ARRANGEMENT OF ELECTRONS IN ATOMS Chapter 4

Visible Light We are all familiar with light but what is “visible” is just a very, very small portion of the electromagnetic spectrum What colors make up the rainbow? Red, Orange, Yellow, Green, Blue, Indigo, Violet (ROYGBIV)

The E-M Spectrum Gamma Rays (Very Harmful / Cancerous) X-rays (Cancerous in large doses; small doses medical scanning) Ultraviolet ( not as harmful - sunburn; black lights) Infrared (Heat, communication) Microwaves (Cooking, communications) Radio Rays (TV, Radio, other communications)

The Development of a New Atomic Model Wavelength ( ) - length of one complete wave Frequency ( ) - # of waves that pass a point during a certain time period  hertz (Hz) = 1/s Amplitude (A) - distance from the origin to the trough or crest

Waves A greater amplitude (intensity) greater frequency (color) crest origin trough A

The Electromagnetic Spectrum AM radio Short wave radio Television channels FM radio Radar Microwave Radio Waves Gamma Rays X- Raysinfrared Increasing photon energy Increasing frequency Decreasing wavelength Red Orange Yellow Green Blue Indigo Violet UV Rays R O Y G B I V VisibleLightVisibleLight

Electromagnetic Spectrum Frequency & wavelength are inversely proportional c = c:speed of light (3.00  10 8 m/s) :wavelength (m, nm, etc.) :frequency (Hz)

Electromagnetic Spectrum GIVEN: = ? = 434 nm = 4.34  m c = 3.00  10 8 m/s WORK : = c = 3.00  10 8 m/s 4.34  m = 6.91  Hz EX: Find the frequency of a photon with a wavelength of 434 nm.

So why is the electromagnetic spectrum so important to chemistry? Why is the steel emitting light when it is heated? We take it for granted that when things get hot they turn red then orange and finally white; but that isn’t good enough any more

Black Body Radiation Colors

So why is the electromagnetic spectrum so important to chemistry? Incandescence is heat made visible – the process of turning heat energy into light energy. Our usage of terms like "red hot," "white hot," and so on, is part of the color sequence black, red, orange, yellow, white, and bluish white, seen as an object is heated to successively higher temperatures.

So why is the electromagnetic spectrum so important to chemistry? The light produced consists of photons emitted when atoms and molecules release part of their thermal vibration energy. For increasing temperatures, the sequence of radiated colors is: black, red, orange, yellow-white, bluish- white.

Heat and Light Planck (1900)  Observed - emission of light from hot objects  Concluded - energy is emitted in small, specific amounts (quanta)  Quantum - minimum amount of energy change

Energy and Light E:energy (J, joules) h:Planck’s constant (  J·s) :frequency (Hz) E = h zThe energy of a photon is proportional to its frequency.

Energy and Light GIVEN: E = ? = 4.57  Hz h =  J·s WORK : E = h E = (  J·s ) ( 4.57  Hz ) E = 3.03  J EX: Find the energy of a red photon with a frequency of 4.57  Hz.

Niels Bohr and the Bohr model of the atom Bohr hypothesized that instead of haphazardly orbiting the nucleus, electrons had clearly defined orbits – very similar to the planetary orbits circling our sun His model is (cleverly) named the Planetary Model

Niels Bohr Bohr Model (1913)

Bohr’s Proof Bohr said this: If you assume that the electrons have clearly defined orbits that are congruent to the energy levels…

Bohr’s Proof … then when an electron gets “excited” it jumps to a higher energy level. When it “relaxes” it emits a certain wavelength of light. Bohr showed the energy of an electron in an atom is quantized, which means it has a particular numerical value, not some arbitrary number.

Excitation of Hydrogen Atoms Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 328

Return to Ground State

Bohr’s Proof n=1 n=2 n=6 n=5 n=4 n=3 n=7 1 = x 10 7 m -1 λ nr2nr2 ne2ne Lyman Series (uv) Balmer Series (vis and uv) Paschen Series (ir)

Emission Spectrum of an Element 1 nm = 1 x m = “a billionth of a meter” 410 nm434 nm486 nm656 nm 1 nm = 1 x m = “a billionth of a meter”

Continuous and Line Spectra

Hydrogen to Steel If Hydrogen emits 4 distinct wavelengths of light when its one electron is excited what can we extrapolate to that of steel which is made mostly of iron? s.html s.html

Flame Emission Spectra Photographs of flame tests of burning wooden splints soaked in different salts. Include link to web page methane gas wooden splintstrontium ioncopper ionsodium ion calcium ion

Fireworks

Composition of Fireworks Gunpowder  Sulfur, charcoal, potassium nitrate (saltpeter) Salts (to give color)  Red = lithium  Green = copper

Good News Bad News Good News Bohr’s Model works and moves us along in the development of the Atomic Theory End of this little unit Bad (Frustrating) News Lots of Math Everything I taught you only works for Hydrogen and therefore is completely wrong and obsolete.

Check for Understanding c= λν E=hν c=3.0 x10 8 m/s h=6.626 x J s What is the frequency of a radar photon with an energy of 7.2 x J? What is the frequency of light having a wavelength of 6.20x10 -7 m?

Models of the Atom Dalton’s model (1803) Thomson’s plum-pudding model (1897) Rutherford’s model (1909) Bohr’s model (1913) Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 125 Greek model (400 B.C.) e e e e ee e e e e

Your Current View of the Atom electrons nucleus

Again… so why is it so important to chemistry? Einstein (1905)  Observed - photoelectric effect

Again… so why is it so important to chemistry? Einstein (1905)  Concluded - light has properties of both waves and particles “wave-particle duality”  Photon - particle of light that carries a quantum of energy

Quantum Mechanical Model Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).

Modern View The atom is mostly empty space Two regions  Nucleus  protons and neutrons  Electron cloud  region where you might find an electron Also called the electron cloud model

Modern View of Atom e-e- e-e- Ground state Excited state Electrons can only be at specific energy levels, NOT between levels.

Models of the Atom Dalton’s model (1803) Thomson’s plum-pudding model (1897) Rutherford’s model (1909) Bohr’s model (1913) Charge-cloud model (present) Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 125 Greek model (400 B.C.) e e e e ee e e e e

C. Johannesson Electrons as Waves Louis de Broglie (1924)  Applied wave-particle theory to e -  e - exhibit wave properties QUANTIZED WAVELENGTHS

C. Johannesson Quantum Mechanics Heisenberg Uncertainty Principle  Impossible to know both the velocity and position of an electron at the same time

C. Johannesson Quantum Mechanics Schrödinger Wave Equation (1926)  finite # of solutions  quantized energy levels  defines probability of finding an e -

Quantum Theory quantum theory-  Describes mathematically the wave properties of electrons and other small particles orbital- a region of an atom in which there is a high probability of finding electrons Today’s atomic model predicts quantized, or particular energy levels for electrons. does not describe the exact path or location electrons take or can be found around the nucleus concerned with the probability, or likelihood, of finding an electron in a certain position Two electrons can occupy each orbital, also called an electron cloud.

Quantum Numbers Four Quantum Numbers:  Specify the “address” or “seat” of each electron in an atom UPPER LEVEL Courtesy Christy Johannesson

Quantum Numbers Principal Quantum Number n 1. Principal Quantum Number ( n )  Energy level (ladder rungs)  Size of the orbital  Positive integer 1s1s 2s2s 3s3s

Quantum Numbers Principal Quantum Number 1. Principal Quantum Number  > number, further away from the nucleus  1- right next to the nucleus  3- further away from nucleus  > number, higher the energy level  n = 2 greater energy level than n = 1  these electrons have more energy than electrons in the n = 1 level 1s1s 2s2s 3s3s

Quantum Numbers Angular Momentum Quantum # l 2. Angular Momentum Quantum # ( l )  Energy sublevel (orbital)  Shape of the orbital  Often represented by letters than numbers s p d f Courtesy Christy Johannesson

Quantum Numbers pxpx pzpz pypy x y z x y z x y z

d-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 336

Quantum Numbers Orbitals combine to form a spherical shape. 2s 2p z 2p y 2p x Courtesy Christy Johannesson

Quantum Numbers Magnetic Quantum Number m l 3. Magnetic Quantum Number ( m l )  Orientation of orbital  Specifies the exact orbital within each sublevel

Shapes of s, p, and d-Orbitals

Quantum Numbers Spin Quantum Number 4. Spin Quantum Number ( m s )  Electron spin  +½ or -½  An orbital can hold 2 electrons that spin in opposite directions. Courtesy Christy Johannesson

Maximum Capacities of Subshells and Principal Shells n n l Subshell designation designation s s p s p d s p d f Orbitals in subshell subshell Subshell capacity capacity Principal shell capacity capacity n 2 Hill, Petrucci, General Chemistry An Integrated Approach  1999, page 320

Filling Rules for Electron Orbitals Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for. Pauli Exclusion Principle: An orbital can hold a maximum of two electrons. To occupy the same orbital, two electrons must spin in opposite directions. Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results. *Aufbau is German for “building up”

Diagonal Rule 1s2s3s4s5s6s7s1s2s3s4s5s6s7s 2p3p4p5p6p 2p3p4p5p6p 3d4d5d6d 3d4d5d6d 4f5f 4f5f

General Rules Pauli Exclusion Principle Pauli Exclusion Principle  Each orbital can hold TWO electrons with opposite spins. Courtesy Christy Johannesson

RIGHT WRONG General Rules Hund’s Rule Hund’s Rule  Within a sublevel, place one electron per orbital before pairing them. Courtesy Christy Johannesson

Orbital Diagrams and Electron Configurations Orbital diagrams  Show how electrons are distributed within sublevels  Electrons represented by an arrow  Orbital is represented by a box  Direction of spin represented by direction of arrow Electron configuration  Abbreviated form of orbital diagram

Orbital Diagrams and Electron Configurations H1 e - Orbital diagram 1s Electron configuration 1s 1 ↑

O 8e - Orbital Diagrams and Electron Configurations Orbital Diagram Electron Configuration 1s 2 1s 2 2s 2 2s 2 2p 4 1s 2s 2p Courtesy Christy Johannesson

Ne 10e - Orbital Diagrams and Electron Configurations Orbital Diagram Electron Configuration 1s 2 1s 2 2s 2 2s 2 2p 6 1s 2s 2p Courtesy Christy Johannesson

neon's electron configuration (1s 2 2s 2 2p 6 ) Noble Gas Configuration [Ne] 3s 1 third energy level one electron in the s orbital orbital shape Na = [1s 2 2s 2 2p 6 ] 3s 1 electron configuration A B C D

Shorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 Electron Configuration Longhand Configuration Courtesy Christy Johannesson S