Gen Chemistry Ch 5 - lecture notes on chapter 5 PDF

Preview

Summary

lecture notes on chapter 5...

Description

Chapter 5: Chemical Bonding I The Continuum of Bonding: - Distinction between covalent and ionic bonding is not always well defined → represent two ends of the spectrum (sharing vs transfer) Electronegativity: - Ability of an atom to attract bonding electrons to itself (don't confuse with electron affinity) - Periodic trend: - Increases left to right across a period - Decreases going down a group - F is most electronegative, then O, then N, and then Cl - H is located between B and C (higher than B but lower than C) - Noble gases are not assigned electronegativities as they don’t readily bond to other atoms (exceptions: Xe, Kr, and Ar) - Main group metals tend to have low electronegativities (left column) - Impacts several behaviors/properties about a molecule - Molecular polarity - Formal charges - Oxidation numbers - Ability to form H bonds - Acid strength Covalent Bonds: Polar vs Nonpolar Bonds: - Electrons shared evenly or unevenly → dictates polarity - Evenness of sharing depends on electronegativity difference (Δ EN) between the atoms participating in the bond - Greater ΔEN = more polar bond - Think abt electronegativity trend - Bonding spectrum can be arbitrary divided into specific regions based on ΔEN

-

Representing Dipole Moments/Polar Bonds: - A dipole moment caused by a charge separation - Magnitude of charges are partial charges - Partial negative and Partial positive VS an arrow - Arrow goes from partial positive atom to partial negative - Electrostatic potential map - Sometimes it's beneficial to picture electrons as cloud of charge smeared around the molecule - More red regions = higher region of electron density Ranking Bond Polarity: rank following bonds in order of increasing polarity and indicate dipole moment of each bond with an arrow: C-F, C-O, C-H Calculating Dipole Moments: - Dipole moment, u, is a measure of bond polarity - Occurs any time there is a charge separation of a (+) and a (-) by a distance - u (dipole moment) = q x r - q = magnitude of charge (in C) - r = distance (in m) between charges - Directly proportional both to the magnitude of the partial charge and to the distance (r) between them - commonly measured in debyes (D) where 1 D = 3.34 x 10^-30 C *m - Large dipole moment or great electronegativity difference = polar Percent ionic Character - Indicates the extent of electron transfer between two atoms within a bond - It is the ratio of a bond’s actual dipole moment to the dipole moment the bond would have if electron was completely transferred - % Ionic Character = (experimentally measured dipole moment of bond)/(dipole moment if electron were completely transferred) * 100 - % ionic character increases as electronegativity differences increase - Greater % ionic character = more polar bond - Bonds with 50% or above ionic character are referred to as ionic bonds Drawing Lewis Structures: - A lewis structure is a simple way of representing the structure of a molecule or polyatomic ion - Bonding electrons shared by two atoms represented by dashes - Unshared, nonbonding electrons shown as dots - Usually central atom is first element listed (or the one with lower electronegativity) - Put brackets and charge outside it

Ranking Bond Polarity - Rank the following bonds in order of increasing polarity and indicate the dipole moment of each bond (if present) with an arrow

Common Oxyanions that may have H+(s) attached to Oxygen(s) - Nitrate ion NO3- Nitrite ion NO2- Carbonate ion C032- Oxalate ion C2O42- Sulfate ion SO42- Sulfite ion SO32-

-

Phosphate ion PO430 Phosphite ion PO33Perchlorate ion ClO4Chlorate ion ClO3Chlorite ion ClO20 Hypochlorite ion ClOAcetate ion CH3COOFormate ion HCOO*highlighted hydrogen in CH3COO- and HCOO- are attached to carbon though any subsequent hydrogens attach to oxygen

Resonance - The structures of some molecules and ions cannot be adequately described by a single valid lewis structure - More than one valid lewis structure differing in only in the position of electrons/multiple-bonds - Resonance structures must have the same order of atom connectivity (i.e., the same skeletal structure - Cyanate ion (NCO-) vs. fulminate ion (CNO-) - The true structure of the species is a mixture/hybrid of the different resonance structures - The actual structure of the species DOES NOT FLIP-FLOP between the different resonance structures - The “better” the resonance structure, the more it contributes to the overall structure of the species Resonance Hybrid - an analogy - Horse + donkey = mule - The mule (the hybrid) does not flip-flop between being a horse and being a donkey Three Resonance Structure for the Nitrate Ion

Formal Charge - Charge assigned to atoms for the purposes of distinguishing and ranking competing lew structures - Fictitious charges used to designate electron ownership of each atom - An atom only “owns” all its nonbonding electrons - An atom only “owns” half of each pair of bonding electrons Some rules regarding formal charges 1. The sum of all formal charges must equal the overall charge of the species 2. Minimal formal charge (+1, -1) are always better than excessive formal charges (±2, ±3…

etc) 3. When unavoidable, negative formal charges should reside on most electronegative atoms (F, O, N..) Exceptions to the Octet Rule - Odd-Numbered Valances - Aka radicals or free radicals - Highly reactive and unstable species - They occur when there is an odd-number of valence electrons due to the presence of an odd number of group 5A and/or 7A elements - Usually place the single electron on the atom that minimizes formal charges - Example : NO, ClO Exceptions to the Octet Rule - Incomplete Valances - Some atoms are satisfied having less than eight electrons in their outer valence

shell - beryllium (Be): for electrons in outer shell - Boron (B): six electrons in outer shell

Exceptions to the Octet Rule - Expanded Octets - Period 3 atoms(P, S, Cl…) and beyond (As, Se, Br, Kr, I, Xe,,,) can expand their octets and accommodate more than eight electrons in their outer shell - Due to easy access to relatively low energy d orbitals - If they can, atoms will expand their octets to reduce formal charges - PERIOD 2 ATOMS (C, N, O, F) NEVER EXPAND THEIR OCTETS - The cardinal sin of Lewis structures! - No d-orbitals for Period 2 Criteria for the “Best” Lewis Structure - A summary 1. The structure must have the proper number of electrons as determined by the total valences of the atoms 2. All atoms that can have complete otets (C, O, N, F) DO have complete octets or expand their octets (in the case of Period 3 atoms and beyond in periodic table) a. Never expand the octet of Period 2 elements (C, O, N, F) b. Remember that Be and B can have incomplete octets 3. Formal charges are minimized a.

4. If formal charges are absolutely necessary, negative formal charges should be on the more/most electronegative atom(s) Average Bond Energy: - Bond energy: amount of energy required to break 1 mole of covalent bonds in the gas phase into separated atoms in the gas phase

-

- Memorize: X-Y (g) → X (g) + Y (g) - Always endothermic (energy is absorbed) - The reverse process (making bonds) is exothermic (energy released) If you know the energy change of one process, you also know the energy change of the reverse process just by changing the sign The stronger the bond, the more stable (less reactive) the bond

-

- N2 (g) (946 kJ/mol) vs O2 (g) (498kJ/mol) Average bond energies: average bc bond energies may change slightly depending on what environment it is in In general, for a given pair of atoms, the bond energies increase with multiple bonds - Triple > double > single

Average Bond Length: trends - Bond length: distance between the nuclei of the two atoms taking part in the covalent bond - Usually in pm (1pm=10^-12m) - An average of the bond length from several different compounds containing a particular bond) - Closely *thought not perfectly* related to atomic radius - F2 vs Cl2 vs Br2 vs I2 (all group 7), bond lengths gets larger - HF vs HCl vs HBr vs HI (all halogens), bond lengths get larger down the group - In general, when comparing the same two atoms, bond length decreases as the number of bonds increase - C-C vs C=C vs C=-C, triple bond will be the shortest but strongest while the single bond is the longest but the weakest *Example: Rank the following species in order of increasing N-O bond length: NO2-, NO+, NO2+ (hint: use lewis structures) Shapes of Molecules: - Shape is responsible for many properties of a molecule - Impacts polarity which impacts intermolecular forces which impacts physical properties - Shape of drug molecules and enzymes are responsible for their bioactivity - Shapes of large, complex molecules can be understood by focusing on the shape around each central atom - Valence shell electron pair repulsion (VSEPR) theory - Electron groups around an atom want to get as far away from one another to minimize repulsion - A single bond is one electron group - A double bond is one electron group - A triple bond is one electron group - A lone pair is one electron group - A single electron (in a radical species, unpaired electron) is an

-

electron group Maximum separation of electron groups around an interior/central atom leads to the shape of a molecule

Electron Group geometry vs molecular geometry - Electron group geometry (e.g.g.): spatial arrangement taken by the electron groups around a central atom - Molecular geometry (m.g.): spatial arrangement of only the bonds around a central atom - Sometimes the e.g.g. And the m.g. Can be the same (CO2 and CH4) - Sometimes they can be different (H2O and SF4) - Electron groups repel one another and try to get as far away from one another to minimize repulsions - Ideal geometry of electron groups depends on the total number of them around the central atom

-

-

-

2 electron groups: - 2 bonds, 0 lone pairs: 180, linear, linear 3 electron groups: - 3 bonds, 0 lone pairs: 120, trigonal planar, trigonal planar - 2 bonds, 1 lone pair: 119, trigonal planar, bent 4 electron groups: - 4 bonds, 0 lone pairs: 109.5, tetrahedral, tetrahedral - 3 bonds, 1 lone pair: 107.8, tetrahedral, trigonal pyramidal - 2 bonds, 2 lone pairs: 104.5, tetrahedral, bent 5 electron groups: - 5 bonds, 0 lone pairs: 90, trigonal bipyramidal, trigonal bipyramidal - Axial and equatorial: 180, 120, and 90 - 4 bonds, 1 lone pair: 87.8, trigonal bipyramidal, seesaw - 3 bonds, 2 lone pairs: 87.5, trigonal bipyramidal, t shape - 2 bonds, 3 lone pairs 180, trigonal bipyramidal, linear

-

- Replace all equatorial positions with only axial 6 electron groups: - 6 bonds, 0 lone pairs: 90, octahedral, octahedral - 5 bonds, 1 lone pair: 84.8, octahedral, square pyramidal - 4 bonds, 2 lone pairs: 90, octahedral, square planar

Deviation from ideal bond angles: - Presence of lone pairs unsymmetrically around the central atom will cause the bond geometries to deviate from their ideal values -

The H-O-H bond angle of water is 104.5 (5 degrees less than 109.5 which is what its predicted to be) → what is

-

causing this compression of bond angles? Lone pairs repel bonding electrons, causing compression of bond angles - lp/lp repulsion > lp/bp repulsion > bp/bp repulsion In general, more lone pairs around the central atom = more angle compression - CH4 vs NH3 vs H2O If lone pairs are symmetrically distributed, then bond angles won’t deviate - XeF2 (180 degrees) and XeF4 (90 degrees)

Shapes of larger molecules: - Molecules with two or more interior/central atoms - Overall shape of the molecule will be determined by the shape around each

interior atom -

Ex: acetic acid (CH3COOH) → tetrahedral + trigonal planar + bent (lone pair not shown)

ICL-1 → 2 lone pairs and 4 single bonds → 6 electron groups → octahedral → square planar → bond angles (symmetric since one is 180 degrees): 90 and 180 Molecular Polarity: - Already learned that bonds can be polar due to uneven sharing of bonding electrons between two atoms - Bond acts similar to a little magnet and will respond to an external electric shield - Molecules also may exhibit polarity (i.e. have a net dipole moment) - Polar molecules can interact with each other much like different poles of magnets interact - Polar molecules will easily mix with other polar molecules to make homogenous solutions - Like dissolves like (polar want to interact with polar and create a homogenous mixture vs nonpolar interacts with nonpolar) Requirements for Molecular Polarity: - In molecules where there are two or more bonds, the entire molecule can be polar and may exhibit a net dipole moment provided some requirements are met

-

-

- Examples: Water (polar molecule) response to an external electric field: acts like a magnet - The water orients itself so that the partial negative charge of the molecule wants to interact with the partial positive charge of the plate 1) contains polar bonds (ignoring exceptions to these) 2) the individual bond dipoles around the central atom do not cancel each out

Geometries that commonly lead to molecular polarity: - This is assuming that all bonds are identical

-

-

-

ex) molecular geometry : bent, trigonal pyramidal, seesaw, t-shaped, square pyramidal It gets more complicated when there is more than one type of bond - ex) FCN (two bond dipoles going in opposite directions but not equal in strengths; net movement is still in one direction) Bond dipoles are vector quantities - They have both magnitude and direction

Examples: - Determine if the following molecules are polar or nonpolar: NF3 and XeF4

-

https://docs.google.com/document/d/1DvuuxvdbPLgCk8ZPLY7rgcpF9rmm5WQxZY0d-i1bpo/edit

file:///Users/minjookang/Downloads/2014%20Purple%20Book%20Answers%20Week %201%20(1).pdf https://courses.lumenlearning.com/introchem/chapter/sp3-hybridization/...