Chapter 10 Lecture Notes - CHEM 1113 Broering
Page 1: Introduction to Covalent Bonding
Covalent Bonds
Definition: A bond created by two atoms sharing one or more pairs of electrons.
Bond Length: The distance between the nuclei of two atoms joined together in a bond.
Bond Energy: The energy needed to break one mole of a covalent bond (also known as bond strength).
Page 2: Lewis Symbols and Structures
Lewis Symbols
Introduced by Gilbert N. Lewis.
Atoms form bonds by sharing pairs of electrons to mimic the electron configuration of noble gases (ns2np6).
Octet Rule: Main group atoms tend to gain, lose, or share electrons so that each atom has eight valence electrons.
Hydrogen: Forms bonds to achieve two valence electrons (1s2), referred to as the duet rule.
Example 10.1: Carbon has four valence electrons and must make four bonds to satisfy the octet rule.
Page 3: Lewis Symbols Continued
Lewis Symbols:
A chemical symbol surrounded by dots representing valence electrons; place dots on four sides of the symbol one at a time before pairing.
Bonding Capacity: Number of covalent bonds an atom can form equals the number of unpaired electrons in its Lewis symbol.
Lewis Structures of Ionic Compounds:
Can represent ionic compounds where cations have no valence electrons and anions gain electrons to fill their valence shell, represented with brackets showing charge.
Page 4: Lewis Structures of Molecular Compounds
Molecular Compounds:
Covalent bonds are pairs of shared valence electrons, counted for both atoms in a bond.
Electrons behave more like waves than particles, existing as overlapping charge clouds accessible to both atoms.
Page 5: Lewis structure of HCl
In HCl, H shares one pair of electrons with Cl.
H has access to two electrons (duet) and Cl has three lone pairs plus one shared pair, forming a covalent bond.
Page 6: Example of Shared and Unshared Electrons
Hydrazine (N2H4): Contains 10 shared electrons (5 single bonds) and has 4 unshared electrons (2 lone pairs).
Hydrogen Cyanide (HCN): Contains 8 shared electrons (1 single bond and 1 triple bond) and 2 unshared electrons (1 lone pair).
Page 7: Valid Lewis Structures
Characteristics: Must equal total number of valence electrons and obey the octet/duet rule.
Example: In CO2, each O and C must have 8 electrons.
Steps for Drawing Lewis Structures:
Count total valence electrons.
Determine the central atom and arrange others around it.
Draw single bonds.
Fill in electrons on the bonded atoms (except hydrogen).
Compare leftover valence electrons to total.
Complete the octet on the central atom if incomplete.
Verify it’s a valid structure.
Page 8: Lewis Structure of CHCl3
Steps:
Sum up valence electrons: 4 (C) + 1 (H) + 3(7) (Cl) = 26.
Select C as central and bond to H and Cl atoms.
Completion of octets results in all atoms fulfilling their valence requirements.
Page 9: Drawing Lewis Structures: Ammonia (NH3)
Total valence electrons = 8.
Nitrogen is the central atom; H atoms complete their valence shell through single bonds.
Page 10: Double and Triple Bonds
Double Bond: Two atoms share two pairs of electrons.
Triple Bond: Two atoms share three pairs of electrons.
Page 11: Practice with Acetylene (C2H2)
Calculate valence electrons for C2H2, bond C atoms with H, leading to a structure with shared pairs that fulfill bonding capacities.
Page 12: Drawing Lewis Structure for CO2
Follow steps similar to previous, modify bonds as needed to ensure that C achieves octet through multiple structures.
Page 13: Lewis Structures of More Complex Molecules
Use molecular formula to guide arrangements (e.g., ethanol CH3CH2OH).
Structure indication: first C bonded to three H, second C attached to first, 2 H, and O.
Page 14: Polyatomic Ions
Polyatomic ions consist of covalently bonded atoms with a net charge.
Adjust valence counts based on charges and represent structures in brackets, indicating overall charge.
Page 15: Resonance Structures for Carbonate Ion
The carbonate ion (CO3^2-) has multiple resonance structures; properties shared equally among all structures.
Page 16: Resonance and Stability
Structures can be drawn that reflect delocalization of electrons, stabilizing the molecule and reducing potential energy.
Page 17: Electronegativity
Definition: The ability of a bonded atom to attract shared electrons.
Trends: Increases left to right across a period and bottom to top in a group.
Page 18: Calculating Formal Charge
Formal Charge: Difference between the number of valence electrons an atom has and its apparent count in a Lewis structure.
Best Lewis structures have charge distribution as close to zero as possible, particularly for central atoms.
Page 19: Example of Formal Charge on Cyanate
Analyzing formal charge across Lewis structures of different configurations to identify most accurate representation.
Page 20: Concept Test on Formal Charge
Understanding concept tests that reinforce learning about charge in Lewis structures.
Page 21: Exceptions to the Octet Rule
Explanation of situations where atoms do not complete an octet (e.g., hypervalent molecules, odd electron species like radicals).
Page 22: Expanded Octets
Nonmetals in period 3 and below can have more than eight electrons in their valence shells.
Characteristics of examples like SF6 and PCl5 where expanded octets occur for stabilization.
Page 23: Concept Test on Lewis Structures with Expanded Octets
Applying understanding of Lewis structures involving central atoms with expanded valence shells.
Page 24: Electronegativity and Bond Polarity
Guiding principles to infer bond characteristics based on electronegativity differences, contributing to classification of bonds.
Page 25: Bond Polarity Explained
Definitions and illustrations of polar and nonpolar covalent bonds.
Page 26: Electronegativity Guidelines
ΔEN values categorize bonds as nonpolar covalent, polar covalent, or ionic.
Page 27: Bonding Continuum
Dipole moments quantify bond characteristics and indicate sharing spectrum from polar to nonpolar.
Page 28: Enthalpy of Reaction and Bond Energies
Bond energy's importance in determining reaction enthalpies, showcasing energy transitions.
Page 29: Example of Calculating ΔH for Methane Combustion
Application of bond enthalpy for estimating reaction energy savings.
Page 30: Bond Length and Strength Relation
Discussion on how bond strength affects bond length and relationship dynamics.
Page 31: Ranking Bond Lengths
Importance of bond order in determining lengths and energies of various nitrogen-oxygen bonds.
Page 32: Practice with Bond Energies
Activities focused on comparing bond energies and valence implications in ions.
Introduction to Covalent Bonding
Covalent Bonds: Formed by sharing one or more pairs of electrons between atoms. Key properties include bond length (distance between nuclei) and bond energy (energy required to break the bond).
Lewis Symbols and Structures
Lewis Symbols: Represent valence electrons around an atom's symbol; based on the octet rule (eight electrons for main group atoms) and the duet rule for hydrogen (two electrons).
Bonding Capacity: Determined by unpaired electrons in Lewis symbols.
Structures for Ionic Compounds: Represent cations and anions with brackets showing charge.
Lewis Structures of Molecular Compounds
Shared electrons form covalent bonds, behaving as overlapping charge clouds.
Valid Structures: Must follow octet/duet rule and use all valence electrons.
Types of Bonds
Double Bonds: Share two pairs of electrons.
Triple Bonds: Share three pairs of electrons.
Electronegativity and Bond Polarity
Electronegativity: Atom's ability to attract shared electrons; increases across periods and up groups.
Bond Polarity: Determined by differences in electronegativity; classified as nonpolar, polar, or ionic.
Exceptions to the Octet Rule
Some atoms exceed eight valence electrons (expanded octets), common with nonmetals in higher periods (e.g., SF6).