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Chapter 5 - Electrons In Atoms 5.1 Light and Quantized Energy 5.2 Quantum Theory and the Atom 5.3 Electron Configurations

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Section 5.1 Light and Quantized Energy Know the names and how to recognize the parameters that describe an electromagnetic wave (period, wavelength, amplitude, frequency, crest, etc.) Perform calculations involving c = Describe the experiments and the interpretation of both Planck’s experiment and the photoelectric effect. Compare the wave and particle natures of light. Light, a form of electronic radiation, has characteristics of both a wave and a particle.

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Section 5.1 Light and Quantized Energy Define a quantum of energy and explain how it is related to an energy change of matter. Contrast continuous electromagnetic spectra and atomic emission spectra and provide examples of each.

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Key Concepts All waves are defined by their wavelengths, frequencies (periods), amplitudes, and speeds. In a vacuum, all electromagnetic waves travel at the speed of light. The relationship between this speed and the wavelength and frequency is c = λν All electromagnetic waves have both wave and particle properties. This is known as wave-particle duality. Matter emits and absorbs energy in quanta. The amount of energy can be calculated using E quantum = hν White light produces a continuous spectrum. An element’s emission spectrum consists of a series of lines of individual colors. Section 5.1 Light and Quantized Energy

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Nuclear Atom - Unanswered Questions Rutherford model said that: Positive charge and virtually all mass concentrated in nucleus Most of volume occupied by electrons But still unknown Spatial arrangement of electrons? Why doesn’t positive nucleus pull negative electrons into itself?

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How to Explain Chemical Behavior? Elements with AN = 17, 18, 19 Chlorine – highly reactive gas Argon – unreactive gas Potassium – very reactive metal

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Wave Nature of Light Electromagnetic “radiation” Exhibits wavelike behavior Consists of oscillating (periodically varying) magnetic and electric fields Oscillations to direction of travel

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Moving Electromagnetic Wave Magnetic Field Electric Field Wavelength ( ) Propagation direction

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Wave Characteristics Wavelength (lambda)m Frequency (nu)Hz Period T s Speed c (for EM waves; m/s speed of light) Amplitude [symbol not used for this course]

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Wavelength Shortest distance between equivalent points on a continuous wave View of wave frozen in space: x axis - distance

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Wavelength Has unit of length m cm nm etc.

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Frequency Number of waves that pass given point per second SI unit: Hz Waves per second In calculations, use 1/s or s -1 Period (T) T = 1/ Units of time (s) Time between crests for wave sampled over time at one point in space

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Wave period View of wave sampled over time at 1 point in space: x axis = time (oscilloscope trace) Period

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EM Wave Speed If electromagnetic (EM) wave in vacuum, travels at speed of light, c 3.00 10 8 m/s All EM waves in vacuum travel at c Doesn’t vary with wavelength, frequency, or amplitude of wave For all practical purposes, air is equivalent to a vacuum for EM wave

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Frequency / Wavelength For EM waves in vacuum c = c constant, so as increases, must decrease and vice versa

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Frequency / Wavelength c = Red light versus blue light Lower Frequency Higher Frequency Longer Shorter

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Amplitude Wave height measured from origin to crest (or trough) Not concerned about units for amplitude in this course

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EM Wave Energy Energy increases with increasing frequency (or decreasing wavelength) High energy Gamma rays X-rays Low energy Radio waves

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Electromagnetic Spectrum Visible Light Increasing Energy Frequency (Hz)

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Calculating Wavelength Practice problem 5.1, page 140 Wavelength ( ) of microwave having frequency ( ) = 3.44 10 9 Hz? c = = c / = 3.00 10 8 m/s 3.44 10 8 s -1 = 8.72 m

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Practice Electromagnetic Waves Practice problems, page 140 Problems 1-4 Chapter assessment page 166 Problems (concepts) Problems 45 – 48 Supplemental Problems 1-2, page 978

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History of Development of Human Understanding of the Atom Dalton Thomson Rutherford BohrQuantum

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Pre 1900 View of Matter & Energy Matter and EM energy are distinct Matter consists of particles Have mass Position in space can be specified Energy in form of EM radiation is wave Massless Delocalized – position in space can’t be specified

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Pre 1900 View of Matter & Energy View of energy also included idea that energy can be gained or lost in continuous manner Heating water up with Bunsen burner – can change temperature by any arbitrary amount simply by changing size of flame and/or heating time

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Pre 1900 View of Matter & Energy Attempt to explain results of 2 key experiments challenged neat particle/wave division and idea of continuous energy transfer Change in intensity and wavelength of radiation emitted by heated objects as a function of temperature Emission of electrons from metals when certain frequencies of light shown on metal

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EM Radiation of Heated Objects Heated objects emit differing wavelengths of light depending upon temperature Max Planck measured (~1900) energy vs wavelength profile at different temperatures

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EM Radiation of Heated Objects Predictions of classical theory based on wave model fails to explain intensity profile

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Quantum Nature of Energy Based on analysis of heated object data, Planck concluded that matter can gain or lose energy only in small increments call quanta Single quantum minimum amount of energy atom can lose or gain

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Quantum Nature of Energy Energy of quantum related to frequency of radiation E quantum = h h = Planck’s constant = x J s If know energy, can get frequency Consistent with idea that high frequency EM radiation has high energy

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Quantum Nature of Energy E quantum = h Above gives magnitude of quantum of energy Planck’s theory: Energy transfer only happens in increments = integer multiple of E quantum E transfer = n h n = 1, 2, …

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Photoelectric Effect Measure number and energy of electrons ejected when light of certain wavelength strikes metal surface Beam of light Metal surface Electrons Nuclei Electron ejected from surface

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Photoelectric Effect

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Classical wave model (continuous energy) prediction: Given enough time, low frequency (= low energy) light will eventually transfer enough energy to metal to eject an electron Analogous to heating water model Actual (quantum) results No electrons unless > threshold

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Photoelectric Effect Metal only ejects electrons when energy (frequency) of light above minimum threshold Even dim light above threshold ejects electrons Increasing intensity of light of given frequency above threshold ejects more electrons but energy of electrons same Increasing light frequency above threshold ejects higher energy electrons but not more electrons

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Photoelectric Effect Basis for Einstein proposal (1905) that photons have both a wavelike and particle nature Einstein extended Planck’s quantum idea to photons E photon = h Photon = particle (packet) of EM energy with no mass that carries a quantum of energy

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Wave/Particle Views Light as a wave phenomenon Light as a stream of photons

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Calculating Energy of Photon Practice problem 5.2, page 143 Energy of photon from violet part of rainbow if = 7.23 s -1 ? E photon = h h = x J s E photon = (6.626 x J s)(7.23 s -1) = 4.79 J

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Practice Photon Energy Practice problems, page 143 Problems Section assessment page 145 Problem 13 Chapter assessment page 166 Problems (concepts) Problems 49 – 53, 55 – 57 Supplemental, page 978 Problems 3 - 4

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Atomic Emission Spectra The fact that only certain colors are seen in fireworks, neon signs, etc. is a further indication that energy can’t come out of an atom in arbitrary amounts

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Atomic Emission Spectra Fact that only certain colors seen means only certain distinct frequencies are possible Energy = n h Strontium Copper

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Discrete Spectrum Continuous Spectrum

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Atomic Emission Spectra Use optical spectroscopes and diffraction glasses to view: Light from fluorescent tubes Incandescent light from overhead projector Emission from spectrum tubes (H, He)

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Atomic Emission Spectra of Hydrogen Prism Slit Hydrogen gas discharge tube Spectrum

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Hg, Sr, H Spectra Comparison

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Emission / Absorption Spectra - Na

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Emission / Absorption Spectra – He Fig 5.9

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5.1 Light and Quantized Energy 5.2 Quantum Theory and the Atom 5.3 Electron Configurations Chapter 5 - Electrons In Atoms

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Section 5.2 Quantum Theory and the Atom Compare the Bohr and quantum mechanical models of the atom. Describe the process of atomic emission and calculate the wavelength of an emitted photon given the energy levels Explain the impact of de Broglie's wave particle duality and the Heisenberg uncertainty principle on the current view of electrons in atoms. Identify the relationships among a hydrogen atom's energy levels, sublevels, and atomic orbitals. Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals.

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Key Concepts Bohr’s atomic model attributes hydrogen’s emission spectrum to electrons dropping from higher-energy to lower-energy orbits. ∆E = E higher-energy orbit - E lower-energy orbit = E photon = hν The de Broglie equation relates a particle’s wavelength to its mass, its velocity, and Planck’s constant. λ = h / mν The quantum mechanical model of the atom assumes that electrons have wave properties. Electrons occupy three-dimensional regions of space called atomic orbitals. Section 5.2 Quantum Theory and the Atom

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Rutherford Model of Atom - Limitations Accelerated charged particles will radiate EM waves Electron in orbit; therefore accelerated Should lose energy and spiral inwards to nucleus EM waves This doesn’t happen!

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Bohr Model of Atom 1913, Niels Bohr proposed that H atom has only certain allowable (quantized) energy states Lowest state = ground state Energy gains promote electrons in atom to excited state Electrons confined to distinct circular orbits Smaller orbits lower energy

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Bohr’s Model (nucleus way out of scale) Nucleus Electron Orbit Energy Levels orbit radius

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Quantum Staircase Emission Absorption

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Bohr Atom Picture of Energy Transfer Photon AbsorptionPhoton Emission Electron: Red Photon: Orange

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Emission Spectrum of Hydrogen

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Bohr Model of Atom 3 Series of Atomic Emission Lines Visible Series (Balmer) final state n=2 UV Series (Lyman) final state n=1 Infrared Series (Paschen) final state n=3

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Bohr Energy Formula (not in book) Energy change = E = E (final) – E(initial) = E photon = h R H = Rydberg constant = 2.18 J E < 0 (Energy at infinite separation = 0)

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Bohr Freq. Formula (not in book) Energy transitions for hydrogen’s four visible spectra lines - Balmer series

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Bohr Model of Atom - Limitations Explained emission spectra of hydrogen very well Failed to explain spectrum of other elements Did not fully account for chemical behavior of atoms

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Bohr Model of Atom Has been shown that this model is fundamentally incorrect – electrons not particles in fixed orbits (classical model)

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Schematic for X-ray Diffraction X-ray beam with continuous range of wavelengths incident on crystal Diffracted radiation intense in certain directions Intense spots correspond to constructive interference from waves reflected from layers of crystal Diffraction pattern detected by photographic film

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Electrons as Waves - Evidence Diffraction observed when electrons with sufficient momentum strike an ordered crystal lattice Electrons Nickel Crystal Detector Screen Diffraction pattern on detector screen

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Comparing X-Ray and Electron Diffraction Patterns in Al Foil X-RayElectron Electrons - particles with mass and charge - create diffraction patterns in a manner similar to electromagnetic waves!

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Electrons as Waves Louis De Broglie (1924) Wavelength associated with particle of mass m moving at velocity v = h/ mv de Broglie equation All moving particles have wave characteristics

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Idea for Particle as Wave Vibrating guitar string Only multiples of /2 allowed Orbiting electron Only multiples of allowed

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Heisenberg Uncertainty Principle 1927 Fundamentally impossible to know precisely both velocity and position of particle at same time Not just a technical limitation

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Finding Out Where an Electron Is Act of measuring changes properties To determine electron location, use light as probe But light moves electron And hitting the electron changes the frequency of the light

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Heisenberg Uncertainty Principle Impact of photon on knowledge of location and velocity of electron in an atom Before collision After collision

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Quantum Mechanical Model of Atom 1926 – Schrödinger wave equation Treated hydrogen atom’s electron as a wave Unlike Bohr model, worked for other atoms besides hydrogen Also limited electron’s energy to quantized values Makes no attempt to describe electron’s path around the nucleus

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Bohr Model According to Bohr’s atomic model, electrons move in definite orbits around nucleus, much like planets circle sun These orbits, or energy levels, are located at certain fixed distances from nucleus

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Wave Model (Electron Cloud) Quantum mechanical (wave / electron cloud) model of atom specifies only probability of finding electron in certain regions of space

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Marble Model Plum Pudding Model Nuclear Model Planetary (Bohr) Model Quantum Mechanical (wave / electron cloud) Model

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Classical to Quantum Theory IndivisibleElectronNucleusOrbitElectron Cloud Greek X Dalton X Thomson X Rutherford X X Bohr X X X Wave X X X

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Summary Major Observations & Theories Leading from Classical to Quantum Theory ObservationScien tist Theory Spectral shape of blackbody radiation Planck 1900 Energy quantized Photoelectric effect Einstein 1905 Light has particle behavior (photons)

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Summary Major Observations & Theories Leading from Classical to Quantum Theory ObservationScien tist Theory Atomic line spectra Bohr 1913 Energy of atoms quantized; photon emitted in orbit energy transition Bohr model works for H de Broglie 1924 All moving particles have wave-like nature

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Summary Major Observations & Theories Leading from Classical to Quantum Theory Observ ation Scien tist Theory ?Schrodinger 1926 Quantum mechanical description of atom using wave function ?Heisenberg 1927 Fundamentally impossible to know precisely both velocity & position at same time

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Wave Function (From Schrödinger Wave Equation) is wave function (solution to wave equation) Square of gives probability of finding electron within particular volume of space around nucleus

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Wave Function and Orbital (wave function) defines atomic orbital – 3D description of electron’s probable location Entire family of wave functions exists, each having particular set of quantum numbers “n” example of a quantum number Quantum numbers determine electron energy and shape/size of probability distribution

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Nucleus found inside blurry “electron cloud” Orbital describes chance of finding electron in a region Draw line/surface at 90% probability Shape may be complex Quantum Mechanical Model

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Electron Density – Hydrogen Atom AB A – likelihood of finding electron at particular point dot density B – Orbital boundary: volume encloses 90% probability of finding electron inside

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Radial Probability Distribution Distance from nucleus (r) Most probable distance

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Hydrogen Atom Schrödinger wave equation can be analytically solved for the H atom Energy levels same as Bohr model – also labeled by “n” Position of electron no longer described by circular orbit Position specified by probability only – details described by orbital

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Electron Density – Hydrogen Atom Some (small) probability electron will be found a large distance from nucleus or very close to nucleus

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Hydrogen’s Atomic Orbitals Quantum mechanical model assigns principal quantum number (n) Indicates relative size and energy of orbitals As n increases, electron spends more time farther from nucleus n=1 is lowest = ground state

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Hydrogen’s Atomic Orbitals Principal energy levels contain energy sublevels Labeled as s, p, d, or f also g, etc but not concerned with in this course Sublevel value determines orbital shape

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Hydrogen’s Atomic Orbitals n = 1 Only one s sublevel with one s orbital n = 2 s and p sublevels; three p orbitals in p sublevel n = 3 s, p, and d sublevels; 5 d orbitals in d sublevel n = 4 s, p, d and f sublevels; 7 f orbitals in f sublevel

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Hydrogen’s Atomic Orbitals For hydrogen only (special case), all sublevels of a given principal quantum number n have the same energy For n=2, the 2s and 2p sublevels have the same energy For n=3, the 3s, 3p, and 3d sublevels have the same energy Special case means energy levels have same pattern as Bohr atom model

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Hydrogen’s Atomic Orbitals For a given sublevel, energy increases with increasing n Energy: 5p > 4p > 3p >2p (no 1p !)

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Orbital Energies for Hydrogen Atom (Aufbau diagram)

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Orbital Notation n sublevel direction (as subscript) n = principal quantum number sublevel = s, p, d, or f direction = not applicable for s x, y, z for p xy, xz, yz, x 2 -y 2, z 2 for d forget it for f

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Orbital Notation n sublevel direction (as subscript) 1s 3s 2p x 3d xy 4p z 4d x 2 -y 2

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s and p Orbitals ( p shape exaggerated)

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Hydrogen 1s, 2s, 3s Orbitals 1s 2s 3s 1s 2s 3s Node

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1s 2s 3s

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2p z

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d Orbitals (also exaggerated) The odd one (different shape)

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3d xz

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3d z 2

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Hydrogen’s Atomic Orbitals A given orbital 2s, 3p x, 4d yx, 5s, etc can be occupied by at most two electrons However, hydrogen has only one electron to worry about

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Summary of Sublevels s p d f # of orbitals Max # electrons Starts at energy level

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Summary by Energy Level 1 st Energy Level n=1 Only s orbital Holds 2 electrons 1s 2 2 total electrons =2n 2 2 d Energy Level n=2 s and p orbitals are available 2 in s, 6 in p 2s 2 2p 6 8 total electrons =2n 2

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By Energy Level 3 d energy level n=3 s, p, and d orbitals 2 in s, 6 in p, and 10 in d 3s 2 3p 6 3d total electrons =2n 2 4 th energy level s,p,d, and f orbitals 2 in s, 6 in p, 10 in d, and 14 in f 4s 2 4p 6 4d 10 4f total electrons=2n 2

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Orbital Summary for Hydrogen

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5.1 Light and Quantized Energy 5.2 Quantum Theory and the Atom 5.3 Electron Configurations Chapter 5 - Electrons In Atoms

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Section 5.3 Electron Configuration Apply the Pauli exclusion principle, the aufbau principle, and Hund's rule to write electron configurations using orbital diagrams and electron configuration notation (including noble gas notation). Define valence electrons, and draw electron-dot structures representing an atom's valence electrons. A set of 3 rules determines the arrangement of electrons in an atom. This arrangement is called the electron configuration.

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Key Concepts The arrangement of electrons in an atom is called the atom’s electron configuration. Electron configurations are defined by the aufbau principle, the Pauli exclusion principle, and Hund’s rule. An element’s valence electrons determine the chemical properties of the element. Electron configurations can be represented using orbital diagrams, electron configuration notation, and electron-dot structures. Section 5.3 Electron Configuration

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Rules for Filling Orbitals (1) Aufbau principle – one by one build up Each electron occupies lowest energy orbital available Energy level order determined from diagram

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Atomic Orbital Energies For a given principal quantum number n, all sublevels of a given type have the same energy (said to be degenerate) All three p orbitals for n=2 are degenerate All five d orbitals for n=3 are degenerate For hydrogen, all sublevels are degenerate Levels are same as Bohr atom levels

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Aufbau Diagram for Hydrogen Atom

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Orbitals in Many-Electron Atoms For n 2, the s- and p-orbitals are no longer degenerate because the electrons interact with each other Unlike case for hydrogen, 3d > 3p >3s Aufbau diagram looks different for many-electron systems No longer follows simple Bohr model

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Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Aufbau Diagram – Multi-Electron Atoms

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Orbital Energies: Multi-Electron Atoms

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Aufbau Diagram Features – Table 5.3

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Orbital Filling Order (see page 160)

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Spin of the Electron Associated with electrons is a property called spin Spin generates a magnetic field which can be oriented up or down Use or to indicate

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Rules for Filling Orbitals (2) Pauli Exclusion Principle A maximum of two electrons may occupy a single atomic orbital provided the electrons have opposite spins

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Rules for Filling Orbitals (3) Hund’s Rule – how to handle degenerate orbitals Single electrons with same spin must occupy each degenerate orbital before additional electrons with opposite spins share an orbital Rule arises because increased electron-electron repulsion occurs when 2 electrons occupy the same orbital

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Hund’s Rule for p Orbitals Second electron goes into empty degenerate orbital with spin in same direction as first All degenerate orbitals filled – can start pairing now

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Electron Configurations Periods 1, 2, and 3 (only) Three rules: Electrons fill orbitals starting with lowest n and moving upwards (Aufbau) No two electrons can fill one orbital with the same spin (Pauli Exclusion Principle) For degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule)

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For 1 electron, box with single arrow Constructing Orbital Diagrams 1s 2s 2p C Use a box for each orbital Label each box with n and sublevel If 2 electrons in 1 orbital, use opposite arrows

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Electron Configuration Notation Indicate n and sublevel for each orbital and the total electron occupancy with a superscript C 1s 2 2s 2 2p 2 Distribution of electrons in the three p orbitals (p x, p y, p z ) not explicit in this notation We write in n order, not energy order Ti 1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 4s 2 (note energy of 3d > 4s) (Textbook uses energy order for configs.)

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1s2s 2p Orbital Diagram/ Electron Configuration H 1s 1 1s He 1s 2 Li 1s 2 2s 1 1s Be 1s 2 2s 2 B 1s 2 2s 2 2p 1 1s2s 1s C 1s 2 2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s 2 2s 2 2p 6

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Electron Configuration Determine the electron configuration for Phosphorus (AN=15) Need to account for 15 electrons

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First two electrons go into the 1s orbital Notice the opposite spins 13 more to go 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Increasing energy

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Next two electrons go into the 2s orbital 11 more to go 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Increasing energy

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Next six electrons go into the 2p orbitals 5 more to go 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Increasing energy

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Next two electrons go into the 3s orbital 3 more to go 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Increasing energy

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Last three electrons go into the 3p orbitals Each go into separate orbitals 3 unpaired electrons 1s 2 2s 2 2p 6 3s 2 3p 3 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Increasing energy

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Using the Aufbau Diagram 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2 electrons

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Fill from bottom up following arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 4 electrons

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1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 12 electrons Fill from bottom up following arrows

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1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 20 electrons Fill from bottom up following arrows

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1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 38 electrons Fill from bottom up following arrows

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1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 56 electrons Fill from bottom up following arrows

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1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 88 electrons Fill from bottom up following arrows

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1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p electrons Maxed out Fill from bottom up following arrows

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Noble Gas Notation 1s 2s 2p 3s1s 2s 2p Na 1s 2 2s 2 2p 6 3s 1 (standard notation) Ne 1s 2 2s 2 2p 6 Na [Ne]3s 1 (noble gas notation)

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Noble Gas Configurations Noble gases always have s and p orbitals completely filled (He has no p) ns 2 np 6 Principal quantum number (n) of these orbitals is same as the period in which the gas is found - p orbitals are the last filled (s for He) He: 1s 2 Ne: 2s 2 2p 6 Ar: 3s 2 3p 6 Kr: 4s 2 4p 6

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Table Noble Gas Configurations How to express noble gas using noble gas configuration

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Determining Electron Configuration Germanium (Ge), a semiconducting element, is commonly used in the manufacture of computer chips. What is the ground state configuration for an atom of germanium using noble gas notation? Atomic number of Ge = 32

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Orbital Filling Order (page 160) [Ar]4s 2 3d 10 4p 2 Above configuration obtained from order of filling. For this class, use principal energy level (n) order [Ar] 3d 10 4s 2 4p 2 32 electrons

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Practice Electron Configurations Practice problems, page 160 Probs 21(a-f), Chapter assessment page 167 Probs 85(a-d), 86(a-d), 87(a-f), 88, 89 Supplemental Problems, page 978 Probs 5(a-d), 6

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Orbital Filling Order Lowest energy to higher energy. Adding electrons can change energy of orbital Half filled sublevels have lower energy Makes them more stable Causes exceptions in the filling order shown on the diagram

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Exceptions to Filling Order Aufbau diagram works to vanadium, AN 23 Half-filled and fully-filled set of d orbitals have extra energy stability, so chromium is Cr [Ar]3d 5 4s 1 (1/2 filled d) Not [Ar]3d 4 4s 2 Next exception is copper: Cu [Ar]3d 10 4s 1 (filled d) Not [Ar]3d 9 4s 2

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Valence Electrons Defined as those electrons in the atom’s outermost orbitals Orbitals with highest n If have a (n-1)d 10 component in the configuration, then ignore these electrons for counting valence electrons, even if higher energy than the ns 2 electrons Zn [Ar]3d 10 4s 2 2 valence electrons, not 12 Br [Ar] 3d 10 4s 2 4p 5 7 valence electrons, not 15 or 17

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Electron-Dot Structures Shorthand notation for indicating valence electrons Write the element’s chemical symbol Add a dot for each valence electron One at a time on all four sides of symbol Then pair them up until all are used Mg [Ne]3s 2 S Mg S [Ne]3s 2 3p 4

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Valence Electrons Determine the chemical (bonding) properties of the element S [Ne]3s 2 3p 4 6 valence electrons Cs [Xe]6s 1 1 valence electron

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Dot Structures for Elements in 2d Period

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Example Problem page 162 What is electron dot structure for tin? Sn: [Kr]4d 10 5s 2 5p 2 4 valence electrons Sn

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Practice Electron-Dot Structures Practice problems, page 162 Probs 26 (a-c), 27, 28 Section 5.3 Assessment, page 162 Prob 33 Chapter assessment pages Probs 81(a-d), 90 (a-e), Supplemental Problems, page 978 Probs 7, 8, 9 (a-d)

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