Term
|
Definition
the energy required to remove an electron from a gaseous atom or ion
trend: increases across period and up group (highest=He) |
|
|
Term
|
Definition
| the energy change associated with the addition of an electron to a gaseous atom trend: becomes more negative across period and up group (changes down group are small and numerous exceptions) |
|
|
Term
|
Definition
| half the distance between the nuclei in a molecule consisting of two identical atoms trend: decrease from left to right and increases down group |
|
|
Term
| Mixture vs. Pure Substance |
|
Definition
| mixture: variable composition (homogeneous (solution) or heterogeneous) pure substance: constant composition (element or compound) |
|
|
Term
| Law of Definite Proportion |
|
Definition
| A given compound always contains exactly the same proportion of elements by mass |
|
|
Term
| Law of Multiple Proportions |
|
Definition
| When two elements form a series of compounds the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers |
|
|
Term
|
Definition
1. each element is made up of tiny particles called atoms
2. the atoms of a given element are identical, the atoms of different elements are different in some fundamental way or ways
3. chemical compounds are formed when atoms of different elements combine with each other. a given compound always has the same relative numbers and types of atoms
4. chemical reactions involve reorganization of the atoms, the atoms themselves are not changed |
|
|
Term
| Empirical Formula vs. Molecular Formula |
|
Definition
empirical: simplest whole number ratio of the atoms in a compound
molecular: exact formula of molecule |
|
|
Term
| Strong Electrolytes vs. Weak Electrolytes vs. Nonelectrolytes |
|
Definition
strong: completely ionize
weak: partially ionize
non: don't produce any ions |
|
|
Term
|
Definition
| moles solute/Liters solution |
|
|
Term
|
Definition
1. nitrate (NO3-) soluble
2. alkali metals and ammonium (NH4+) soluble
3. chloride, bromide, iodide soluble (except Ag+, Pb2+, Hg2 2+)
4. sulfate (SO4 2-) soluble (except Ba, Pb, Hg2, Ca)
5. hydroxides (OH-) slightly soluble (except NaOH, KOH)
6. sulfide (S2-), carbonate (CO3 2-), chromate (CrO4 2-), and phosphate (PO4 3-) slightly soluble |
|
|
Term
| Types of Chemical Equations |
|
Definition
formula equation: reactants and products in undissociated form
complete ionic equation: all strong electrolytes shown as ions
net ionic equation: spectator ions eliminated |
|
|
Term
| Bronstead Acids and Bases |
|
Definition
Acid: proton donator
Base: proton acceptor |
|
|
Term
| Arrenhius Acids and Bases |
|
Definition
Acid: produces H+ ions (H3O+) in water
Bases: produces OH- ions in water |
|
|
Term
| Oxidation-Reduction Reactions (Redox) |
|
Definition
electrons are transferred
oxidation: increase in oxidation state, loss of electrons
reduction: decrease, gain
oxidizing agent: electron acceptor
reducing agent: electron donator |
|
|
Term
|
Definition
Boyle: PV=k
Charles: V=bT
Avogadro: V=an
Ideal Gas Law: PV=nRT
molar mass=(dRT)/P
Dalton's Law of Partial Pressures: Ptotal=P1+P2+P3...
|
|
|
Term
| Kinetic Molecular Theory Gases |
|
Definition
1. volume is negligible
2. constant random motion
3. no repulsive or attractive forces
4. average kinetic energy in directly proportional to Kelvin temp
5. collisions with each other and container are elastic
followed at low pressure and high temperatures |
|
|
Term
| Endothermic vs. Exothermic |
|
Definition
Exothermic: potential energy stored in bonds is converted to thermal energy, reactants have higher potential energy, bonds in products are stronger than reactants, more energy is released when new bonds in products are formed than is consumed to break bonds in reactants
Endothermic: energy flows into system as heat increases potential energy, products have higher potential energy and weaker bonds |
|
|
Term
| Thermodynamic Law-thingies |
|
Definition
sum of kinetic and potential energies
ΔE=q + w
q: + energy added, - energy out
w: + surroundings do work on system, - system does work on surroundings
for gas work: w=-PΔV |
|
|
Term
|
Definition
H=E + PV
ΔH=qp (at constant pressure)
Hess's Law: enthalpy same regardless of steps
from bond energies: ΔH= Σn x D(bonds broken)-Σn x D(bonds formed) |
|
|
Term
|
Definition
constant pressure: energy released by reaction=specific heat capacity of solution x mass of solution x increase in temperature= s x m x ΔT
constant volume: ΔT x heat capacity of container |
|
|
Term
| Standard Enthalpy of Formation |
|
Definition
| change in enthalpy that accompanies the formation of one moles of a compound from its elements with all substances in their standard states Compound: gas at 1 atm, pure substance in condensed state is pure substance, solution is 1M Element: form at 1 atm and 25C ΔHf reaction=ΔHf products - ΔHf reactants |
|
|
Term
|
Definition
λv=c
Ephoton=hv=(hc)/λ
de Broglie's Equation: λ=h/(mv) to calculate wavelength of particle |
|
|
Term
|
Definition
n: principle quantum number
l: angular momentum, 0 to (n-1), shape 0=s, 1=p, 2=d, 3=f, 4=g
ml: magnetic, -l to l, orientation
ms: electron spin, -1/2 or 1/2 |
|
|
Term
| Writing Electron Configurations Like a Boss |
|
Definition
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
number of orbitals per subshell: s=1, p=3, d=5, f=7, g=9 |
|
|
Term
|
Definition
the ability of an atom in a molecule to attract shared electrons to itself
trend: fluorine highest, increases left to right, decreases down |
|
|
Term
|
Definition
same number of electrons
more protons=smaller |
|
|
Term
|
Definition
| the change in energy that takes place when separated gaseous ions are packed together to for an ionic solid, negative, exothermic |
|
|
Term
|
Definition
1. Things can exceed the octet rule (I know, say whaaa?) like third row+ elements because of valence d orbitals
2. B and Be often have fewer than eight electrons, very reactive compounds
C, N, O, F always assumed to obey octet rule
Formal Charge: valence e- on free atom - valence e- assigned to atom in molecule, close to zero as possible and negative on most electronegative atom |
|
|
Term
| VSEPR (valence shell electron-pair repulsion) Model |
|
Definition
| 2 effective pairs: linear, 180 3: trigonal planar, 120 4: tetrahedral, 109.5 (3: trigonal pyramidal, 2: bent) 5: trigonal bipyramidal, 120 & 90 (4: see-saw, 3: t-shaped, 2: linear) 6: octahedral, 90 (4: square planar) |
|
|
Term
|
Definition
2: sp, 3: sp2, 4: sp3, 5: dsp3, 6: d2sp3
sigma bonds: between atoms, pi bonds: above and below atoms, double bond: one sigma + one pi |
|
|
Term
|
Definition
bonding: lower in energy, antibonding: higher in energy
bond order: (number of bonding e - number of antibonding e)/2, larger=stronger, larger bond energy, smaller length
paramagnetism=unpaired, dimagnetism=paired
B, C, N: sp mixing changes order O,F: no sp mixing |
|
|
Term
|
Definition
1. Lewis structure
2. Arrangement of electron pairs according to VSEPR
3. Specify hybridization |
|
|
Term
|
Definition
dipole-dipole attraction: polar molecules
hydrogen bonding: strong dipole-dipole forces, when hydrogen is bound to highly electronegative atom (N, O, F)
london dispersion:induced dipole, polarizability: large atoms, more electrons higher
electrostatic attraction: ionic compounds |
|
|
Term
|
Definition
surface tension: resistance of a liquid to an increase in its surface area, high IM forces=high surface tension
capillary action: cohesive forces (IM between liquid molecules) and adhesive forces (between liquid and container, when container is polar)
viscosity: resistance to flow |
|
|
Term
|
Definition
crystalline (regular arrangement, lattice) vs. amorphous (disorder)
types of crystalline: atomic, ionic, molecular
X-ray diffraction according to nλ=2dsinθ |
|
|
Term
|
Definition
| Simple Cubic: 1 net sphere, side= 2r Body-Centered Cubic: 2, side= r4/sqrt(3) Face-Centered Cubic: 4, side=rsqrt(8), 8 tetrahedral holes, 4 octahedral holes |
|
|
Term
|
Definition
| closest packing: aba=hexagonal closest packed (hcp) structure, abc= cubic closest packed (ccp) structure, face-centered |
|
|
Term
|
Definition
| endothermic enthalpy of vaporization: energy required to vaporize 1 mol liquid at 1 atm higher IM forces, lower vapor pressure larger molar mass, lower vapor pressure (dispersion forces) higher temp, higher vapor pressure melting point: liquid and solid have identical vapor pressures |
|
|
Term
|
Definition
solid/liquid line sloped negative: solid less dense than liquid, melting point decreases as pressure increases
positive: solid more dense than liquid, melting point increases with pressure |
|
|
Term
| Henry's Law: gas solubility and pressure |
|
Definition
C=kP
solubility increases when pressure increases
C=concentration of dissolved gas
P=partial pressure of gas solute above solution
k=constant for solution |
|
|
Term
|
Definition
|
|
Term
| Temperature and Gas Solubility |
|
Definition
| gas solubility in liquid typically decreases when temp increases |
|
|
Term
| Vapor Pressure of Solutions |
|
Definition
nonvolatile solute lowers vapor pressure of solvent
Raoult's Law: Psolution=Xsolvent x Pinitial of solvent (modified when both components are volatile) |
|
|
Term
| Colligative Properties: depend only on number of solute particles |
|
Definition
Boiling Point Elevation: ΔT=Kb x molality of solute
kb=molal boiling-point elevation constat
Freezing Point Depression: ΔT=Kf x molality of solute
Osmotic Pressure: Π=MRT, M=molarity, T=kelvin temp, isotonic=identical, hyper=more concentrated solvent goes out, hypo=less concentrated, solvent comes in |
|
|
