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Arrangement of Electrons in Atoms (Chapter 4) Notes Part 1 Electromagnetic Radiation

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I. Properties of Light-Different types of electromagnetic radiation (x-rays, radio waves, microwaves, etc…) SEEM to be very different from one another. Yet they share certain fundamental characteristics. All types of electromagnetic radiation, also called radiant energy, move through a vacuum at a speed of 3.00 x l0 8 meters per second.

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A. Wavelength – distance between identical points on successive waves; may be measured in any length unit but is usually dependent on how long the wave is (X- rays are usually measured in nanometers or Angstroms while the very long radio waves might be measured in meter. The Greek letter lambda,, is used to depict wavelength (pg 92)

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B. Frequency – the number of complete wave cycles that pass a given point in one second: the unit is cycles/second but is written as sec -1, or Hertz. The Greek letter nu,, is used to depict frequency.

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If the frequency and wavelength are known then the product of the two (wavelength x frequency) is always equal to the same speed. It is known as the speed of light or c. c = speed of light = 3.00 x l0 8 m/s c = = wavelength (in m) = frequency (in Hz)

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1. What is the wavelength of radiation whose frequency is 6.24 x l0 13 sec -1 ? A: 4.81X10 - 6 m 2. What is the frequency of radiation whose wavelength is 2.20 x l0 -6 nm? (1 m = 1,000,000,000 nm) A: 1.36X10 23 s -1 or Hz

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II. The Photoelectric Effect (pg 93) – refers to the emission of electrons from a metal when light shines on the metal.

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The wave theory of light (early 1900) could not explain this phenomenon. The mystery of the photoelectric effect involved the frequency of the light striking the metal. For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum – regardless of how long the light was shone. Light was known to be a form of energy, capable of knocking loose an electron from a metal. But the wave theory of light predicted that light of any frequency could supply enough energy to eject an electron. Scientists couldn’t explain why the light had to be of a minimum FREQUENCY in order for the photoelectric effect to occur.

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Let’s Review Energy – Radiation of different wavelengths affect matter differently – certain wavelengths (near infrared) may burn your skin with a heat burn, overexposure to X radiation causes tissue damage. These diverse effects are due to differences in the energy of the radiation. Radiation of high frequency and short wavelength are more energetic than radiation of lower frequency and longer wavelength.

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THE QUANTITATIVE RELATIONSHIP BETWEEN FREQUENCY AND ENERGY WAS DEVELOPED THROUGH THE QUANTUM THEORY OF MAX PLANCK. The explanation of the photoelectric effect dates back to 1900 when Max Planck revised classical ideas of light by proposing that light, which before was thought of as a collection of waves, consisted of BUNDLES OF ENERGY called QUANTA. A quantum is the minimum quantity of energy that can be lost or gained by an atom. Max Planck

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Planck proposed the following relationship between a quantum of energy and the frequency of radiation: E = h h = Planck’s constant = 6.63 x l0 -34 Joules sec E = energy (in Joules) = frequency (in Hz)

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Examples: 1. If a certain light has 7.18 x l0 -19 J of energy what is the frequency of this light? A: 1.08X10 15 s -1 or Hz b. What is the wavelength, in nm, of this light? A: 278 nm 2. If the frequency of a certain light is 3.8 x l0 14 Hz, what is the energy of this light? A: 2.5X10 -19 J

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Albert Einstein expanded on Planck’s theory by explaining that electromagnetic radiation has a dual wave-particle nature. While light exhibits many wavelike properties, it can also be thought of as a stream of particles. Each particle of light carries a quantum of energy. Einstein called these particles PHOTONS. A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy. Albert Einstein

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Einstein explained the photoelectric effect by proposing that electromagnetic radiation is absorbed by matter only in whole numbers of photons. In order for an electron to be ejected from a metal surface, the electron must be struck by a single photon possessing at least the minimum energy (E photon = hv) required to knock the electron loose, this minimum energy corresponds to a minimum frequency. If a photon’s frequency is below the minimum, then the electron remains bound to the metal surface. Electrons in different metals are bound more or less tightly, so different metals require different minimum frequencies to exhibit the photoelectric effect.

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Example from problem 4: An atom or molecule emitting or absorbing radiation whose wavelength is 589 nm cannot lose or gain energy by radiation except in MULTIPLES OF 3.37x l0 -19 J. It cannot, for example, gain 5.00 x l0 -19 J from this radiation because this amount is not a multiple of 3.37 x l0 -19.

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5. In astronomy, it is often necessary to be able to detect just a few photons because the light signals from distant stars are so weak. A photon detector receives a signal of total energy 4.05 x l0 -18 J from radiation of 540 nm wavelength. How many photons have been detected? A: 11 photons 6. Excited chromium atoms strongly emit radiation of 427 nm. What is the energy in kilojoules per photon? A: 4.66X10 -22 kJ 7. Light hitting certain chemical substances may cause rupture of a chemical bond. If a minimum energy of 332 kJ is required to break a carbon-chlorine bond in a plastic material, what is the longest wavelength of radiation that possesses the necessary energy? A: 5.99X10 -31 m

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III. The Hydrogen-Atom Line-Emission Spectrum When investigators passed an electric current through a vacuum tube containing hydrogen gas at low pressure, they observed the emission of a characteristic pinkish glow. When a narrow beam of the emitted light was shined through a prism, it was separated into a series of specific frequencies (and therefore specific wavelengths, c = ) of visible light. The bands of light were part of what is known as hydrogen’s LINE- EMISSION SPECTRUM. (page 95) The lowest energy state of an atom is its ground state. A state in which an atom has a higher amount of energy is an excited state. When an excited atom returns to its ground state, it gives off energy.

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IV. Bohr’s Model of Hydrogen – Neils Bohr incorporated Planck’s quantum theory to explain line-emission spectra. Bohr said the absorptions and emissions of light by hydrogen corresponded to energy changes within the atom. The fact that only certain frequencies are absorbed or emitted by an atom tells us that only certain energy changes are possible. Bohr’s model incorporated (l) Rutherford’s Experiment, which established a nucleus and (2) Einstein’s theory that used Planck’s quantum theory to determine that light is discrete bundles of energy.

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V. Bohr’s Theory of the Atom: Electrons cannot have just any energy; only orbits of certain radii having CERTAIN energies are permitted. Thus, when an electron absorbs quanta of energy, it will cause them to jump away from the nucleus to a higher orbit (energy level or n) and when the electron falls from a high orbit to a lower one, a photon of a particular wavelength is released, and a particular color will be given off. Bohr was able to calculate a set of allowed energies. Each of these allowed energies corresponds to a circular path of a different radius. Thus the larger the value of n, the farther the electron is from the nucleus and the higher energy it possesses. The success of Bohr’s model of the hydrogen atom in explaining observed spectral lines led many scientist to conclude that a similar model could be applied to all atoms. It was soon recognized, however, that Bohr’s approach did not explain the spectra of atoms with more than one electron. Nor did Bohr’s theory explain the chemical behavior of atoms.

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V. Bohr’s Theory of the Atom:

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Electrons in Atoms (Chapter 4) Notes Part 2 Quantum Model of the Atom

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So where are the electrons of an atom located? A. Various Models of the Atom Dalton’s Model Thomson’s Plum Pudding Model Rutherford’s Model Bohr’s ‘Solar System’ Model – electrons rotate around the nucleus Quantum Mechanics Model – modern description of the electron in atoms, derived from a mathematical equation (Schrodinger’s wave equation)

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B. In 1926, the Austrian physicist Erwin Schrodinger used the hypothesis that electrons have a dual wave/particle nature (developed by Louis de Broglie in 1924) to develop an equation that treated electrons in atoms as waves. Erwin Schrodinger

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Electrons as Waves Louis de Broglie (1924) Applied wave-particle theory to electrons electrons exhibit wave properties QUANTIZED WAVELENGTHS Adapted from work by Christy Johannesson www.nisd.net/communicationsarts/pages/chem Standing Wave 200 150 100 50 0 - 50 -100 -150 -200 0 50 100 150 200 Second Harmonic or First Overtone 200 150 100 50 0 - 50 -100 -150 -200 0 50 100 150 200 Fundamental mode 200 150 100 50 0 - 50 -100 -150 -200 0 50 100 150 200 Louis de Broglie ~1924

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Electrons as Waves QUANTIZED WAVELENGTHS Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem n = 4 n = 6 Forbidden n = 3.3 n = 5

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Schrodinger’s equation results in a series of so called wave functions, represented by the letter (psi). Although has no actual physical meaning, the value of 2 describes the probability distribution of an electron. (Same concept covered in Algebra II when dealing with linear regressions and finding best fit lines.) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Electron Probability vs. Distance Electron Probability (%) Distance from the Nucleus (pm) 100150200250500 0 10 20 30 40 Orbital 90% probability of finding the electron

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We cannot know both the location and velocity of an electron (Heisenberg’s uncertainty principle), thus Schrodinger’s equation does not tell us the exact location of the electron, rather it describes the probability that an electron will be at a certain location in the atom. Here is an overview of electron properties: 1. Waves are confined to a space and can only have certain frequencies. 2. Electrons are considered waves confined to the space around an atomic nucleus. Electrons can only exist at specific frequencies. And according to E=hv (Planck’s hypothesis), these frequencies correspond to specific energies (or quantified amounts of energy.) 3. Electrons, like light waves, can be bent or diffracted. Orbital 90% probability of finding the electron

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C. Heisenberg’s Uncertainty Principle says that there is a fundamental limitation on just how precisely we can hope to know both the location and the momentum of a particle. It turns out that when the radiation used to locate a particle hits that particle, it changes its momentum. Therefore, the position and momentum cannot both be measured exactly. As one is measured more precisely, the other is known less precisely. Today we say that the electrons are located in a region outside the nucleus called the electron cloud. Werner Heisenberg

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Heisenberg Uncertainty Principle Impossible to know both the velocity and position of an electron at the same time Microscope Electron Werner Heisenberg ~1926

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I. Electron Cloud – Energy Levels Electrons are found in various energy levels around the nucleus. The energy levels are analogous to the rungs of a ladder. The lowest rung of the ladder corresponds to the lowest energy level. A person can climb up or down a ladder by going from rung to rung. Similarly, an electron can jump from one energy level to another. A person on a ladder cannot stand between the rungs; similarly, the electrons in an atom cannot exist between energy levels.

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A. Quantum: To move from one rung to another, a person climbing a ladder must move just the right distance. To move from one energy level to another, an electron must gain or lose just the right amount of energy. The exact amount of energy required to move from one energy level to another is called a quantum of energy. B. Photon: When electrons move from one energy level to another energy level we see light – going from one energy level to another energy level gives off an exact amount of light (called a photon). Electron Absorbing Energy (Photon) Electron will move from a ground state to an excited state. Electron Emitting Energy (Photon) Electron will move from an excited state to a ground state.

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II. Quantum Mechanics Model of the Atom and Quantum Numbers Periodic Table with predicted ending electron configurations. Quantum Numbers – a series of numbers which describe several properties of an energy level (or orbit)

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Quantum Numbers UPPER LEVEL Four Quantum Numbers: Specify the “address” of each electron in an atom Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

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Quantum Numbers Principal Quantum Number n Principal Quantum Number ( n ) Angular Momentum Quantum # l Angular Momentum Quantum # ( l ) Magnetic Quantum Number m l Magnetic Quantum Number ( m l ) Spin Quantum Number Spin Quantum Number ( m s )

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A. Principal Quantum Number, “n” (Energy Levels): energy levels (represented by the letter n) are assigned values in order of increasing energy: n=1,2,3,4, and so forth…. which correspond to the periods in the periodic table. The principle q. n. is related to the size and energy of the orbital. n=1, n=2, n=3, n=4, n=5, etc… Which energy level is furthest away from the nucleus and has electrons with the highest energy - 1, 2,3, or 4?

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Relative Sizes 1s and 2s 1s 2s Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 334

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Quantum Numbers Principal Quantum Number n Principal Quantum Number ( n ) Energy level Size of the orbital n 2 = # of orbitals in the energy level Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 1s1s 2s2s 3s3s

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B. Angular Momentum or Azimuthal Quantum Number, “l” (Sublevels): Within each energy level, the electrons are located in various sublevels – there are 4 different sublevels s, p, d, and f. “l” defines the shape of the orbital (s, p, d, & f). The possible values of “l” are limited by the value for “n”. If n = 3, “l” can be 0, 1, or 2, but not 3 or higher. This q.n. is related to the shape of the orbital.

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Shapes of s, p, and d-Orbitals s orbital p orbitals d orbitals

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p-Orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 335 pxpx pypy pzpz

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d-orbitals Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 336

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Atomic Orbitals

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Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

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l = 0, is referring to the s sublevel l = 1, is referring to p sublevel l = 2, is referring to d sublevel l = 3, is referring to f sublevel 1s 2s 2p 3p

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Quantum Numbers s p d f Angular Momentum Quantum # l Angular Momentum Quantum # ( l ) Energy sublevel Shape of the orbital Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

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Quantum Numbers Orbitals combine to form a spherical shape. 2s 2p z 2p y 2p x Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

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Quantum Numbers Magnetic Quantum Number m l Magnetic Quantum Number ( m l ) Orientation of orbital Specifies the exact orbital within each sublevel Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

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Feeling overwhelmed? Read chapter 4.2! Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem "Teacher, may I be excused? My brain is full." Chemistry

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C. Orbitals: Where are the electrons in the various sublevels located in relation to the nucleus? Electrons are NOT confined to a fixed circular path, they are, however, found in definite regions of the atoms – these regions are called atomic orbital’s! Each orbital can only hold 2 electrons at a time (Pauli exclusion principle).

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Within the s sublevel (l=0) there is only 1 orbital (which is spherical) it is called the s orbital. http://www.shef.ac.uk/chemistry/orbitron/AOs/1s/index.html http://www.shef.ac.uk/chemistry/orbitron/AOs/1s/index.html Within the p sublevel (l=1) there are 3 orbital’s (which are dumbbell shaped) called the p x, p y, p z orbital’s. Within the d sublevel (l=2) there are 5 orbital’s (4 of which are cloverleaf shaped) called the d xy, d xz, d yz, d x 2 -y 2, d z 2 orbital’s. Within the f sublevel (l=3) there are 7 orbital’s - which are too complex to draw

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The magnetic quantum number, m l, refers to the position of the orbital in space relative to other orbital’s. It may have integral numbers ranging from 0 in the s sublevel, 1 to –1 in the p sublevel, 2 to –2 in the d sublevel and 3 to –3 in the f sublevel.

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m l = 0, is referring to the s orbital m l = -1, 0, +1, are referring to the three p orbital’s (px, py, and pz) m l = -2, -1, 0, +1, +2, are referring to the five d orbitals m l = -3, -2, -1, 0, +1, +2, +3, are referring to the seven f orbitals

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D. D. How many electrons can go into each energy level? Each orbital can hold two electrons. (2n 2 = number of electrons per energy level) The 1st energy level (n=1) only has 1 sublevel called 1s. s only has 1 orbital called the s orbital, so only 2 electrons will be found in the 1st energy level. (2n 2 = 2)

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The 2nd energy level (n=2) has 2 sublevels called 2s and 2p. s only has 1 orbital called the s orbital, p has 3 orbital’s called p x, p y, and p z orbitals, so 8 electrons will be found in the 2nd energy level. (2n 2 = 8)

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The 3rd energy level (n=3) has 3 sublevels called 3s, 3p, and 3d. s only has 1 orbital called the s orbital, p has 3 orbital’s called p x, p y, and p z orbitals, and d has 5 orbital’s, so 18 electrons will be found in the 3rd energy level. (2n 2 = 18)

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How about the 4th energy level? It has 4 sublevels called 4s, 4p, 4d, and 4f. s only has 1 orbital, p has 3 orbital’s, d has 5 orbital’s, and f has 7 orbitals, so 32 electrons will be found in the 4th energy level. (2n 2 = 32)

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E. Lets put it all together: Example of neon atom:

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Fourth Quantum Number, m s, refers to the magnetic spin of an electron within an orbital. Each orbital can hold two electrons, both with different spins. Clockwise spin is represented with a value of +1/2 and counterclockwise spin is represented with a value of –1/2. Electrons fill the orbital’s one at a time with the same spin (+1/2), then fill up the orbital(s) with electrons of the opposite spin (-1/2). m s = +1/2 or –1/2

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Quantum Numbers Spin Quantum Number 4. Spin Quantum Number ( m s ) Electron spin +½ or -½ An orbital can hold 2 electrons that spin in opposite directions. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

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Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.

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Quantum Numbers Analogy Energy Levels (n) or Principal Q.N. n=1 (Weir) n=2 (Liberty Hill) n=3 (Georgetown)n=4 (Austin) Sublevels (l) or Azimuthal Q.N. l=0 – s shape 1 bedroom l=1 – p shape 3 bedroom l=2 – d shape 5 bedroom l=3 – f shape 7 bedroom Orbitals (m l ) or Magnetic Q.N. If l=0 then m l =0 (Represents the 1 bed/orbital in the s sublevel) If l=1 then m l = -1, 0, 1 (Represents the 3 bed’s/orbital’s in the p sublevel) If l=2 then m l = -2, -1, 0, 1, 2 (Represents the 5 bed’s/orbital’s in the p sublevel) If l=3 then m l = -3, -2, -1, 0, 1, 2, 3 (Represents the 7 bed’s/orbital’s in the p sublevel) Magnetic Spin – Fourth Q.N. (m s ) m s = +1/2 - 1 st electron in orbital m s = -1/2 – 2 nd electron in orbital

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Level n 1 2 3 Sublevel l Orbital m l Spin m s 00 00 1 0 0 1 0 21 0 -2 2101 = +1/2 = -1/2 Allowed Sets of Quantum Numbers for Electrons in Atoms

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Maximum Number of Electrons In Each Sublevel Maximum Number of Electrons In Each Sublevel Maximum Number SublevelNumber of Orbitals of Electrons s 1 2 p 3 6 d 5 10 f 7 14 LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 146

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Quantum Numbers n shell l subshell m l orbital m s electron spin 1, 2, 3, 4,... 0, 1, 2,... n - 1 - l... 0... +l + 1 / 2 and - 1 / 2

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Electrons In Atoms Notes (Chapter 4) Part 3 Electron Configurations

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I. Electron Configuration: It should be obvious to you now that it is very difficult to draw a representation or model of atom showing where the electrons are located, so what we do instead is write electron configurations for elements. Definition of electron configuration: An electron configuration is a written representation of the arrangement of electrons in an atom.

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When constructing orbital diagrams and electron configurations, keep the following in mind: Aufbau Principle – electrons fill in order from lowest to highest energy. The Pauli Exclusion Principle – An orbital can only hold two electrons. Two electrons in the same orbital must have opposite spins. You must know how many electrons can be held by each angular momentum number, l. (ie; s can hold 2, 6 for p, l0 for d, 14 for f) Hund’s rule – the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons for a set of degenerate orbitals. By convention, all unpaired electrons are represented as having parallel spins with spin “up”.

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Filling Rules for Electron Orbitals Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for. Pauli Exclusion Principle: An orbital can hold a maximum of two electrons. To occupy the same orbital, two electrons must spin in opposite directions. Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results. *Aufbau is German for “building up”

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Quantum Numbers 1. Principal # 2. Ang. Mom. # 3. Magnetic # 4. Spin # energy level sublevel (s,p,d,f) orbital electron Pauli Exclusion Principle No two electrons in an atom can have the same 4 quantum numbers. Each electron has a unique “address”: Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Wolfgang Pauli

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II. What? How do we write an electron configuration? A. 1st rule - electrons occupy orbitals that require the least amount of energy for the electron to stay there. So always follow the vertical rule (Aufbau Principle): You notice, for example, that the 4s sublevel requires less energy than the 3d sublevel; therefore, the 4s orbital is filled with electrons before any electrons enter the 3d orbital!!!! (So just follow the chart and you can’t go wrong!!!!)

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II. What? How do we write an electron configuration? 1st rule - electrons occupy orbitals that require the least amount of energy for the electron to stay there. So always follow the vertical rule (Aufbau Principle):

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You notice, for example, that the 4s sublevel requires less energy than the 3d sublevel; therefore, the 4s orbital is filled with electrons before any electrons enter the 3d orbital!!!! (So just follow the chart and you can’t go wrong!!!!)

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B. 2nd rule – only 2 electrons can go into any orbital, however, you must place one electron into each orbital in a sublevel before a 2nd electron can occupy an orbital. Orbital’s with only 1 electron in the orbital are said to have an unpaired electron in them.

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III. Writing Electron Configurations (3 ways): A. Orbital Notation: an unoccupied orbital is represented by a line______, with the orbitals name written underneath the line. An orbital containing one electron is written as __ ___, an orbital with two electrons is written as __ __. The lines are labeled with the principal quantum number and the sublevel letter.

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Examples: (Remember that you must place one electron into each orbital before a second electron in placed into an orbital.) Hydrogen _ ___Helium _ _ 1s 1s Lithium _ __ _ ___ 1s 2s Carbon __ __ __ __ __ __ __ __ _____ 1s 2s 2px 2py 2pz You try to write the notation for Titanium

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H = 1s 1 1s1s He = 1s 2 1s1s Li = 1s 2 2s 1 1s1s 2s2s Be = 1s 2 2s 2 1s1s 2s2s C = 1s 2 2s 2 2p 2 1s1s 2s2s 2px2px 2py2py 2pz2pz S = 1s 2 2s 2 2p 4 1s1s 2s2s 2px2px 2py2py 2pz2pz 3s3s 3px3px 3py3py 3pz3pz THIS SLIDE IS ANIMATED IN FILLING ORDER 2.PPTFILLING ORDER 2.PPT

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H = 1s 1 1s1s He = 1s 2 1s1s Be = 1s 2 2s 2 1s1s 2s2s +1 e-e- +2 e-e- e-e- +4 e-e- e-e- e-e- e-e-

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B. Electron Configuration Notation: eliminates the lines and arrows of orbital notation. Instead, the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation. The superscript indicates the number of electrons present in that sublevel.

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Examples: Hydrogen: 1s 1 Helium: 1s 2 Lithium: 1s 2 2s 1 Carbon: 1s 2 2s 2 2p 2 You try to write the notation for Titanium

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Fe = 1s 1 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 1s1s 2s2s 2px2px 2py2py 2pz2pz 3s3s 3px3px 3py3py 3pz3pz +26 e-e- e-e- e-e- e-e- 4s4s 3d3d 3d3d3d3d 3d3d Iron has ___ electrons. 26 3d3d Arbitrary Energy Scale 18 32 8 8 2 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e-

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Orbital Filling Element 1s 2s 2p x 2p y 2p z 3s Configuration Orbital Filling Element 1s 2s 2p x 2p y 2p z 3s Configuration Electron Configurations Electron H He Li C N O F Ne Na 1s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 5 1s 2 2s 2 2p 4 1s 2 2s 2 2p 3 1s 2 2s 2 2p 2 1s 2 2s 1 1s 2 NOT CORRECT Violates Hund’s Rule Electron Configurations Electron H He Li C N O F Ne Na 1s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 5 1s 2 2s 2 2p 4 1s 2 2s 2 2p 3 1s 2 2s 2 2p 2 1s 2 2s 1 1s 2

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C. Short Hand or Noble Gas Notation: Use the noble gases that have complete inner energy levels and an outer energy level with complete s and p orbital’s. Use the noble gas that just precedes the element you are working with.

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Boron is ls 2 2s 2 2p 1 The noble gas preceding Boron is He, so the short way is [He]2s 2 2p 1. Sulfur is ls 2 2s 2 2p 6 3s 2 3p 4 Short way: [Ne]3s 2 3p 4 Example: Titanium

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More Practice Problems: Write electron configurations for each of the following atoms: 1. boron 2. sulfur 3. vanadium 4. iodine Draw orbital diagrams for these: 5. sodium 6. phosphorus 7. chlorine Write shorthand electron configuration for the following: 8. Sr 9. Mo 10. Ge

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Irregular Electron configurations – sometimes the electron configuration is NOT what we would predict it to be. Sometimes electrons are moved because (l) it will result in greater stability for that atom or (2) for some unknown reason??

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It is very important to define “stable” here. STABLE means: 1. all degenerate (equal energy) orbital’s are FULL 2. all degenerate orbital’s are half-full 3. all degenerate orbital’s are totally empty.

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Examples – draw the orbital’s (lines or boxes) and fill each orbital with the predicted number of electrons. Predict the electron configuration for Cr #24: [Ar]4s 2 3d 6 However, the real E. C. is [Ar]4s 1 3d 5. The 4s 1 electron has been moved to achieve greater stability. ALWAYS USE THE ACTUAL E. C. AND NOT THE PREDICTED ONE. YOU WILL HAVE THESE ATOMS WITH IRREGULAR E. C. HIGHLIGHTED OR MARKED ON YOUR PERIODIC TABLE.

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Electron configurations for Ions-First, determine if the element will lose or gain electrons. Secondly, what number of electrons will be gained or lost? It is recommended that you write the e.c. for the atom and then determine what will happen.

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For cations (positive ions) – look at the element and decide how many electrons will be lost when it ionizes and keep that in mind when writing the E. C. The last number in the E. C. will now be LESS than what is written on your periodic table. Ex. Write the electron configuration for magnesium ion: [Ne]3s 2 is for the atom. Mg is a metal and will lose its valence (outer) electrons, so the e.c. for Mg 2+ is 1s 2 2s 2 2p 6 Practice: 1. #3 2. #12 3. #19 4. #13

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For anions (negative ions) – look at the element and decide how many electrons that element will GAIN when it ionizes. The last number in the E. C. will be MORE than what is written on the periodic table. Ex. Sulfide ion: Sulfur atom is 1s 2 2s 2 2p 4. Sulfur is a nonmetal with 6 valence electrons (2s 2 and 2p 4 ) and will gain 2 electrons: 1s 2 2s 2 2p 6 is for the sulfide ion. Practice: 1. #17 2. #7 3. #16 4. #30

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