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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**

Light, a form of electronic radiation, has characteristics of both a wave and a particle. 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.

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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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**Section 5.1 Light and Quantized Energy**

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 Equantum = hν White light produces a continuous spectrum. An element’s emission spectrum consists of a series of lines of individual colors.

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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 Propagation direction Wavelength (l)

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

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EM Wave Speed If electromagnetic (EM) wave in vacuum, travels at speed of light, c 3.00 108 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**

Red light versus blue light Longer l Shorter l Lower Frequency Higher Frequency

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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 104 1022 Frequency (Hz) Increasing Energy

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**Calculating Wavelength**

Practice problem 5.1, page 140 Wavelength (l) of microwave having frequency () = 3.44109 Hz? c = = c / = 3.00108 m/s 108 s-1 = 8.7210-2 m

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**Practice Electromagnetic Waves Practice problems, page 140**

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**

+ - Bohr Quantum + - + - Thomson + - Rutherford Dalton

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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 Equantum = h h = Planck’s constant = x10-34 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**

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

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**Electron ejected from surface**

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

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Photoelectric Effect

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Photoelectric Effect 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 n > 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 Ephoton = 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 = 7.231014 s-1? Ephoton= h h = x10-34 J s Ephoton= (6.626 x10-34 J s)(7.231014 s-1) = 4.7910-19 J

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**Practice Photon Energy Practice problems, page 143 Problems 5 - 7**

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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Continuous Spectrum Discrete 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**

Hydrogen gas discharge tube Prism Slit 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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**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 5.2 Quantum Theory and the Atom**

Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals. 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.

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**Section 5.2 Quantum Theory and the Atom**

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.

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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)**

Electron Orbit Energy Levels orbit radius

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Quantum Staircase Absorption Emission

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**Bohr Atom Picture of Energy Transfer**

Electron: Red Photon: Orange Photon Absorption Photon Emission

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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 = D E = E (final) – E(initial) = E photon = h RH = Rydberg constant = 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**

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 l 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 l/2 allowed Orbiting electron Only multiples of l 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 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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**Planetary (Bohr) Model**

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

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**Classical to Quantum Theory**

Indivisible Electron Nucleus Orbit Electron Cloud Greek X Dalton Thomson Rutherford Bohr Wave

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**Summary Major Observations & Theories Leading from Classical to Quantum Theory**

Scien 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**

Scien 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**

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)**

Y is wave function (solution to wave equation) Square of Y gives probability of finding electron within particular volume of space around nucleus

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**Wave Function and Orbital**

Y(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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**Quantum Mechanical Model**

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

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**Electron Density – Hydrogen Atom**

A – likelihood of finding electron at particular point µ dot density B – Orbital boundary: volume encloses 90% probability of finding electron inside A B

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**Radial Probability Distribution**

Most probable distance Distance from nucleus (r)

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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**

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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**n sublevel direction (as subscript)**

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, x2-y2, z2 for d forget it for f

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**n sublevel direction (as subscript)**

Orbital Notation n sublevel direction (as subscript) 1s 3s 2px 3dxy 4pz 4dx2-y2

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**s and p Orbitals (p shape exaggerated)**

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

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1s 2s 3s

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2pz

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**d Orbitals (also exaggerated)**

The odd one (different shape)

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3dxz

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3dz2

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**Hydrogen’s Atomic Orbitals**

A given orbital 2s, 3px, 4dyx, 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 # of orbitals Max # electrons**

Starts at energy level s 1 2 1 p 3 6 2 d 5 10 3 7 f 14 4

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**Summary by Energy Level**

1st Energy Level n=1 Only s orbital Holds 2 electrons 1s2 2 total electrons =2n2 2d Energy Level n=2 s and p orbitals are available 2 in s, 6 in p 2s22p6 8 total electrons =2n2

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**By Energy Level 3d energy level n=3 s, p, and d orbitals**

2 in s, 6 in p, and 10 in d 3s23p63d10 18 total electrons =2n2 4th energy level s,p,d, and f orbitals 2 in s, 6 in p, 10 in d, and 14 in f 4s24p64d104f14 32 total electrons=2n2

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**Orbital Summary for Hydrogen**

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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 5.3 Electron Configuration**

A set of 3 rules determines the arrangement of electrons in an atom. This arrangement is called the 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.

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**Section 5.3 Electron Configuration**

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.

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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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**Aufbau Diagram – Multi-Electron Atoms**

7p 6d 7s 5f 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s Increasing energy 3p 3s 2p 2s 1s

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**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**

1 2 3 4 5 6 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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**Constructing Orbital Diagrams**

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

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**Electron Configuration Notation**

Indicate n and sublevel for each orbital and the total electron occupancy with a superscript C 1s22s22p2 Distribution of electrons in the three p orbitals (px, py, pz) not explicit in this notation We write in n order, not energy order Ti 1s22s22p63s23p63d24s (note energy of 3d > 4s) (Textbook uses energy order for configs.)

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**Orbital Diagram/ Electron Configuration**

1s He 1s2 1s H 1s1 1s 2s 2s 1s Li 1s22s1 Be 1s22s2 1s 2s 2p F 1s22s22p5 Ne 1s22s22p6 O 1s22s22p4 N 1s22s22p3 C 1s22s22p2 B 1s22s22p1

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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 **

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

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**Next two electrons go into the 2s orbital 11 more to go**

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

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**Next six electrons go into the 2p orbitals 5 more to go**

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

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**Next two electrons go into the 3s orbital 3 more to go**

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

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**Last three electrons go into the 3p orbitals **

6d 7s 5f 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s Increasing energy 3p Last three electrons go into the 3p orbitals Each go into separate orbitals 3 unpaired electrons 1s22s22p63s23p3 3s 2p 2s 1s

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**Using the Aufbau Diagram**

2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s2 2 electrons

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**Fill from bottom up following arrows**

2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s2 2s2 4 electrons

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**Fill from bottom up following arrows**

2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s2 2s2 2p6 3s2 12 electrons

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**Fill from bottom up following arrows**

2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s2 2s2 2p6 3s2 3p6 4s2 20 electrons

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**Fill from bottom up following arrows**

2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 38 electrons

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**Fill from bottom up following arrows**

2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 56 electrons

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**Fill from bottom up following arrows**

2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 88 electrons

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**Fill from bottom up following arrows**

2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 108 electrons Maxed out

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**Noble Gas Notation Na 1s22s22p63s1 (standard notation) Ne 1s22s22p6**

Na [Ne]3s1 (noble gas notation)

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**Noble Gas Configurations**

Noble gases always have s and p orbitals completely filled (He has no p) ns2np6 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 Ne: 2s22p6 Ar: 3s23p Kr: 4s24p6

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**Table 5.5 - 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)**

32 electrons [Ar]4s23d104p2 Above configuration obtained from order of filling. For this class, use principal energy level (n) order [Ar] 3d104s24p2

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**Practice Electron Configurations Practice problems, page 160**

Probs 21(a-f), 22-25 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]3d54s1 (1/2 filled d) Not [Ar]3d44s2 Next exception is copper: Cu [Ar]3d104s1 (filled d) Not [Ar]3d94s2

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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)d10 component in the configuration, then ignore these electrons for counting valence electrons, even if higher energy than the ns2 electrons Zn [Ar]3d104s valence electrons, not 12 Br [Ar] 3d104s24p5 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]3s2 Mg S [Ne]3s23p4 S

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Valence Electrons Determine the chemical (bonding) properties of the element S [Ne]3s23p valence electrons Cs [Xe]6s valence electron

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**Dot Structures for Elements in 2d Period**

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**Example Problem 5.3 - page 162 What is electron dot structure for tin?**

Sn: [Kr]4d105s25p2 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 167-8 Probs 81(a-d), 90 (a-e), 91-93 Supplemental Problems, page 978 Probs 7, 8, 9 (a-d)

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END

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Chapter 4 Arrangement of Electrons in Atoms. 4-1 The Development of the New Atomic Model Rutherford’s atomic model – nucleus surrounded by fast- moving.

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