General and Organic Chemistry Flashcards
1
Q
Atoms
A
tiny particles that makes up all mass
2
Q
nucleus
A
consists of protons and neutrons which are collectively called nucleons
3
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Strong nuclear force
A
force that holds protons and neutrons together
4
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Binding energy
A
the energy that would be required to break the nucleus into individual protons and neutrons; a measure of the stability of the nucleus
5
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Proton
A
positively charged, same mass as neutron but much greater than electron
6
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Neutron
A
neutrally charged, same mass as proton but much larger than electron
7
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Electron
A
negatively charged, much smaller than proton and neutron
8
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Element
A
building blocks of compounds and cannot be decomposed into simpler substances by chemical means
9
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Atomic number
A
indicates the number of protons and provides the identity of the element
10
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Mass number
A
number of protons plus neutrons
11
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Atomic weight
A
molar mass; units of amu (atomic mass unit or g/mol)
12
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Isotopes
A
two or more atoms of the same element that contain different number of neutrons
13
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Ion
A
when then number of electrons in an atom does not match the number of protons and so the atom carries a charge
14
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Cation
A
positively charged ion, smaller than neutral compound
15
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Anion
A
negatively charged ion, larger than neutral compound
16
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Periodic table
A
lists the elements from left to right in order of their atomic number
17
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Period
A
horizontal row of periodic table
18
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Group/family
A
vertical column of periodic table; elements in the same group share some similar chemical and physical properties
19
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Metals
A
large atoms that tend to lose electrons to form positive ions and positive oxidation states; have a loose hold on their outer electrons which causes their characteristics: lustrous, ductile, malleable, and thermally and electrically conductive
20
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Nonmetals
A
form covalent bonds with one another and generally speaking have lower melting points than metals; tend to form anions
21
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Metalloids
A
have some metallic and some non-metallic characteristics
22
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Representative elements
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groups 1-2, 13-18; make ions by forming the closest noble gas electron configuration
23
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Transition metals
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groups 3-12, can form multiple ions with various charges
24
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Elements with half filled or filled orbital
A
groups 1-2: half-filled and completely filled s orbitals; groups 7 and 12: half-filled and completely filled d orbitals; groups 15 and 18: half-filled and completely filled p orbitals `
25
Valence electrons
electrons in the outermost shell; contribute most to an element's chemical properties; in most cases, only electrons from the s and p subshells are considered valence electrons
26
Alkali metals
Group 1 metals; soft metallic solids with low densities and low melting points that easily form 1+ cations; highly reactive with nonmetals to form ionic compounds and react with H to form hydrides; react exothermically with H20 to produce metal hydroxide and H2 gas
27
Alkaline earth metals
Group 2 metals; harder, more dense, and melt at higher temperatures than alkali metals; form 2+ cations and are less reactive than alkali metals because their highest energy electron completes the s orbital
28
Oxygen group/chalcogens
know oxygen and sulfur; oxygen is 2nd most electronegative element, it is divalent and can form strong pi bonds to make double bonds; typically reacts with metals to form metal oxides; akali metals form peroxides and super oxides; sulfur can form many more bonds than oxygen because it has access to the 3d orbital; S also has ability to form double bonds
29
Halogens
F, Cl, Br, I; highly reactive; like to gain electrons; react with metals to form ionic halides; can combine with hydrogen to form gaseous hydrogen halides which are soluble in water, forming hydrohalic acids
30
Noble gases
inert gases that nonreactive
31
Atomic radius
distance from the center of the nucleus to the outermost electron
32
Atomic radius trend
increases moving down a group and across a period from right to left
33
Electrostatic force
the force between charged objects following coulomb's law F = kq1q2/r^2
34
Effective nuclear Charge (Zeff)
the amount of charge felt by the most recently added electron
35
Zeff trend
increases going left to right and from top to bottom
36
Ionization energy
the energy needed to detach an electron from an atom
37
First ionization energy
the energy necessary to remove an electron from a neutral atom in its gaseous state to form a +1 cation
38
Second ionization energy
the energy required for the removal of a second electron from the same atom to form a +2 cation; always greater than the first ionization energy
39
Ionization energy trend
increases along the periodic table from left to right and from bottom to top
40
Electronegativity
the tendency of an atom to attract electrons shared in a covalent bond
41
Electronegativity trend
increases from left to right and from bottom to top
42
Electron affinity
the willingness of an atom to accept an additional electron
43
Electron affinity trend
increases from left to right and from bottom to top
44
Quantum mechanics
elementary particles can only gain or lose energy in discrete units
45
Bohr atom
represents the atom as a nucleus surrounded by electrons in discrete electron shells
46
orbital structure of the H atom
a single e- orbits the hydrogen's nucleus in an electron shell
47
Pauli Exclusion Principle
no 2 electrons in the same atom can have the same 4 quantum numbers
48
Principal quantum number, n
the first quantum number which designates the shell level of the electron, with low numbers closest to the nucleus
49
Subshell, L
shape of orbital; L= 0 is s subshell, L =1 is p subshell etc.; range from 0 to n-1
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Magnetic quantum number, Ml
specific orbital within a subshell, value from -L to L
51
Electron spin quantum number, Ms
distinguishes between 2 electrons in the same orbital; one is spin +1/2 and the other is spin -1/2
52
Heisenberg Uncertainty Principle
there is an inherent uncertainty in the product of the position of a particle and its momentum; xp >= h/2
53
Aufbau principle
with each new proton added to create a new element, the new electron is added to maintain neutrality will occupy the lowest energy level available
54
Exceptions for electron configuration
Elements in Group 6 (Cr) and Group 11 (Cu), borrow 1 electron from the highest s subshell so they end up with a half-filled s subshell and a half-filled or filled d subshell
55
Electron configuration
lists the shell and the subshells of an element's electrons in order typically from lowest to highest energy level
56
Ground state
lowest energy level
57
Excited state
when at least one e- has moved from a lower energy subshell to a higher energy subshell
58
Hund's rule
electrons will not fill any orbital in the same subshell until all orbitals in that subshell contain at least one electron and that the unpaired electrons will have parallel spins
59
Paramagnetic elements
elements with unpaired electrons that align with an external magnetic field
60
Diamagnetic elements
elements with no unpaired electrons and are unresponsive to an external magnetic field
61
Emission line spectrum
spectrum that is characteristic of a given element from energy released when excited electrons fall from a higher energy state to a lower energy state
62
Absorption line spectrum
measures the radiation absorbed when electrons absorb energy to move to a higher energy state
63
Photoelectric effect
the emission of electrons when light falls on a material; KE = hf - work function
64
Covalent bond
electrons are shared between atoms
65
Ionic bond
electrons are transferred from one atom to another
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Molecule
atoms held together by only covalent bonds
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Bond length
the distance between the nuclei of 2 atoms when they are at their lowest possible energy state
68
Bond dissociation energy/bond energy
the energy necessary for a complete separation of the bond
69
Partial ionic character
when the difference in electronegativity is significant
70
Dipole moment
occurs when the center of positive charge in a bond does not coincide with the center of negative charge
71
Intermolecular attractions
attractions between separate molecules due to dipole moments
72
Hydrogen bond
the strongest type of dipole-dipole interaction; occurs between a H that is covalently bound to a F, O, or N and a F, O, or N from another molecule
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Induced dipole
occurs when dipole moment is momentarily induced in an otherwise nonpolar molecule or bond by a polar molecule, ion, or electric field
74
Instantaneous dipole moment
arise spontaneously and occur because the electrons move about and at any given moment they may not be distributed exactly between the 2 bonding atoms even when the atoms have equivalent electronegativity
75
London dispersion forces/Van der Waals forces
weakest dipole-dipole force between 2 instantaneous dipoles
76
Empirical formula
the smallest ratio of whole numbers that can be used to represent these proportions
77
Molecular formula
represents the exact number of elemental atoms in each molecule
78
Physical reaction
when a compound undergoes a reaction and maintains its molecular structure ie. melting, evaporation, dissolution, and rotation of polarized light
79
Chemical reaction
when a compound undergoes a reaction and changes its bonding or structure to form a new compound ie. combustion, metathesis, and redox
80
Combination Reaction
A + B --> C
81
Decomposition reaction
C --> A + B
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Single Displacement/Replacement
A + BC --> B + AC
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Double Displacement/Replacement or Metathesis
AB + CD --> AD + BC; often occurs between ionized salts dissolved in water
84
runs to completion
means that the reaction generates products until the supply of at least one reactant is fully depleted
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Limiting reactant
the first reactant to be used up first
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Theoretical yield
the amount of product that should be created when a reaction runs to completion based on the stiochiometry
87
Percent yield
actual yield/theoretical yield x 100%
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Actual yield
the amount of product created by a real experiment
89
mole
6.022 x 10^23 of something; grams/mw
90
Avogadro's number
6.022 x 10^23
91
Radioactive decay
atoms that spontaneously break apart
92
Half-life
the length of time necessary for one half of a given substance to decay
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Type of decay of radioactive decay
exponential decay
94
Semi-log plot
plotting the logarithm of amount of atoms as a function of time would produce a straight line for something that exponentially decays
95
4 variables of a half-life
1. initial amount of substance 2. final amount of substance 3. length of the half life 4. the number of half lives (often give as a time period in which you divide by the length of the half-life)
96
Alpha decay
the loss of an alpha particle/helium nucleus
97
Beta decay
the breakdown of a neutron into a proton and electron and the expulsion of the newly created electron
98
Neutrino
also emitted during beta decay; virtually a massless particle
99
Positron emission
the emission of a positron when a proton becomes a neutron; type of beta decay; neutrino also emitted
100
Electron capture
the capture of an electron and the merging of that electron with a proton to create a neutron
101
Gamma ray
a high frequency photon that has no mass or charge and does not change the identity of the atom from which it is given off
102
Gamma decay
often accompanies the other types of radioactive decay; can occur when an electron and positron collide
103
Valence
the number of bonds an atom usually forms
104
Formal charge
the number of valence electrons of an atom, minus the number of bonds it is a part of, minus the number of nonbonding electrons it has
105
Dash formula
shows the bonds between each atom of a molecule but does not show lone pairs nor the three dimensional structure of the molecule
106
Condensed formula
shows neither the bonds nor the 3D structure
107
Bond-line formula
line intersections, corners and endings represent a carbon atom unless another atom is drawn in; hydrogens are not drawn
108
Fischer projection
vertical lines are assumed to be oriented into the page and horizontal lines are assumed to be oriented out of the page
109
Newman projection
a view straight down the axis of one of the sigma bonds which gives information about steric hindrance with respect to a particular sigma bond
110
Dash-line wedge formula
the solid black wedges represent bonds coming out of the page and the dashed wedges represent bonds going into the page
111
Space-filling model
a 3D representation of a molecule with spheres of various colors representing different elements with respect to their relative sizes
112
Ball and sticks model
bond lengths are drawn to approximately twice their length so that the atoms are clearly visible; give information about relative size of atoms
113
Sigma bond
forms when the bonding pair of electrons are localized to the space directly between the two bonding atoms; lowest energy, strongest and most stable covalent bond
114
Pi bind
created by overlapping p orbitals; double and triple bonds are made by adding pi bonds to a sigma bond
115
How to determine number and type of hybrid orbital
count the number of sigma bonds and lone pairs of electrons on that atom, match this number to the sum of the superscripts in a hybrid name
116
Valence shell electron pair repulsion (VSEPR)
the electrons in an orbital seek to minimize their energy by moving as far away from other electron pairs as possible, minimizing the repulsive forces between them
117
Shape/ bond angle of sp
linear, 180
118
Shape/ bond angle of sp2
trigonal planar, 120
119
Shape/ bond angle of sp3
tetrahedral, trigonal pyramidal or bent, 109.5
120
Shape/ bond angle of sp3d
trigonal-bipyramidal, see-saw, t-shaped, or linear; 90, 120
121
Shape/ bond angle of sp3d2
octahedral, square pyramidal, or square planar, 90
122
Delocalized electrons
bonding electrons which are spread out over 3 or more atoms
123
Resonance structures
molecules containing delocalized electrons can be represented by a combination of 2 or more alternative Lewis structures called resonance structures; weighted average represents the actual molecule
124
Aromaticity
the increased stability of a cyclic molecule due to electron delocalization; must follow Huckel's rule
125
Huckel's rule
planar monocyclic rings whose number of pi electrons can be described with the equation 4n + 2 will be aromatic (lone pairs count as pi electrons)
126
Nucelophilic functional groups
have a partial negative charge and seek positively charged nuclei ; they donate electrons and attack functional groups with partial positive charges ie. amines
127
Electrophilic functional groups
have a partial positive charge and seek electrons; usually get attacked by electrons from other functional groups
128
Stereochemistry
three-dimensional structure of a molecule
129
Isomers
molecules that share the same molecular formula
130
Structural isomer
have the same molecular formula but different bond-to-bond connectivity and thus different chemical properties due to differences in functional groups
131
Conformational isomers/conformers
different spatial orientations of the same molecule
132
Stereoisomers
2 unique molecules with the same molecular formula and the same bond-to-bond connectivity
133
Enantiomers
non-superimposable mirror images of one another; same molecular formula and connectivity but are not the same molecule because they differ in their configuration; have the same chemical and physical characteristics except for interactions with other chiral molecules and interactions with polarized light
134
Chiral molecule
carbon bonded to four different substituents that rotate plane-polarized light
135
Determine absolute configuration
atoms attached to chiral center are number from highest to lowest priority; highest priority is given to atom with the largest atomic weight,turn molecule so that lowest priority molecule is into the page, draw a circle from highest to lowest priority; clockwise = R, counterclockwise = S
136
Relative configuration
two molecules ave the same relative configuration about a chiral carbon if they differ by only one substituent and the other substituents are oriented identically about the carbon; still may have different absolute configurations however
137
Determine absolute configuration for Fischer projection
assign priorities, ignore lowest, draw circle from highest to lowest, if 4 is into the page (vertical), keep configuration, if 4 is out of the page reverse the configuration
138
Plane-polarized light
photons in which their electric fields are oriented in the same direction
139
Specific rotation
a standardized form of observed rotation that is calculated from the observed rotation and experimental parameters
140
Dextorotatory
rotates plane-polarized light clockwise; d or +
141
Levorotary
rotates plane-polarized light counterclockwise; l or -
142
Observed rotation
the direction and number of degrees that the electromagnetic field is rotated when it passes through a compound
143
Racemic mixture
when enantiomers are present in equal concentration and so it does not rotate plane-polarized light
144
Diastereomers
have same molecular formula and same bond-to-bond connectivity but are not mirror images of each other and are not the same compound; have same absolute configuration at one or more of their chiral centers
145
Maximum number of optically active isomers
2^n where n is the number of chiral centers
146
Meso compound
compound that has multiple chiral centers but is optically inactive as they have a plane of symmetry through their center causing the chiral centers to offset each other
147
Epimers
diastereomers that differ in configuration at only one chiral carbon
148
Anomers
cyclic diastereomers that are formed when a ring closure occurs at an epimeric carbon
149
Anomeric carbon
chiral carbon of an anomer
150
Cis-isomer
molecules with substituents on same side of double bond or ring structure; have dipole moments causing a higher boiling points but lower melting points since they cannot form crystals
151
Trans-isomer
molecules with substituents on opposite sides of double bond or ring structure
152
Steric hindrance
when substituents in the cis position crowd each other which raises their energy levels, decreasing stability
153
E Diastereomer
when the higher priority substituents are on opposite sides of the locked bond
154
Z Diastereomer
when the higher priority substituents are on the same side of the locked bond
155
Substitution reactions
one group leaves and is replaced with another
156
SN1 reaction
has a rate that is dependent on only the substrate/electrophile, leaving group leaves spontaneously to form a carbocation; carboskeleton rearrangement can occur
157
SN2 reaction
occurs in a single step, rate dependent on concentration of nucleophile and substrate, inversion of configuration
158
Substrate of SN1/SN2
SN1: secondary or tertiary; SN2: don't occur with sterically hindered substrate, requires methyl, primary, or secondary substrate
159
Solvent of SN1/SN2
highly polar solvent (polar protic) increases reaction rate of SN1 by stabilizing the carbocation but slows down sped of SN2 by stabilizing the nucleophile; SN2: polar aprotic solvent
160
Speed of SN1/SN2
SN1: concentration of substrate; SN2: concentration of substrate and nucleophile
161
Stereochemistry of SN1/SN2
SN1: creates both enantiomers; SN2: inverts sterepchemistry about the chiral center
162
Carbon skeleton rearrangment SN1/SN2
SN1: may occur; SN2: never occurs
163
Leaving group
a good leaving group is defined by its stability or solubility in a solution: if it is stable, it will not return to attack the molecule
164
Alcohols
consist of an oxygen bound to a hydrogen and an R group and undergo three major types of reactions: nucleophile, acid, leaving group
165
Trend of alcohol acidity
from strongest to weakest acid: methyl > primary > secondary > tertiary; reverse for nucleophilicity trend
166
Formation of sulfonates
nucleophlic substitution where alcohol acts as the nucleophile
167
Tosylates/mesylates
commonly used sulfonates that are widely used in the protection of alcohols: the conversion to a sulfonate prevents the alcohol from acting as an acid or nucleophile, or from undergoing undesirable reactions
168
Ethers
relatively unreactive and are a comment solvent for organic reactions since organic compounds tend to be more soluble in ethers than alcohols since no hydrogen bonds need to be broken
169
Cleavage of an ether with strong acid
R2O + HBr --> ROH + RBr
170
Nitrogen
acts as a nucleophile where the lone pair of electrons attacks a positive charge or it can take on a fourth bond and become positively charged
171
Electron donating groups
increase the basicity/nucleophilicity of the molecule
172
Electron withdrawing groups
increase the acidity of the molecule (decrease its nucleophilicity)
173
Properties that make carbonyls good electrophiles
planar stereochemistry and partial positive charge on carbon
174
Planar stereochemistry
leaves open space above and below a molecule reducing steric hindrance and making it more receptive to a nucleophilic attack
175
Racemix mixture
equal amounts of R and S
176
Stereoselective
a reaction prefers either R or S
177
Stereospecific
when only the R or S is formed
178
Trend of reactive of carbonyls
electron donating groups reduce carbonyl reacitivty to a nucleophile attack while electron withdrawing groups increase this reactivity
179
Carboxylic acid derivatives
the less reactive derivatives can be synthesized from the more reactive derivatives but the more reactive derivatives cannot be synthesized from the less reactive ones
180
Hydrolysis of amides
only possible under extreme chemical conditions that are unlikely to occur in biological systems
181
Anhydrides
group of molecules in which the leaving group is a carboxylate ion
182
Esters
formed from an alcohol undergoing nucleophilic substitution with a carboxylic acid (esterification)
183
Transesterification
alcohols react with esters in a reaction called transesterification where one alkoxy group is substituted for another in the ester
184
Amides
synthesized when an amine, acting as a nucleophile, substitutes at the carbonyl of a carboxylic acid or one of its derivatives; found in peptide bounds between amino acids
185
Lactams
cyclic amides which are unstable
186
Tautomerization
shift from a carbonyl to an alkene with an alcohol; is a reaction at equilibrium and NOT a resonance
187
alpha carbon
any carbon attached directly to a carbonyl carbon
188
Kinetic enolate
the less substituted enolate; it is formed quickly by the removal of one of three possible by the removal of an alpha hydrogen
189
Thermodynamic enolate
the more substituted double bond; it is more difficult to form because it has a higher activation energy but it is the more thermodynamically stable product
190
Hemiacetals/hemiketals
aldehydes and ketones react with alcohols to form hemiacetals/hemiketals; have one bond to an OH group and one bond to an OR group
191
acetal/ketal
synthesized when another alcohol molecule reacts with a hemiacetal/ketal where two OR groups are attached to what was previously the carbonyl carbon; good protecting groups ie. monosaccharides
192
Imine
looks like a carbonyl as there is a carbon-nitrogen double bond; formed when aldehyde/ketone reacts with amine
193
Enamine
an alkene with an amine substituent; Enamines are not stable due to the electron donating nitrogen off of the pi bond. Because nitrogens are less electronegative and more electron donating than oxygen, enamines are better nucleophiles than enols; forms when aldehyde/ketone reacts with amine
194
Grignard synthesis of an alchool
organometallic reagent + aldehyde/ketone --> alcohol with R group from organometallic reagent
195
Reduction synthesis of an alcohol
use NaBH4 or LiAlH4; only LiAlH4 is strong enough to fully reduce carboxylic acids and esters and acetates to alcohol
196
Cyanohydrins
a nitrile and alcohol attached to the same carbon; synthesized when nitriles attack carbonyls
197
Oxidation
increase in bonds to oxygen or halogen, loss of C-H bonds
198
Reduction
increase in bonds to H or R groups, loss of bonds to oxygen or halogen
199
Neither oxidation nor reduction
addition or loss of H+, water, HX etc.
200
Decarboxylation
the oxidation of a carboxylic acid to carbon dioxide gas ; the activation energy is lower then the beta carbon is a carbonyl because either the anion intermediate is resonance stabilizied or the acid forms a more stable cyclic intermediate
201
Aldol condensation
a carbonyl nucleophile attacks another carbonyl; the alpha carbon (carbon adjactent to carboynl carbon) acts as a nucleophile to form an aldol (aldehyde and alcohol) and then a condensation reaction turns it into an enal (an aldehyde having a neighboring double bond)
202
Carboyhdrates
carbon chains with an alcohol on each carbon except for one which either has an aldehyde or ketone attached in straight chain form; ring is formed when an alcohol group on a chiral carbon far from the carboynl (in glucose C5) acts as a nucleophile attacking the carboynyl forming a hemiacetal; most naturally occurring ones are D epimers
203
Aldose
when carbohydrate contains an aldehyde
204
Ketose
when carbohydrate contains a ketone
205
Epimers
carbohydrates that have the same structure except for the configuration around a single chiral center
206
Anomeric carbon
formerly the carbonyl carbon of a straight chain carbohydrate and can be identified as the only carbon in the sugar attached to 2 oxygens
207
Anomer
the alcohol group of the anomeric carbon may point upwards or downwards on the ring structure resulting in the alpha or beta anomer (alpha- opposite; beta- same)
208
Furanose
five member ring carbohydrate
209
Pyranose
six member ring carbohydrate
210
Glucopyranose
glucose ring
211
Disaccharides/polysachharides
glycosides where the aglycone is another sugar and are connected by glycosidic linkages
212
Aglycone
the group attached the the anomeric carbon of a glycoside
213
Glycoside
sugar that is formed when a sugar is attacked by an alcohol to create an acetal
214
Sucrose
1,1 glycosidic linkage: glucose and fructose, alpha with respect to glucose and beta with respect to fructose
215
Maltose
alpha 1-4 glycosidic linkage: 2 glucose molecules
216
Lactose
beta 1,4 galactosidic linkage: galactose and glucose
217
Amylose (starch)
alpha 1,4 glycosidic linkage: chain of glucose molecules
218
Amylopectin
alpha 1,4 glycosidic linkages:a branched chain of glucose molecules with alpha 1,6 glucosidic linkages forming the branches
219
Glycogen
alpha 1,4 glycosidic linkage: a branched chain of glucose molecules with alpha 1,6 glucosidic linkages forming the branches
220
Gabriel synthesis
way to form amino acids
221
Strecker synthesis
way to form amino acids
222
Lipogenesis
the formation of lipids
223
Lipolysis
lipid breakdown
224
Saponification
when lipid breakdown is base catalyzed where a soap forms (fatty acid salts and glycerol)
225
Acidity of hydrogens in carbonyl
alpha carbon is most acidic
226
Changing the reactivity of aromatic rings
electron withdrawing groups (meta, NO2, halides, -SO2R, -CO2R) deactivate while electron donating groups (ortho, para, H, alkyl groups, -OH, -OR, -NH2, -NR2)
227
Thermodynamics
study of energy and its relationship to macroscopic properties of chemical systems
228
Extensive property
proportional to the size of the system ie. volume and number of moles
229
Intensive property
independent of the size of the system ie. pressure and temperature
230
Temperature
represents the amount of molecular movement in a substance
231
Thermal energy
the sum of translational, rotational, and vibrational energies
232
Absolute zero
0K or -273 C
233
Reaction rate
how quickly the concentration of the reactants or products are changing over the course of the reaction
234
Elementary reaction
reaction that occurs in a single step
235
Intermediates
species that are products of one step and reactants of a later step in a multistep reaction; often present in low concentrations; not shown in overall equation
236
Zero order reactant
reaction rate does not depend on concentration
237
1st order reactant
reaction rate directly proportional to concentration
238
2nd order reactant
reaction rate is proportional to the square of the concentration
239
Rate-determining step
slowest elementary step; rate law derived from this step
240
Catalyst
substance that increases the rate of reaction without being consumed or permanently altered; increase rate of both forward and reverse reaction; may lower the activation energy or increase the steric factor
241
Steric factor
fraction of collisions having the effective spatial orientation
242
Homogeneous catalyst
is in the same phase as the reactants and prodcts
243
Heterogeneous catalyst
is in a different phase than the reactants or products such as when gas or aqueous particle reacts on a solid
244
State
the physical condition of a system as described by a specific set of thermodynamic properties
245
State functions
properties that describe the current state of a system; do not depend on the path taken
246
Path functions
properties that do not describe the state of a system but rather depend on the pathway used to achieve that state ie. work and heat
247
Internal energy
the collective energy of molecules measured on a microscopic scale; includes vibrational, rotational, translational, electronic, intermolecular and rest mass energies; does not include macroscopic mechanical energies
248
Vibrational energy
created by the back and forth motion of atoms within a molecule
249
Rotational energy
created by the rotation of a molecule around its center of mass
250
Translational energy
created by movement of the center of mass of a molcule
251
Electronic energy
the potential electrical energy created by the attractions between electrons and their nuclei
252
Intermolecular potential energy
created by intermolecular forces between molecular dipoles
253
Rest mass energy
energy described by E =mc^2
254
Heat
spontaneous transfer of energy from a warmer body to a cooler body
255
Zeroth law of thermodynamics
two systems in thermal equilibrium with a third system are in thermal equilibrium with each other; two bodies in thermal equilibrium share a thermodynamic property - temperature, a state function
256
Conduction
thermal energy transfer via molecular collisions; requires direct physical contact
257
Convection
thermal energy transfer via fluid movements; differences in pressure or density drive warm fluid in the direction of cooler fluid
258
Radiation
thermal energy transfer via electromagnetic waves
259
Newton's law of cooling
a body's rate of cooling is proportional to the temperature difference between the body and its environment
260
Emissivity
a constant between 0 and 1 and specific to that object; higher emissivity indicates that a higher amount of radiation energy is absorbed
261
First law of thermodynamics
the total energy of the system and surroundings is always conserved
262
Second law of thermodynamics
the net entropy or disorder of the universe is always increasing
263
Third law of thermodynamics
absolute zero (zero energy) can never be achieved; assigns a zero entropy value to any pure element or compound in its solid form at absolute zero and in internal equilibrium
264
Standard state
the reference form of a substance at any chosen temperature and a pressure of 1 bar `s
265
Standard enthalpy of formation
the change in enthalpy for a reaction that creates one mole of that compound from its raw elements in their standard states
266
Naught
indicates standard state conditions
267
Endothermic reaction
a reaction with a positive enthalpy change
268
Exothermic reaction
a reaction with a negative enthalpy change
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Entropy
nature's tendency to create the most probably arrangement that can occur within a system; state function; extensive property
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Gibbs free energy
extensive property; state function; represents the maximum work available/free for a reaction
271
Endergonic reaction
positive delta G; non-spontaneous
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Exergonic reaction
negative delta G; spontaneous
273
Hess's law of heat summation
the sum of the enthalpy changes for each step is equal to the total enthalpy change regardless of the path chose; also indicates that a forward reaction has exactly the opposite change in enthalpy as the reverse reaction
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Transition state
peak of reaction coordinate where old bonds are breaking and new bonds are forming
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Intermediate location on reaction coordinate
would show multiple peaks (transition states) and the intermediates would be the troughs lying between these peaks
276
Chemical equilibrium
when the forward reaction rate equals the reverse reaction rate and there is no net change in the concentration of the products or reactants
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Law of Mass Action
the mathematical relationship between a chemical equation and the associated equilibrium constant
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Le Chatelier's principle
when a system at equilibrium is stressed, the system will shift in a direction that will reduce that stress ; do not use for inert gases added or solvation
279
Relationship between K and G
if K = 1, G = 0; if K >1, G < 0; if K <1, G > 0
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Ideal gas model/kinetic molecular theory
1. gas molecules have no size; 2. no attractive/repulsive forces 3. completely elastic collisions; 4. average kinetic energy of gas molecules directly proportional to the temperature of the gas
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Simple mercury barometer
a tube of mercury that is closed at one end is inverted and placed in an uncovered mercury bath that is open to the atmosphere; the amount of mercury left in the tube is related to the atmospheric pressure pushing down on the mercury bath
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Boyle's law
pressure and volume inversely related
283
Charles' law
volume of a gas is directly proportional to temperature
284
Avogadro's law
volume of a gas is directly proportional to the number of moles of that gas
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Adiabatic process
no heat transfer
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Isothermal process
no change in internal energy
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Isovolumetric process
no change in volume, therefore no work
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Partial pressure
the total pressure of the gaseous mixture multiplied by the mole fraction of the particular gas
289
Real gases
volume is larger, pressure is smaller (due to intermolecular attractive forces)
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Deviations from ideal gas law
if pv/rt > 1, deviation due to molecular volume; of pv/rt < 1, deviation due to attractive intermolecular forces
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Heat capacity
the added energy required to increase the temperature of a given substance by one K (or degree C)
292
Constant volume heat capacity
system cannot due PV work and thus all input energy must be in the form of heat
293
Constant pressure heat capacity
added energy can be lost by work done by the system by expanding its volume therefore it is typically greater than Cv
294
Calorimeter
measures heat change; holds a liquid with a thermometer placed inside to measure any changes in temperature
295
Coffee cup calorimeter
constant pressure calorimeter that measures energy change at atmospheric pressure (open top); used to measure heats of reactions
296
Bomb calorimeter
measures energy change at constant volume and thus indicates the internal energy change in a reaction; reaction takes place in a rigid container
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Heat of fusion
the enthalpy change associated with melting
298
Heat of vaporization
enthalpy change associated with boiling
299
Sublimation
solid to gas
300
Deposition
gas to solid
301
Condensation
gas to liquid
302
Phase diagram
indicates the phases of a substance at different pressures and temperatures
303
Critical temperature
the temperature above which a substance cannot be liquefied regardless of the pressure applied
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Critical pressure
the pressure required to produce the liquid phase when the substance is at the critical temperature
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Triple point
the only point where a substance can exist in equilibrium between the solid, liquid, and gas phases
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Solution
homogeneous mixture of 2 or more compounds in a single phase
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Solvent
the compound of which there is more of in a solution
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Solute
the compound of which there is less of in a solution
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Solvation
when ionic compounds dissolve, the cations and anions break apart and are surrounded by the oppositely charged ends of the polar solvent
310
Hydration
solvation in water; something that is hydrated is said to be in an aqueous phase
311
Hydration number
the number of water molecules that must surround an ion for hydration to occur; commonly 4-6
312
Electrolyte
a compound that forms ions in aqueous solutions; conduct electricity in solution
313
Parts per million
10^6 multiplied by the ratio of the mass of solute to the total mass of the solution
314
Normality
measures the number of equivalents per liter of solution; ex. a 1 molar solution of sulfuric acid is called a 2 normal solution because it can donate 2 protons for each sulfuric acid molecule
315
Solution formation
1. breaking of intermolecular bonds between solute molecules 2. breaking of intermolecular bonds between solvent molecules 3. formation of intermolecular bonds between solvent and solute molecules; first 2 steps are endothermic, last step is exothermic
316
Vapor pressure
the pressure created by the molecules in the open space at equilibrium
317
Evaporation
a compound evaporates when the vapor pressure of its liquid phase is greater than the partial pressure of its gaseous phase
318
Condensation
a compound condenses when the vapor pressure of its gas phase is greater than the partial pressure of its liquid phase
319
Raoult's law for nonvolatile solutes
if 97% of the solution is solvent, the vapor pressure will be 97% of the vapor pressure of the pure solvent
320
Raoult's law for volatile solutes
if 97% of the solution is solvent, the vapor pressure will be 97% of the vapor pressure of the solvent plus 3% of the vapor pressure of the pure solute
321
Solubility
quantifies a solute's tendency to dissolve in a solvent
322
Saturated solution
the concentration of dissolved salt has reached a maximum
323
Spectator ions
ions that have no effect on a reaction
324
Henry's law
the solubility of a gas is proportional to its vapor partial pressure
325
Redox titration
used to find the molarity of a reducing agent, an oxidizing agent of known molarity is slowly added
326
Half reaction
no half reaction can occur by itself; any reduction half reaction must be accompanied by an oxidation half reaction
327
Electric potential
the ability of a molecule to be reduced, more positive = more readily reduced
328
Galvanic/voltaic cell
generates a current from one phase to another in a conversion of chemical to electrical energy
329
Terminals
conductors such as metal wires
330
Electrodes
conductors
331
Electromotive force (emf)
the potential difference between the terminals when they are not connected
332
Salt bridge
a type of liquid junction that minimizes the potential difference caused by ions moving across a liquid junction
333
Concentration cell
limited form of a galvanic cell in which a reduction half reaction takes place in one half cell while the exact reverse of that half reaction is taking place in the other half cell
334
Electrolytic cell
created by hooking up a power source across the resistance of a galvanic cell and forcing the reactions to run in reverse; will have a negative emf; cathode is marked negative and anode is marked positive
