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General Chemistry Flashcards

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1
Q

Rutherford Model

A
  1. electrons surround nucleus
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2
Q

Bohr Model

A 
1913. describes orbit in much more detail:
AHED:
- absorb light 
- higher potential
- excited 
- distant from nucleus

photons emitted when n is lower and photons absorbed when n is higher; further orbitals, the more energy

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3
Q

Heinsberg Uncertainty

A

it is impossible to know the momentum and position at the same time

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4
Q

Hund’s Rule

A

electrons only double up in orbitals if all orbitals first have one electron

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5
Q

Pauli’s Exclusion Principle

A

paired electrons must be +/- 1/2

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6
Q

Aufbau Principle

A

if you fill up all the lower energy levels, then you can work your way up

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7
Q

Diamagnetic vs Paramagnetic

A

Dia- all electrons are paired (one up one down), so repelled by an external magnetic force
Para- 1 or more unpaired electrons, so there is a pull from an external magnetic force

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8
Q

Atom Size

A

cation

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9
Q

Electronegativity for Covalent Bonds

A
  • nonpolar covalent bonds = EN1.7
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10
Q

Coordinate

A

a single atom provides both bonding electrons

ex. B12

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11
Q

Covalent bonds are most often found in…

A

Lewis Acid- Base Chemistry

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12
Q

Van der Waals encompasses both…

A

dipole and London dispersion forces

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13
Q

Crystalline Lattice

A

large, organized arrays of ions

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14
Q

Equivalents

A

For ions, it is the number of moles needed to balance the charge of one mole of the oppositely charged monovalent
ex. N^-3 / H+ 1 mol of N^-3 = 3 Eq

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15
Q

Equivalent Mass

A

the mass of an acid that yields 1 mol of H+ or the mass of a base that reacts with 1 mol of H+

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16
Q

Gram Equivalent Weight (GEW)

A

the mass of a substance that can deliver 1 Eq of the species of interest (molar mass)/(mol H+ or electrons)

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17
Q

Combustion Reaction

A

always results to CO2 and H2O

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18
Q

Neutralization Reaction

A

Acid+ Base = H2O + salt

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19
Q

Hydrolysis Reaction

A

uses water to break the bonds in a molecule

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20
Q

Reaction Order and Michaelis- Menten Curve

A

at low substrate concentration, the reaction is in first order and at very high concentrations, the reaction is in zeroth order since the reaction ceases to depend on substrate concentration

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21
Q

Difference Between Electronic Geometry and Molecular Shape

A

Electronic Geometry- arrangement of electron groups

Molecular shape- arrangement of atoms, excluding lone pairs

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22
Q

Stability Constant (Kf)

A

the equilibrium constant for complex formations; usually much higher than Ksp

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23
Q

Common Ion Effect

A

decrease in solubility of a compound in a solution that already contains one of the ions in the compound; the presence of the ion shifts the dissolution reaction the left; decreasing its dissociation (compound and solution share similar ion)

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24
Q

Chelation

A

when a central cation is bonded to the same ligand in multiple places; chelation therapy sequesters toxic metals from body

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25
Colligative Properties
physical properties of solutions that depend on concentration of dissolved particles, but not on their chemical properties nor identities
26
Raoult's Law
vapor pressure depression; the presence of other solutes decrease evaporation rate of solvent, causing vapor pressure to decrease Pa= XaPa deg
27
Boiling Point Elevation
delta Tb= -iKbm Kb- .51 C/m m - molality of solution i- Van't Hoff factor
28
Freezing Point Depression
delta Tb= iKfm Kf- -1.86 C/m m - molality of solution i- Van't Hoff factor
29
Osmolarity
the number of individual particles in solution | ex. 1 M of NaCl = 2 osmol/ L
30
Osmotic Pressure
sucking pressure generated by solutions in which water is drawn into solution pi= iMRT. (i= Van't Hoff factor)
31
Arrhenius and Bronsted Acid
produces or donates H+
32
Arrhenius and Bronsted Base
produces OH- or accepts H+
33
Lewis Acid
accepts electron pair
34
Lewis Base
donates electron pair
35
Amphoteric Species
species have both acid and base; specifically behaves as a Bronsted Acid/ Base
36
Polyprotic Acid
an acid with multiple ionized H atoms
37
Buffers
weak acid/weak base + conjugating salt, which differs by 1 H+
38
Buffering Capacity
the ability of a buffer to resist change in pH. Maximum buffering capacity is within 1 pH point of the pKA, and when [A-]=[HA] at the half equivalence point, then pH=pKa - at the half equivalence point, have of the titrant has been protonated or deprotonated
39
Polyvalent
can donate or accept multiple equivalents
40
Equivalence Point
the point at which the equivalent amounts of acid and base have reacted N1V1=N2V2
41
pH Based on Strength of Acids and Base for Titrations
- Strong acid and base pH=7 - Weak acid and strong base pH>7 - Weak base and strong acid pH<7 - Weak acid and weak base depends on strength of both so can be either > or < than pH of 7
42
Kinetic Products
HIGHER in free energy than thermodynamic products and can form at lower temperatures."Fast” products because they can form more quickly under such conditions.
43
Thermodynamic Products
LOWER in free energy than kinetic products, more stable. Slower but more spontaneous (more negative delta G)
44
pH in comparison to respiration of CO2
- When pH decreases, there is an increase in respiration to take out the CO2 - When pH increases, the respiration levels decrease to trap CO2
45
Isolated System
Exchange neither matter nor energy with the environment.
46
Closed System
Can exchange energy but not matter with the environment.
47
Open System
Can exchange BOTH energy and matter with the environment.
48
Isothermal Process
Constant temperature
49
Adiabatic Process
Exchange no heat with the environment
50
Isobaric Process
Constant pressure
51
Isovolumetric or Isochoric
Constant volume
52
Standard Conditions
- 298 K, 1 atm, 1 M - Note that in gas law calculations, Standard Temperature and Pressure (STP) is 0°C, 1 atm. - 1 mol of gas= 22.4 L - Units: 1 atm = 760 mmHg = 760 torr = 101.3 kPa = 14.7 psi
53
Fusion
Solid to Liquid
54
Sublimation
Solid to Gas
55
Deposition
Gas to Solid
56
Supercritical Fluid
Density of gas = density of liquid, no distinction between those two phases.
57
Temperature
Scaled measure of average kinetic energy of a substance.
58
Room Temp
25°C = 75°F
59
Body Temp
37°C = 98.6°F
60
Enthalpy (H)
A measure of the potential energy of a system found in intermolecular attractions and chemical bonds.
61
Phase changes
S-L-G is endergonic because Gas has the most heat energy, so progressively stores more heat G-L-S is exergonic since these reactions release heat
62
Hess's Law
Enthalpy changes are additive - Enthalpy from heat of formations is product minus reactants - Enthalpy from bond dissociation energies is reactants- products
63
Ideal Gas
Theoretical gas whose molecules occupy negligible space and whose collisions are perfectly elastic. Gases behave ideally under reasonably ­increased temperatures and decreased pressures.
64
Dalton's Law
Total pressure is equal to all the partial pressures or Partial Pressure= XPtotal X is the mol fraction
65
Henry's Law
C=kP C - concentration of dissolved gas k- constant P- partial pressure of gas
66
Effusion
The movement of gas from one compartment to another | through a small opening under pressure
67
Solvent particles surround solute particles via what kind of interactions?
Electrostatic
68
Solvation or Dissolution
The process of dissolving a solute in solvent. Most | dissolutions are endothermic, although dissolution of gas into liquid is exothermic.
69
Molar Solubility
Molarity of the solute at saturation
70
Complex ions
Cation bonded to at least one ligand which is the e- | pair donor. It is held together with coordinate covalent bonds. Formation of complex ions increase solubility.
71
Indicators
Weak acids or bases that display different colors in the | protonated and deprotonated forms. The indicator’s pKa should be close the pH of the equivalence point.
72
Ways to Test pH
- Litmus: Acid = red; Base = blue; Neutral = purple - Phenolphthalein: pH < 8.2 = colorless; pH > 8.2 = purple - Methyl Orange: pH < 3.1 = red; pH > 4.4 = yellow - Bromophenol Blue: pH < 6 = yellow; pH > 8 = blue
73
Endpoint
When indicator reaches full color.
74
Oxidation and Reduction regarding Oxygen Transfer
Oxidation is gain of oxygen and reduction is loss of oxygen
75
Disproportionation Reactions | dismutation
A type of REDOX reaction in which one element is both oxidized and reduced, forming at least two molecules containing the element with different oxidation states 2H2O2 (-1) = 2H2O (-2) + O2 (0)
76
Oxidation Rules
- any free element or diatomic species = 0 - When in compounds, group 7A elements = -1, unless combined with an element of greater EN - H = +1 unless it is paired with a less EN element, then = -1 - O = -2 except in peroxides, when it = -1, or in compounds with more EN elements - The sum of all oxidation numbers in a compound must = overall charge
77
Redox Titration
Similar in methodology to acid-base titrations, however, | these titrations follow transfer of charge
78
Potentiometric Titration
A form of REDOX titration in which a voltmeter measures the electromotive force of a solution. No indicator is used, and the equivalence point is determined by a sharp change in voltage
79
Anode vs Cathode
Anode is cite for oxidation and Cathode is always site of reduction Electrons flow from anode to cathode Current flows from cathode to anode
80
Galvanic Cells (Voltaic)
- A spontaneous redox reaction (transfer of e-) is used to produce electric current House spontaneous reactions. - delta G, +Emf, +E cell Anode = NEG, Cathode = POS
81
Electrolytic Cells
- an electric current is used to drive a non spontaneous redox reaction House non spontaneous reactions. + delta G, -Emf, -E cell Anode = POS, Cathode = NEG
82
Concentration Cells
Specialized form of galvanic cell in which both electrodes are made of the same material. It is the concentration gradient between the two solutions that causes movement of charge.
83
Rechargeable Batteries
Can experience charging (electrolytic) and discharging (galvanic) states.
84
Lead-Acid
Discharging: Pb anode, PbO2 cathode in a concentrated | sulfuric acid solution. Low energy density.
85
Ni-Cd
Discharging: Cd anode, NiO(OH) cathode in a concentrated KOH solution. Higher energy density than lead-acid batteries.
86
NiMH
More common than Ni-Cd because they have higher energy density.
87
Reduction Potential
Quantifies the tendency for a species to gain e- and be reduced. More positive Ered = greater tendency to be reduced.
88
Standard Reduction Potential
``` E red= Calculated by comparison to the standard hydrogen electrode (SHE). ```
89
Standard Electromotive Force (Emf)
``` E cell= The difference in standard reduction potential between the two half-cells E red (cathode minus the E red (anode) ```

MCAT (4 decks)

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