Chemistry 1A03 Introductory Chemistry I PDF

Chemistry 1A03 Introductory Chemistry I Chemistry in the context of health, energy and the environment ©2008 – 2023 McMaster University Unit 5 Chemical Bonding Ch....

Chemistry 1A03 Introductory Chemistry I Chemistry in the context of health, energy and the environment ©2008 – 2023 McMaster University Unit 5 Chemical Bonding Ch.10: Chemical Bonding Chem 1 1A03  Bonding Involves transfer or sharing of outer electrons, usually to acquire a stable configuration (Lewis) Ionic bonding (transfer of electrons) ©2008 – 2023 McMaster University Usually between a metal and non-metal Na [Ne]3s1 becomes Na+ [Ne] Cl [Ne]3s23p5 becomes Cl− [Ar] Chem 2 © 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 – 2023 McMaster University Bonds depicted as lines : : H : Cl : H Cl : : : Chem 3 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 – 2023 McMaster University Chem 4 © 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 – 2023 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 5 1A03 Electrostatic Potential Maps ©2008 – 2023 McMaster University Effect of EN on charge distribution 6 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 – 2023 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 7 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 – 2023 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) 8 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 – 2023 McMaster University Initial e─: 24 Bonds: −6 18 Outer e─: −18 0 Chem 9 © 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 – 2023 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 10 © 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 – 2023 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 11 © 1A03  iClicker #1 Select the most appropriate Lewis structure for chlorate anion (ClO3–) that shows all lone pairs of electrons. A B C ©2008 – 2023 McMaster University D E Chem 12 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 – 2023 McMaster University D E Chem 13 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 – 2023 McMaster University A B C D E Chem 14 © 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 – 2023 McMaster University A B C D E Chem 15 © 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 – 2023 McMaster University average of all equivalent resonance states Most polyatomic anions have resonance structures Chem 16 © 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: =− 4 4 Bond order: single (1), double (2), triple (3) ©2008 – 2023 McMaster University 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 17 © 1A03  iClicker #3 Rank the following molecules in increasing average formal charge on terminal O’s (from negative to positive): HCO3−, CO2, CO32− ©2008 – 2023 McMaster University A) HCO3− < CO2 < CO32− B) HCO3− < CO32− < CO2 C) CO32− < CO2 < HCO3− D) CO32− < HCO3− < CO2 Chem 18 1A03  iClicker #4 Rank the following molecules in increasing terminal C-O bond order: HCO3−, CO2, CO32− A) HCO3− < CO2 < CO32− ©2008 – 2023 McMaster University B) HCO3− < CO32− < CO2 C) CO32− < CO2 < HCO3− D) CO32− < HCO3− < CO2 Chem 19 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 – 2023 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 20 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 – 2023 McMaster University Repulsion increases: bond pair/bond pair < bond pair/lone pair < lone pair/lone pair Note: double bonds occupy slightly more space than a lone pair and have more repulsive power towards other electrons. 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 21 © 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 – 2023 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 22 © 1A03 2 Electron Groups Electron Group Geometry - Linear VSEPR class AX2 Molecular Geometry Linear Angles 180° ©2008 – 2023 McMaster University Symmetry Symmetrical Example BeCl2, CO2 Chem 23 © 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 – 2023 McMaster University Chem 24 1A03 4 Electron Groups Electron Group Geometry - Tetrahedral VSEPR class AX4 AX3E AX2E2 Molecular Geometry Tetrahedral Trigonal Pyramidal Bent Angles 109° < 109.5° <