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Chapter 10

Lecture and Textbook Collaborated Notes - Chapter 10 - CHEM 1A03

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Aadil Merali Juma

Chem 1A03 Chapter 10: Chemical Bonding I: Basic Concepts 10.1 Lewis Theory: An Overview  Lewis Theory: 1. Electrons, especially those of the outermost electronic shell, play a fundamental role in chemical bonding 2. Ionic Bonds – Electrons are transferred from one atom to another a. Positive ions and negative ions are formed and attract each other through electrostatic forms b. Usually between metal and non-metal 3. Covalent Bonds – One or more pairs of electrons are shared between atoms a. A bond formed and attract each other through electrostatic forces b. Often between two non-metals 4. Electrons are transferred or shared in such a way that each atom acquires an especially stable electronic configuration – usually a noble gas configuration, one with eight outer electrons, or an octet  Lewis Symbols and Lewis Structures o Lewis Symbol – consist of a chemical symbol to represent the nucleus and core (inner-shell) electrons of an atom, together with dots placed around the symbol to represent the valence (outer-shell) electrons o Lewis Structure – combination of Lewis symbols that represent either the transfer or the sharing of electrons in a chemical bond, designate electrons involved in bond formation from one atom as an X, and a  from the other 10.2 Covalent Bonding: An Introduction  Coordinate Covalent Bond – an atom contributes both of the electrons to a shared pair o Eg/ NH + H  NH 4+ 10.3 Polar Covalent Bonds and Electrostatic Potential Maps  Polar Covalent Bond – covalent bond in which electrons are not shared equally between two atoms o Electrons are displaced toward the more nonmetallic element – leads to a partial negative charge on the more nonmetallic element (δ-) and a partial positive charge on the more metallic element (δ+)  Electrostatic Potential Map – a way to visualize the charge distribution within a molecule o Electrostatic Potential – the work done in moving a unit of positive charge at a constant speed from one region of a molecule to another o Obtained by hypothetically probing an electron density surface with a positive point charge  Positive point charge will be attracted to an electron-rich region – electrostatic potential will be negative  If the point charge is placed in an electron-poor region, the point charge will be repelled – electrostatic potential will be positive o Red, the low energy end of the spectrum – most negative electrostatic potential Ionic Bond – typically one atom holds more charge than the other Polar Covalent Bond Nonpolar Covalent Bond – all diatomic molecules Chem 1A03 Electronegativity  Electronegativity – describes an atom ability to compete for electrons with other atoms to which it is bonded  Increases up a group and across a period  EN A (I A EA A  The lower the EN, the more metallic it is  The higher the EN, the more nonmetallic it is  As the ionization energy increases across the period we expect the electronegativity to increase  The distinction between electron affinity and electronegativity -1 o EA Cl349 kJmol ) is somewhat more negative than EA F-328 kJmol ) but the EN iCl3.0 which is significantly lower than EAF, which is 4.0  because the decreased IE Cl251 kJ/mol) relative toEA F1681 kJ/mol)  ΔEN – electronegativity difference, the absolute value of the difference in EN values of the bonded atoms ΔEN Bonding Example Large Ionic NaCl Intermediate Polar Covalent PCl5 Small Pure Covalent Cl2  Percent Ionic Character o Decreases across a period 10.4 Writing Lewis Structures  Lewis Structures o Show boding and non-bonding electrons and formal charges  Octet can be achieved by combination of bonding and nonbonding electrons o Not all atoms have an octet  Bonding electrons can be single, double or triple bonds  Steps 1. Count total # of electrons (include charge) a. Add electrons for negative charge b. Subtract electrons for positive charge 2. Draw skeletal structure (central and terminal atoms) a. The least electronegative atom is generally the central atom b. H and F are always terminal 3. Deduct two electrons for each single bond of skeleton 4. Use remaining electrons to complete octet of terminal atoms a. Only 2 electrons for H 5. Subtract electrons used for terminal octets 6. If electrons remain, place on central atom 7. Do all atoms have octet? a. If not: Use lone pairs to form multiple bonds b. Note: Often Group 2 and 3 elements have < 8 electrons 8. Calculate Formal Charges on each atom a. Formal Charge – apparent charges on certain atoms in a Lewis Structure that arise when atoms have not contributed equal numbers of electrons to the covalent bonds joining them b. FC = (Group # - # non-binding electrons - # of bonds) Chem 1A03 c. All formal charges must add up to total charge 9. Minimize formal charges to zero where possible, by creating multiple bonds  Octet is not exceed in Period 2 o C, N, O and F “Can Not Over Fill” o Cannot exceed the octet rule even if there are adjacent formal charges that cold be minimized  For elements in Periods 3, 4, etc. – minimizing formal charges is the priority, even if it means breaking the octet rule  Usually don’t have adjacent atoms with same formal charge  Negative formal charges usually appear on the most electronegative atoms, positive charge son the most electropositive atoms Porphryn Ring  O2molecules bind to iron of the heme unit  Note how the geometry around iron changes as 2 s bound vs. free  Carboxyhemoglobin: The Silent Killer o CO binds competitively to the heme unit and displaces O 2 o CO poisoning can result o < 50 ppm: Upper safety limit o 50 – 200 ppm: Headache/nausea o > 200 ppm: Dizziness /convulsions o Treatment: Saturation with 2 can reverse the process – the binding is an equilibrium process  Earth CO Atmospheric Conditions Chem 1A03 10.5 Resonance  Resonance – the situation in which two or more plausible Lewis Structures contribute to the “correct” structure  Same skeletal structure  Differ only in how electrons are distributed within the structure  Eg/ Average Formal Charge on O = 0 + (-1) + (-1) + (-= -3/4 4 Average P-O Order = 1 + 1 + 1 + 2 = 5/4 = 1.25 4  Double headed arrow between different structures  Average Formal Charge for an atom = total charges on atom total # of that atom  Average Bond Order = total number of 1 type of bond # of places where the bond is found 10.6 Exceptions to the Octet Rule Odd-Electron Species  Lewis theory deals with electron pairs and does not tell us where to put the unpaired electron 
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