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Electrons in Atoms Greek Idea Democritus and Leucippus Matter is made up of indivisible particles Dalton - one type of atom for each element Thomson’s Model Discovered electrons Atoms were made of positive stuff Negative electron floating around “Plum-Pudding” model Rutherford’s Model Discovered dense positive piece at the center of the atom Nucleus Electrons moved around Mostly empty space Bohr’s Model Why don’t the electrons fall into the nucleus? Move like planets around the sun. In circular orbits at different levels. Amounts of energy separate one level from another. Bohr’s Model Nucleus Electron Orbit Energy Levels Bohr’s Model Increasing energy Fifth Fourth Third Second First Nucleus Further away from the nucleus means more energy. There is no “in between” energy Energy Levels The Quantum Mechanical Model Energy is quantized. It comes in chunks. Quanta - the amount of energy needed to move from one energy level to another. Since the energy of an atom is never “in between” there must be a quantum leap in energy. The Quantum Mechanical Model Schrodinger derived an equation that described the energy and position of the electrons in an atom Heisenberg Uncertainty Principle – it is impossible to know both the position and energy (momentum) of an electron at the same time. Photons – light particles that allow us to make observations The Quantum Mechanical Model Things that are very small behave differently from things big enough to see. Wave-Particle Duality of Nature – an electron can behave like a wave or a particle depending on what is being studied. The Quantum Mechanical Model Describes the location of an electron by using four quantum numbers: 1st - Principle Energy Level (n) 2nd – Sublevel (l) • 3rd - Angular Momentum (m) – 4th – Spin (s) Each level is getting more specific for the location of the electron. 1st Quantum Number Principal Energy Level (n) = the energy level of the electron. There are 7 main energy levels for all known elements today. The value equals the row of the periodic table (exception is the d and f sublevels) 2nd Quantum Number Within each energy level the complex math of Schrodinger’s equation describes several shapes. Sublevel – describes the shape of the orbital. Orbital – probable region where there is an electron. The value is associated with the type of orbital. S = 0; p = 1; d = 2; f = 3 s orbitals Spherical shaped Start at first energy level Each s orbital can hold 2 electrons p orbitals Start at the second energy level 3 different directions dumbbell shaped Each can hold 2 electrons – 6 total p Orbitals d orbitals Start at the third energy level 5 different shapes Each can hold 2 electrons – 10 total f orbitals Start at the fourth energy level 7 different shapes Each can hold 2 electrons – 14 total f orbitals Summary # of Max shapes electrons Starts at energy level S 1 2 1 p 3 6 2 d 5 10 3 f 7 14 4 Electron Configurations Use the 1st and 2nd quantum numbers to describe an atoms electrons. Aufbau principle- electrons enter the lowest energy first. This causes difficulties because of the overlap of orbitals of different energies. The easy way to remember 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s • 2 1s • 2 electrons Fill from the bottom up following the arrows 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s • 2 2 1s 2s • 4 electrons Fill from the bottom up following the arrows 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s • 2 2 6 2 1s 2s 2p 3s • 12 electrons Fill from the bottom up following the arrows 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s • 2 2 6 2 1s 2s 2p 3s 6 2 3p 4s • 20 electrons Fill from the bottom up following the arrows 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s • 2 2 6 2 1s 2s 2p 3s 6 2 10 6 3p 4s 3d 4p 5s2 • 38 electrons Fill from the bottom up following the arrows 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s • 2 2 6 2 1s 2s 2p 3s 6 2 10 6 3p 4s 3d 4p 5s2 4d10 5p6 6s2 • 56 electrons Fill from the bottom up following the arrows 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s • 2 2 6 2 1s 2s 2p 3s 6 2 10 6 3p 4s 3d 4p 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 • 88 electrons Fill from the bottom up following the arrows 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s • 2 2 6 2 1s 2s 2p 3s 6 2 10 6 3p 4s 3d 4p 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 • 108 electrons Orbital Diagram A visual representation of the electrons orbital arrangement. Write the electron configuration for P. 1s2 2s2 2p6 3s2 3p3 1s 2s 2p 3s 3p 4s 3d Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to . 1s2 1s 2s 2p 3s 3p 4s 3d Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to . 1s2 2s2 2p6 1s 2s 2p 3s 3p 4s 3d • 1s22s22p63s23p3 • 3 Unpaired electrons - paramagnetic 1s 2s 2p 3s 3p 4s 3d 3rd Quantum Number Angular Momentum (m) - describes the orientation The value is (– l) to (+l) Each orbital within the sublevel is assigned a number s = 0 p = -1; 0; +1 d = -2; -1; 0; +1; +2 f = -3; -2; -1; 0; +1; +2; +3 3rd Quantum Number The value is (– l) to (+l) Each orbital within the sublevel is assigned a number s = 0 p = -1; 0; +1 d = -2; -1; 0; +1; +2 f = -3; -2; -1; 0; +1; +2; +3 0 0 -1 0 +1 0 -1 0 +1 0 1s 2s 2p 3s 3p 4s -2 -1 0 +1 +2 3d 4th Quantum Number Spin (s) – each orbital contains two electrons with different spins. Pauli Exclusion Principle- at most 2 electrons per orbital – must have different spins Value is either +½ or -½ Indicated by an up or down arrow 4th Quantum Number Spin – Up arrow = +½ Down arrow = -½ 0 0 -1 0 +1 0 -1 0 +1 0 1s 2s 2p 3s 3p 4s -2 -1 0 +1 +2 3d All Four Quantum Numbers 1st – Energy Level = 3 2nd – Sublevel Value (p) = 1 3rd – Angular Momentum Value = +1 4th – Spin – Up arrow = +½ 0 0 -1 0 +1 0 -1 0 +1 0 1s 2s 2p 3s 3p 4s -2 -1 0 +1 +2 3d Exceptions to Electron Configuration Orbitals fill in order Lowest energy to higher energy. Adding electrons can change the energy of the orbital. Half filled orbitals have a lower energy. Makes them more stable. Changes the filling order Write these electron configurations Titanium - 22 electrons 1s22s22p63s23p64s23d2 Vanadium - 23 electrons 1s22s22p63s23p64s23d3 Chromium - 24 electrons 1s22s22p63s23p64s23d4 is expected But this is wrong!! Chromium is actually 1s22s22p63s23p64s13d5 Why? This gives us two half filled orbitals. Slightly lower in energy. The same principal applies to copper. Copper’s electron configuration Copper has 29 electrons so we expect 1s22s22p63s23p64s23d9 But the actual configuration is 1s22s22p63s23p64s13d10 This gives one filled orbital and one half filled orbital. Remember these exceptions Light The study of light led to the development of the quantum mechanical model. Light is a kind of electromagnetic radiation. Electromagnetic radiation includes many kinds of waves All move at 3.00 x 108 m/s ( c) Parts of a wave Crest Wavelength Amplitude Orgin Trough Parts of Wave Orgin - the base line of the energy. Crest - high point on a wave Trough - Low point on a wave Amplitude - distance from origin to crest Wavelength - distance from crest to crest Wavelength - is abbreviated l Greek letter lambda. Frequency The number of waves that pass a given point per second. Units are cycles/sec or hertz (hz) Abbreviated n the Greek letter nu c = ln Frequency and wavelength Are inversely related As one goes up the other goes down. Different frequencies of light is different colors of light. There is a wide variety of frequencies The whole range is called a spectrum High Low energy energy Radio Micro Infrared Ultra- XGamma waves waves . violet Rays Rays Low High Frequency Frequency Long Short Wavelength Wavelength Visible Light Atomic Spectrum How color tells us about atoms Prism White light is made up of all the colors of the visible spectrum. Passing it through a prism separates it. If the light is not white By heating a gas with electricity we can get it to give off colors. Passing this light through a prism does something different. Atomic Spectrum Each element gives off its own characteristic colors. Can be used to identify the atom. How we know what stars are made of. • These are called discontinuous spectra • Or line spectra • unique to each element. • These are emission spectra • The light is emitted given off. Light is a Particle Energy is quantized. Light is energy Light must be quantized These smallest pieces of light are called photons. Energy and frequency are directly related. Energy and frequency E=hxn E is the energy of the photon n is the frequency h is Planck’s constant h = 6.6262 x 10 -34 Joules sec. joule is the metric unit of Energy The Math in Chapter 11 Only 2 equations c = ln E = hn Plug and chug. Examples What is the wavelength of blue light with a frequency of 8.3 x 1015 hz? What is the frequency of red light -5 with a wavelength of 4.2 x 10 m? What is the energy of a photon of each of the above? An explanation of Atomic Spectra Where the electron starts When we write electron configurations we are writing the lowest energy. The energy level and electron starts from is called its ground state. Changing the energy Let’s look at a hydrogen atom Changing the energy Heat or electricity or light can move the electron up energy levels Changing the energy As the electron falls back to ground state it gives the energy back as light Changing the energy May fall down in steps Each with a different energy Ultraviolet Visible Infrared Further they fall, more energy, higher frequency. This is simplified the orbitals also have different energies inside energy levels All the electrons can move around. What is light Light is a particle - it comes in chunks. Light is a wave- we can measure its wave length and it behaves as a wave If we combine E=mc2 , c=ln, E = 1/2 mv2 and E = hn We can get l = h/mv The wavelength of a particle. Matter is a Wave Does not apply to large objects Things bigger that an atom A baseball has a wavelength of about 10-32 m when moving 30 m/s An electron at the same speed has a wavelength of 10-3 cm Big enough to measure. The physics of the very small Quantum mechanics explains how the very small behaves. Classic physics is what you get when you add up the effects of millions of packages. Quantum mechanics is based on probability because Heisenberg Uncertainty Principle It is impossible to know exactly the speed and velocity of a particle. The better we know one, the less we know the other. The act of measuring changes the properties. More obvious with the very small To measure where a electron is, we use light. But the light moves the electron And hitting the electron changes the frequency of the light. Before Photon Moving Electron After Photon changes wavelength Electron Changes velocity