1 Electron Dot Diagrams G.N. Lewis idea (UC Berkeley) –Elegantly simple idea, but very instructive –Show each bonding electron as a dot As elements brought together, dots merge Most stable configuration is filled shell –2 dots for Hydrogen (2s 2 or [He] configuration) –8 dots for most others (s 2 +p 6, Octet rule) Methane example C(4dot) + 4*H (1dot) –Can have more electron pairs than bonds “lone pairs” are non-bonding electrons Lone pairs occupy a geometrical position –Are part of molecular shape consideration

2 Rules of Lewis-Dot diagrams Only “valence electrons” considered –Inner core (noble gas configs.) ignored They play no part in chemical reactions –Octet rule applies for >90% of atoms Hydrogen is main exception with 2 electrons A few other exceptions –Starting point is an element’s electrons Pure element where electrons = protons Not elements in oxidized or reduced state “AE” or “Available Electrons”

3 Rules of Lewis-Dot diagrams Electron Counting –Add up electrons in all elements involved If product is an ion, must add or subtract electrons Sulfur (6) + 2 oxygens (6) = 3*6=18 electrons All require 8 shared electrons 3*8=24 Sharing is 24-18=6 –Each electron is a dot, each pair CAN be a line Dot plot must meet all the rules (one of these 2 does NOT)

4 Octet Rule 2 electrons per outer shell for Hydrogen –Hydrogen starts out with single electron - 1s –Diatomic hydrogen provides 2 shared electrons 8 electron outer shell for most elements –Some gain electrons, such as Cl(7e) to Cl - (8e) –Others lose electrons, as Na(11e) to Na + (10e) Polyatomic Ions share 8 electrons/element –Nitrate, sulfate, phosphate, etc.

5 Diatomic Hydrogen Formation of H 2 via sharing 2 outer electrons, emulating He

6 Lewis Structure (electron dot diagram) for ammonia Each of the 3 hydrogen atoms will share its electron with nitrogen to form a bonding pair of electrons (covalent bond) so that each hydrogen atom has a share in 2 valence electrons (electronic configuration of helium) and the nitrogen has a share in 8 valence electrons (electron configuration of neon)

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8 Lone Pair in Ammonia NH 3 has 8 electrons around the N = ok! NH 3 has 2 electrons around each H = ok! –Fits the octet and 2 electron rules What happens to 2 unbonded N electrons ? –3 of 5 AE are bonded, 2 are “leftovers” –These 2 are “lone pairs” Fulfills the electrical requirements Non-bonding, help define shape due to repulsion

9 Lone Pairs Having more electrons than bonds = leftover “lone pair” –Ammonia example Nitrogen (5 electrons to share) Hydrogen (total of 3 electrons to share) After N-H Bonding, 2 nitrogen electrons “left over” –Lone pairs They complete octets of 8 electrons per element They are part of electronic structure They are NOT part of molecular structure

10 3 of Nitrogen’s 5 valence electrons shared with 3 Hydrogen atoms in Ammonia. “Lone Pair” electrons attract Hydrogen ion Result is formation of the Ammonium ion

11 Each oxygen will share 2 of its valence electrons in order to form 2 bonding pairs of electrons (a double covalent bond) so that each oxygen will have a share in 8 valence electrons (electronic configuration of neon). Lewis Structure (electron dot diagram) for the oxygen molecule

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13 Examples across the chart

14 Electronic Configurations of Elements

15 Nitrate Formation from elements 1 additional electron needed to fill shells, meeting octet rule provides (-1) charge to anion

16 Alternative Nitrate Representations “Ball & Stick” easy to model and understand “Bubbles” represent electron clouds, size

17 Sulfate Formation from elements 2 additional electrons needed to fill shells, meeting octet rule provides (-2) charge to anion

18 Electron Dot Diagrams Lines between atoms are 2-electrons –One line equivalent to 2 dots 2 lines (double bond) equivalent to 4 dots 3 lines (triple bond) equivalent to 6 dots –Can rotate around one line (no interference) 2 lines (double bond) restricts rotation, planar 3 lines (triple bond), no rotation, linear

19 (Almost) all bonding can be represented by lines and dots

20 Each of 4 carbon valence electrons shares orbitals with 1 from Chlorine

21 Important Exception Ionic bonds  electron transfer –Each ion is free agent, unattached to the other If not shared, no “double counting” possible –Na + Cl  NaCl(s)  Na + + Cl - –Ions are “Isoelectronic” with [Ne] and [Ar] Core looks like noble gas, but net charge on ions –A similar question in the experiment Metal halides and oxides generally not covalent

22 Lewis Diagram for Chloride Ion

23 Lewis diagram for Sodium Ion