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Electrons in an Atom • According to the Bohr Model, electrons (e-) can only orbit the nucleus in specific, allowed pathways. • They move toward and away from the nucleus by “steps” or discrete amounts of energy that are released or absorbed. • e- farther from the nucleus have more energy. Those closer to the nucleus have less energy. The Problem (s)… • If electrons orbited the nucleus, they would constantly be accelerating • Accelerating charges give off all colors of light at the same time • White light is made up of every visible color of light • There’s nothing stopping the electrons from crashing into the nucleus? (+ attract -) Solution: Very similar to a ladder. Just as you cannot step on the air between the rungs, an electron cannot exist between the levels. Light is going to help us understand the atom! We need to know… • What is light? • How does light behave? • What does light have to do with electrons? • Can we explain/predict the kind of light that is given off by different elements? The Electromagnetic Spectrum All the types of light! ROY G. BIV Think… Slinky Light is a Wave Light – as a Wave (How do we know this?) Can interfere with itself (two waves crash together and cancel each other out) Can refract (be bent as it passes through different materials) Parts of a Wave Light – as a Wave Wavelength (λ – “lambda”): the distance between the peaks Frequency (ν – “nu”): number of waves that pass a given point in a specific amount of time λ crest trough Light as a wave speed of light = frequency x wavelength c=νλ (where c is always equal to 3.00 x 108 m/sec) Longer wavelength → _________ frequency Shorter wavelength → _________frequency In the visible spectrum, color is associated with different wavelengths (ROY G. BIV) Longer λ and lower ν: towards Red Shorter λ and higher ν: towards Violet Einstein’s Theory To make sense of the fact that light can eject electrons even if its intensity (brightness) is low… Light is both waves and particles –think “stream of particles” Particle: The photon – it has zero mass and carries a quantum (piece) of energy If the energy of the light particle is too low, no matter how many particles hit the metal, no electrons will pop off! New Ideas about Light Prior to the 1900s – light was thought of as only a wave … Now there is a duality of light “Duality” means that there are two ways to represent it Wave Particle Properties of both are present This also means that… all particles have a wave nature, all waves have a particle nature. Turn in to me: 1. Rank the colors of light from low to high energy. 2. Rank the colors of light from small to large wavelength Violet Blue Green Yellow Orange Red Particle of Light Emission of light by hot objects happens in small specific amounts (packets)– quanta Each particle carries with it an amount of energy which depends on the frequency of light Quantum – that small bundle of energy Energy = Planck’s Constant * frequency emitted E=hν (where h = 6.626 x 10-34 J*s) Spectroscopic analysis of the visible spectrum… …produces all of the colors in a continuous spectrum Spectroscopic analysis of the hydrogen spectrum… …produces a “bright line” spectrum Ground State – lowest energy state This means that e- are found in shells closer to the nucleus. Excited State – higher potential energy of an atom A form of heat, light, electrical, or mechanical energy is needed to go from the ground to an excited state As electrons increase in energy, they move away from the nucleus and into outer shells. Atoms and Light Absorption (take in) of energy moves electrons from a ground state to a higher energy state Heat, light, electrical, mechanical energy Emission (or give off) of energy lets electrons fall back down to a lower energy state Always light We don’t know why this happens Excited State Energy Going In Ground State Energy Coming Out “Color” The kind of light given off (emitted) by each element is UNIQUE because each element has a UNIQUE arrangement of electrons! (Think fingerprint) (It’s not just about electrons in “shells”… it’s more complicated than that! Hydrogen’s Line Emission Spectrum Excited State Energy (Light) Ground State This led to the quantum model of the atom! Electrons are in orbitals, not orbits Not Come from a mathematical equation Orbitals (3-D Fuzzy Shapes) Three dimensional space that electrons most probably occupy The math equation treats electrons like waves You can solve the equation to get the shape of space in which electrons are Shapes look like “clouds” of probability What is the “address” of an electron? Where do you live? - Where on the block - Style of house - Which way is your house facing Electrons are identified the same way … - Principle Energy Level (identified by n) - Sublevel (identified by l) - Orbital Orientation (identified by ml) - Spin State (identified by ms) Four Quantum Numbers… 1. Principle Quantum Number (n ) • Indicates the main energy level occupied by the e- (distance from the nucleus) • Shell Number (1st shell is closest to nucleus, 2nd is further, and so on …) • Come from the Bohr Model • Values of n can only be positive integers (1, 2, 3, etc.) Four Quantum Numbers…cont. 2. Angular Momentum Quantum Number ( l ) • Indicates the shape of the orbitals (orbital = individual 3D shape in space where the probability of finding e- is greatest) • Nickname is subshell of n • Designated s, p, d, f • Values of l are zero and all positive integers less than equal to n - 1 The number of subshells in a shell is equal to the shell number 1st shell – 1 subshell 2nd shell – 2 subshells 3rd shell – 3 subshells s subshell Spherical shaped One orbital, 2 e- p subshell Dumbbell shaped Three orbitals, 6 ed subshell Double peanut shaped Five orbitals, 10 ef subshell Flower shaped Seven orbitals, 14 e- What are the two quantum numbers learned so far? 1. Principal Quantum Number shell number 1-7 distance from nucleus 2. Angular Momentum Quantum Number subshell s p d f actual space where the probability of finding an e- pair is greatest (orbital) Four Quantum Numbers…cont. 3. Magnetic Quantum Number (ml ) • Indicates the orientation of an orbital (the plane that the orbital is in) • Because an s orbital is spherical, it only has one orientation (ml = 0) • p orbitals can have three different orientations, one along the x-axis, one along the y-axis, and one along the z-axis (ml = -1, ml = 0, ml = +1) • (just an illustration: I don’t need you to memorize) 4. Spin Quantum Number ( s Indicates the two spin states of an e- in an orbital Only 2 e- fit in each orbital, and they spin in opposite directions (up and down!) Possible ms values are - ½ , + ½ Spin is represented by dashes inside circles Empty m) Half-Filled Filled Half-Filled Filled or arrows Empty What are the four quantum numbers? Principal Quantum Number shell number 1-7 distance from nucleus 2. Angular Momentum Quantum Number subshell s p d f actual space where the probability of finding an e- pair is greatest (orbital) 3. Magnetic Quantum Number the # of directions your shape (orbitals) faces 4. Spin Quantum Number clockwise, counterclockwise 1. Bohr Models Represented with pictures… Mathematical Models Represented with… ? Summary Table Shells Subshell Orbitals 1 2 3 4 s 1 s p s 1 3 1 p d s 3 5 1 p d 3 5 f 7 Total # orbitals in shell 1 4 9 16 #e- Total #ein shell 2 2 2 6 2 6 10 2 6 10 14 8 18 32 It is a bit complicated… shells don’t always get filled from 1 to 2 to 3 etc. because some subshells overlap Start End Nucleus 1 Start Start 2 End 3 End Start Start End 5 4 End What order are they filled in? 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d ? 4f 5f ? ? ? ? ? ? ? Electron Configurations Orbital Notations The Arrangement of electrons in atoms Aufbau Principle “building up” An electron occupies the lowest energy possible The levels follow a pattern of increasing energy Fill starting at nucleus (Bohr Models!) p subshell 3 orbitals Fill left to right not Hund’s Rule Electrons do not pair up until there are no more empty orbitals in that subshell 2e3e- NOT 4e- OR Pauli Exclusion Principle No two electrons in the same atom will have the same set of four quantum numbers No two electrons have the same address An orbital can only hold two electrons because there are only two different spin states (and the must be opposite) To represent e- configurations, you need to know: 1. The atomic # (# of e-) 2. Order in which the e- fill 3. The number of orbitals in each sublevel (s, p, d, f) (remember – they want a private room) What to write: 1. Shell # (1, 2, 3, 4, …) 2. Sublevel letter (s, p, d, f) 3. # of e- orbital notation (Remember how many e- are in each sublevel) Electron configuration and the periodic table What do elements in column 1 have in common? Column 2? Column 3? What do elements in each of the rows have in common? Periodic Table Different blocks on the periodic table (shaded in different colors here) correspond to different types of orbitals. The lazy electron configuration Start with the previous noble gas in brackets, and then finish the rest of the configuration “Condensed” electron configuration Long way is Sodium = 1s2 2s2 2p6 3s1 Condensed Sodium = [Ne] 3s1 “valence” shell electrons are the ones listed as following the noble gas (aka the outside shell electrons) Element Lithium Configuration notation 1s22s1 [He]2s1 ____ 1s Beryllium ____ ____ 2p ____ ____ 2s ____ ____ 2p ____ [He]2s2p2 ____ 2s ____ ____ 2p ____ 1s22s2p3 [He]2s2p3 ____ 2s ____ ____ 2p ____ 1s22s2p4 [He]2s2p4 ____ 2s ____ ____ 2p ____ 1s22s2p5 [He]2s2p5 ____ 1s Neon ____ 2s 1s22s2p2 ____ 1s Fluorine ____ [He]2s2p1 ____ 1s Oxygen ____ 2p 1s22s2p1 ____ 1s Nitrogen ____ [He]2s2 ____ 1s Carbon ____ 2s 1s22s2 ____ 1s Boron Noble gas notation Orbital notation ____ 2s ____ ____ 2p ____ 1s22s2p6 [He]2s2p6 ____ 1s ____ 2s ____ ____ 2p ____ Valence Shell Electrons Valence = outermost Count the total electrons in the highest shell number Do not count electrons in d subshells How many valence shell electrons are in the following elements? H He O S Br