BONDING Bond types bond energies

Types of Chemical Bonds (what holds atoms together): Covalent – sharing electrons between non metal atoms to form molecules. Nonpolar – equal sharing Polar – unequal sharing Ionic – giving or taking electrons between atoms to form compounds. Metallic – a “sea of electrons” among metal atoms to form the metal solid.

Types of Bonds IONIC COVALENT Bond Formation e- are transferred from metal to nonmetal e- are shared between two nonmetals Type of Structure crystal lattice true molecules Physical State solid liquid or gas Melting Point high low Solubility in Water yes usually not Electrical Conductivity yes (solution or liquid) no Other Properties odorous

METALLIC e- are delocalized among metal atoms “electron sea” solid Bond Formation e- are delocalized among metal atoms Type of Structure “electron sea” Physical State solid Melting Point very high Solubility in Water no yes (any form) Electrical Conductivity Other Properties malleable, ductile, lustrous

Ionic Bonding - Crystal Lattice Types of Bonds Ionic Bonding - Crystal Lattice

Covalent Bonding - True Molecules Types of Bonds Covalent Bonding - True Molecules Diatomic Molecule

Metallic Bonding - “Electron Sea” Types of Bonds Metallic Bonding - “Electron Sea”

NaCl CO2 CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular

Chemical Bond attractive force between atoms or ions that binds them together as a unit bonds form in order to… decrease potential energy (PE) increase stability

Bond Polarity Most bonds are a blend of ionic and covalent characteristics. Difference in electronegativity determines bond type.

Bond Polarity Electronegativity Attraction an atom has for a shared pair of electrons. higher e-neg atom  - lower e-neg atom +

Bond Polarity Electronegativity Trend Increases up and to the right.

Bond Polarity Nonpolar Covalent Bond e- are shared equally symmetrical e- density usually identical atoms

+ - Bond Polarity Polar Covalent Bond e- are shared unequally asymmetrical e- density results in partial charges (dipole) + -

Bond Polarity Nonpolar (Pure) Polar Ionic View Bonding Animations.

Bond Polarity 3.0-3.0=0.0 Nonpolar 3.0-2.1=0.9 Polar 3.0-0.9=2.1 Ionic Examples: Cl2 HCl NaCl 3.0-3.0=0.0 Nonpolar 3.0-2.1=0.9 Polar 3.0-0.9=2.1 Ionic

MOLECULAR POLARITY When does a MOLECULE become polar?? Write some examples:

ATTRACTIONS BETWEEN ATOMS INTRAMOLECULAR FORCES Covalent Bonds Ionic (Electrostatic) Bonds Metallic Bonds Examples:

8.4 in TEXTBOOK: Attractions between Molecules INTERMOLECULAR FORCES Van der Waals Forces – weakest attractions Result of dipole interactions – when one polar molecule “lines up” to another molecule Dispersion forces – occur between ALL molecules (polar or not) and are momentary Hydrogen Bonds – strongest type of intermolecular force Always involves hydrogen An attraction to a hydrogen atom ALREADY bonded to another strongly electronegative atom Hydrogen becomes strongly polarized with a bond and will therefore try to compensate when close to another molecule

BOND ENERGIES

Energy of Bonding CD + 100kJ  C + D Reaction 1(endo): Endothermic – heat is a reactant heat is absorbed breaking a bond Reaction 1(endo): CD + 100kJ  C + D

Reaction 2 (exo): A + B  AB + 100kJ Energy of Bonding Exothermic – heat is a product heat is released making a bond less potential energy bond is more stable Reaction 2 (exo): A + B  AB + 100kJ Reaction 3: E + B  EB + 100kJ Reaction 4: F + G  FG + 400kJ

Which bond is more stable, EB or FG? Reaction 1(endo): CD + 100kJ  C + D Reaction 2 (exo): A + B  AB + 100kJ Reaction 3: E + B  EB + 100kJ Reaction 4: F + G  FG + 400kJ Which bond is more stable, EB or FG? FG, because it takes more energy (300kj more) to break the bond than EB.

Energy of Bond Formation Bond Energy Short bond = high bond energy