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- 1. Electrons in Atoms<br />1<br />
- 2. GPS Standards<br />SC3a – Discriminate between the relative size, charge, and position of protons, neutrons, and electrons in the atom.<br />Identify the inadequacies in the Rutherford atomic model.<br />Identify the new proposal in the Bohr model of the atom.<br />Describe the energies and positions of electrons according to the quantum mechanical model.<br />Describe how the shapes of orbitals related to different sublevels differ.<br />2<br />
- 3. Essential Question <br />How are Rutherford’s, Bohr’s, and the quantum mechanical models related to each other?<br />3<br />
- 4. Notes<br />Inadequacies in Rutherford’s Model<br />Could not explain why metals and metal compounds give off characteristic colors when heated in a flame<br />Could not explain why heated metals glow red, then yellow, then white <br />Could not explain the chemical properties of elements<br />Treated the electron as a particle<br />4<br />
- 5. The Bohr Model<br />Revised Rutherford’s model to include information about how the energy of an atom changes when it absorbs or emits light<br />Proposed that an electron is found only in specific circular paths, or orbits, around the nucleus<br />Each proposed orbit has a fixed energy called an energy level<br />Higher the energy of an electron, the farther it is from the nucleus<br />Quantum – the amount of energy required to move an electron from one energy level to another energy level<br />Gave results in agreement with experiments for the hydrogen atom<br />Failed to explain the energies absorbed and emitted by atoms with more than one electron<br />Treated the electron as a particle<br />5<br />
- 6. Quantum Mechanical Model<br />Schrodinger<br />Devised a mathematical equation describing electron as a wave<br />Quantum mechanical model <br />modern description of the electrons around an atom<br />based on mathematical solutions to Shrödinger’s equation<br />Based on the probability of finding an electron within a particular volume of space around the nucleus<br />By treating the electron as an electron wave instead of a particle, most of the problems associated with Bohr’s model were alleviated. There are still some problems that we will look at later. The model is still a work in progress.<br />6<br />
- 7. September 7, 2011<br />Essential Question<br />How are quantum numbers used to describe electrons?<br />7<br />
- 8. Quantum Numbers<br />Each electron around an atom has a set of 4 quantum numbers which describe the “energy address” of the electron.<br />Principal quantum number (n)<br />First quantum number<br />Represents the energy level in which the electron is found (larger value of n = higher energy)<br />Determines the size of an orbital (larger value of n = larger orbital size)<br />The values of n are successive integers beginning with 1 (n = 1, 2, 3, 4, …., )<br />Each energy level represents 1 period on the periodic table.<br />Maximum number of orbitals in an energy level = n2<br />Maximum number of electrons in an energy level = 2n2<br />8<br />
- 9. Angular momentum quantum number (l)<br />Designates the shape of the orbital in which the electron is found<br />Indicates the sublevel of the electron<br />Values of l = successive integers from zero to n-1 (l = 0, 1, 2, …., n-1)<br />Each energy level has a number of sublevels equal to the value of n.<br />Energy level n=1 has 1 sublevel (l=0)<br />Energy level n=2 has 2 sublevels (l=0 and l=1)<br />Energy level n=3 has 3 sublevels (l=0, l=1, and l=2)<br />Energy level n=4 has 4 sublevels (l=0, l=1, l=2, l=3)<br />Commonly used labels of the sublevels<br />l=0 is the s-sublevel<br />l=1 is the p-sublevel<br />l=2 is the d-sublevel<br />l=3 is the f-sublevel<br />9<br />
- 10. 10<br />
- 11. Magnetic quantum number (ml)<br />Determines the orientation of the orbital within the sublevel<br />Each energy level has an s-sublevel that contains 1 s-orbital<br />Beginning with the 2nd energy level, each energy level has a p-sublevel containing 3 p-orbitals.<br />Beginning with the 3rd energy level, each energy level has a d-sublevel containing 5 d-orbitals.<br />Beginning with the 4th energy level, each energy level has an f-sublevel, containing 7 f-orbitals<br />Values of ml are integers from –l to +l<br />Orbital – a region in the space surrounding the nucleus where the probability of finding an electron is above 90%<br />11<br />
- 12. Spin quantum number (ms)<br />Each orbital can hold a maximum of 2 electrons.<br />Spin makes the electron act like a tiny magnet<br />Values of ms are +1/2 or -1/2<br />12<br />
- 13. Orbital filling diagrams<br />Show all 4 quantum numbers for each electron surrounding an atom<br />V: 23e-<br /><br /><br /><br /><br /><br /><br /><br /><br /><br /><br /><br /><br /><br />1s<br />2s<br />2p<br />3s<br />3p<br />4s<br />3d<br />13<br />
- 14. Aufbau Principle<br />Electrons occupy the orbitals of least energy first.<br />Always fill one sublevel before adding electrons to a higher energy sublevel.<br />Hund’s Rule<br />Electrons occupy orbitals of the same energy level in a way that makes the number of electrons with the same spin direction as large as possible.<br />Always add 1 electron to each orbital in a sublevel before adding a second electron to any orbital in that sublevel.<br />Pauli’s Exclusion Principle<br />No 2 electrons in the same atom can have the exact same four quantum numbers<br />When 2 electrons occupy the same orbital, they must have opposite spins.<br />14<br />
- 15. Complete electron configuration<br />Shows the energy level (principal quantum number), sublevel (angular momentum quantum number) and the number of electrons in that sublevel.<br />Coefficient = energy level<br />Letter = sublevel<br />Superscript = number of electrons<br />The sum of the superscripts should equal the atomic number of the element.<br />V: 23e-<br />1s22s22p63s23p64s23d3<br />15<br />
- 16. Noble gas configuration<br />Uses the symbol of the previous noble gas in brackets to represent the configuration of the inner energy levels<br />Vanadium: [Ar]4s23d3<br />16<br />
- 17. Electron-dot diagrams<br />Shows only the electrons in the outermost energy level<br />For elements in the s-block and p-block, the number of dots equals the last number in the group number<br />For transition elements, the number of dots is 2 for all elements other than those in groups 6 and 11. These two groups will exhibit 1 dot.<br />The symbol of the element represents all inner electrons.<br />17<br />
- 18. Physics & the Quantum Mechanical Model<br />Light<br />Sir Isaac Newton<br />Tried to explain light behavior by assuming that light travels as a particle but other evidence convinced scientists that light travels as a wave<br />18<br />
- 19. Wave properties of light<br />Amplitude – height from the equilibrium position to the crest or trough<br />Wavelength () – distance between two crests<br />Frequency () <br />Number of waves that pass a given point per second<br />Measured in hertz (Hz)<br />1 Hz = 1 wave per second<br />Speed of light (c)<br />A constant (2.998 x 108 m/s in a vacuum)<br />c = <br />speed of light(m/s) = wavelength(m) x frequency(Hz)<br />Wavelength and frequency are inversely proportional (seesaw relationship)<br />19<br />
- 20. Electromagnetic radiation<br />Includes visible light as well as infrared, ultraviolet, gamma, x-rays, radio waves, etc. (see p. 139)<br />Continuous spectrum<br />All of the different frequencies of light coming from light source as seen through a prism<br />Sunlight contains all of the frequencies of light<br />Each color blends into the next as in a rainbow<br />20<br />
- 21. Atomic Spectra<br />When atoms absorb energy, electrons move into higher energy levels<br />When electrons lose energy, they fall into lower energy levels by emitting the same amount of energy as light.<br />21<br />
- 22. Atomic emission spectrum<br />Each fall of an electron to a lower energy orbital represents a specific frequency of light which corresponds to a particular color.<br />When light from an excited atom is passed through a prism, the frequencies represented by the changes in energy of the electrons are separated into distinct lines.<br />Each line represents a single movement of an electron to a lower energy orbital.<br />22<br />
- 23. 23<br />
- 24. Emission spectrum of an element is like a fingerprint for that element and can be used to identify the element.<br />Explanation of Atomic Spectra<br />Ground state<br />Electrons are in their lowest possible energy states<br />Excited state<br />Electrons have moved into higher energy orbitals by absorbing energy<br />Max Planck<br />Determined the relationship between the energy of a quantum (photon) and the frequency of light.<br />E = h <br />h = Planck’s constant = 6.626 x 10-34 Js<br />24<br />
- 25. 25<br />
- 26. Quantum Mechanics<br />Einstein<br />Revisited the concept of light as a particle<br />Called a quantum of light a photon<br />Won the nobelprize for his explanation of the photoelectric effect<br />26<br />
- 27. de Broglie<br />Based on the dual wave/particle nature of light, proposed a similar duality for the electron, calling the wavelike behavior of particles matter waves<br />All moving objects exhibit wavelike behavior; however the mass must be very small in order for the wavelength to be large enough to observe.<br />Davison and Germer<br />Found experimental evidence to support de Broglie’s claim that electrons travel as waves<br />Heisenberg<br />Heisenberg Uncertainty Principle<br />It is impossible to know exactly both the velocity and position of a particle at the same time.<br />27<br />
- 28. 28<br />
- 29. 29<br />
- 30. 30<br />
- 31. 31<br />

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