Chemistry 1A03 Introductory Chemistry I PDF

Chemistry 1A03 Introductory Chemistry I Chemistry in the context of health, energy and the environment ©2008 – 2024 McMaster University Unit 5 Chemical Bonding Ch....

Chemistry 1A03 Introductory Chemistry I Chemistry in the context of health, energy and the environment ©2008 – 2024 McMaster University Unit 5 Chemical Bonding Ch.10: Chemical Bonding Chem 1 1A03  This Week at ChemClub! We’re headed to the McMaster ©2008 – 2024 McMaster University Museum of Art! Meet in front of the MMA at 11:25 am! Chem 2 1A03  Bonding Involves transfer or sharing of outer electrons, usually to acquire a stable configuration (Lewis) Ionic bonding (transfer of electrons) ©2008 – 2024 McMaster University Usually between a metal and non-metal Na [Ne]3s1 becomes Na+ [Ne] Cl [Ne]3s23p5 becomes Cl− [Ar] Chem 3 © 1A03 Covalent Bonding Covalent bonding Sharing of electrons Often to attain an octet of electrons Often between 2 non-metals Lewis structure shows all electrons as equivalent ©2008 – 2024 McMaster University Bonds depicted as lines : : H : Cl : H Cl : : : Chem 4 1A03 Electronegativity (EN) – The Trend Atom’s ability to compete for e− in a bond Trend: EN increases across a period and up a group Pauling scale: F 4.0 (highest EN) ©2008 – 2024 McMaster University Chem 5 © 1A03 Bond Polarity Polar covalent bonds Unequal sharing of e− + - Indicated by polar arrow and partial charges H Cl Dictated by the difference in electronegativity (EN) ©2008 – 2024 McMaster University between atoms EN Bonding Example Large (> 1.9) Ionic NaCl (EN ≈ 2.23) Intermediate Polar Covalent PCl5 (EN ≈ 0.97) (0.5 - 1.9) Small (< 0.5) Pure Covalent Cl2 (EN ≈ 0) Chem 6 1A03 Electrostatic Potential Maps ©2008 – 2024 McMaster University Effect of EN on charge distribution 7 Chem © 1A03 Lewis Structures Show bonding (b) and non-bonding (nb) e−, and formal charges 'Complete shells’ can be achieved by combination of bonding and nonbonding e− (lone pairs) A complete shell is typically an octet with the following ©2008 – 2024 McMaster University exceptions Hydrogen is satisfied with 1 electron pair Beryllium can be satisfied with 2 electron pairs Boron/Aluminum can be satisfied with 3 electron pairs Elements in periods 3 and beyond can have expanded octets if involved as the central atom Bonding e− can be involved in single, double, triple bonds Chem 8 1A03 Drawing Lewis Structures 1. Count total # of valence e- including charge of structure Add e- for negative charge, subtract e- for positive charge 2. Draw skeletal structure (central and terminal atoms) Least electronegative atom is usually the central atom Hydrogen and Fluorine are always terminal 3. Use remaining e- to complete octet of terminal atoms (or 2 e- for hydrogen) ©2008 – 2024 McMaster University 4. Subtract all e- used in previous steps and place any remaining e- on the central atom. Sometimes leads to expanded octets 5. Calculate formal charges (FC) on each atom FC = (Valence e- - ½ bonding e- - nonbonding e-) 6. Minimize formal charges by creating multiple bonds using nonbonding electrons Typically happens when neighbouring atoms have opposite charges 7. Ensure all atom have an allowed electron count (C, N, O, F must obey octet rule; some elements need to have at least or exactly 8 electrons in valence shell) 9 Chem 1A03  BF3 B: 1x3= 3 F: 3x7= 21 Formal Charge: Total e─= 24 F: 7−½(2)−6= 0 B: 3−½(6)−0= 0 ©2008 – 2024 McMaster University Initial e─: 24 Bonds: −6 18 Outer e─: −18 0 Chem 10 © 1A03  NO3− N: 1x5= 5 Formal Charge: O: 3x6= 18 O: 6−½(2)−6= −1 charge= 1 N: 5−½(6)−0= +2 Total e─= 24 ©2008 – 2024 McMaster University Initial e─: 24 Formal Charge: Bonds: −6 O: 6−½(2)−6= −1 18 O: 6−½(4)−4= 0 Outer e─: −18 N: 5−½(8)−0= +1 0 If a molecule is charged: Negative formal charge typically on the most electronegative atom Positive formal charge typically on the least electronegative atom Chem 11 © 1A03  BrOF2+ Br: 1x7= 7 O: 1x6= 6 F: 2x7= 14 charge= −1 Formal Charge: Total e─= 26 F: 7−½(2)−6= 0 O: 6−½(2)−6= −1 Br: 7−½(6)−2= +2 ©2008 – 2024 McMaster University Initial e─: 26 Bonds: −6 Formal Charge: 20 F: 7−½(2)−6= 0 Outer e─: −18 O: 6−½(4)−4= 0 2 Br: 7−½(8)−2= +1 Center e─: −2 0 If a molecule is charged: Negative formal charge typically on the most electronegative atom Positive formal charge typically on the least electronegative atom Chem 12 © 1A03  iClicker #1 Select the most appropriate Lewis structure for chlorate anion (ClO3–) that shows all lone pairs of electrons. A B C ©2008 – 2024 McMaster University D E Chem 13 1A03  iClicker #1 – solution Select the most appropriate Lewis structure for chlorate anion (ClO3–) that shows all lone pairs of electrons. A B C ©2008 – 2024 McMaster University Not charge-minimized Not charge-minimized Incorrect FC on Cl and O Missing a lone pair on Cl Missing a lone pair on Cl D E Incorrect FC on Cl Violates octet rule on O Chem 14 1A03  iClicker #2 Select the Lewis structure that is most appropriate for CO and shows all lone pairs of electrons. The octet rule is more important for C, N, O, F. Minimizing formula charges is more important for other elements in the periodic table. ©2008 – 2024 McMaster University A B C D E Chem 15 © 1A03  iClicker #2 – solution Select the Lewis structure that is most appropriate for CO and shows all lone pairs of electrons. The octet rule is more important for C, N, O, F. Minimizing formula charges is more important for other elements in the periodic table. ©2008 – 2024 McMaster University A B C Wrong FC Only 6 e– on C Violates octet Missing lone rule for O pair on C D E Wrong FC (not indicated) Chem 16 © 1A03 Resonance Structures for PO43- For each PO43- there are 4 equivalent charge-minimized structures (resonances structures) Molecule exists as a hybrid of all formal resonance structures, known as the resonance hybrid Resonance hybrid bond lengths, orders and charges are the ©2008 – 2024 McMaster University average of all equivalent resonance states Most polyatomic anions have resonance structures Chem 17 © 1A03 Resonance Structures for PO43- Average formal charge for an atom: total charges on atoms total # of that atom 0+ −1 + −1 +(−1) 3 Average formal charge on O: =− One of the 4 4 4 resonance Bond order: single (1), double (2), triple (3) structures for ©2008 – 2024 McMaster University PO43– Average bond order: total number of bond orders total # of places the bond is formed 2+1+1+1 5 1 Average P-O bond order = = = 1 4 4 4 Chem 18 © 1A03  iClicker #3 Rank the following molecules in increasing average formal charge on terminal O’s (from negative to positive): HCO3−, CO2, CO32− ©2008 – 2024 McMaster University A) HCO3− < CO2 < CO32− B) HCO3− < CO32− < CO2 C) CO32− < CO2 < HCO3− D) CO32− < HCO3− < CO2 Chem 19 1A03  iClicker #4 Rank the following molecules in increasing terminal C-O bond order: HCO3−, CO2, CO32− A) HCO3− < CO2 < CO32− ©2008 – 2024 McMaster University B) HCO3− < CO32− < CO2 C) CO32− < CO2 < HCO3− D) CO32− < HCO3− < CO2 Chem 20 1A03 Bond Order, Length & Energy Covalent bond length Approximately distance between 2 nuclei involved in covalent bond Bond dissociation energy (homolysis) Approximate energy required to break 1 mol of bonds in gas phase As bond order increases, bond length decreases ©2008 – 2024 McMaster University As bond length decreases, bond energy increases Length Energy Bond Order (pm) (kJ mol-1) C-C 1 154 347 C=C 2 134 611 CC 3 120 837 NN 3 109.8 946 21 Chem 1A03  Molecular Shape VSEPR (valence shell electron pair repulsion) Theory AKA Gillespie-Nyholm theory Ron Gillespie, McMaster Chemistry! Electron pairs repel one another ©2008 – 2024 McMaster University Repulsion increases: bond pair/bond pair < bond pair/lone pair < lone pair/lone pair More generally: We should use electron groups instead of electron pairs. An electron group includes bond pair, lone pair, a single electron, double or triplet bond. Chem 22 © 1A03  A note about VSEPR Theory In the OpenStax textbook, the authors claim that the order of repulsive power, from largest to smallest, is: lone pair > triple bond > double bond > single bond McMaster researchers have shown that double bonds ©2008 – 2024 McMaster University occupy slightly more space than a lone pair. (See this reference by Gary Schrobilgen.) Correct order of repulsive power, from largest to smallest: triple bond > double bond > lone pair > single bond Chem 23 1A03 VSEPR Classes AXnEm A = central atom X = atoms bonded to the central atom E = lone electron pairs Electron group geometry dictates the observed ©2008 – 2024 McMaster University molecular shape (watch where they are different!) Table 10.1 – know the shapes! Note: Table 10.1 gives only ideal angles; know which ones are non-ideal also! (class notes) Chem 24 © 1A03 2 Electron Groups Electron Group Geometry - Linear VSEPR class AX2 Molecular Geometry Linear Angles 180° ©2008 – 2024 McMaster University Symmetry Symmetrical Example BeCl2, CO2 Chem 25 © 1A03 3 Electron Groups Electron Group Geometry – Trigonal Planar VSEPR class AX3 AX2E Molecular Geometry Trigonal Planar Bent Angles 120° < 120° (non-ideal) Symmetry Symmetrical Asymmetrical Example BF3 BH2- ©2008 – 2024 McMaster University Chem 26 1A03 4 Electron Groups Electron Group Geometry - Tetrahedral VSEPR class AX4 AX3E AX2E2 Molecular Geometry Tetrahedral Trigonal Pyramidal Bent Angles 109° < 109.5° <