Gen Chem Flashcards
1
Q
effective nuclear charge (Zeff)
A
a measure of the net positive charge experienced by the outermost electrons (how close valence e- are to the nucleus), increases from left to right for elements in the same period
2
Q
Principal quantum number
A
increases by one each time as one moves down the elements of a given group
3
Q
Atomic radius
A
decreases from left to right across a period, increases down a group
4
Q
Ionization energy
A
the energy required to remove an electron from a gaseous species, increases from left to right across a period and from bottom to top in a group
5
Q
Electron affinity
A
the energy dissipated by a gaseous species when it gains an electron, increases across a period from left to right and decreases in a group from top to bottom
6
Q
Electronegativity
A
a measure of the attractive force that an atom will exert on an electron in a chemical bond, increases across a period from left to right and decreases in a group from top to bottom (trends are the same as ionization energy)
7
Q
Isomers
A
molecules w/ the same molecular formulas but different structures
8
Q
Stereoisomers
A
have the same chemical formulas and the same atomic connectivity (the same structural backbone), they differ in how these atoms are arranged in space (their wedge-and-dash pattern)
9
Q
Conformational isomers
A
differ in rotation around single (σ) bonds,
10
Q
Anti Staggered conformation
A
substituents are at the maximum distance from each other (no overlap of atoms along line of sight), also called anti coformation
11
Q
Gauche staggered conformation
A
the two largest groups are 60 degrees apart
12
Q
Eclipsed conformation
A
when the two groups are 120 degrees apart and overlap w/ hydrogens on the adjacent carbon, when two groups atoms and/or groups directly overlap each other with 0 degrees separation (totally eclipsed)
13
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Cis
A
both groups located on the same side
14
Q
Trans
A
groups are on opposite sides
15
Q
Configurational isomers
A
can only change from one form to another by breaking and reforming covalent bonds, 2 categories: enantiomers and diastereomers
16
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Chiral
A
mirror image cannot be superimposed on the original object; the molecule lacks an internal plane of symmetry (ex: left hand cannot fit in right-handed glove even though they are essentially identical)
17
Q
Enantiomers
A
two molecules that are non superimposable mirror images of each other, have the same connectivity but opposite configurations at every chiral center in the molecule
18
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Diastereomers
A
molecules are chiral and share the same connectivity but are not mirror images of each other, have different physical properties
19
Q
Meso compound
A
a molecule w/ chiral centers that has an internal plane of symmetry (two halves that are mirror images)
20
Q
Cahn-Ingold-Prelog priority rules
A
priority is assigned based on the atom bonded to the double-bonded carbons: the higher the atomic number, the higher the priority
21
Q
S
A
when assigning priority the circle is drawn counterclockwise (lowest priority group projects into the page- dashes, #4 doesn’t count when determining the order– true designation is the opposite of what you find when initially assigning priority)
22
Q
R
A
when assigning priority the circle is drawn clockwise
23
Q
Incomplete octet
A
elements that are stable w/ fewer that 8 electrons in their valence shell and include hydrogen (2), helium (2), lithium (2), beryllium (4), and boron (6)
24
Q
Expanded octet
A
any element in period 3 and greater can hold more than 8 electrons, including phosphorus (10), sulfur (12), chlorine (14), and many others
25
Odd numbers of electrons
any molecule w/ an odd number of valence electrons cannot distribute those electrons to give eight to each atom; for example, nitric oxide (NO) has eleven valence electrons
26
Ionic bonding
one or more e- from an atom w/ low ionization energy, typically a metal, are transferred to an atom w/ high electron affinity, typically a nonmetal
--atom that loses e- (cation)
--atom that gains e- (anion)
--difference in electronegativity must be greater than 1.7 on the Pauling scale
27
Covalent bonding
an electron pair is shared btwn 2 atoms, typically nonmetals, that have relatively similar values of electronegativity
--if electron pair is shared equally = nonpolar
--if electron pair is shared unequally = polar
28
Ionic compound properties
-have high melting and boiling points
-many dissolve readily in water and other polar solvents
-in the molten or aqueous state they are good conductors of electricity
-in the solid state they ionic constituents of the compound form a crystalline lattice
29
Covalent compound properties
-have lower melting and boiling points
-bc they do not break down into constituent ions, they are poor conductors of electricity in the liquid state or in aqueous solutions
-the # of shared electron pairs btwn 2 atoms = bond order
-bond length= avg distance btwn 2 nuclei of atoms in a bond (decreases as # of shared e- pairs increases)
-bond energy= the energy required to break a bond by separating its components (the greater the # of pairs of e- shared, the more energy required)
30
Lewis acid
any compound that will accept a lone pair of electrons
31
Lewis base
any compound that will donate a pair of electrons
32
Formal electrons
= valence electrons - dots (lone electron) - sticks (bond)
33
Electronic geometry
describes the spatial arrangement of all pairs of electrons around the central atom, including both the bonding and the lone pairs
34
Molecular geometry
describes the spatial arrangement of only the bonding pairs of electrons
35
Hydrogen bonds
- H bonds to three highly electronegative atoms-- nitrogen, oxygen, or fluorine
-substances that display H bonding have unusually high boiling points
36
Formula for determining the number of moles of a sample substance
Moles = Mass of sample (g)/ Molar mass (g/mol)
37
Percent composition formula
Percent composition = mass of element in formula/ molar mass x 100%
38
Combination reaction
has two or more reactants forming one produce (*they have more reactants than products)
A + B -> C
39
Decomposition reactions
a single reactant breaks down into two or more products, usually the result of heating , high-frequency radiation, or electrolysis (*have more products than reactants)
A -> B + C
40
Combustion reaction
a special type of reaction that involves a fuel- usually a hydrocarbon- and an oxidant (normally oxygen), commonly these reactants form the two products of carbon dioxide and water (*involves oxidation of a fuel)
41
Percent yield
Percent yield = actual yield/ theoretical yield x 100%
42
Rate-determining step
the slowest step in any proposed mechanism
-the rate of the whole reaction is only as fast as the rate-determining step
43
The collision theory of chemical kinetics
the rate of a reaction is proportional to the number of collisions per second between the reacting molecules
-suggests that not all collisions result in a chem rxn
-an effective collision occurs only if the molecules collide with each other in the correct orientation and with sufficient energy to break their existing bonds and form new ones
44
Collision theory equations
- rate = Z (# of collisions) x f (fraction of collisions that are effective)
- Arrhenius equation: k = Ae x (-Ea/RT) where k is the rate constant, A is the frequency factor, Ea is the activation energy, R is the ideal gas constant, and T is the temp
45
Free energy change of the reaction (ΔGrxn)
the difference between the green energy of the products and the free energy of the reactants
+ΔG = exergonic = energy absorbed
-ΔG = endergonic = energy given off
46
rate law
rate = k[A] ^x [B]^y
--the values of k, x, and y must be determined experimentally for a given rxn at a given temp
47
Zero-order reaction
one in which the rate of formation of product C is independent of changes in concentrations of any of the reactants
- rate = k[A]^0[B]^0 = k, where k has the units M/s
-it is possible to change the rate by changing the temp or by the addition of a catalyst which lowers the activation energy thereby increasing the value of K
-results in a linear graph
48
First-order reaction
has a rate that is directly proportional to only one reactant, such that doubling the concentration of that reactant results in a doubling of the rate of formation of the product
- rate = k[A]^1 or rate = k[B]^1, where k has the units s^-1
-ex: process of radioactive decay
-results in a nonlinear graph
49
Second-order reaction
has a rate that is proportional to either the concentrations of two reactants or to the square of the concentration of a single reactant
- rate = k[A]^1[B]^1 or rate = K[A]^2 or rate = k[B]^2, where k has the units of M^-1x^-1
-suggest a physical collision between two reactant molecules
-results in a nonlinear graph
50
Reaction quotient
QKeq: ΔG>0, reaction proceeds in the reverse direction
51
Le Chatelier's Principle
states that if a stress is applied to a system, the system shifts to relieve that applied stress
-change in concentration: the system will always react in the direction away from the added species or toward the removed species
-change in pressure (and volume): if the pressure is increased the system will react in the direction that produces fewer moles of gas, and if the pressure is decreased the system will react in the direction that produces more moles of gas
-change in temp: results in a change in Keq, if a rxn is endothermic heat functions as a reactant, if a rxn is exothermic heat functions as a product (the equilibrium shifts in the direction that consumes energy)
52
First law of thermodynamics
ΔU (change in internal energy) = Q (heat added to the system) - W (work done by the system)
53
Isothermal process
the system's temperature is constant
-implies that the total internal energy of the system (U) is constant, ΔU= 0 and the first law simplifies to Q = W
54
Adibatic process
no heat is exchanged between the system and the environment, the thermal energy of the system is constant throughout the process
-when Q = 0, the first law of thermodynamics simplifies to ΔU = -W (the work done on the system)
55
Isobaric process
the pressure of the system is constant
-do not alter the first law of thermodynamics
56
Isovolumetric (isochoric) process
experience no change in volume
-bc the gas neither expands nor compresses, no work is performed in the process
-the first law of thermodynamics simplifies to ΔU = Q
57
Standard condition
for measuring enthalpy, entropy, and Gibbs free energy changes of a reaction
--- 25 degrees celsius (298 K), 1 atm pressure, and 1 M concentrations
-used for kinetics, equilibrium, and thermodynamics problems
58
Standard temperature and pressure (STP)
the temperature is 0 degrees celsius (273 K) and pressure is 1 atm
-used for ideal gas calculations
59
Fusion or melting
the transition from solid to liquid
60
Solidification, crystallization, or freezing
the transition from liquid to solid
61
Sublimation
solid goes directly into the gas phase
62
Deposition
from the gaseous to the solid phase
63
Triple point
the point at which the three phase boundaries (solid-liquid, liquid-gas, and solid-gas) meet
64
Critical point
where the phase boundary between the liquid and gas phases terminates (there is no distinction between the two phases)
65
Endothermic
processes in which the system absorbs heat (ΔQ >0)
66
Exothermic
process in which the system releases heat (ΔQ <0)
67
The equation for heat transfer
q = mcΔT (m= mass, c= specific heat of the substance)
68
Enthalpy
heat changes at constant pressure
69
Hess's law
enthalpy changes of reactions are additive
-applies to any state function, including entropy and Gibbs free energy
70
The second law of thermodynamics
states that energy spontaneously disperses from being localized to becoming spread out if it is not hindered from doing so
71
Entropy
is a measure of the spontaneous dispersal of energy at a specific temperature: how much energy is spread out, or how widely spread out energy becomes in a process
ΔS = Q rev (heat that is gained or lost in a reversible process) /T (temp in kelvins), units are J/ mol x K
72
The change in Gibbs free energy
the maximum amount of energy released by a process-- occurring at constant temperature and pressure--that is available to do work
ΔG = ΔΗ - ΤΔS
73
Exergonic
movement toward the equilibrium position, when a system releases energy (ΔG <0)
74
Endergonic
movement away from the equilibrium position, (ΔG>0)
75
ΔG
-if negative, the reaction in spontaneous
-if positive, the reaction is non spontaneous
-if zero, the system is in a state of equilibrium; ΔH = TΔS
76
Variables of gaseous samples
-pressure, volume, temperature, and number of moles (n)
1 atm = 760 mmHg = 760 torr = 101.325 kPa
77
Ideal gas
represents a hypothetical gas with molecules that have no intermolecular forces and occupy no volume
78
The ideal gas law
shows the relationship among four variables that define a sample of gas:
PV = nRT
-R represents the ideal gas constant, which has a value of 8.21 x 10^-2 L x atm / mol x K, may also be given as 8.314 J/ K x mol which is derived when the SI units of pascal
79
Combined gas law
can be used to relate changes in temperature, volume, and pressure of a gas:
P1V1/T1 = P2V2/T2
80
Avogadro's principle
states that all gases at a constant temperature and pressure occupy volumes that are directly proportional to the number of moles of gas present
n/V = k or n1/V1 = n2/V2
81
Boyle's law
for a given gaseous sample held at constant temperature (isothermal conditions), the volume of the gas is inversely proportional to its pressure:
PV = k or P1V1 = P2V2
82
Charles's Law
at constant pressure, the volume of a gas is proportional to its absolute temperature, expressed in kelvin
V/T = k or V1/T1 = V2/T2
83
Gay-Lussac's law
is complementary to Charles's law, it utilizes the same derivation from the ideal gas, but it relates pressure to temperature instead
P/T = k or P1/T1 = P2/T2
84
Dalton's law of partial pressures
states that the total pressure of a gaseous mixture is equal to the sum of the partial pressures of the individuals
PT= PA + PB + PC +...
-The partial pressure of a gas is related to its mole fraction: PA = XAPT, XA = moles of gas A/ total moles of gas
85
Solvation
the electrostatic interaction between solute and solvent molecules
-exothermic = new interactions are stronger than original ones
-endothermic= new interactions are weaker than original ones
86
General solubility rules in aqueous solutions
1. all salts containing ammonium and alkali metal cations are water-soluble
2. all salts containing nitrate and acetate anions are water-soluble
3. Halides excluding fluorides, are water-soluble, with the exceptions of those formed with Ag+, Pb2+, and Hg 2+
4. All salts of the sulfate ion are water-soluble, with the exceptions of those formed with Ca2+, Sr2+, and Ba2+
5. all metal oxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and CaO, SrO, and BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides
6. All hydroxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and Ca2+, Sr2+, and Ba 2+
7. All carbonates, phosphates, sulfides, and sulfites are insoluble with the exception of those formed with the alkali metals and ammonium
87
Percent composition by mass
mass of solute/ mass of solution x 100%
88
mole fraction (X)
XA = moles of A/total moles of all species
*the sum of the moles fractions in a system will always equal 1
89
molarity (M)
moles of solute/liters of solution
90
colligative properties
physical properties of solutions that are dependent on the concentration of dissolved particles but not on the chemical identity of the dissolved particles
-include: vapor pressure depression, boiling point elevation, freezing point depression, and osmotic pressure
91
Raoult's law
accounts for vapor pressure depression caused by solutes in solution
-as solute is added to a solvent, the vapor pressure of the solvent decreases proportionately
-as more solute is dissolved into solvent, the vapor pressure of the solvent decreases
PA = XAPA (PA = vapor pressure of solvent A when solutes are present, XA= the mole fraction the solvent A in the solution, PA is the vapor pressure of solvent A in its pure state)
92
The extent to which the boiling point of a solution is raised relative to that of the pure solvent is given by the formula
ΔTb = iKbm (where Tb is the increase in boiling point, I is the Van't Hoff factor, Kb is a proportionality constant characteristic of a particular solvent, and m is the molality of the solution)
93
The formula for calculating the freezing point depression for a solution is
ΔTf = iKfm (Tf is the freezing point depression)
94
Osmotic pressure
a sucking pressure generated by solutions in which water is drawn into a solution, it is the amount of pressure that must be applied to counteract this attraction of water molecules for the solution
- π = iMRT
95
Arrhenius acids & bases
-acid: will dissociate to form an excess of H+ in solution
-base: will dissociate to form an excess of OH- in solution
---*these behaviors are generally limited to aqueous acids and bases
---acids contain H at the beginning of their formula, bases contain OH at the end of their formula
96
Bronsted-Lowry acids & bases
-acid: a species that donates hydrogen ions (H+)
-base: a species that accepts hydrogen ions
-not limited to aqueous solutions
97
Lewis acids & bases
-acid: an electron pair acceptor
-base: an election pair donor
98
pH
pH = -log[H+] = log 1/[H+]
99
pOH
-log[OH-] = log 1/[OH-]
100
Henderson-Hasselbach equation
used to estimate the pH or pOH of a buffer solution
pH = pka + log [A-]/[HA]
101
dehydrogenases
enzymes that catalyze oxidations in which electrons are lost in the form of hydrogen (dehydrogenation)
102
Oxidizing agents
*almost all contain oxygen or another strongly electronegative element (such as a halogen)
- 02, H202, F2, Cl2, Br2, I2, H2SO4, HNO3, NaClO, KMn04, Cr03, Na2Cr2O7, pyridinium chlorochromate (PCC), NAD+, FADH
103
Reducing agents
often contain metal ions or hydrides (H+)
-CO, C, B2H6, Sn2+ and other pure metals, Hydrazine, Zn(Hg), Lindlar's catalyst, NABH4, LiAlH4, NADH, FADH2
104
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