Modern Atomic Theory - Mrs. Ellis' Science Class!

Modern Atomic Theory - Mrs. Ellis' Science Class!

Modern Atomic Theory Chapter 11 The ELECTRON: Wave Particle Duality No No familiar familiar conceptions conceptions can can be be woven woven around around the the electron.

electron. Something Something unknown unknown is is doing doing we we dont dont know know what. what. -Sir -SirArthur ArthurEddington Eddington The

TheNature Natureof ofthe thePhysical PhysicalWorld World(1934) (1934) The Dilemma of the Atom Electrons outside the nucleus are Electrons outside the nucleus are attracted attracted to to the the protons

protons in in the the nucleus nucleus Charged particles moving in curved Charged particles moving in curved paths paths lose lose energy energy What keeps the atom from collapsing? What keeps the atom from collapsing? Wave-Particle Duality JJ Thomson won the Nobel prize for describing the electron as a particle.

His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave! Electromagnetic Electromagnetic Radiation Radiation Most subatomic particles behave

as PARTICLES (photons) and obey the physics of waves. The Wave-like Electron The The electron electron propagates propagates through through space space as as an an energy energy wave. wave. To To

understand understand the the atom, atom, one one must must understand understand the the behavior behavior of of electromagnetic electromagnetic waves. waves. Louis deBroglie

Electromagnetic radiation propagates through space as a wave moving at the speed of light. c = c = speed of light, a constant (3.00 x 108 m = frequency, in units of hertz (hz, sec-1) = wavelength, in meters The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation. E = h

E = Energy, in units of Joules (kgm2/s2) h = Plancks constant (6.626 x 10-34 Js) = frequency, in units of hertz (hz, sec-1) Often times in calculations, you will use the following formula (remember c=v):): E = hc ): Long Wavelength = Low Frequency = Low ENERGY

Short Wavelength = High Frequency = High ENERGY Answering the Dilemma of the Atom Treat electrons as waves Treat electrons as waves As the electron moves toward the As the electron moves toward the

nucleus, nucleus, the the wavelength wavelength shortens shortens Shorter wavelength = higher Shorter wavelength = higher energy energy Higher energy = greater distance Higher energy = greater distance from from the the nucleus nucleus

Electromagnetic radiation (light) is divided into various classes according to wavelength. Tell me what you know Which color has the greatest wavelength? Which color has the shortest wavelength? Electromagnetic Electromagnetic Spectrum Spectrum Long wavelength --> small frequency --> LOW ENERGY Short wavelength --> high frequency --> HIGH ENERGY increasing

frequency increasing wavelength Electromagnetic Spectrum In increasing energy (from low to high), ROY G BIV Back to Light as energy Photons carry energy E=hc This energy can mean different things

Excited State atom with excess energy Ground State lowest possible energy state Wavelengths of light carry different amounts of energy per photon Only certain types of photons are produced (see only certain colors) Quantized only certain energy levels (and therefore colors) are

allowed Emission of Energy by Atoms When atoms receive energy from some source, they become excited, and they can release this energy by emitting light by releasing a photon. The energy of the photon corresponds exactly to the energy change experienced by the emitting atom Red=lower energy=longer wavelength Violet=higher energy=shorter wavelength Some Definitions

Ground State: Lowest possible energy state of an atom Excited States: Possible Higher energy states. Only certain states are allowed for certain atoms. Figure 11.8: An excited lithium atom emitting a photon of red light to drop to a lower energy state. Copyright Houghton Mifflin

Company Flame Tests Many elements give off characteristic light which can be used to help identify them. strontium sodium lithium potassium copper

Emission Spectra Atoms in the excited state are unstable. The electrons will fall back to the ground state. When an electron moves from a higher energy level to a lower energy level, energy is released and light is emitted. This energy is emitted in a form of light. The energy (color) of the light depends on how far the electron falls. 6 5

4 3 2 1 nucleus As e- fall from higher orbitals energy is given off. Amount of energy given off = to the distance of the fall Atomic Fingerprints

Each element has a unique atomic line emission spectrum. Used to identify the element Continuous and Line Spectra Visible Light Spectrum Na H Ca Hg o 4000 A

5000 6000 7000 Flame Emission Spectra Photographs of flame tests of burning wooden splints soaked in different salts. methane gas wooden splint sodium ion

calcium ion copper ion strontium ion Lets look more at hydrogen Figure 11.9: A sample of H atoms receives energy from an external source. Copyright Houghton Mifflin Company Figure 11.9: The excited atoms

release energy by emitting photons. Excited atom can release some or all of its energy by emitting a photon (electromagnetic radiation particle) Copyright Houghton Mifflin Company Figure 11.10: An excited H atom returns to a lower energy level. Energy contained in photon = change in energy of atom Copyright Houghton Mifflin Company

Figure 11.11: Colors and wavelengths of photons in the visible region. Visible light photons emitted by Hydrogen always the same Because only certain photons are emitted, only certain energy changes are occurring Hydrogen atom has certain discrete energy levels Energy levels of Hydrogen are quantized only certain Copyright Houghton Mifflin Company values allowed Figure 11.12: The color of the photon emitted depends on the energy change that produces it.

Copyright Houghton Mifflin Company Figure 11.13: Each photon emitted corresponds to a particular energy change. Copyright Houghton Mifflin Company Figure 11.14: Continuous (a) and discrete (b) energy levels. Quantized nature of energy

surprised scientists (b) Previously assumed atom could exist at any energy level (a) Copyright Houghton Mifflin Company Figure 11.15: The difference between continuous (a) and quantized (b) energy levels. Ramp can be at any elevation Staircase can move from one step to another or even skip, but must be on a step

Copyright Houghton Mifflin Company Lets look at Old Atom Models and Compare them to Newer Ones OLD Rutherford Bohr NEW de Broglie

Schrodinger Figure 11.1: The Rutherford atom. Copyright Houghton Mifflin Company Rutherford Atom Review Alpha particle/Gold foil experiment Nuclear Atom Nucleus composed of protons & neutrons Nucleus small compared to atomic size Electrons account for rest of atom

Unanswered questions: What are electrons doing? How are they arranged & how do they move? Thought electrons revolved around nucleus like planets orbit the sun Couldnt explain why electrons arent attracted to protons causing atom to collapse Copyright Houghton Mifflin Company

Limitations of Rutherfords Atomic Model It explained only a few simple properties of atoms. It could not explain the chemical properties of elements. For example, Rutherfords model could not explain why an object such as the iron scroll shown here first glows dull red, then yellow, and then white when heated to higher and higher temperatures. The Bohr Model In 1913, Niels Bohr (18851962), a

young Danish physicist and a student of Rutherford, developed a new atomic model. He changed Rutherfords model to incorporate newer discoveries about how the energy of an atom changes when the atom absorbs or emits light. The Bohr Model According to Bohrs atomic model, electrons move in definite orbits around the nucleus, much like planets circle the sun. These orbits, or energy levels, are located

at certain distances from the nucleus. The energy levels are also called electron shells. Figure 11.17: The Bohr model of the hydrogen atom. Electrons moved in circular orbits like planets Electrons could jump from one orbit to another by emitting/absorbing a photon Only worked for H,

didnt work for other atoms Showed experimentally to be incorrect Copyright Houghton Mifflin Company Paved way for other theories The Modern Model of the Atom Accredited to Schrodinger,

According to the wave mechanic model, electrons do not move about an atom in a definite path, like the planets around the sun. Electrons are NOT in fixed paths Electrons are in orbitals which are nothing like Bohrs orbits The electrons move constantly throughout the energy levels forming an electron cloud

Electron Cloud The space in which electrons are most likely to be found. The cloud is more dense where the probability of finding the electron is high. We cant know exactly where an electron is (Heisenberg Uncertainty Principle) Electron cloud Heisenberg Uncertainty Principle One

One cannot cannot simultaneously simultaneously determine determine both both the the position position and and momentum momentum of of an an electron. electron. You can find out where the electron is,

but not where it is going. OR Werner Heisenberg You can find out where the electron is going, but not where it is! Figure 11.18: A representation of the photo of the firefly experiment (lightning bugs). Shows probability (or likelihood) of where firefly will be found

Usually near the center, but can be found in any of the shaded areas at any time Copyright Houghton Mifflin Company Figure 11.19: The orbital that describes the hydrogen electron in its lowest

possible energy state. Darker pink = greater probability Copyright Houghton Mifflin Company Drawbacks of wave mechanical model: Gives no information about when the electron occupies a certain point in space or how it moves We will probably never know the details of

electron motion Confident that Bohr model is incorrect Copyright Houghton Mifflin Company 11 50 Quantum Mechanical Model Quantum Mechanical Model Parts of the Wave Mechanical Model

1. Principle Energy Level (n) energy level designated by numbers 1-7. -called principle quantum numbers 1 2 3 4 5 6 7 2. Sublevel exist within each principle energy level -the energy within an energy level is slightly different -each electron in a given sublevel has the same

energy s p d -lowest sublevel = s, then p, then d, then f f Parts of the Wave Mechanical Model cont. 3. Orbital region within a sublevel or energy level where electrons can be found s sublevel 1 orbital p sublevel 3 orbitals d sublevel 5 orbitals f sublevel 7 orbitals

- ** No more than two electrons can occupy an orbital** -an orbital can be empty, half-filled, filled Shapes of orbitals All s orbitals are spherical as the principle energy level increases the diameter increases. All p orbitals are dumbbell or figure-8 shaped all have the same size and shape within an energy level 4 of the d orbitals are 4-leaf clover shaped and the last is a figure-8 with a donut all have the same size and shape within an energy level

f orbitals are complicated!!!!! Summary Principle Energy Level # of sublevels # of orbitals present s p d f

Total # of orbitals Maximum # of electrons 1 1 1- - - 1

2 2 2 13- - 4 8 3 3

9 18 4 4 135 135 7 16 32 Sublevels in Each Energy Level

Principal Energy Level # of Sublevels 1 1 (s) 2 2 (s,p) 3 3 (s,p,d) 4 4 sublevels (s,p,d,f)

Orbitals in Each Sublevel Sublevel s No. of Orbitals 1 p d f 3 5 7

No. of Electrons 2 6 10 14 Electron Configuration arrangement of the electrons among the various orbitals of the atom Ex: 1s22s22p6 = Neon Sulfur 1s22s22p6 3s23p4 = Cd =1s22s22p6 3s23p6 4s2 3d10 4p6 5s24d10 Na =1s22s22p6 3s1 N

a N e Orbital filling table Electron configuration of the elements of the first three series *Examples: Write the electron configuration using sublevels. 1. Hydrogen: ______________________ 2. Carbon: _________________________ 3. Phosphorous: ______________________________________ 4. Potassium: _______________________________________

5. Calcium:___________________________________________ 6. Iodine: __________________________________________ **NOTE 1: Superscripts should ALWAYS add up to the # of electrons the atom has! (This is a good way of checking yourself!!!** **NOTE 2: When electrons are in an excited state, they will jump to a higher energy level before all the orbitals in the lower energy level are completely filled. Sample Problems Which is the electron configuration of an atom in the excited state? a. 1s22s22p2 b. 1s22s22p1 c.

1s22s22p53s2 d. 1s22s22p63s1 An atom in the excited state can have an electron configuration of? a. 1s22s2 b. 1s22p1 c. 1s22s22p5 d. 1s22s22p6 Which electron configuration represents a potassium atom in the excited state?

e. 1s22s22p63s23p3 f. 1s22s22p63s13p4 g. 1s22s22p63s23p64s1 h. 1s22s22p63s23p54s2 Electron Spin Spin motion that resembles earth rotating on its axis clockwise or counterclockwise Pauli Exclusion Principle two electrons in the same orbital must have opposite spins Hunds Rule All orbitals within a sublevel must contain at least one electron before any orbital can have two

Orbital Diagram describes the placement of electrons in orbitals use arrows to represent electrons with spin line represents orbital (s=1, p=3, d=5, f=7) ____ full ____ half-full ____ empty Aufbau Order Aufbau Order Tool to predict the order in which sublevels will fill. Aufbau Principle:

Must fill lower energy levels first OR use order on Periodic Table Orbital Diagrams Neon = 1s__2s__ 2p__ __ __ Carbon = 1s__2s__ 2p__ __ __ Zinc =1s__2s__ 2p__ __ __ 3s__ 3p__ __ __ 4s__ 3d__ __ __ __ __ Gallium = 1s__2s__ 2p__ __ __ 3s__ 3p__ __ __ 4s__ 3d__ __ __ __ __ 4p__ __ __ Atomic # (= # of electrons)

Symbol Orbital Filling Diagram Electron Configuratio n 10 Ne Atomic # (= # of electrons)

Symbol Orbital Filling Diagram Electron Configuratio n 5 B Atomic # (= # of electrons)

Symbol Orbital Filling Diagram Electron Configuratio n 7 N Atomic # (= # of electrons)

Symbol Orbital Filling Diagram Electron Configuratio n 8 O Noble Gas Configuration Shorthand configuration that substitutes a noble gas for

electrons. Ex: Na = 1s 2s 2p 3s or [Ne]3s 2 2 6 1 1 Sn = 1s22s22p63s23p64s23d104p65s24d105p2

or [Kr]5s24d105p2 Valence Electrons Electrons in the outermost (highest) principle energy level in an atom, electrons use in bonding (outside of brackets) Core Electrons innermost electrons not involved in bonding (inside brackets) Valence Configuration shows just the valence electrons Na = 3s 3 Shell/1valence electron Ex: Sn = 5s 5p 5 Shell/4 valence electrons

1 2 rd 2 th Element Lithium Configuration notation

Orbital notation 1s22s1 [He]2s1 ____ 1s Beryllium Oxygen ____ 2s

____ ____ 2p ____ ____ 2s ____ ____ 2p ____

1s22s22p2 [He]2s2p2 ____ 2s ____ ____ 2p ____ 1s22s22p3

[He]2s2p3 ____ ____ 1s 2s ____ ____ ____ 2p

1s22s22p4 [He]2s2p4 ____ 2s ____ ____ 2p ____ 1s22s22p5

[He]2s2p5 ____ 1s Neon ____ [He]2s2p1 ____ 1s Fluorine ____ 2p

1s22s22p1 ____ 1s Nitrogen ____ [He]2s2 ____ 1s Carbon ____

2s 1s22s2 ____ 1s Boron Noble gas notation ____ 2s ____

____ 2p ____ 1s22s22p6 [He]2s2p6 ____ 1s ____ 2s ____

____ 2p ____ Figure 11.34: Periodic table with atomic symbols, atomic numbers, and partial electron configurations. Copyright Houghton Mifflin Company 11

76 Ion Ion Configurations Configurations To form anions from elements, add 1 or more efrom the highest sublevel. P [Ne] 3s2 3p3 + 3e- ---> P3- [Ne] 3s2 3p6 or [Ar] 3p 3p 3s 3s 2p

2p 2s 2s 1s 1s Ions Cations: Take away that many electrons as the number on the charge (+1 means take away highest energy electron in configuration) Anions: Add that many electrons as the number on the charge. Abide by Aufbaus

Principle

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