Presentation on theme: "Chapter 6 Electronic Structure of Atoms"— Presentation transcript:
1 Chapter 6 Electronic Structure of Atoms or“How I Learned to Stop Worrying and Love the Electron”
2 Problems with the Rutherford Model Classical physics says atoms should emit light and destroy themselves - they don’tAtoms can be induced to emit light, but they give off a line spectrum, rather than a continuous spectrum.Every atom gives off different colours of light.No explanation of why different atoms have different properties, or the same properties.
4 What is Light? light is radiant energy. a better term is electromagnetic energy (EMR)includes not only visible light, but many forms of EMR we cannot see directly:heatmicrowavesx-rayslight has properties of both waves and particles.
5 Wave model of LightThe distance between corresponding points on adjacent waves is the wavelengthThe symbol is the Greek letter lamda ().The height of the wave is the amplitude. It corresponds to the intensity of the light
6 WavesThe number of waves passing a given point per unit of time is the frequency,The symbol is the Greek letter nu ().The longer the wavelength, the smaller the frequency.
7 Electromagnetic Radiation EMR is a continuous spectrum of wavelengths.
8 Using λ and ν to determine “colour” “Colour” is a term used to describe visible light.Visible light is a very small part of the electromagnetic spectrum:radio waves -- ultravioletmicrowaves -- x-raysinfrared -- gamma raysvisible
10 Colour of Light we can identify light by its wavelength or frequency: a wave 5.00 x 10-7 m isgreena wave 6.80 x 10-7 m isreda wave 2.60 x 10-5 m isinfrareda wave 7.80 x m isin the x-ray region.
11 Speed of Lightis the one thing that is constant in the universe (sort of).in a vacuum, the speed of light is 3.00 x 108 m/sthere is a relationship between wavelength and frequency:wherec = speed of lightλ = wavelengthν = frequency
13 Quantized Energy and Photons The wave model of light is good, but it does not explain:how an object can glow when its temperature increases. (Blackbody radiation)emission of electrons by shining light on the surface of a metal. (photoelectric effect)emission of light from electrons of excited gas atoms. (emission spectra)
14 Blackbody Radiation Heated solids emit radiation (blackbody radiation) The wavelength distribution depends on the temperature (i.e., “red hot” objects are cooler than “white hot” objects).Why does wavelength or frequency depend on temperature?Max Planck suggested a way out by assuming that energy comes in packets called quanta.
15 as energy increases, so does frequency. Planck proposed a relationship between energy and the frequency of light quanta:where:E = energy in Joulesh = Planck’s constant ( x J·s)ν = frequency, in Hertz (1/s, s-1)as energy increases, so does frequency.
18 Photons Einstein used the quantum to explain the photoelectric effect: Light comes in particles, called photons.The energy of each photon is determined by Planck’s equation.Light shining on the surface of a metal can cause electrons to be ejected from the metal.
19 Light has wave-like AND particle-like properties. The electrons will only be ejected if the photons have sufficient energy:Below the threshold frequency no electrons are ejected.Above the threshold frequency, the excess energy appears as the kinetic energy of the ejected electrons.Light has wave-like AND particle-like properties.Complete question 6.30
20 Line Spectra and the Bohr Model Another mystery involved the emission spectra observed from energy emitted by atoms and molecules.
21 The Nature of EnergyA white light source produces a continuous spectrum, like a rainbow.When elements are excited, only a line spectrum of discrete wavelengths is observed.
22 Hydrogen Spectrum has line spectra in 3 regions of the EM Spectrum: Ultraviolet – Lymann SeriesVisible – Balmer SeriesInfrared - Paschen Series
23 The Nature of EnergyNiels Bohr adopted Planck’s ideas about the quantum and applied them to the electrons around a nucleus.
24 The Nature of Energy Bohr’s model is based on three postulates: Electrons in an atom can only occupy certain permitted orbits, or quanta.Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.Energy is only absorbed or emitted to move an electron from one “allowed” energy state to another; the energy is emitted by a photon:E = h
25 The Nature of EnergyAn electron in its lowest permissible energy is at ground stateIf an electron accepts a quantum of energy it will move to a higher energy level, or excited state.When the electron moves back down to ground state it emits a photon of light of a frequency which correlates to the energy of the quantum.
26 The Nature of EnergyThe energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation:E = −RH( )1nf2ni2-where RH is the Rydberg constant, 2.18 10−18 J, and ni and nf are the initial and final energy levels of the electron.
27 this explains the line spectra of hydrogen: Lyman series. High energy, in the UV range. Represents energy transition from higher quanta to ground state.Balmer series. Intermediate energy, in the visible range. Represents energy transition from higher quanta to quantum 2.Paschen series. Low energy, in the IR range. Represents energy transition from higher quanta to quantum 3.
28 Limitations of Bohr Model cannot explain the spectra of atoms other than hydrogen.However, the model introduces two important ideas:The energy of an electron is quantized: electrons exist only in certain energy levels described by quantum numbers.Energy gain or loss is involved in moving an electron from one energy level to another.
29 The Wave Nature of Matter Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties.He demonstrated that the relationship between mass and wavelength was =hmv
30 What does this mean? electrons have wave properties. the orbits of electrons are multiples of the electron wavelengththe first orbit has a circumference of 1 λ, the second is 2 λ , etc.
31 The Uncertainty Principle Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known.Our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!
32 Quantum MechanicsErwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.It is known as quantum mechanics.
33 The wave function describes the electron’s matter wave; it gives the probability of finding the electron.Electron density is another way of expressing probability.A region of high electron density is one where there is a high probability of finding an electron.
37 Orbitals and Quantum Numbers Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies.Each orbital describes a spatial distribution of electron density.An orbital is described by a set of three quantum numbers.
38 Principal Quantum Number, n describes the energy level on which the orbital resides.The values of n are integers ≥ 0.correspond to the periods on the Periodic Table.
39 Azimuthal Quantum Number, l defines the shape of the orbital.Allowed values of l are integers ranging from 0 to n − 1.letter designations communicate the different values of l and, therefore, the shapes and types of orbitals.
40 Azimuthal Quantum Number, l Value of l123Type of orbitalspdfTheoretical g, h, i, etc. orbitals exist, but no atoms have been created to use them.
41 Magnetic Quantum Number, ml three-dimensional orientation of the orbital.Values are integers ranging from -l to l :−l ≤ ml ≤ lon any given energy level, there can be up to:1 s orbital3 p orbitals5 d orbitals7 f orbitals
42 Magnetic Quantum Number, ml Orbitals with the same value of n form a shell.Different orbital types within a shell are subshells.Each subshell is designated by a number and a letter.For example, 3p orbitals have n = 3 and l = 1.
44 s Orbitals Value of l = 0. Spherical in shape. Radius of sphere increases with increasing value of n.
45 p Orbitals Value of l = 1. Have two lobes with a node between them. The letters correspond to allowed the values of ml of –1, 0, and +1.
46 d OrbitalsValue of l is 2.Four of the five orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.The letters correspond to allowed the values of ml of -2, –1, 0, +1 and +2.
47 f orbitalsValue of l is 3.7 possible shapes, including 8 lobes and 2 doughnuts.The letters correspond to allowed the values of ml of -3, -2, –1, 0, +1, +2 and +3.
48 Energies of OrbitalsOrbitals of the same energy are degenerate.
49 Spin Quantum Number, mstwo electrons in the same orbital do not have exactly the same energy.The “spin” of an electron describes its magnetic field, which affects its energy.
50 Spin Quantum Number, ms There is a spin quantum number, ms. has only 2 allowed values: +1/2 and −1/2.
51 Pauli Exclusion Principle No two electrons in the same atom can have exactly the same energy.Therefore, no two electrons in the same atom can have identical sets of quantum numbers.
52 Electron Configurations Distribution of all electrons in an atomTell us how the electrons are distributed among the various orbitals of an atom.Consist ofNumber denoting the energy level
53 Electron Configurations Distribution of all electrons in an atomTell us how the electrons are distributed among the various orbitals of an atom.Consist ofNumber denoting the energy levelLetter denoting the type of orbital
54 Electron Configurations Distribution of all electrons in an atom.Tell us how the electrons are distributed among the various orbitals of an atom.Consist ofNumber denoting the energy level.Letter denoting the type of orbital.Superscript denoting the number of electrons in those orbitals.
55 Electron Configurations Electrons tend to occupy the lowest energy orbitals at ground state.Thus the electron configuration of an atom is the arrangement of the electrons from the lowest energy level to the highest.
57 Electron Configurations For any atom or ion:find total number of electronsstart with the 1s electrons and fill each orbital in order of energywatch the order of filling; energy is all importantFor instance:Iron (Fe) – contains 26 electrons1s22s22p63s23p64s23d6
58 Electron Configuration of Period 2 Elements Element Electron configuration lithium 1s22s1 beryllium 1s22s2 boron 1s22s22p1 carbon 1s22s22p2 nitrogen 1s22s22p3 oxygen 1s22s22p4 fluorine 1s22s22p5 neon 1s22s22p6
62 Electron Promotionmovement of an outer ‘s’ electron to the adjacent ‘p’ orbital at the same energy level.turns non-bonding electrons into bonding electronsallows atoms to make more chemical bonds and achieve a lower energyapplies to elements from groups 2, 13 and 14 onlyfor these elements promotion is the rule
63 Electron Promotion Element Unhybridized Hybridized beryllium 1s22s2 1s22s12p1boron 1s22s22p1 1s22s12p2carbon 1s22s22p2 1s22s12p3
64 Orbital Diagramsare another way to illustrate the position of electrons.They are best learned by comparison with electron configuration:Na (11 protons, 11 electrons)electron configuration: 1s22s22p63s1orbital diagram:1s 2s 2p 3s↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑
65 Orbital Diagrams Each line represents one orbital. Arrows represent the electrons.The direction of the arrow represents the spin of the electron.
66 Hund’s Rule“In any orbital with more than 1 sub- orbital electrons will half-fill each sub- orbital before pairing up.”not
67 Orbital Diagrams Representative Group Element Electron configuration Orbital Diagram1s 2s p1 lithium 1s22s1 ↑↓ ↑2 beryllium 1s22s12p1 ↑↓ ↑ ↑13 boron 1s22s12p2 ↑↓ ↑ ↑ ↑14 carbon 1s22s12p3 ↑↓ ↑ ↑ ↑ ↑15 nitrogen 1s22s22p3 ↑↓ ↑↓ ↑ ↑ ↑16 oxygen 1s22s22p4 ↑↓ ↑↓ ↑↓ ↑ ↑17 fluorine 1s22s22p5 ↑↓ ↑↓ ↑↓ ↑↓ ↑18 neon 1s22s22p6 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓Note the electron promotion for the elements from groups 2, 13 & 14
68 Repeat the last assignment, giving the orbital diagrams for the elements.
73 Condensed Electron Configurations Looks mainly at electrons with the highest energy:find element on the Periodic Table.find noble gas with lower atomic numberwrite electron configuration from that point.Example:P is 1s22s22p63s23p3,but Ne is 1s22s22p6.Therefore, P is [Ne]3s23p3.
74 Electron Configuration and the Periodic Table The periodic table can be used as a guide for electron configurations.The period number is the value of n.Groups 1 & 2 - s blockGroups 13 to 18 - p blockGroups 3 – 12 – d blockThe lanthanides and actinides - f blockWatch the order of filling!