Chem 101 Lecture 10: Lewis Bonding PDF

Chapter 10 Chemical Bonding 1: The Lewis Model  I. Chemical Bonds A. Types of Chemical Bonds Electrons transferred Electrons shared  I. Chemical Bonds B. Valence Electrons as Dots Oxygen: Lewis Symbol ...

Chapter 10 Chemical Bonding 1: The Lewis Model  I. Chemical Bonds A. Types of Chemical Bonds Electrons transferred Electrons shared  I. Chemical Bonds B. Valence Electrons as Dots Oxygen: Lewis Symbol  B. Valence Electrons as Dots Filled valence shell Filled valence shell Octet of electrons Duet of electrons Octet rule – atoms transfer or share electrons to gain octets of electrons  II. Ionic Bonding A. Lewis Symbols Lewis symbol  II. Ionic Bonding B. Lattice Energy 1 𝑁𝑎 𝑠 + 𝐶𝑙2 𝑔 → 𝑁𝑎𝐶𝑙 𝑠 Δ𝐻𝑓𝑜 = −411 𝑘𝐽/𝑚𝑜𝑙 2 𝑁𝑎 𝑔 → 𝑁𝑎+𝑔 + 𝑒 − Δ𝐻 𝑜 = +496 𝑘𝐽/𝑚𝑜𝑙 𝐶𝑙 𝑔 + 𝑒 − → 𝐶𝑙 −𝑔 Δ𝐻 𝑜 = −349 𝑘𝐽/𝑚𝑜𝑙 Δ𝐻𝑓𝑜 = +147 𝑘𝐽/𝑚𝑜𝑙  B. Lattice Energy Lattice energy – energy associated with the formation of a crystalline lattice of alternating cations and anions from the gaseous ions Coulomb’s Law Energy decreases when cations and anions are close together Calculated via Born-Haber Cycle  B. Lattice Energy Born-Haber Cycle Hypothetical formation of ionic compound from constituent elements Enthalpy of all steps known except lattice energy Hess’s Law  III. Covalent Bonding A. Single Bonds Covalent bond – sharing of electrons between two atoms Water Lewis symbols Shared electrons count for both atoms  III. Covalent Bonding A. Single Bonds Covalent bond – sharing of electrons between two atoms Water Lewis symbols Shared electrons count for both atoms  III. Covalent Bonding A. Single Bonds Covalent bond – sharing of electrons between two atoms Water Lewis structure Lewis symbols Shared electrons count for both atoms  A. Single Bonds Why do halogens form diatomic molecules? or 7 valence electrons 8 valence electrons  III. Covalent Bonding B. Double and Triple Bonds Oxygen molecule Double bond Incomplete Complete octet octet Double bonds are shorter and stronger than single bonds  B. Double and Triple Bonds Nitrogen molecule Incomplete Incomplete Complete octet octet octet Triple bonds are shorter and Triple bond stronger than double bonds  B. Double and Triple Bonds Bonds generally get weaker as they get longer  III. Covalent Bonding C. Electronegativity and Bond Polarity Equal sharing of electrons + - Polar covalent bond – positive pole and negative pole Dipole moment – separation of positive and negative charge C. Electronegativity and Bond Polarity Unequal sharing of electrons Fluorine is more electronegative than hydrogen Electronegativity – ability of an atom to attract electrons in a chemical bond C. Electronegativity and Bond Polarity Increase across row Decrease down column C. Electronegativity and Bond Polarity Covalent bond – small difference in electronegativities Ionic bond – large difference in electronegativities Polar covalent bond – intermediate difference in electronegativities C. Electronegativity and Bond Polarity  III. Covalent Bonding D. Drawing Lewis Structures Draw the Lewis structure of CH2O 1. Draw skeletal structure. Hydrogen atoms are always terminal More electronegative atoms terminal Least electronegative atoms central  D. Drawing Lewis Structures Draw the Lewis structure of CH2O 2. Calculate the total number of valence electrons  D. Drawing Lewis Structures Draw the Lewis structure of CH2O 3. Complete octets of terminal atoms 1 bond = 2 electrons Complete as many octets/duets as possible  D. Drawing Lewis Structures Draw the Lewis structure of CH2O 4. If there are electrons are left over - distribute them to the central atom 5. If there are no electrons left over – move lone pairs of electrons to complete the octets of the central atom  D. Drawing Lewis Structures Ex. Draw the Lewis structure for HCN.  III. Covalent Bonding E. Resonance Ozone (O3) Resonance – more than one valid Lewis structure Same skeletal formula; different electron arrangement  E. Resonance Electrons delocalized over all resonance structures – resonance hybrid Delocalization stabilizes molecules  E. Resonance Ex. Draw the three resonance structures for the nitrate ion. Sketch the resonance hybrid structure  III. Covalent Bonding F. Formal Charge Formal charge – charge an atom would have if all bonding electrons were equally shared (ignore effects of electronegativity) 1 𝐹𝑜𝑟𝑚𝑎𝑙 𝑐ℎ𝑎𝑟𝑔𝑒 = # 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒 − # 𝑛𝑜𝑛𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝑒 + # 𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝑒 − − − 2  F. Formal Charge Used to distinguish between preferred Lewis/resonance structures Rules: 1. Sum of all formal charges in a neutral molecule equals zero 2. Sum of all formal charges in an ion equals the ionic charge 3. Small formal charges on individual atoms are better than large ones 4. Negative formal charges should reside on the most electronegative atom  F. Formal Charge Which is the preferred Lewis structure?  F. Formal Charge Ex. Assign formal charges to each atom in the resonance forms of the cyanate ion (OCN-). Which resonance form is likely to contribute the most to the correct structure?  III. Covalent Bonding G. Exceptions to the Octet Rule Odd number of electrons Free Radicals Can never complete octet of all atoms Draw structures to minimize formal charge 11 valence electrons Assign unpaired electrons to least electronegative element  G. Exceptions to the Octet Rule Incomplete Octets Typically observed in B, and Be compounds Why not form double bonds to complete octets?  G. Exceptions to the Octet Rule Formal charges: 0 +1 0 0 0 0 -1 0  G. Exceptions to the Octet Rule Expanded Octets Only occurs for 3rd row and higher elements Accessible d-orbitals accommodate Energy extra electrons  G. Exceptions to the Octet Rule Expand octets in order to minimize formal charge  G. Exceptions to the Octet Rule Ex. Write the Lewis structure for H3PO4. If necessary, expand octet(s) to minimize formal charges.  III. Covalent Bonding H. Bond Energy Bond energy – energy required to break 1 mole of the bond in the gas phase 𝐶𝑙2 𝑔 → 2 𝐶𝑙 𝑔 Δ𝐻 = 243 𝑘𝐽/𝑚𝑜𝑙 H-Cl bond is stronger 𝐻𝐶𝑙 𝑔 →𝐻 𝑔 + 𝐶𝑙 𝑔 Δ𝐻 = 431 𝑘𝐽/𝑚𝑜𝑙 Higher bond energy → stronger bond → more chemically stable (less reactive)  H. Bond Energy Bond energies are molecule dependent 𝐻3 𝐶 − 𝐻 𝑔 → 𝐻3 𝐶 𝑔 +𝐻 𝑔 Δ𝐻 = 438 𝑘𝐽/𝑚𝑜𝑙 𝐹3 𝐶 − 𝐻 𝑔 → 𝐹3 𝐶 𝑔 +𝐻 𝑔 Δ𝐻 = 446 𝑘𝐽/𝑚𝑜𝑙 𝐵𝑟3 𝐶 − 𝐻 𝑔 → 𝐵𝑟3 𝐶 𝑔 +𝐻 𝑔 Δ𝐻 = 402 𝑘𝐽/𝑚𝑜𝑙 𝐶𝑙3 𝐶 − 𝐻 𝑔 → 𝐶𝑙3 𝐶 𝑔 +𝐻 𝑔 Δ𝐻 = 401 𝑘𝐽/𝑚𝑜𝑙 Average C-H bond energy : 422 kJ/mol  H. Bond Energy 𝐸𝑠𝑖𝑛𝑔𝑙𝑒 𝑏𝑜𝑛𝑑 < 𝐸𝑑𝑜𝑢𝑏𝑙𝑒 𝑏𝑜𝑛𝑑 < 𝐸𝑡𝑟𝑖𝑝𝑙𝑒 𝑏𝑜𝑛𝑑  H. Bond Energy Δ𝐻𝑟𝑥𝑛 can be estimated from bond energies 𝐻3 𝐶 − 𝐻 𝑔 + 𝐶𝑙 − 𝐶𝑙 𝑔 → 𝐻3 𝐶 − 𝐶𝑙 𝑔 + 𝐻 − 𝐶𝑙 𝑔 Bonds broken Bonds formed (requires energy) (releases energy)  H. Bond Energies Exothermic reaction −Δ𝐻𝑟𝑥𝑛 : weak bonds break; strong bonds form Endothermic reaction +Δ𝐻𝑟𝑥𝑛 : strong bonds break; weak bonds form