The Periodic Table with Electron Configuration Unit 3

Unit 3 – The Periodic Table, Electron Configuration, &Periodic Trends Mendeleev’s Table Dmitry Mendeleev He created the first periodic table based on the properties of the elements ChemCatalyst: This document was created in 1889, when chemists only knew of 63 different elements. How do you think the elements are organized? What do you think the numbers represent? Mendeleyev's Table of the Elements- 1889 Group I H= 1 2 3 4 5 6 7 8 9 10 11 12 Li= Group II Group IV Group V Group VI Group VII _ Group VIII 0 = 16 F= l Be= 9 7 Na= QS ]\·Ib..,. = o-,1- • Ca= -H> K = :;9 Cu= 63 Hb = 8."> Group III Zn = 65 Sr= 87 B= 1 I Al== (' = 27 IQ N = lt Si= 28 - = 44 Ti= 48 P = 31 V = 51 = 68 -·-- = 72. /\� = 7J Yt = 88 Zr = 90 Nb= 94 s = jQ Cr = 52 Se - - - Er= 178 - 78 :Mo= 96 - - Cl= 35 Mn= 55 Ag= 108 Cd = 112 :u = l t {j Sn= 11 s Sb= 1�2 Te= 125 Cs= 133 Ba= 137 Di= 138 Ce = 140 - 19 - La= 180 Ta= 182 W= 184 Br = so Fe= 56 Co= 59 Ni= 59 Cu= 63 _= 100 Ru=104 Rh =106 Pd =106 Ag=108 l= 127 - - - 0s = 195 Ir =197 Pt =198 Au= 199 Au= 199 Hg= 200 TI= 204 Pb= 207 Bi= 208 U= 240 Th= 231 Note: Mendeleyev's symbol for iodine, "J", has been changed to "I" to match modern symbols. -- Create a Periodic Table Activity Activity Compared to Real Table Transition metals missing! Electron Configuration Electrons are found in energy levels Fill the lowest possible energy levels 1st and move outward as the energy levels fill up Energy levels called “rings” in lower level classes n=1 is lowest energy level (closest to the nucleus) Each Energy Level is Different…  Each energy level has specific orbitals (3- D pathways) where electrons can be found Each energy level can hold a specific # of electrons 1st Energy Level- can hold 2 electrons - Puts them in “S” shaped orbital (always Eonly hold 2 electrons). E- 1s2 1st energy level Sphere shape # of electrons This is how you write out the location of the electrons 2nd Energy Level   Can hold up to 8 electrons Has an s-shaped orbital & a p-shaped orbital  The s-orbital always fills up 1st! ee- 2s2 ee- e- e- 2p6 ee- 3rd Energy Level   Can hold up to 18 Electrons Has s, p and d orbitals ee- 3s2  e- e- e- e- 3p6 e- ee- e-e- e- e-ee e e-e 3d10 NOTE: The 4s orbital is actually at a lower energy, so electrons will fill it before the 3d orbital! 4th Energy LevelHas 4 different types of orbitals (s,p,d & f)  f- orbitals are the most complicated  7 possible “f” orientations  14 electrons can fit in f orbitals (2 electrons x 7 orientations = 14) One little trick… Here’s the order that an atom will fill it’s electrons. Starts with the easiest, lowest energy level to put electrons and moves up. Diagram on pg 150 of your text book  Whoa… skips from 3p to 4s to 3d? Based on energy, it’s easier to fill the “s” orbital on 4th energy level then the complicated “d” on the 3rd. Look further up the fill chart and you’ll see more of this. Order that Orbitals Fill Up… don’t memorize!!! •You will learn to use your periodic table to figure this out! n=7 n=6 n=5 n=4 n=3 n=2 n=1 7p6 6d10 5f14 7s2 6p6 5d10 4f14 6s2 5p6 4d10 5s2 4p6 3d10 4s2 3p6 3s2 2p6 2s2 1s2 Writing Electron Configurations Orbital Notation & Electron Configuration Notation Ex: E. Config. for Fluorine (9 electrons) 1s2 2s2 2p6 Write the order that they fill the electrons More Practice (a harder one) Ex: Titanium (22 electrons) E. Configuration Notation 1s2 2s2 2p6 3s2 3p6 4s2 3d2 Remember the fill order, 4s before 3d! Using the Periodic Table for E. Config. S Block p Block d Block (n-1) (Energy level -1) Noble Gases (last column) f block (n-2) E. Configuration with Periodic Table Practice Writing E. Configs. Carbon  Magnesium  Iron  Iodine Challenging:  Gold Even more challenging:  Plutonium   1s22s22p2 1s22s22p63s2 1s22s22p63s23p64s23d6 1s22s22p63s23p64s23d104p65s24d105p5  1s22s22p63s23p64s23d104p65s24d105p66s24f145d9  1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p 67s26d15f5    Noble Gas Notation- the “short cut” For Na E. config: 1s2 2s2 2p6 3s1 Noble Gas Not.: [Ne] 3s1 For Cl E. config: 1s2 2s2 2p6 3s1 Noble Gas Not.: [Ne] 3s2 3p5 *Start at the noble gas ABOVE the element and do the configuration from there. Noble Gas Notations through the D and F Blocks Noble Gas Notation for Br [Ar] 4s2 3d10 4p5 Remember d block is n-1 (row 4 -1 =3) Noble Gas Notation for Pb [Xe] 6s2 4f14 5d10 6p2 Remember f block is n-2 (row 6 -2 =4) Using Noble Gas Notation  The noble gas notation can tell you the identity of an element Period # [He]2s22p2 Which element in the block block Element identity = Carbon Decoding Noble Gas Notation Period Block Group Identity [Ne]3s2 [Ar]4s23d8 [Xe]6s24f145d106p2 3 S 2 Mg 4 d 10 Ni 6 p 14 Pb Electrons History Behind Electron Configuration Certain elements emit distinct, visible light when heated in a flame. But why? Strontium Copper What is light?  It’s a form of energy The Electromagnetic Spectrum (see image) shows other types of energy in our environment. Red Light = lowest frequency Violet Light = highest frequency Visible light makes up only a small portion of the spectrum Spectroscopic Analysis  When white light is scattered through a prism (or spectroscope), all of the colors of the visual spectrum can be seen.  This is seen as a “continuous spectrum” (without breaks) Electron Transitions  Electrons can “jump” to higher energy levels when atoms are exposed to an energy source   This is known as the “excited state” When the electrons fall back down, they release that energy in the form of light  This is known as the “ground state” Energy Light Energy Transitions for Hydrogen Shows how the 1 electron in hydrogen can go up and down in different energy levels. Emission Spectroscopy  Because atoms have different numbers of electrons, different types of atoms emit specific wavelengths and have a different pattern of spectral lines  This is the “line-emission spectrum” Spectroscopy Elements have a unique set of spectral lines that allows us to identify them This is how we know the sun contains H and He, even though we’ve never been there. Argon Hydrogen Valence Electrons, Octet Rule, and Ions These 3 atoms have similar reactivity and chemical behavior. A) where are they located on the periodic table? B. What do you think might be responsible for their similar properties? Valence Electrons   Valence Electrons electrons in the outermost energy level These are the electrons that interact with other atoms  They determine an atom’s chemical reactivity Valence Electrons & E. Congfig.  1s2 2s22p6  In the electron configuration, the valence electrons are found in the s & p orbitals of the highest energy level. Examples:  3s23p7 Cl: [Ne]3s23p5  Chlorine  Fe: [Ar]4s23d6   Has 7 valence electrons Has 2 valence electons Sn: [Kr]5s24d105p2  Has 4 valence electrons Valence Electrons and the Periodic Table Figuring out # of Valence Electrons Using the Periodic Table (Short Cut) The column that they are in is the number of valence electrons an atom has. (EXCEPTION: This does not work for the D Block) 2 Valence Electrons The Octet Rule   Atoms tend to gain, lose, or share electrons to “fill” their valence shell. Exceptions: H & He abide by the “duet” rule.  They only need 2 electrons in their valence shell because the 1st energy level only holds 2 electrons Ions   Ions are charged particles or atoms that have gained or lost electrons to “fill” their octet. Anions have a negative charge.    They have gained electrons & electrons are negative. They have more electrons than protons. Cations have a positive charge.   They have lost electrons. They have more protons than electrons Ion Examples   Potassium has 1 valence electron, what ion will it form? It will lose 1 electron and form the ion K+   Sulfur has 6 valence electrons, what ion will it form? It will gain 2 electrons and form S2- Joke joke  Two atoms walk into a bar. One atom stops and says to the other, "I think I just lost an electron."  The second atom asks "Are you sure?"  The first atom replies, "I'm positive!" Organization of the Periodic Table Characteristics of the Periodic Table    Elements are arranged in order of increasing atomic number Elements with similar properties appear at regular intervals (“periods” or rows) Elements with similar properties fall in the same column (“group” or “family”) Families/ groups Periods (rows) Organization of the Periodic Table    Metals – excellent conductors of heat & electricity  Alkali metals – Group 1  Alkaline-earth metals – Group 2  Transition metals – Groups 3-12 Metalloids – properties of metals & non-metals (along zigzag) Non-Metals- poor conductors of heat & electricity. Usually brittle solids or gases.  Halogens – Group 17  Noble gases – Group 18  Other solid non-metals – above metalloids Alkali Metals Group 1 of the Periodic Table  All have 1 valence electron  e- Highly reactive (with water)  Silvery in appearance  Soft enough to be cut with a knife  Alkaline-Earth Metals   Group 2 of the Periodic Table All have 2 valence electrons ee-   ee- Harder, denser, stronger the alkali metals Also reactive, but not as much as alkali metals Transition Metals Groups 3-12 of the periodic table  All have 2 valence electrons  ee- ee- Halogens   Group 17 of the Periodic Table All have 7 valence electrons ee-    ee- Despite chemical similarities, some are solids, liquids, and gases Most reactive non-metals React with metals to make salts Noble Gases   Group 18 of the Periodic Table All have 8 valence electrons, a complete octet ee-  e- e- Total lack of reactivity, inert  “too noble to react with anyone else” PERIODIC TRENDS Periodic Trends    Characteristics of elements are predictable based on their location on the Periodic Table. These characteristics are dependent on the structure of the atom and the location of its electrons. Periodic Trends include: Reactivity  Atomic radius (size)  Ionization energy  Electronegativity  Atomic Radii   Size of the atom – measured using half the distance between the nuclei of two identical atoms bonded together Trend – decreases across a period, increases down a group •Largest atom = Francium •Smallest atoms = Helium Atomic Radii   As you move across a period, you are adding electrons to the same energy level, but also adding more protons  These protons attract the electron cloud closer, decreasing the atomic radius. As you move down a group, you add energy levels.  Each energy level is farther away from the nucleus, increasing the atomic radius. Ionization Energy (IE)  The energy required to remove one electron from a neutral atom Ex: Noble Gases require a ton of energy to lose an electron because they are “happy” with their full shells Ex: Alkali Metals have low ionization energies because they want to lose their outer electron. Trend – increases across a period, decreases down a group Ionization Energy Increase Decrease Electronegativity  The ability of an atom to attract electrons (how “greedy” it is) Ex: Flourine really wants another electron to get to the octet rule so it has a very high electronegativity. Anything close to Fluorine will have a high electronegativity Trend – increase across a period, decrease down a group Electronegativity Increase Decrease Heavy metal (joke)