Unit 3 Light and Electronic Configuration 24-25 HC PDF

UNIT 3 – Light and Electronic Configuration Enloe Magnet High School Mr. Lamberth 3.1 Wave Nature Of Light Electromagnetic Radiation Radiant energy Form of energy that has wave-like characteristics. Speed in a vacuum is 3.00 x 108 m/s Electromagnetic Spectrum North Carolina C...

UNIT 3 – Light and Electronic Configuration Enloe Magnet High School Mr. Lamberth 3.1 Wave Nature Of Light Electromagnetic Radiation Radiant energy Form of energy that has wave-like characteristics. Speed in a vacuum is 3.00 x 108 m/s Electromagnetic Spectrum North Carolina Chemistry Reference Tables  The Wave Nature of Light c = speed of light = 3.00 x 108m/s Wavelength (λ) - distance between troughs Uses the units m (SI unit) or nm Frequency (ν) - the number of complete wavelengths that pass a given point each second. Uses the units cycles per second; s -1 = /s (SI unit) or hertz (Hz) 1 Hz = 1 s -1 Related by c = λν Frequency and wavelength are ____ proportional. 3.2 Quantized Energy and Photons A revolution is born The wave nature of light does not explain three phenomena that are specifically related to how atoms and light interact. Blackbod Photoelectr Emission y ic Effect Spectrum Radiation  Max Planck Explained how radiation is emitted by hot objects (black- body radiation) Assumed energy can be released (or absorbed) by atoms or molecules only in discrete quantities Quantum (plural quanta) – discrete chunk of energy, a fixed amount. E = hν Planck’s constant, h = 6.63 x 10-34 J s The chunks of energy would have values of hν, or 2hν, 3hν, etc. Albert Einstein Credit: CK-12 Foundation - Christoph Auyeung;Raymond Chou Source: CK-12 Foundation License: CC-BY-NC-SA 3.0 Explained photoelectric effect (When electrons are ejected off of the surface of a metal with light with a certain minimum frequency shines on it. Extended Planck’s theory Photons - packets of energy that behave like particles; particles of light If photons have enough energy when they strike the metal, they will pass that energy to the electrons causing the electrons to fly off! Is light composed of waves or particles? What do you think? 3.3 Emission Spectrum and The Bohr Model Spectrum Continuous Spectrum - Continuous range of colors Line (Emission) Spectrum – A spectrum containing radiation of specific wavelengths emitted by a substance. Absorption Spectrum – A spectrum containing radiation of all wavelengths except those absorbed by a substance. Who was the first to shine light through a prism and study it? Hydrogen Spectra Absorption spectrum Line/Emission spectrum Emission Spectrum of Elements  Bohr’s Model Three postulates: Assumed that electrons move in circular orbits around the nucleus. These orbits correspond to certain definite amounts of energy. An electron in a permitted orbit has a specific energy and is in an allowed energy state. It will not spiral into the nucleus. Energy is emitted or absorbed by an electron as it changes from one energy state to another. This energy exists as a photon. Ground State: Electrons are as close to the nucleus as they can be; they are in the lowest energy level Excited State: Electrons are not as close to the nucleus as they can be; they are 3.4 The Quantum Mechanical Model  Louis de Broglie ▪ Dual nature of the electron - suggests that if light can behave like a stream of particles then electrons may possess wave properties. The Wave-Particle Duality. ▪ Soon after this was published it was experimentally demonstrated.  Werner Heisenberg ▪ A wave extends into space, therefore its exact location can not be found. ▪ Also because photons are used to detect electrons and the energy of each of these particles is similar, then any attempt to locate an electron with a photon will knock the elctron off its course. Heisenberg Uncertainty Principle ▪ It is impossible to know the momentum and position of a particle with certainty. ▪ Because of this we know the electron does not orbit the nucleus in a well defined path (as Bohr thought)  Quantum Mechanical Model The dual nature of the electron and Heisenberg’s Uncertainty Principle led to the Quantum Mechanical Model The Schrodinger Equation incorporates both the wave and particle behavior of electrons. The location of an electron cannot be described so simply. Launched a field of physics called quantum mechanics. Quantum Mechanical Model ▪ Quantum mechanics mathematically defines a region in space where the electron has a high probably of being at a given instant. ▪ Regions where there is a high probability of finding the electron are said to be regions of high Electron Density Electron Density  Orbitals ▪ Orbital – a three-dimensional region around the nucleus that indicates the probable location of an electron. ▪ Electrons do not travel around the nucleus, instead they exist in regions called orbitals. ▪ Region in the atom where there is a high probability of finding an electron. ▪ Each has a characteristic  Quantum Numbers Principal Quantum Number (n) ▪ Describes the distance the electron is from the nucleus. ▪ Identifies the energy level of the electron. ▪ Allowed values: n = 1, 2, 3, etc… ▪ In an energy level, n , there are ▪ n sublevels ▪ n2 orbitals and ▪ A maximum of 2n2 electrons that an occupy it.  Quantum Numbers Azimuthal (angular momentum) Quantum Number (l ) ▪ Describes the shapes of orbitals ▪ Identifies the sublevel (subshell) in which the n Value for l electron is located. 1 0 2 0, 1 ▪ Allowed values: l = 0 to n-1 3 0, 1, 2 ▪ In a subshell, l , there are 4 0, 1, 2, 3 5 0, 1, 2, 3, 4 ▪ 2l + 1 orbitals Type of l ▪ Maximum of 2(2l + 1) electrons sublevel s 0 ▪ sharp, principal, diffuse, and fundamentald p 1 2 ▪ Used to describe spectra before quantum f 3  Quantum Numbers Azimuthal (angular momentum) Quantum Number (l ) ▪ Continued… ▪ A collection of orbitals with the same value of n is called a shell (or energy level) ▪ A collection of orbitals with the same value of n and l is known as a subshell (or sublevel) ▪ https://www.chemtube3d.com/orbitals-p/ Orbitals Artistic Picures Realistic Pictures  Quantum Numbers Magnetic quantum number (ml) ▪ Describes the orientation of the orbital around the x, y, and z axes ▪ Identifies the specific orbital in which an electron could be located. ▪ Allowed values: -l to +l l value Possible values # of Orbitals for m orbitals l 0 0 1 s 1 -1, 0, 1 3 px py pz 2 -2, -1, 0, 1, 2 5 d x2 – y2 , dz2 , dxy , dxz , dyz 3 -3, -2, -1, 0, 1, 2, 7 3  Quantum Numbers Spin magnetic quantum number (ms) ▪ Explains the splitting of emission spectra of H when magnetic field is applied. ▪ Electrons are assumed to act like tiny magnets that spin on their own aces. ▪ The electron has two opposite spins, clockwise and counterclockwise. ▪ Allowed values: or ▪ The value has no effect on energy, size, shape, or orientation of an orbital, but it determines how electrons are arranged in orbitals of equal energy.  Summary of Quantum ▪Numbers Principal: Gives Shell (n = 1,2,3, etc.) ▪ Azimuthal: Gives Subshell (l = 0 to n) corresponds to s, p, d, f. ▪ Magnetic: Gives orbital, orientation on x, y, z axis. ml = -l to +l ▪ Spin: two electrons per orbital, values of +½ and -½ ▪ All four quantum numbers enable us to label an electron in any orbital in any atom. 3.5 Electron Configuration Electronic Structure - Inroduction electrons, electrons, electrons Knowing the arrangement, number of electrons, and energy of the electrons (electronic structure) in an atom is the key to understanding the physical and chemical properties of an element. What does the study of light have to do with understanding electronic structure?  Energy Level and Sub ▪Levels n Value for l Energy Levels (or Shell) have 1 0 any positive integer value 2 0, 1 ▪ Determines the size of the 3 0, 1, 2 orbital and the distance it is from the nucleus. 4 0, 1, 2, 3 ▪ Sub levels (or subshells) are 5 0, 1, 2, 3, 4 represented by a letter. In the ground state, this letter can Type of l sublevel only be s, p, d, or f. s 0 ▪ Determines the shape of the p 1 orbital, how many orbitals there d 2 are, and thus how many electrons can exist in the f 3  Orbitals ▪ 3-D region in space of where the electron is most likely to exist. ▪ We are going to represent this with a box or line. l Possible values for ml 0 0 1 -1, 0, 1 2 -2, -1, 0, 1, 2 3 -3, -2, -1, 0, 1, 2, 3  Electron Spin ▪ There can only be two electrons in each orbital. ▪ These two electrons must be different. ▪ We will use arrows to represent these electrons, and up arrow and a down arrow in a box or on a line. Energy Level Types of Electrons in Total # of Sublevels each electrons sublevel 1 s 2 2 2 s, p 2+6 8 3 s, p, d 2 + 6 + 10 18 2 + 6 + 10 + 4 s, p, d, f 14 32  Electron Configuration H: 1s1 1 (the coefficient) – gives energy level s – gives type of subshell -orbital shape -# of orbitals 1s – gives subshell Superscript – gives the number of electrons in that subshell Example Problems Write the electron configuration for S. Write the electron configuration for V. Write the electron configuration for Rb. 3.6 Orbital Notation Orbital Notation Pauli’s Exclusion Principle ▪ No two electrons in an atom can have the same four quantum numbers. They will have different spins. Hund’s Rule ▪ The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins. ▪ Place electrons in each orbital one at a time, then go back and place two in them if necessary. Aufbau Principle ▪ As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to the atomic orbitals. ▪ Electrons will fill lower energy subshells first then move on to higher energy subshells. ▪ 1s 2s 2p 3s 3p 4s 3d 4p etc… Example Problems Write the Orbital Notation for S, V, and Rb from the previous examples. 3.7 Noble Gas Core Configuration Noble Gas Configuration Go to the Noble Gas that precedes the element (by atomic number.) Write that Noble Gas Symbol in Brackets. Then write the rest of the configuration starting after that noble gas. Example: Rb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 Rb: [Kr] 5s1 Write the Noble Gas Configuration for S and V from the previous Examples. 3.8 Valence Electrons Valence Electrons Valence electrons are the electrons in the outermost shell (energy level) Notice that the outermost shell for any atom is always going to be the s and sometimes p subshell. Looking at the electron configurations for S, V, and Rb, how many valence electrons does each have? 3.9 Understanding the Quantum Mechanical Model; Electron Configuration and Orbital Notation Energy Levels, Sublevels, and Orbitals REMEMBER: Energy Level = Shells are represented by the coefficient. Sublevel = Subshell The type is given by the letter The exact one is given by the coefficient with letter Give shape, number of orbitals and thus number of electrons. Boxes(or lines) = orbital Arrow = electron Unpaired electron = only one arrow in a box 3.10 Excited States Excited State Excited State - An electron that has gained energy and moved into a higher energy orbital. The energy gained can come from heat, light, or electricity. Electrons do not stay in this excited state for long. They soon return to their ground state or some other lower energy state by emitting a photon of light that corresponds to that energy change. Some examples of this happening are in chemical reactions such as photosynthesis or when light bulbs have electricity pass through them. https://slideplayer.com/slide/13853867/