Lewis Structure Reminders. Ch 2 Structure & Bonding. BTW: A Note on Formal Charges. Hints on Lewis Dot Structures
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1 Ch 2 Structure & Bonding An amazing thing about the universe - It works in a way that sometimes when things come together, they stick Protons and neutrons in a atomic nuclei Atoms in molecules spin-pairing pairing Lewis Structure Reminders Lots of practice! Work all the examples you can find! Optimal resonance structures: Never violate octet rule for n=2 Satisfy octet rule whenever possible Can violate octet rule for n > 2 Reduce formal charge if possible No negative charge on low c atoms ints on Lewis Dot Structures 1. Octet rule is the most useful guideline. 2. Carbon forms 4 bonds. 3. ydrogen typically forms one bond to other atoms. 4. When multiple bonds are forming, they are usually between C, N, O or S. 5. Nonmetals can form single, double, and triple bonds, but not quadruple bonds. 6. Always account for single bonds and lone pairs before forming multiple bonds. 7. Look for resonance structures. BTW: A Note on Formal Charges #Valence e - s - # Actual e - s 1 FC = G.N. - # BE + # LPE 2 If Step 4 leads to a positive formal charge on an inner atom beyond the second row, shift electrons to make double or triple bonds to minimize formal charge, even if this gives an inner atom with more than an octet of electrons. 1
2 All resonance structures are not created equal!!! Chemical Bonds: attractive force holding two or more atoms together TWO EXTREME CASES Covalent bonding: when atoms share electrons (e.g. organic carbon-based based molecules) Ionic bonding: when Atom A transfers an electron to Atom B (e.g. salts such as NaCl) Polar-covalent bonding: everything in between these two extremes (a good portion of reality KEY POINT: Bonding often happens so all bonding partners acquire a noble-gas electron configuration Covalent Bonding: Considerations Each electron in bonding pair has greater space available than in the unbonded individual atoms, and each gets to feel the positive charge of both nuclei. Covalent Bonding: Considerations Electrons & nuclei balance balance all interactions to give the molecule stability. Balance achieved when electrons are concentrated between the nuclei (electron sharing) A covalent bond forms IF: the energy *released* from sharing electrons is greater than (>) the energy *spent* in partially removing one or more electrons from each of contributing atom 2
3 Atomic size & shape affect bonding! & symmetry!!! Unequal Electron Sharing A pure covalent bond occurs only when two identical atoms are bonded: N 2 Polar Covalent Bond: : Unequal sharing between two dissimilar atoms Therefore, the electrons are nearer to one of the atoms, and that atom acquires a partial negative charge (δ ).( And consequently the other atom has a partial positive charge (δ+). Bond is referred to as polar & the molecule can be called a dipole (having two poles KEY QUESTION: ow do you determine which atom has the partial negative charge and which atom has the partial positive charge? 3
4 Electronegativity is the Answer! p. 58 All atoms are not created equal! Defintion: Electronegativity (χ) The extent to which an element attracts an electron toward itself The bigger the EN difference, the more polar the bond. Trends similar to electron affinity Ionic vs. Covalent Use definition of IE to define ionic χ > 1.7 : Ionic χ < 1.7 : < 1.7 : Polar covalent ( is never ionic) e.g. : C - Cl versus K - Cl χ = 0.5 χ = 2.2 What affects bond length? 1. The smaller the principle quantum numbers of the valence orbitals, the shorter the bond. 2. The higher the bond multiplicity, the shorter the bond. 3. The higher the effective nuclear charge of the bonded atoms, the shorter the bond. 4. The larger the electronegativity difference, the shorter the bond. 4
5 Bond Energy Oxidation Numbers 1. Bond strength increases as more electrons are shared between the atoms 2. Bond strength increases as the electronegativity difference ( χ)( between bonded atoms increases. 3. Bond strength decreases as bonds become longer. Valence Bonding (Localized) vs. Molecular Orbital Theory (Delocalized) Orbital Overlap (Localized Bonding) Bonding orbitals are constructed by combining atomic orbitals from adjacent atoms. From Quantum Mechanics: orbitals can add or subtract; therefore constructive or destructive interference is possible 5
6 Localized bonding 2 Orbital Overlap: As two atoms approach, the overlap of their 1s atomic orbitals increases. The wave amplitudes add, generating a new orbital with high electron density between the nuclei. P 3 Phosphine is a colorless, highly toxic gas with bond angles of Describe the bonding in P 3. Oh shoot What about methane? s and p hybridization What is the electron configuration of methane? 6
7 Methane hybridization General Features of ybridization 1. The # of valence orbitals generated by hybridization equals the # of valence AOs participating in hybridization. 2. The steric number of an inner atom uniquely determines the number and type of hybrid orbitals. 3. ybrid orbitals form localized bonds by overlap with atomic orbitals or with other hybrid orbitals. 4. There is no need to hybridize orbitals on outer atoms, because atoms do not have limiting geometries. The bonds formed by all other outer atoms can be described using valence p orbitals. 7
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