CHEMICAL BONDING

CHEMICAL BONDING

Overview of bonding • Chemical bond: attractionbetween the nuclei and valence electrons of different atoms that binds those atoms together. • Why do atoms bond? Most atoms are more stable when they are bonded (lower potential energy)

Types of chemical bonds: • Ionic – stealelectrons metal—nonmetal • Covalent – shareelectrons nonmetal—nonmetal • Metallic – seaofelectrons between metals

IONIC BONDING • Results from the attraction between cations and anions • “Steal” electrons • Metal—nonmetal

Common properties of ionic compounds: • Solid at room temperature • Very high melting points

Positive and negative charges must cancel each other out EXAMPLES: • Na and Cl • Na and O • Ca and O • Ca and Cl • Al and O

COVALENT BONDING • Covalent bond: results from the sharing of electron pairs between two atoms

COVALENT BONDING • Common properties of covalent compounds: • Low melting/boiling points (many are liquids or gases at room temperature)

Covalent or ionic? Determined by electronegativity difference

Covalent or ionic?

Practice: Ionic or Covalent? • NaCl • H2O • CaCl2 • CBr4 • NO2 • CH4 • BaF2 • MgO

METALLIC BONDING • Sea of electrons between metal atoms • This is why metals conduct electricity, have luster, and are ductile/malleable • Metals—Metals

more COVALENT BONDING • Nature favors lowest energy • In covalent bonds the atoms share electrons to achieve lowest energy

Lewis Structures • Lewis structures are used to represent molecules with covalent bonds. (Covalent bonds are shared electrons.)

Lewis Structures:Used to represent molecules with covalent bonds (shared electrons) 1. Count the valence electrons from eachatom 2. Add up the total number of valence electrons • For anions (negatively charged) add e- • For cations (positively charged) take away e-

3. Arrange atoms and connect with single bonds. • Each bond represents two electrons • Nature likes symmetry! • Halogens (F, Cl, Br, I) are not usually center atoms • Hydrogen is NEVER center • Carbon is always center, or least electronegative atom if no carbon

4. Use “left over” electrons and add unshared pairs (lone pairs) of electrons to atoms to give full shells • This is called the octet rule because everything needs 8 electrons for a full shell • Exceptions to the octet rule: • Hydrogen has a full shell when it has two electrons • Do not draw too many electrons!

5. If there are not enough electrons to give everything a full shell, try multiple bonds.

Molecular Geometry • Applies to any Lewis diagram (covalent bonding)

VSEPR Theory • Valence Shell Electron Pair Repulsion • Each bonding group or lone pair of electrons spreads out as far as possible from every other bonding group or lone pair of electrons. • Think 3D!

Polarity • Separation of charge • Occurs because of difference in electronegativity between bonded atoms. • To draw individual bond polarity: • Compare only two atoms at a time • The more electronegative atom attracts electrons more • Draw arrow towards the more electronegative atom

Bond Polarity – Examples CH4 NH3 H2O

Are there lone pairs on the center atom? No Yes Polar Are the peripheral (outer) atoms different? No Yes Polar Nonpolar

Intermolecular Forces • Intermolecular Forces: attraction between molecules

Intermolecular Forces WEAK STRONGER STRONGEST • London Dispersion Forces: • Attraction from constant motion of electrons • Exist between ALL molecules • Dipole-dipole: • Attraction between polar molecules • Hydrogen Bonding: • Special (stronger) type of dipole-dipole interaction • Occurs only between molecules with lone pairs on center atom AND hydrogen bonded to F, O, or N

Determining Intermolecular Forces Is the molecule polar? No Yes London Dispersion Forces Does the molecule have lone pairs on the center atom, with H bonded to F, O, or N? Yes No London Dispersion Forces Dipole-dipole Hydrogen bonding London Dispersion Forces Dipole-dipole

CHEMICAL BONDING