| Flap 1 | Flap 2 |
| Dipole-Dipole Bonding | occurs between polar molecules; (+) end of 1 molecule is attracted to the (-) end of another molecule; 1-2% of the strength of an ionic bond; larger molecule = higher boiling point |
| Hydrogen Bonding | special type of dipole-dipole (stronger than regular); between hydrogen and "highly electronegative elements" (usually F, N, O); larger molecule = higher boiling point |
| London Dispersion Forces | occurs between nonpolar molecules; size of molecule and strength of force are directionally proportional (larger molecule= higher boiling point) |
| Surface Tension | resistance of liquid to increase its surface area; large intermolecular forces = high surface tension |
| Capillary Action | the spontaneous rising of a liquid in a narrow tube (forming meniscus); when substance is polar, meniscus is concave, when substance is nonpolar, its convex |
| Cohesive Forces | forces among the molecules that can cause capillary action |
| Adhesive Forces | forces between the molecules and the container that can cause capillary action |
| Viscosity | measure of a liquid's resistance to flow; large intermolecular forces = highly viscous |
| Crystalline Solids | solids with a highly regular arrangement of their components |
| Amorphous Solids | solids with considerable disorder in their structures |
| Ionic Solids | type of crystalline solid; ions are at the points of the lattice |
| Molecular Solids | type of cyrstalline solid; have covalently bonded molecules at each lattice point |
| Atomic Solids | type of crystalline solid; have atoms at the lattice points |
| Metallic Solids | subgroup of atomic solids; a special type of delocalized nondirectiona covalent bonding occurs |
| Network Solids | subgroup of atomic solids; atoms bond to each other with strong directional covalent bonds that lead to networks of other atoms |
| Group 8A Solids | subgroup of atomic solids; noble gases are attracted to each other with weak London dispersion forces |
| Hydration | process by which ionic substances dissociate into ions. cations are attracted to the (-) end of H2O (O) and anions are attracted to the (+) end (H2) |
| Properties of Electrolytes | dissociate in solution; solution conducts electricity |
| Strong Electrolytes | completely dissociate in H2O |
| Weak Electrolytes | dissociate to a small degree in H2O |
| Non-Electrolytes | may dissolve in H2O, but do not dissociate at all |
| Strong Acids | HCl, H2SO4, HNO3, HClO4, HBr, HI |
| Strong Bases | NaOH, KOH, LiOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 |
| Weak Acids | HC2H3O2, H3PO4, HNO2, HOCl, C6H5COOH |
| Weak Bases | NH3 |
| Molarity | moles of solute / liters of solution |
| Molality | moles of solute / kg of solvent |
| Mole Fraction | moles of solute / moles of solution |
| Mass Percent | ( g solute / g solution ) x 100 |
| Acids | releases H+ into solution; proton donor; electron acceptor |
| Bases | releases OH- into solution; proton acceptor; electron donor |
| Properties of Acids | dissociate into ions in solution; solution conducts electricity; low pH (below 7); litmus turns red; phenolphthalein remains colorless; if reacted w/ active metal, produces H2 gas; if reacted with carbonates, will form CO2 |
| Equilibrium | equal rates of reaction; lies far to the right for strong acids, far to the left for weak acids |
| Ka | equilibrium expression; [concentration of products] / [concentration of reactants] |
| Amphoteric | can act as both an acid and as a base, e.g., water |
| Evaporation | the process by which molecules go from liquid to gaseous state; endothermic because gas has more energy than liquid, so energy must go into the system |
| Boiling | gas is able to come up from the bottom; occurs throughout entire liquid |
| Vapor Pressure | ability of molecules to enter vapor stage |
| Boiling Point | temperature at which vapor pressure of liquid is equal to the atmospheric pressure |
| Normal Boiling Point | temperature at which the vapor pressure of the liquid is equal to the normal atomospheric pressure |
| Temperature | a measure of the average kinetic energy |
| Relationship Between Temperature and Vapor Pressure | as temperature increases, energy of molecules increase. higher energy means less energy has to be added to become gaseous, so it is easier to become gaseous. higher ability to become gaseous = higher vapor pressure.as temp increases so does vapor pressure |
| Why does evaporation lower a liquid's temperature? | evaporation is endothermic, so energy enters the system and increases the molecules' energy. the energetic molecules evaporate, leaving low-energy molecules (not enough energy to become gas). lower energy is reflected in a lower temperature. |
| Melting Point | point where the liquid and solid have identical vapor pressures |
| Normal Melting Point | temperature at which the solid and liquid states have the same vapor pressure under conditions where the total pressure is one atmosphere |
| Raolt's Law (Vapor Pressure of Solution Formula) | P(soln) = X(solv)P(solv) (vapor pressure of solution = (mole fraction of solvent)(vapor pressure of pure solvent) |
| Sublimation | process by which molecules of a solid become gaseous without passing through the liquid state |
| Colloids | suspension of tiny particles in some medium; particles are single large molecules or aggreates of molecules/ions; also called colloidal dispersions |
| Tyndall Effect | suspended particles scatter light, making the beam of light visible from the side; can be used to distinguish between a suspension and a true solution |
| Coagulation | destruction of a colloid; can be caused by heating or adding an electrolyte |
| Van't Hoff Factor (i) | i = ( moles of particles in solution / moles of solute dissolved ); number of particles a substance dissociates into when in solution |
| Ion Pairing | some ions remain paired instead of dissociating; leads to deviation from expected value of "i" |
| Osmotic Pressure (formula) | osmotic pressure = iMRT |
| Gas Law Constant (R) | 0.0826 |
| Change in Freezing Point (formula) | Tf = imTf ; Tf = temperature of pure solvent in K |
| Change in Boiling Point | Tb = imTb ; Tb = boiling point of pure solvent in K |
| Triple Point | only point at which all three phases can coexist |
| Critical Point | end of the line between vapor and liquid phases; shows the critical temperature and critical pressure |
| Critical Temperature | highest temperature at which vapor can be liquified |
| Critical Pressure | the pressure needed to liquify vapor at the critical temperature |
| Ion Dipole | ionic bonding; occurs between two polar ionic molecules; strongest type of intermolecular bond |
| Mixture | two or more elements physically sharing the same space; the elements' appearances may change, but their chemical properties won't; a mixture can be separated into its component parts with relative ease |
| Compound | two or more elements chemically combined; a new substance is formed and the individual elements lose their original properties |
| Homogenous | consistent throughout; components are evenly dispersed |
| Solution | a homogenous mixture |
| Solubility | the ability to dissolve |
| Temperature's Effect on Solubility | increases solubililty in solids and liquids; decreases solubility in gases |
| Pressure's Effect on Solubility | increases solubility in gases; no effect in solids/liquids |
| Why is it beneficial to have measures of concentration that are Temperature Independent? | temperature affects volume, but not mass. molality, mole fraction, & mass percent are mass-based, but molarity is volume-based. M of a solution may change as the temp. changes, but m, x, and m% won't change |
| Explain nature of solute and solvent in relation to (delta)T. | (delta)T doesn't depend on the nature of the solute, just on the number of solute particles (i); it does however depend on the nature of the solvent, because Kb and Kf are specific to each individual solvent. |
| Colligative Properties | properties that change when a solute dissolves in a solvent; vapor pressure, osmotic pressure, freezing point, boiling point |
| Enthalpy of Vaporization | (delta)Hvap; energy required to vaporize one mole of molecules at standard atmospheric pressure |
| Enthalpy of Fusion | (delta)Hfus; energy required to melt one mole of molecules at standard atmospheric pressure |
| Conjugate Base | everything that remains of the acid molecule after a proton is lost |
| Conjugate Acid | formed when the proton is transferred to the base; essentially base + H+ |
| Conjugate Acid-Base Pair | two substances related to each other by the donating and accepting of a single proton; two substances that differ only by the presence or lack of H+; e.g., acid and conjugate base |
| Acid-Base Reaction | H+ and OH- combine to form H2O; also called neutralization reaction; enough base is added to react exactly with the acid in a solution |
| Precipitation Reaction | a reaction in which an insoluble substance forms and separates from the solution |
| Oxidation-Reduction (Redox) Reaction | a reaction in which one or more electrons are transferred |
| Equivalence Point | point in titration where enough titrant has been added to neutralize the base |
| Endpoint | point where indicator changes color |
