Genchem Chapter 10 PDF

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Summary

Chemical Bonding...

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Chapter 10 : Chemical Bonding I Chemical Bonding - Chemical bond - Holds atoms together - Lower energy (stable) -

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Electrons on each atom - Feel (+) of other nucleus - New orbitals (allowed states) - Electron near both nuclei -

lower s total energy - Bond energy

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Most stable configuration - Full octet (octet rule) - 8 electrons, full s and p orbitals of valence level

How

Covalent Bonds -

Atoms share electrons - Octet (valence) Covalent bond - Bond (total energy decreases) with sharing - Orbitals (wave functions)

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NB: covalent bond

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- Shared electrons “count” for both atoms - Electrons: pairs Two types - Equal sharing (non polar)

- Unequal sharing (polar) Bond Polarity in Chemical Bonds

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 Nonpolar covalent bond - Electrons shared equally Polar covalent bond - Electrons shared unequally

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Ionic bond - Electrons transferred Electronegativity - Attraction for shared electrons -

Property of each element

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Number (EN): polar character of atom’s bonds

Electronegativity - Polar vs nonpolar - the difference in EN of 2 atoms - Type of bond

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Normally need only TRENDS in EN to solve problems

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Metal - nonmetal (Na, Cl) - Ionic - Unlike nonmetals (C, O) - Polar covalent - Identical nonmetals (N, N in N2) - Nonpolar covalent Dipoles in polar covalent bonds -

Dipole - Two equal, opposite charges separated by a distance Dipole moment (magnitude of dipole) - Two equal (opposite) charges Q separated R has dipole - μ = Qr

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Dipole moment - Amount of (+/-) separation

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Dipole moment arrow

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- Direction of dipole moment - Points to negative end of bond Dipole moment - Vector quantity, both magnitude and direction

Lewis Structures: drawings that show bonding in molecules - Bonding - Covalent molecules - Atoms share electrons : octet (duet for H) - Shared electrons count: both atoms

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 Lewis structure - Diagram (electron dot notation) - Valence electron arrangement - How many bonds - Number of valence electrons needed for octet Valence (outer shell) electrons - Electrons in highest occupied energy level - Generally s and p electrons only

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Filled s and p valence shell by sharing electrons Main group number - Valence electrons

 To draw a proper lewis structure - 1) add up valence electrons of all atoms, arrange in pairs -

2) draw a skeletal structure by using pairs of electrons to make bonds - Guidelines for connectivity of atoms in molecules - Minority atom in center, majority atoms on sides - LEAST electronegative atom in center - H - atoms always on sides, never center - Hydrocarbons (C and H) - C atoms center, H atoms outside

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- 3) complete octets (duets for H) for all atoms, outer atoms first, using the remaining valence electrons as lone pairs - 4) if octets are not produced, have atoms with octets share more electron pairs with atoms without octets Lewis Structures for Polyatomic Ions -

Must include ion’s charge in total count of electrons -

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Resonance structures

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 Resonance - Multiple equivalent structures ( resonance structures) Real structure - Average of all resonance structures Real O3 molecule - A single structure, average of its resonance structures (NOT two separate structures)

Formal charge in covalent bonds -

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Formal charge: “net charge” atom (gained/lost valence electrons) Difference: (starting electrons) - (final electrons) - Each bonding electron - Half its charge to each atom (shared equally) Starting electrons = valence electrons Lewis structure - Electrons assigned to atom - All lone pair electrons (2 per pair) - One electron for each covalent bond (2 x ½) - If starting does not equal final - Atom has acquired a net charge - Formal charge

Lewis structures and formal charges - Dominant lewis structure - Formal charges closest to zero - Negative formal charge on most electronegative atom Ionic Bonds - Metal and nonmetal

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- metals : lose electrons (+ ions) - Nonmetals : gain electrons (- ions) Ionic bond - Attraction between (+) and (-) The octet rule : ionic compounds -

Atoms >>> ions : octet of valence electrons

Ionic Compounds properties

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Lattice energy

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 ΔHLat : binding energy of compound => Properties (melting) ΔHLat depends on: - 1) Ion’s sizes - smaller => closer together => lower energy (greater ΔHLat)

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ΔHLat depends on:

 - Ions’ charge: greater charge => greater attraction => greater ΔHLat Resonance structures

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 Resonance - Multiple equivalent structures Real structure - Average of all resonance structures Real 03 molecule - A single structure, average of its resonance structures (NOT two separate structures)

 Exceptions to the octet rule - Three cases - Odd number electrons - Important molecules with odd number of electrons - Very reactive

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Less than octet -

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2nd row before carbon (b, be) - Compounds with < 8 electrons on central atom Bf3: - Boron with fill octet > negative formal on B, positive on F

 >> 8 valence electrons (expanded octet) - Row 3 or higher: - > 8 electrons in valence shell - Can use d-orbitals to make more than four bonds - Important cases - 1) row 3+ central atom in oxides/anions reduce their formal charge - Form more bonds with their O atoms - Examples: PO43- , SO4 2- , ClO4- - 2) Heavy atom (row 3+) halides (especially Fluorides) - NB: -

Covalent bond Strength -

Energy to break a bond - Bond enthalpy For Cl—Cl, D(Cl—Cl), = 242 kJ/mole: Cl—Cl → 2 Cl• ΔH = +242 kJ/mole Positive, energy input (endothermic).

Only period 3 and higher Never period 2 (C, N, O, F)

Lewis Structure of Carbon Monoxide - Valence electrons - 4+6=10 > 5 pairs, 4 pairs - Formal charges

Using Bond enthalpies to estimate enthalpy of reaction -

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Estimate DHrxn: D(bonds broken) and D(bonds formed): - => DHrxn = S(D(broken)) − S(D(formed)). - Example: Calculate ΔHrxn for this reaction using bond enthalpies

Condensed Structural formula notation

Bond Enthalpy and bond length - Measured average bond length for different bond types -

Number of bonds (bond order) increases > bond length decreases

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The higher the bond order (triple > double > single) - The shorter the bond length - The stronger the bond

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