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9/17/14 By convention there is color, By convention sweetness, By convention bitterness, But in reality there are atoms and space. -Democritus (c. 400 BCE) Development of Early Atomic Theory Democritus Through Rutherford Democritus (460-370 BC) • Greek Philosopher • Pondered the composition of matter • Matter is made up of tiny indivisible particles called ________________ Law of Conservation of Mass • Mass is neither created or destroyed in chemical reactions. The atoms are _________________. Hg(NO3)2 + 2KI → HgI2 + 2KNO3 1 9/17/14 Which diagram best represents the LOCM? Heat Heat Law of Definite Proportions • Formulated by Joseph Proust • Different samples of a pure substance always contain the ________________ proportion of elements by ____________ • Water is 11% H and 89% O regardless of its source. • Elements combine in ___________________ proportions, not in random proportions. Dalton’s Atomic Theory • Elements are made of ______________________________________. • Atoms of the same element are the ______________ and atoms of different elements are ______________________. • Atoms combine in simple ____________________________________________. • Atoms are not _____________________________________________ in a chemical reaction. They are just rearranged. 2 9/17/14 Radioactivity • Becquerel, 1896 • Alpha-α, ________________, heaviest, lowest energy,positive charge, _________________________________________ • Beta-β, ___________________, light weight, higher energy, negative charge, _____________________ • Gamma-γ, ___________________, highest energy, _________________________________________ http://highered.mcgraw-hill.com/sites/0072512644/student_view0/chapter2/ animations_center.html# J.J. Thomson and Cathode Ray Tubes-1897 • Objects cast shadows. • Rays can be formed with different metals. • The rays exert a ___________________. • Rays are deflected by ___________________ and magnetic fields. • J.J. Thomson was the first person to deflect the rays. • Called the particles corpuscles. • The term ______________________ coined in 1891 by G. Johnstone Stoney. http://highered.mcgraw-hill.com/sites/ 0072512644/student_view0/chapter2/ animations_center.html# Paddle Wheel Maltese Cross Simulation Plum Pudding Model 3 9/17/14 The Oil Drop Experiment • ____________________ (1868-1953) • Suspended a charged oil drop between two plates • The _________________ of the electron was determined to be 1.61 x10-19 C. • The mass of the electronhttp://highered.mcgraw-hill.com/sites/0072512644/ student_view0/chapter2/animations_center.html# explanation simulation Rutherford Gold Foil Experiment http://cameron4gfs.glogster.com/Gold-foil-experiment/ Summary • Experiment performed by Ernest ____________________ • Most of the particles passed straight through as expected— atoms are mostly ________________________! • However, some were _________________ at angles and some bounced back— atom contains small, dense, positively charged ___________________ • Called the planetary or ________________ model 4 9/17/14 The Atom Three Major Subatomic Particles Nuclear Symbols Isotopes Average Atomic Mass Three Major Subatomic Particles Particle Symbol Relative Charge Mass Number Rel. Mass Act. Mass Electron 0.0005u 9.11•10-28 g Proton 1.007 u 1.637•10-24 g Neutron 1.009 u 1.675•10-24 g 1 u = 1.66•10-24 g Carbon has 6 p+ and 6no in its nucleus. By international agreement, 1 amu is 1/12 the mass of a carbon atom. Symbols on the Periodic Table 5 9/17/14 Isotopes and Mass Number • Isotopes - Atoms of the _________ ___________that have different masses. - Mass difference is due to the number of ______________ in the nucleus. - To identify an isotope, you must know the ______________ of the element and the _________ of the isotope - A _________ is an individual isotope. Mass Number, A • The total number of protons and neutrons in the nucleus of an atom. Mass Number, A = _______________+_______________ • Uranium that has 92 protons and 143 neutrons. • Uranium that has 92 protons and 146 neutrons. Representing Isotopes • Individual isotopes can be represented using _____________ ____________ or the name of the element followed by the mass number of the isotope. 14 6 C Carbon-14 6 9/17/14 Nuclear Symbols • A = • Z = A Z X Charge • Charge = • Positive • Neutral • Negative Examples # 1 – 4 • Write the nuclear symbol for an atom that has 15 protons, 16 neutrons and 15 electrons. • Write the nuclear symbol for an element that has 30 protons, 28 electrons and a mass number of 65. • Calculate the number of protons, neutrons, and electrons represented by the nuclear symbol on the 27 3+ right. 13 • An element has an atomic number of 12, 13 neutrons in the nucleus, and 10 electrons. Write the nuclear symbol for this ion. Al Isotopes of Hydrogen MCMURRAY, J., & Fay, R. C. (2001). Isotopes of Hydrogen. In Chemistry (3rd ed., p. 46). Upper Saddle River, NJ: Prentice Hall. 7 9/17/14 Average Atomic Mass • The average atomic mass is a ____________average of the masses of all of the isotopes for a given element. • Average atomic mass is determined using the following formula: A.A.M. = (Mass I #1)(Fractional Abundance #1)+(Mass I#2)(Fractional Abundance#2)+… Isotope Problem #1 • Calculate the average atomic mass for Ne. Isotope Atomic Mass Percent Abundance Ne-20 20.00 amu 90.92 Ne-21 21.00 amu 0.26 Ne-22 22.00 amu 8.82 Fractional Abundance Isotope Problem #2 • Natural chlorine is a mixture of isotopes. Determine its atomic weight if 75.53% of the naturally occurring element is chlorine 35, and 24.47% is chlorine 37. Fogiel, M., Dr. (Ed.). (1996). Problem 708. In REA's Problem Solvers Chemistry (p. 672). Piscataway, NJ: Research and Education Association. 8 9/17/14 Isotope Problem #3 • Only two isotopes of copper occur naturally, copper-63 and copper-65. Their masses are 62.9298 amu and 64.9278 amu respectively. If the average atomic mass for copper is 63.546 amu calculate the relative abundance of each isotope. The Periodic Table Periods / Series • ______ on the periodic table • Represent the shell that the valence electrons occupy. • __________Electrons – ____________ electrons in an atom. – The electrons that are ________, _________, or ________ in chemical reactions. – The _____________ of valence electrons in an element's outer shell determines the ______________ properties of that element. 9 9/17/14 Groups/Families • ____________ • All elements in the same group have the same number of _________ electrons, and therefore have similar __________properties. • Some groups have special names Special Groups • • • • Group 1 Group 2 Group 7 Group 8 - • ___________ Elements • In addition some regions of the periodic table have special names. Metals • Located on the _____ side of the periodic table. • ______ valence electrons in chemical reactions. • ________,________, ________,________ • Transition metals. 10 9/17/14 Nonmetals • Located on the ______ side of the periodic table. • _______,__________ • ______ electrons in chemical reactions. • __________ ______: Every atom wants ___ _________electrons. Metalloids • Located on the _________-_____ line. • Have ______ and _______properties. Electronic Structure of Atoms II Chapter 13 11 9/17/14 Wave nature of light • ___________________________________— carries energy through space – Examples: visible light, x-rays, radio waves – Speed of light in a vacuum c = 3.00x108 m/s • ____________________ (λ)—distance between two adjacent crests or troughs • ____________________ (ν)—number of complete wavelength cycles that pass a given point each second (units: Hz or s-1) Wave nature of light • Wavelength and frequency vary inversely – νλ = c • high frequency, short wavelength • low frequency, longer wavelength What is the amplitude of a wave? http://www.qrg.northwestern.edu/projects/vss/docs/Communications/1-what-is-wavelength.html Electromagnetic spectrum Unit/symbol Length (m) Type of radiation Angstrom (Å) 10-10 X-ray nanometer (nm) 10-9 Ultraviolet; visible micrometer (µm) 10-6 Infrared Other: mm—infrared cm—microwave m—TV/radio http://www.antonine-education.co.uk/physics_gcse/Unit_1/Topic_5/topic_5_what_are_the_uses_and_ha.htm 12 9/17/14 http://micro.magnet.fsu.edu/primer/java/wavebasics/index.html Example problems • Yellow light from a sodium vapor lamp has a wavelength of 589 nm. What is the frequency of this light? – A: • An FM radio station broadcasts electromagnetic radiation at a frequency of 103.4 MHz. Calculate the wavelength of this radiation – A: Relationship between Energy and Frequency • E = hν • h = Planck’s Constant, ____________ J*s • Energy and frequency are ______________________ related. • Energy and wavelength are ______________________ related. 13 9/17/14 Example problems • Calculate the energy electromagnetic radiation that has a frequency of 2.44 x 1015 Hz. – A: • Calculate the frequency of electromagnetic radiation that has an energy of 2.99x10-19J. – A: • Calculate the energy of electromagnetic radiation that has a wavelength of 750 nm. – A: So light is a wave! • Light Diffracts • Light Refracts • Light has a wavelength and a frequency http://lukewest.edublogs.org/ The wave model of light can’t explain everything... I. Blackbody radiation—emission of light from hot objects) II. Photoelectric effect—electron emission from metals when excited by light III. Emission spectra of light for excited gas atoms • Let’s look at these more closely... 14 9/17/14 I. Blackbody radiation • If you heat a solid, it emits radiation – Stove burner; light bulb • When metals are heated they change color from red to yellow and from yellow to white. • As the wavelength decreases, the intensity increases. I. Blackbody radiation I. Blackbody radiation • In 1900, German physicist Max Planck determined that energy was only released/ absorbed in discrete chunks – He called the smallest quantity of energy a quantum • Energy equals a constant times the frequency – E=hν h = 6.626x10-34 J*s – Can only have E in multiples of hν—2hν, 3hν, 4hν, etc – Energy is _________________________—can only have certain values! 15 9/17/14 Quantized Energy • ___________________the smallest amount of energy that can be emitted or absorbed as electromagnetic radiation. Plural-______________. • An elevator • A ramp vs. stairs • So why does energy seem continuous to us? II. Photoelectric effect • 1905—Einstein realized that when light struck a metal, the metal emitted electrons • __________________ of the light was important – minimum needed to get emission – Increasing beyond this minimal affect • Einstein theorized that light was behaving more like an energy “packet” than a wave • Called this “packet” a _______________, which behaves like a tiny particle http://www.cem.msu.edu/~harrison/cem483/ Example problem • A laser emits light with a frequency of 4.69x1014 s-1. What is the energy of one photon? – A: • If the laser pulse has 5.0x1017 photons, what is the total energy? – A: • If the laser emits 1.3x10-2 J of energy during a pulse, how many photons are emitted during the pulse? – A: 16 9/17/14 Behavior of light • Light behaves like a wave (c=νλ) • Energy of light depends on frequency (E=hν) • Light behaves like a particle (photoelectric effect) • So is light a particle or a wave? – Both!!! Wave-particle duality III. Spectra • _______________ light—light with only one wavelength • ________________—light separated into different wavelength components – ____________________ spectrum— contains light of all wavelengths – ________________ spectrum—contains light of a few distinct wavelengths http://www.upscale.utoronto.ca/IYearLab/Intros/Spectra/Spectra.html Emission Spectrum • Atoms ______________ radiation when they are excited. • Light given off consists of discrete wavelengths called the line or emission spectrum. • What is special about an element’s line spectrum? 17 9/17/14 Atomic Absorption Spectrum • Atoms will also ________ discrete wavelengths of light. • The absorptions are observed as _________________ in a continuous spectrum. • How does an element’s emission spectrum relate to its absorption spectrum? Hydrogen spectrum • Mid-1800’s, only 4 distinct lines seen in the hydrogen gas spectrum http://faculty.sdmiramar.edu/fgarces/LabMatters/Instruments/AA/AAS_Theory/AtomicLineOrigins.htm http://faculty.colostate-pueblo.edu/linda.wilkes/111/3c.htm • 1885 Johann Balmer developed an equation for the wavelengths of these 4 lines Bohr’s model: three postulates • Only orbits of certain radii, corresponding to certain _____________, are permitted. http://www.goiit.com/posts/list/community-shelf-the-bohr-atom-917720.htm 18 9/17/14 Bohr’s model: three postulates • Electron in permitted orbit has specific E, and is “allowed”. It does not radiate E, and won’t spiral into the nucleus. http://www.goiit.com/posts/list/community-shelf-the-bohr-atom-917720.htm Bohr’s model: three postulates • E is emitted or absorbed by electron only if electron changes from one “allowed” orbit to another. E is emitted or absorbed as a photon. Lowest energy, n=1 is called the ________________ When electron is in a higher energy orbit (n=2 or higher), it is in an ___________________ http://www.goiit.com/posts/list/community-shelf-the-bohr-atom-917720.htm Principal quantum number • Bohr calculated energy for each orbit: E = (−hcRH )( 1 1 ) = (−2.18 x10 −18 J )( 2 ) n2 n • h = Planck’s constant, c = speed of light, RH= Rydberg constant • The integer n is called the _________________________________ What happens as n becomes infinitely large? State when electron is removed is called the _____________________ 19 9/17/14 Changing energy • Electrons jumping from n to n change E • Calculate the change in energy of an electron ΔE = E f − Ei = E photon = hν • Recall that c=νλ and the Rydberg equation..... 1 1 ΔE = (−2.18 x10 −18 J )( n 2f − ni2 ) This explains line spectra... If nf is smaller than ni, the electron moves __________ to the nucleus and E is released • Visible hydrogen line spectrum comes from transition from excited states to the _______ level http://faculty.colostate-pueblo.edu/linda.wilkes/111/3c.htm So why not Bohr 4-eva? UNFORTUNATELY...this model ONLY works for hydrogen... BUT...we get two important ideas • Electrons exist ONLY in discrete energy levels described by __________________ • ________________ is involved in moving electrons between levels • Where to next? 20 9/17/14 Wave behavior of matter • Louis de Broglie wrote a (very short) PhD thesis in 1925 • Proposed that if radiation (light) is a wave that can act like a particle, matter can act like a wave – Wave-particle duality – Equation for the de Broglie wavelength of matter: λ= h mv m = mass v = velocity Example problems • A major league pitcher can throw a baseball 90 miles per hour. If the average baseball weighs 145 g, what is the de Broglie wavelength of the baseball? – A: • What is the de Broglie wavelength of an electron moving at a speed of 5.97x106 m/s? The mass of an electron is 9.11x10-28 g. – A: New model for the atom! • Electron is a particle with some wave-like properties (de Broglie) • Schrodinger proposed a model that abandoned the “electron in an orbit” notion. • New model was based on the wavelike properties of the electron. 21 9/17/14 Heisenberg Uncertainty Principle (HUP) • German physicist Werner Heisenberg determined that the dual nature (wave & particle) limits how precisely the momentum and location of a particle can be known ΔxΔ(mv) ≥ h 4π Change in position times change in momentum (mass times velocity) Heisenberg Uncertainty Principle (HUP) • Typically only comes into affect for small masses (ie, the position or momentum of an electron) • We know me = 9.11x10-31 kg and the average speed is 5x106 m/s. If the uncertainty of the speed is 1%, what is the uncertainty in the position of the electron? – A: 1x10-9 m • But the diameter of one H atom is 1x10-10! Let’s try the calculation for a baseball (mass 145g, speed 90 mph) A:9.1x10-34 m Quantum mechanics • QM is a new way of describing the behavior of subatomic particles (mostly electrons) • 1926, Austrian physicist Erwin Schrödinger developed an equation (the wave function or Ψ) to describe both the wave and particle properties of the electron – Solving this equation is complicated and requires advanced calculus...so we’ll skip that step. We will, however, look at the results! • The square of the equation Ψ2 (electron density) tells us the probability of where the electron will be – Remember, we can’t know for certain because of HUP 22 9/17/14 Orbitals and Orbits • So Ψ gives us an idea of where the electron is allowed to be in space—called an _________________ • Note that QM orbital _____________ the same as a Bohr orbit http://www.goiit.com/posts/list/community-shelf-the-bohr-atom-917720.htm https://www2.bwdsb.on.ca/ ~f_schlenker/4U/4U %20quantum%20chemistry/ university%20website/ imgres_files/a.htm http://www.sparknotes.com/chemistry/ organic1/atomicstructure/section1.rhtml Quantum numbers • Describe an orbital & where electrons are • Bohr’s model gave us the primary quantum number, n • QM has ___________ quantum numbers, _______ ________ _______ ________ • Let’s take a closer look at these numbers... First quantum number, n • ____________________ quantum number: positive integer (1, 2, 3, ...) – Also defined as the ____________________ – As n increases, the size of the orbital increases and the electron spends more time away from the nucleus – Higher n means higher E – Corresponds to periods on PT 23 9/17/14 Sizes of orbitals • As n ______________, so does the size of the orbital • As orbital size increases, the probability of the electron being found near the nucleus decreases http://chemistry.umeche.maine.edu/~amar/fall2007/orbitals.html http://physchem.ox.ac.uk/~hill/tutorials/qm1_tutorial/atomorb/index.html Second quantum number, l • _____________________ quantum number : integral values from 0, 1, ..., n-1 – Represents the _______________ or __________________ of the electron shell – Defines the _____________ of the orbital – Also have ______________ designations Value of l 0 1 2 3 Letter used Azimuthal Quantum Number, l • l =0 s-sublevel - with spherical orbital • l =1 p-sublevel - with dumbbell orbitals • l =2 d-sublevel - with cloverleaf orbitals • l =3 f-sublevel - with eightlobed orbitals • Corresponds to the blocks on the periodic table. 24 9/17/14 Shapes of orbitals • Second quantum number, l, tells us the shape of the orbital • l = 0 (also called s) is a ________________ orbital – There is __________________ l = 0 orbital in each electron shell (for each n) http://www.sparknotes.com/chemistry/fundamentals/atomicstructure/section1.rhtml Shapes of orbitals • l = 1 (or p) orbitals are bow-tie shaped – Remember for l = 1 there are _____ orbitals – Same shape, oriented differently in space… Which quantum number is different with these orbitals? http://chemistry.umeche.maine.edu/~amar/fall2007/orbitals.html Shapes of orbitals • l = 2 (or d) orbitals are complicated – 4 are clover leaf shaped – Remember l = 2 there are ____ orbitals l = 3 (or f) are waaaaaaay too complicated to worry about the shape…just know that they’re there and how many there are! http://www.chem.ufl.edu/~itl/2045_s00/lectures/lec_10.html 25 9/17/14 f orbital shapes http://www.chem.tamu.edu/rgroup/hughbanks/courses/673/handouts/handouts.html Third quantum number, ml • _________________ quantum number ml : have integral values between -l and l, including zero – Represents the ____________ in the subshell – Describes the orientation in space – The total number of orbitals in all sublevels of a shell is _____ The Fourth Quantum Number, ms • The line spectra of many electron atoms show each line as a closely spaced pair of lines. • Stern and Gerlach experiment placed a beam of atoms through a slit into a magnetic field. • Two spots were found: one with electrons spinning in one direction and one with electrons spinning in the other direction. 26 9/17/14 ________ Quantum Number, ms • Electrons are free to spin clockwise or counterclockwise about their axis. • ms = ±1/2 • ms is independent of the other three quantum numbers. Pauli Exclusion Principle • PEP: No two electrons can have the exact same quantum numbers • Since electrons can only spin two ways, (CW or CCW), then two electrons will fit in one orbital and will have ______________ spins – Designate one as +1/2 and the other as -1/2 Quantum Model in a Nutshell • Principal quantum number, n, gives the electrons average distance from the nucleus and the energy level. • Azimuthal quantum number, l, describes the sublevels. One sublevel for each value. • Magnetic quantum number, ml , describes the orientation of the orbital in space. One orbital for each possible value. • Spin quantum number, ms, describes the spin of the electron in a magnetic field. 27 9/17/14 Interpreting Quantum Numbers 2, 1, -1, +1/2 Quantum Number Principle Symbol Describes Possible Values n Energy Level Whole #s ≥ 1 Azimuthal l Sublevel Whole #s 0 to n-1 Magnetic ml Orbital Spin ms Spin - l to +l ±1/2 Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning Center, 2008. Determining Quantum Numbers 5d5 Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning Center, 2008. Examples • Ex: Give the quantum numbers for the selected arrow. • Ex: Give the quantum numbers for the selected arrow. Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning Center, 2008. 28 9/17/14 Incorrect Quantum Numbers • Identify what is wrong with the following sets of quantum numbers. 1, 1, 0, 1/2 2, 1, -2, -1/2 3, 2, -1, 1 2, 1, 1, 1/2 Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning Center, 2008. What order do we fill the orbitals? Want to have ________________possible energy! • As n increases, E _____________ (E1s<E2s<E3s) • As l increases, E ______________ SLIGHTLY (E3s<E3p<E3d) – Orbitals within the same subshell l (example, px, py, and pz) have the same energy and are called ______________________ • Fill all degenerate orbitals in a subshell before moving on to the next subshell Order of filling • Electrons are arranged based on the energy of each orbital. • The order of filling from lowest to highest E is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d • Notice that the d-orbitals fill ____ energy level later than expected and that f-orbitals fill ____ energy levels later than expected 29 9/17/14 Energy Levels in Single and Multi-electron Atoms Electron Configurations • Tell us how the electrons are distributed among the various orbitals of an atom. » Li – 1s22s1 » Na – 1s22s22p63s1 » K – 1s22s22p63s23p64s1 » He - 1s2 » Ne - 1s22s22p6 » Ar - 1s22s22p63s23p6 • The most stable configuration or ground state is that in which the electrons are in the lowest possible energy state. Rules for Writing Electron Configurations • Electrons fill orbitals of increasing energy with no more than two electrons per orbital. (___________________) • For degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (__________). – Another way to think of this: in degenerate orbitals, electrons prefer to be unpaired. • No two electrons can fill each orbital with the same spin (________________). • How do we show spin? – An arrow pointing upward has ms = +1/2 (spin up). – An arrow pointing downward has ms = -1/2 (spin down). 30 9/17/14 Writing electron configurations • Determine the number of electrons. • Fill orbitals in order of increasing energy – s - ____e- maximum – p – ____e- maximum – d – ____e- maximum – f – ____e- maximum • d-orbitals fill 1 level later, f-orbitals fill 2 levels later • Continue filling until all electrons are used. Electron configuration and the periodic table Writing Configurations Using the Periodic Table • Trace the path you take to arrive at the element. • d-sublevels fill one period behind • f-sublevels fill two periods behind 31 9/17/14 Examples • O • Mg • Ga • Ru • V Configurations Within a Group • Li • Na • K • F • Cl • Br • Core electrons vs. valence electrons • Elements in the same group have the same outer valence shell configuration. Exceptions to Aufbau Ordering • Several elements do not follow the Aufbau order of filling. • _________ and __________ also follow copper pattern. • ________ also follows Cr pattern. 32 9/17/14 Exceptions Continued • Because of the stability of certain configurations. • Full sublevel > Half-filled > No particular organization • A half-filled d sublevel is more stable than a full s sublevel. Abbreviated or Noble Gas Configurations • Way to describe the core electrons and only show valence electrons • Cl • Write the symbol for the Noble gas that precedes the element. • Trace the path to the element • Sn More Examples of Abbreviated Configurations • Ca • S • Rb • Ne 33 9/17/14 Orbital Configurations • Orbital diagrams represent the detail lost by E.C. • Each orbital within a sublevel is represented by a box, line, or circle. • Electrons are represented by half-arrows. • Used primarily for valence electrons, since all inner core electrons will be paired. Hund’s Rule • Within a given sublevel, each orbital must contain one electron before they are allowed to pair up. • Orbitals in the same sublevel are said to be degenerate. Examples of orbital configurations • O • Fe • K • P • Ar 34 9/17/14 Orbital Configuration Questions • From the orbital diagrams select an example which demonstrates – a violation of Hund's rule – a violation of the Pauli exclusion principle – a ground state orbital diagram – an excited state orbital diagram – a violation of the Aufbau principle Lewis Dot Diagrams • Introduced by Gilbert N. Lewis in 1916 • Use to represent _________________ electrons and bonding between atoms • Used with the representative elements (s- and p-block) Whitley, Kelley. "The Periodic Table." ChemProfessor. 1 Oct. 2008. 30 Nov. 2008 <http://www.chemprofessor.com/aboutme.htm>. Dot Diagram Examples • Lithium – Config: – Dot Diagram: • Aluminum – Config: – Dot Diagram: • Arsenic – Config: – Dot Diagram: Whitley, Kelley. "The Periodic Table." ChemProfessor. 1 Oct. 2008. 30 Nov. 2008 <http://www.chemprofessor.com/aboutme.htm>. 35 9/17/14 Configuration Comparison Pro Con Standard Quicker than orbital Does not show pairing Abbreviated (Noble Gas) Shows only valence e- No core eQuick No pairing Orbital Shows pairing of electrons Longest Dot Shows only valence e- No core eVery quick Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning Center, 2008. Review Questions • Se: 1s22s22p63s23p64s23d104p4 1. Atom that has 2 unpaired electrons. • Mg: 1s22s22p63s2 • Al: 1s22s22p63s23p1 • N: 1s22s22p3 • F: 1s22s22p5 2. Atom with exactly 2 electrons in the highest sublevel. 3. The atom that must obtain 3 electrons to fill a sublevel. Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning Center, 2008. Review Questions A. 1s22s22p63s23p64s23d104p4 B. 1s22s22p63s23p2 C. 1s22s22p63s23p24s1 D. 1s22s22p63s23p64s23d8 E. 1s22s22p63s23p64s23d104p6 F. 1s22s22p6 • Which atom is in an excited state? • Which configuration represents the element Ni? • Which configuration(s) represents a noble gas? • Which configuration represents a sodium ion? Huang, Wayne, PhD, Kelly Deters, MA, and Debbie Bilyen, MA. Teach Yourself AP Chemistry in 24 Hours. CD-ROM. Premium ed. Yorba Linda: Rapid Learning Center, 2008. 36 9/17/14 The End For now… 37