Chap 10 Liquids & Solids

Key terms Molecules – atoms joined by covalent bonds (molecular compounds) Condensed states – solid and liquid Intramolecular forces – within the molecule Intermolecular forces – between molecules

Intermolecular Forces Known as IMFs Arise from unequal distribution of electrons in atoms and the attraction of opposites Contribute to properties of the molecules o Melting & boiling points (lower mp or bp = weaker IMFs) o Bonding (type of bond affects IMFs) o Vapor pressures (higher IMFs = lower vapor pressure) o Dissolving process – making solutions strongest = H Bonding; weakest = London Dispersion strongest = solids; weakest = gases

Types of Intermolecular Forces Dipole-Dipole Forces – forces between polar molecules molecules with dipoles attract each other line up so positive and negative ends are close molecules arrange to maximize interactions Hydrogen bonding – strong dipole-dipole o H bonded to highly electronegative atom London Dispersion Forces – forces between nonpolar molecules and Noble gases Induce a dipole – polarizability – temporary dipole Weak forces Increase as size of atom increases

Interactions Dipole - dipole Ion – dipole o H-Bonding Ion – ion o Ions only Induced dipole – induced dipole o London dispersion only

Liquids General Properties: Definite volume; indefinite shape Molecules in constant motion – ability to flow Low compressibility Surface tension = resistance of a liquid to increase its surface area Capillary action = spontaneous rising of liquid up narrow tube o Cohesion – intermolecular forces among molecules within the liquid o Adhesion – forces between liquid molecules and container Viscosity = measure of a liquid’s resistance to flow o Higher IMF = higher viscosity

Solids General properties: Definite shape; definite volume Rigid structure/shape Particles vibrate in place 2 Types of solids: o Crystalline solids – ionic & molecular compounds Highly regular pattern/arrangement Lattice = 3D arrangement showing positions of atoms Unit cell = smallest repeating unit of the lattice Determined through X-Ray diffraction o Amorphous solids – mostly covalent Random order/arrangement EX: glass

Network solids – large “giant” molecules Called atomic solids Tend to be brittle – due to directional bonds; arrangement/structure o Rings = networks o Layers of atoms are weaker – gaps between layers o Stronger structure = all atoms bonded Refer to diamond vs graphite (p471) Molecular SolidsIonic Solids Strong covalent bonding within molecule but weak between molecules Low melting points Soft Made of nonmetals Strong electrostatic forces between ions High melting points Hard Made of oppositely charged ions

Metals General Properties: Solids; good conductors; malleable; ductile; high melting points Metallic crystals o Bonding is strong but non directional Ionic bond when metal to non-metal “sea” of electrons when metal to metal Mostly form alloys when combined

Phase Changes aka changes of state Vaporization or evaporation o Enthalpy/heat of vaporization (∆H vap ) – energy needed to vaporize 1 mole of liquid at 1 atm. o Endothermic process o Vapor pressure Condensation Sublimation Boiling Melting o Enthalpy/heat of fusion (∆H fus ) – energy change occurs at the melting point of a solid Freezing o Melting point and freeing point same Requires energy change for all - E = q=mC∆T

Vapor Pressure

Phase Diagram Can show phases of materials as function of temperature and pressure