Bonding: General Concepts

Chapter 8 Notes - Bonding: General Concepts

Types of Chemical Bonds

• Ionic Bonding

• electrons are transferred
• Metals react with nonmetals
• Ions paired have lower energy (greater stability) than separated ions

• Coulomb's Law

• E = 2.31x10-19 J x nm (Q1Q2/r)

• E = energy in joules
• Q1 and Q2 are numerical ion charges
• r = distance between ion center in nanometers
• negative sign indicates an attractive force

• Bond Length (covalent)

• Distance at which the system energy is at a minimum
• Forces at work

• Attractive forces (proton-electron)
• Repulsive forces (electron-electron, proton-proton)

• Energy is given off (bond energy) when two atoms achieve greater stability together than apart

• Covalent Bonds

• Electrons are shared by nuclei
• Pure covalent (non-polar covalent)

• electrons are shared evenly

• Polar covalent bonds

• Electrons are shared unequally
• Atoms end up with fractional changes (δ+ or δ-)

Electronegativity

• Electronegativity - the ability of an atom in a molecule to attract shared electrons to istelf
• Electronegativity trends generally increase up and to the right, because more protons to the right hold electrons more tightly, and atoms going up have weaker shielding effect from having less electrons
• Characterizing bonds-

• Greater eletronegativity difference between two elements means less covalent character and greater ionic character
• Any compound that conducts an electrical current when melted is an ionic compound

Bond Polarity and Dipole Moments

• Dipolar molecules

• molecules with a somewhat negative end and a somewhat positive end (a dipole moment)
• Molecules with preferential orientation in an electrical field
• All diatomic molecules with a polar covalent bond are dipolar

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• Molecules with polar bonds but no dipole moment

• linear, radial, or tetrahedral symmetry of charge distribution

• CO2 - has polar bonds, but is linear, so charges cancel out
• CCl4 - has polar bonds, but is tetrahedral, so charges cancel out

Ions: Electron Configuration and Sizes

• Bonding and Noble Gas electron configurations

• ionic bonds

• electrons are transferred until each species attains a noble gas electron configuration

• Covalent bonds

• electrons are shared in order to complete the valence configurations of both atoms

• Predicting formulas of ionic compounds

• Placement of elements on the periodic table suggest how many electrons are lost or gained to achieve a noble-gas configurations

• Group 1 loses one electron, group two loses two, group six gains two, etc.

• Formulas for compounds are balanced so that the total positive ionic charge is equal to the total negative ionic charge

• Al2+3O3-2

• Total positive = +6
• Total negative = -6

• Sizes of Ions

• 1. Anions are larger than the parent atom
• 2. Cations are smaller than the parent atom
• 3. Ion size increases within a family
• 4. Isoelectronic ions

• Ions with the same number of electrons
• Size decreases as the nuclear charge Z increases

Formation of Binary Ionic Compounds

• Lattice Energy

• The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid

• M+ (g) + X- (g) → MX (s)

• Energy change is exothermic (negative sign)

• Example: Formation of Lithium Fluoride

• Process Description Energy Change (kJ)
• L(s) → Li(g); Sublimation energy 161
• Li(g) → Li+(g) + e-; Ionization energy 520
• 1/2F2 → F(g); Bond energy (1/2 mole) 77
• F(g) + e- → F-(g); Electron affinity -328
• Li+(g) + F-(g) → LiF(s); Lattice energy -1047
• Li(s) + 1/2F2(g) → LiF(s); ΔH -617

• The formation of ionic compounds is endothermic until the formation of the lattice
• The lattice formed by alkali metals and halogens (1:1 ratio) is cubic except for cesium salts

• Lattice Energy Calculations

• Lattice Energy = k (Q1Q2/r)

• k = a proportionality constant dependent on the solid structure and the electron configuration
• Q1 and Q2 are the charges on the ions
• r = shortest distance between centers of the cations and the anions

• Lattice energy increases as the ionic charge increases and the distance betwen anion and cations decreases

The Covalent Chemical Bond: A Model

• Strengths of the Bond Model

• Associates quantities of energy with the formation of bonds between elements
• Allows the drawing of structures showing the spatial relationship between atoms in a molecule
• Provides a visual tool to understanding chemical structure

• Weaknesses of the Bond Model

• Bonds are not actual physical structures
• Bonds cannot adequately explain some phenomena

•  resonance

8.8 Covalent Bond Energies and Chemical Reactions

• Average Bond Energies

Process                         Energy Required (kJ/mol)

CH4(g) → CH3(g) + H(g)                 435

CH3(g) → CH2(g) + H(g)                453

CH2(g) → CH(g) + H(g)                 425

CH(g) → C(g) + H(g)                 339

Total                                         1652