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Bohr’s Model of the Atom

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Bohr’s Model Why don’t the electrons fall into the nucleus? e- move like planets around the sun. They move in circular orbits at different levels. Amounts of energy separate one level from another.

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Bohr postulated that: Fixed energy related to the orbit Electrons cannot exist between orbits The higher the energy level, the further it is away from the nucleus An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive )

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How did he develop his theory? He used mathematics to explain the visible spectrum of hydrogen gas

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Radio waves Micro waves Infrared. Ultra- violet X- Rays Gamma Rays Low energy High energy Low Frequency High Frequency Long Wavelength Short Wavelength Visible Light Electromagnetic Spectrum

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The line spectrum Electricity passed through a gaseous element emits light at a certain wavelength The colors can be seen when passed through a prism Every gas has a unique pattern (color)

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Further away from the nucleus means more energy. There is no “in between” energy First Second Third Fourth Fifth Increasing energy }

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Line spectrum of various elements

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Electrons orbiting closest to the nucleus are said to be in the lowest energy state called the ground state Atoms can absorb an amount of energy This promotes an electron to a higher energy level called the excited state This energy level is unstable and so the electron will fall back to its ground state When it does this, the excess energy will be emitted as light

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When the e- falls from one energy level to another, an amount of energy is emitted as light This light emitted at specific wavelengths, which corresponds to our atomic spectra Each atom will have different electron “jumps” therefore emitting different amounts of energy as light This creates different line spectra for various elements

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Let’s watch this video http://www.mhhe.com/physsci/ch emistry/chang7/esp/folder_struct ure/pe/m1/s3/

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More videos (view OYO) http://www.mhhe.com/physsci/astronomy/applets/Bohr/appl et_files/Bohr.html http://www.mhhe.com/physsci/astronomy/applets/Bohr/appl et_files/Bohr.html http://highered.mcgraw- hill.com/sites/0072482621/student_view0/interactives.html# http://highered.mcgraw- hill.com/sites/0072482621/student_view0/interactives.html#

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Bohr’s Triumph His theory helped to explain periodic law Halogens are so reactive because it has one e- less than a full outer orbital Alkali metals are also reactive because they have only one e- in outer orbital

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Drawback Bohr’s theory did not explain or show the shape or the path traveled by the e-. His theory could only explain hydrogen and not the more complex atoms

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The Quantum Mechanical Model Energy is quantized – meaning it comes in chunks. A quanta is 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. An equation has been developed that described the energy and position of the electrons in an atom

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Atomic Orbitals Principal Quantum Number (n) = the energy level of the electron. Within each energy level the complex math equation describes several shapes. These are called atomic orbitals Orbitals are regions where there is a high probability of finding an e-

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S sublevel 1 s orbital for every energy level 1s 2s 3s Spherical shaped Each s orbital can hold 2 electrons Called the 1s, 2s, 3s, etc.. orbitals

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P sublevel Start at the second energy level 3 different directions 3 different shapes Each orbital can hold 2 electrons The p Sublevel has 3 p orbitals

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The D sublevel contains 5 D orbitals The D sublevel starts in the 3 rd energy level 5 different shapes (orbitals) Each orbital can hold 2 electrons

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The F sublevel has 7 F orbitals The F sublevel starts in the fourth energy level has seven different shapes (orbitals) 2 electrons per orbital

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Summary Sublevel# of Orbitals# e- in sublevel Starts in what energy level s121 p362 d5103 f7144

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Electron Configurations The way electrons are arranged in atoms. Aufbau principle- e- enter the lowest energy first. This causes difficulties because of the overlap of orbitals of different energies. Pauli Exclusion Principle- at most 2 e-per orbital with different spins Hund’s Rule- When e- occupy orbitals of equal energy they don’t pair up until they have to.

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

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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

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Electron Configuration for phosphorus The first two electrons go into the 1s orbital Notice the opposite spins only 13 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

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The next electrons go into the 2s orbital only 11 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

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The next electrons go into the 2p orbital only 5 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

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The next electrons go into the 3s orbital only 3 more Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

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Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into separate shapes 3 unpaired electrons 1s 2 2s 2 2p 6 3s 2 3p 3

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Write these electron onfigurations Titanium Vanadium Chromium Copper

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Electron Configurations Titanium - 22 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 Vanadium - 23 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 Chromium - 24 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 Expected. But, this is wrong!! 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Copper – 29 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 Actual configuration is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 Why are chromium and copper configurations different? This gives two half filled orbitals for chromium and one completely full and one half filled orbital for copper. Slightly lower in energy. Remember these exceptions

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Practice

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