numbers

1] Principle Quantum Number (n)First number Represents the energy level of the electronValues range from 1 to 7

2] Azimuthal Spin Number (l)Second NumberRepresents the sublevelDescribes the shape of the orbitalValues from 0 to 3

Quantum Numbers

3] Magnetic Spin Magnetic Spin Number (Number (mmll ))

Tells the orientation of the orbital along x, y, z axesValues for:

l = 0, ml = 0

l = 1, ml = +1, 0, -1

l = 2, ml = +2, +1, 0, -1, -2 l = 3, ml =+3, +2, +1, 0, -1, -2, -

3

4] Spin Number (Spin Number (mmss))

Tells if the electrons spin clockwise, or counterclockwiseValues:

+1/2 [spin up] or –1/2 [spin down]