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MCAT > General and Organic Chemistry > Flashcards

General and Organic Chemistry Flashcards

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1
Q

Atoms

A

tiny particles that makes up all mass

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2
Q

nucleus

A

consists of protons and neutrons which are collectively called nucleons

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3
Q

Strong nuclear force

A

force that holds protons and neutrons together

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4
Q

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

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5
Q

Proton

A

positively charged, same mass as neutron but much greater than electron

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6
Q

Neutron

A

neutrally charged, same mass as proton but much larger than electron

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7
Q

Electron

A

negatively charged, much smaller than proton and neutron

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8
Q

Element

A

building blocks of compounds and cannot be decomposed into simpler substances by chemical means

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9
Q

Atomic number

A

indicates the number of protons and provides the identity of the element

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10
Q

Mass number

A

number of protons plus neutrons

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11
Q

Atomic weight

A

molar mass; units of amu (atomic mass unit or g/mol)

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12
Q

Isotopes

A

two or more atoms of the same element that contain different number of neutrons

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13
Q

Ion

A

when then number of electrons in an atom does not match the number of protons and so the atom carries a charge

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14
Q

Cation

A

positively charged ion, smaller than neutral compound

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15
Q

Anion

A

negatively charged ion, larger than neutral compound

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16
Q

Periodic table

A

lists the elements from left to right in order of their atomic number

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17
Q

Period

A

horizontal row of periodic table

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18
Q

Group/family

A

vertical column of periodic table; elements in the same group share some similar chemical and physical properties

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19
Q

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

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20
Q

Nonmetals

A

form covalent bonds with one another and generally speaking have lower melting points than metals; tend to form anions

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21
Q

Metalloids

A

have some metallic and some non-metallic characteristics

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22
Q

Representative elements

A

groups 1-2, 13-18; make ions by forming the closest noble gas electron configuration

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23
Q

Transition metals

A

groups 3-12, can form multiple ions with various charges

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24
Q

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 `

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25
Q

Valence electrons

A

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

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26
Q

Alkali metals

A

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

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27
Q

Alkaline earth metals

A

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

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28
Q

Oxygen group/chalcogens

A

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

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29
Q

Halogens

A

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

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30
Q

Noble gases

A

inert gases that nonreactive

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31
Q

Atomic radius

A

distance from the center of the nucleus to the outermost electron

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32
Q

Atomic radius trend

A

increases moving down a group and across a period from right to left

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33
Q

Electrostatic force

A

the force between charged objects following coulomb’s law F = kq1q2/r^2

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34
Q

Effective nuclear Charge (Zeff)

A

the amount of charge felt by the most recently added electron

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35
Q

Zeff trend

A

increases going left to right and from top to bottom

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36
Q

Ionization energy

A

the energy needed to detach an electron from an atom

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37
Q

First ionization energy

A

the energy necessary to remove an electron from a neutral atom in its gaseous state to form a +1 cation

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38
Q

Second ionization energy

A

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

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39
Q

Ionization energy trend

A

increases along the periodic table from left to right and from bottom to top

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40
Q

Electronegativity

A

the tendency of an atom to attract electrons shared in a covalent bond

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41
Q

Electronegativity trend

A

increases from left to right and from bottom to top

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42
Q

Electron affinity

A

the willingness of an atom to accept an additional electron

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43
Q

Electron affinity trend

A

increases from left to right and from bottom to top

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44
Q

Quantum mechanics

A

elementary particles can only gain or lose energy in discrete units

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45
Q

Bohr atom

A

represents the atom as a nucleus surrounded by electrons in discrete electron shells

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46
Q

orbital structure of the H atom

A

a single e- orbits the hydrogen’s nucleus in an electron shell

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47
Q

Pauli Exclusion Principle

A

no 2 electrons in the same atom can have the same 4 quantum numbers

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48
Q

Principal quantum number, n

A

the first quantum number which designates the shell level of the electron, with low numbers closest to the nucleus

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49
Q

Subshell, L

A

shape of orbital; L= 0 is s subshell, L =1 is p subshell etc.; range from 0 to n-1

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50
Q

Magnetic quantum number, Ml

A

specific orbital within a subshell, value from -L to L

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51
Q

Electron spin quantum number, Ms

A

distinguishes between 2 electrons in the same orbital; one is spin +1/2 and the other is spin -1/2

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52
Q

Heisenberg Uncertainty Principle

A

there is an inherent uncertainty in the product of the position of a particle and its momentum; xp >= h/2

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53
Q

Aufbau principle

A

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

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54
Q

Exceptions for electron configuration

A

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

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55
Q

Electron configuration

A

lists the shell and the subshells of an element’s electrons in order typically from lowest to highest energy level

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56
Q

Ground state

A

lowest energy level

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57
Q

Excited state

A

when at least one e- has moved from a lower energy subshell to a higher energy subshell

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58
Q

Hund’s rule

A

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

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59
Q

Paramagnetic elements

A

elements with unpaired electrons that align with an external magnetic field

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60
Q

Diamagnetic elements

A

elements with no unpaired electrons and are unresponsive to an external magnetic field

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61
Q

Emission line spectrum

A

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

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62
Q

Absorption line spectrum

A

measures the radiation absorbed when electrons absorb energy to move to a higher energy state

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63
Q

Photoelectric effect

A

the emission of electrons when light falls on a material; KE = hf - work function

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64
Q

Covalent bond

A

electrons are shared between atoms

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65
Q

Ionic bond

A

electrons are transferred from one atom to another

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66
Q

Molecule

A

atoms held together by only covalent bonds

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67
Q

Bond length

A

the distance between the nuclei of 2 atoms when they are at their lowest possible energy state

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68
Q

Bond dissociation energy/bond energy

A

the energy necessary for a complete separation of the bond

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69
Q

Partial ionic character

A

when the difference in electronegativity is significant

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70
Q

Dipole moment

A

occurs when the center of positive charge in a bond does not coincide with the center of negative charge

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71
Q

Intermolecular attractions

A

attractions between separate molecules due to dipole moments

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72
Q

Hydrogen bond

A

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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73
Q

Induced dipole

A

occurs when dipole moment is momentarily induced in an otherwise nonpolar molecule or bond by a polar molecule, ion, or electric field

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74
Q

Instantaneous dipole moment

A

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

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75
Q

London dispersion forces/Van der Waals forces

A

weakest dipole-dipole force between 2 instantaneous dipoles

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76
Q

Empirical formula

A

the smallest ratio of whole numbers that can be used to represent these proportions

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77
Q

Molecular formula

A

represents the exact number of elemental atoms in each molecule

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78
Q

Physical reaction

A

when a compound undergoes a reaction and maintains its molecular structure ie. melting, evaporation, dissolution, and rotation of polarized light

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79
Q

Chemical reaction

A

when a compound undergoes a reaction and changes its bonding or structure to form a new compound ie. combustion, metathesis, and redox

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80
Q

Combination Reaction

A

A + B –> C

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81
Q

Decomposition reaction

A

C –> A + B

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82
Q

Single Displacement/Replacement

A

A + BC –> B + AC

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83
Q

Double Displacement/Replacement or Metathesis

A

AB + CD –> AD + BC; often occurs between ionized salts dissolved in water

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84
Q

runs to completion

A

means that the reaction generates products until the supply of at least one reactant is fully depleted

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85
Q

Limiting reactant

A

the first reactant to be used up first

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86
Q

Theoretical yield

A

the amount of product that should be created when a reaction runs to completion based on the stiochiometry

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87
Q

Percent yield

A

actual yield/theoretical yield x 100%

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88
Q

Actual yield

A

the amount of product created by a real experiment

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89
Q

mole

A

6.022 x 10^23 of something; grams/mw

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90
Q

Avogadro’s number

A

6.022 x 10^23

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91
Q

Radioactive decay

A

atoms that spontaneously break apart

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92
Q

Half-life

A

the length of time necessary for one half of a given substance to decay

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93
Q

Type of decay of radioactive decay

A

exponential decay

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94
Q

Semi-log plot

A

plotting the logarithm of amount of atoms as a function of time would produce a straight line for something that exponentially decays

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95
Q

4 variables of a half-life

A
  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)
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96
Q

Alpha decay

A

the loss of an alpha particle/helium nucleus

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97
Q

Beta decay

A

the breakdown of a neutron into a proton and electron and the expulsion of the newly created electron

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98
Q

Neutrino

A

also emitted during beta decay; virtually a massless particle

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99
Q

Positron emission

A

the emission of a positron when a proton becomes a neutron; type of beta decay; neutrino also emitted

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100
Q

Electron capture

A

the capture of an electron and the merging of that electron with a proton to create a neutron

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101
Q

Gamma ray

A

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

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102
Q

Gamma decay

A

often accompanies the other types of radioactive decay; can occur when an electron and positron collide

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103
Q

Valence

A

the number of bonds an atom usually forms

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104
Q

Formal charge

A

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

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105
Q

Dash formula

A

shows the bonds between each atom of a molecule but does not show lone pairs nor the three dimensional structure of the molecule

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106
Q

Condensed formula

A

shows neither the bonds nor the 3D structure

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107
Q

Bond-line formula

A

line intersections, corners and endings represent a carbon atom unless another atom is drawn in; hydrogens are not drawn

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108
Q

Fischer projection

A

vertical lines are assumed to be oriented into the page and horizontal lines are assumed to be oriented out of the page

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109
Q

Newman projection

A

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

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110
Q

Dash-line wedge formula

A

the solid black wedges represent bonds coming out of the page and the dashed wedges represent bonds going into the page

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111
Q

Space-filling model

A

a 3D representation of a molecule with spheres of various colors representing different elements with respect to their relative sizes

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112
Q

Ball and sticks model

A

bond lengths are drawn to approximately twice their length so that the atoms are clearly visible; give information about relative size of atoms

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113
Q

Sigma bond

A

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

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114
Q

Pi bind

A

created by overlapping p orbitals; double and triple bonds are made by adding pi bonds to a sigma bond

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115
Q

How to determine number and type of hybrid orbital

A

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

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116
Q

Valence shell electron pair repulsion (VSEPR)

A

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

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117
Q

Shape/ bond angle of sp

A

linear, 180

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118
Q

Shape/ bond angle of sp2

A

trigonal planar, 120

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119
Q

Shape/ bond angle of sp3

A

tetrahedral, trigonal pyramidal or bent, 109.5

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120
Q

Shape/ bond angle of sp3d

A

trigonal-bipyramidal, see-saw, t-shaped, or linear; 90, 120

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121
Q

Shape/ bond angle of sp3d2

A

octahedral, square pyramidal, or square planar, 90

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122
Q

Delocalized electrons

A

bonding electrons which are spread out over 3 or more atoms

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123
Q

Resonance structures

A

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

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124
Q

Aromaticity

A

the increased stability of a cyclic molecule due to electron delocalization; must follow Huckel’s rule

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125
Q

Huckel’s rule

A

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)

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126
Q

Nucelophilic functional groups

A

have a partial negative charge and seek positively charged nuclei ; they donate electrons and attack functional groups with partial positive charges ie. amines

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127
Q

Electrophilic functional groups

A

have a partial positive charge and seek electrons; usually get attacked by electrons from other functional groups

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128
Q

Stereochemistry

A

three-dimensional structure of a molecule

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129
Q

Isomers

A

molecules that share the same molecular formula

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130
Q

Structural isomer

A

have the same molecular formula but different bond-to-bond connectivity and thus different chemical properties due to differences in functional groups

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131
Q

Conformational isomers/conformers

A

different spatial orientations of the same molecule

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132
Q

Stereoisomers

A

2 unique molecules with the same molecular formula and the same bond-to-bond connectivity

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133
Q

Enantiomers

A

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

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134
Q

Chiral molecule

A

carbon bonded to four different substituents that rotate plane-polarized light

135
Q

Determine absolute configuration

A

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
Q

Relative configuration

A

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
Q

Determine absolute configuration for Fischer projection

A

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
Q

Plane-polarized light

A

photons in which their electric fields are oriented in the same direction

139
Q

Specific rotation

A

a standardized form of observed rotation that is calculated from the observed rotation and experimental parameters

140
Q

Dextorotatory

A

rotates plane-polarized light clockwise; d or +

141
Q

Levorotary

A

rotates plane-polarized light counterclockwise; l or -

142
Q

Observed rotation

A

the direction and number of degrees that the electromagnetic field is rotated when it passes through a compound

143
Q

Racemic mixture

A

when enantiomers are present in equal concentration and so it does not rotate plane-polarized light

144
Q

Diastereomers

A

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
Q

Maximum number of optically active isomers

A

2^n where n is the number of chiral centers

146
Q

Meso compound

A

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
Q

Epimers

A

diastereomers that differ in configuration at only one chiral carbon

148
Q

Anomers

A

cyclic diastereomers that are formed when a ring closure occurs at an epimeric carbon

149
Q

Anomeric carbon

A

chiral carbon of an anomer

150
Q

Cis-isomer

A

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
Q

Trans-isomer

A

molecules with substituents on opposite sides of double bond or ring structure

152
Q

Steric hindrance

A

when substituents in the cis position crowd each other which raises their energy levels, decreasing stability

153
Q

E Diastereomer

A

when the higher priority substituents are on opposite sides of the locked bond

154
Q

Z Diastereomer

A

when the higher priority substituents are on the same side of the locked bond

155
Q

Substitution reactions

A

one group leaves and is replaced with another

156
Q

SN1 reaction

A

has a rate that is dependent on only the substrate/electrophile, leaving group leaves spontaneously to form a carbocation; carboskeleton rearrangement can occur

157
Q

SN2 reaction

A

occurs in a single step, rate dependent on concentration of nucleophile and substrate, inversion of configuration

158
Q

Substrate of SN1/SN2

A

SN1: secondary or tertiary; SN2: don’t occur with sterically hindered substrate, requires methyl, primary, or secondary substrate

159
Q

Solvent of SN1/SN2

A

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
Q

Speed of SN1/SN2

A

SN1: concentration of substrate; SN2: concentration of substrate and nucleophile

161
Q

Stereochemistry of SN1/SN2

A

SN1: creates both enantiomers; SN2: inverts sterepchemistry about the chiral center

162
Q

Carbon skeleton rearrangment SN1/SN2

A

SN1: may occur; SN2: never occurs

163
Q

Leaving group

A

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
Q

Alcohols

A

consist of an oxygen bound to a hydrogen and an R group and undergo three major types of reactions: nucleophile, acid, leaving group

165
Q

Trend of alcohol acidity

A

from strongest to weakest acid: methyl > primary > secondary > tertiary; reverse for nucleophilicity trend

166
Q

Formation of sulfonates

A

nucleophlic substitution where alcohol acts as the nucleophile

167
Q

Tosylates/mesylates

A

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
Q

Ethers

A

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
Q

Cleavage of an ether with strong acid

A

R2O + HBr –> ROH + RBr

170
Q

Nitrogen

A

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
Q

Electron donating groups

A

increase the basicity/nucleophilicity of the molecule

172
Q

Electron withdrawing groups

A

increase the acidity of the molecule (decrease its nucleophilicity)

173
Q

Properties that make carbonyls good electrophiles

A

planar stereochemistry and partial positive charge on carbon

174
Q

Planar stereochemistry

A

leaves open space above and below a molecule reducing steric hindrance and making it more receptive to a nucleophilic attack

175
Q

Racemix mixture

A

equal amounts of R and S

176
Q

Stereoselective

A

a reaction prefers either R or S

177
Q

Stereospecific

A

when only the R or S is formed

178
Q

Trend of reactive of carbonyls

A

electron donating groups reduce carbonyl reacitivty to a nucleophile attack while electron withdrawing groups increase this reactivity

179
Q

Carboxylic acid derivatives

A

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
Q

Hydrolysis of amides

A

only possible under extreme chemical conditions that are unlikely to occur in biological systems

181
Q

Anhydrides

A

group of molecules in which the leaving group is a carboxylate ion

182
Q

Esters

A

formed from an alcohol undergoing nucleophilic substitution with a carboxylic acid (esterification)

183
Q

Transesterification

A

alcohols react with esters in a reaction called transesterification where one alkoxy group is substituted for another in the ester

184
Q

Amides

A

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
Q

Lactams

A

cyclic amides which are unstable

186
Q

Tautomerization

A

shift from a carbonyl to an alkene with an alcohol; is a reaction at equilibrium and NOT a resonance

187
Q

alpha carbon

A

any carbon attached directly to a carbonyl carbon

188
Q

Kinetic enolate

A

the less substituted enolate; it is formed quickly by the removal of one of three possible by the removal of an alpha hydrogen

189
Q

Thermodynamic enolate

A

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
Q

Hemiacetals/hemiketals

A

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
Q

acetal/ketal

A

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
Q

Imine

A

looks like a carbonyl as there is a carbon-nitrogen double bond; formed when aldehyde/ketone reacts with amine

193
Q

Enamine

A

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
Q

Grignard synthesis of an alchool

A

organometallic reagent + aldehyde/ketone –> alcohol with R group from organometallic reagent

195
Q

Reduction synthesis of an alcohol

A

use NaBH4 or LiAlH4; only LiAlH4 is strong enough to fully reduce carboxylic acids and esters and acetates to alcohol

196
Q

Cyanohydrins

A

a nitrile and alcohol attached to the same carbon; synthesized when nitriles attack carbonyls

197
Q

Oxidation

A

increase in bonds to oxygen or halogen, loss of C-H bonds

198
Q

Reduction

A

increase in bonds to H or R groups, loss of bonds to oxygen or halogen

199
Q

Neither oxidation nor reduction

A

addition or loss of H+, water, HX etc.

200
Q

Decarboxylation

A

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
Q

Aldol condensation

A

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
Q

Carboyhdrates

A

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
Q

Aldose

A

when carbohydrate contains an aldehyde

204
Q

Ketose

A

when carbohydrate contains a ketone

205
Q

Epimers

A

carbohydrates that have the same structure except for the configuration around a single chiral center

206
Q

Anomeric carbon

A

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
Q

Anomer

A

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
Q

Furanose

A

five member ring carbohydrate

209
Q

Pyranose

A

six member ring carbohydrate

210
Q

Glucopyranose

A

glucose ring

211
Q

Disaccharides/polysachharides

A

glycosides where the aglycone is another sugar and are connected by glycosidic linkages

212
Q

Aglycone

A

the group attached the the anomeric carbon of a glycoside

213
Q

Glycoside

A

sugar that is formed when a sugar is attacked by an alcohol to create an acetal

214
Q

Sucrose

A

1,1 glycosidic linkage: glucose and fructose, alpha with respect to glucose and beta with respect to fructose

215
Q

Maltose

A

alpha 1-4 glycosidic linkage: 2 glucose molecules

216
Q

Lactose

A

beta 1,4 galactosidic linkage: galactose and glucose

217
Q

Amylose (starch)

A

alpha 1,4 glycosidic linkage: chain of glucose molecules

218
Q

Amylopectin

A

alpha 1,4 glycosidic linkages:a branched chain of glucose molecules with alpha 1,6 glucosidic linkages forming the branches

219
Q

Glycogen

A

alpha 1,4 glycosidic linkage: a branched chain of glucose molecules with alpha 1,6 glucosidic linkages forming the branches

220
Q

Gabriel synthesis

A

way to form amino acids

221
Q

Strecker synthesis

A

way to form amino acids

222
Q

Lipogenesis

A

the formation of lipids

223
Q

Lipolysis

A

lipid breakdown

224
Q

Saponification

A

when lipid breakdown is base catalyzed where a soap forms (fatty acid salts and glycerol)

225
Q

Acidity of hydrogens in carbonyl

A

alpha carbon is most acidic

226
Q

Changing the reactivity of aromatic rings

A

electron withdrawing groups (meta, NO2, halides, -SO2R, -CO2R) deactivate while electron donating groups (ortho, para, H, alkyl groups, -OH, -OR, -NH2, -NR2)

227
Q

Thermodynamics

A

study of energy and its relationship to macroscopic properties of chemical systems

228
Q

Extensive property

A

proportional to the size of the system ie. volume and number of moles

229
Q

Intensive property

A

independent of the size of the system ie. pressure and temperature

230
Q

Temperature

A

represents the amount of molecular movement in a substance

231
Q

Thermal energy

A

the sum of translational, rotational, and vibrational energies

232
Q

Absolute zero

A

0K or -273 C

233
Q

Reaction rate

A

how quickly the concentration of the reactants or products are changing over the course of the reaction

234
Q

Elementary reaction

A

reaction that occurs in a single step

235
Q

Intermediates

A

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
Q

Zero order reactant

A

reaction rate does not depend on concentration

237
Q

1st order reactant

A

reaction rate directly proportional to concentration

238
Q

2nd order reactant

A

reaction rate is proportional to the square of the concentration

239
Q

Rate-determining step

A

slowest elementary step; rate law derived from this step

240
Q

Catalyst

A

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
Q

Steric factor

A

fraction of collisions having the effective spatial orientation

242
Q

Homogeneous catalyst

A

is in the same phase as the reactants and prodcts

243
Q

Heterogeneous catalyst

A

is in a different phase than the reactants or products such as when gas or aqueous particle reacts on a solid

244
Q

State

A

the physical condition of a system as described by a specific set of thermodynamic properties

245
Q

State functions

A

properties that describe the current state of a system; do not depend on the path taken

246
Q

Path functions

A

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
Q

Internal energy

A

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
Q

Vibrational energy

A

created by the back and forth motion of atoms within a molecule

249
Q

Rotational energy

A

created by the rotation of a molecule around its center of mass

250
Q

Translational energy

A

created by movement of the center of mass of a molcule

251
Q

Electronic energy

A

the potential electrical energy created by the attractions between electrons and their nuclei

252
Q

Intermolecular potential energy

A

created by intermolecular forces between molecular dipoles

253
Q

Rest mass energy

A

energy described by E =mc^2

254
Q

Heat

A

spontaneous transfer of energy from a warmer body to a cooler body

255
Q

Zeroth law of thermodynamics

A

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
Q

Conduction

A

thermal energy transfer via molecular collisions; requires direct physical contact

257
Q

Convection

A

thermal energy transfer via fluid movements; differences in pressure or density drive warm fluid in the direction of cooler fluid

258
Q

Radiation

A

thermal energy transfer via electromagnetic waves

259
Q

Newton’s law of cooling

A

a body’s rate of cooling is proportional to the temperature difference between the body and its environment

260
Q

Emissivity

A

a constant between 0 and 1 and specific to that object; higher emissivity indicates that a higher amount of radiation energy is absorbed

261
Q

First law of thermodynamics

A

the total energy of the system and surroundings is always conserved

262
Q

Second law of thermodynamics

A

the net entropy or disorder of the universe is always increasing

263
Q

Third law of thermodynamics

A

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
Q

Standard state

A

the reference form of a substance at any chosen temperature and a pressure of 1 bar `s

265
Q

Standard enthalpy of formation

A

the change in enthalpy for a reaction that creates one mole of that compound from its raw elements in their standard states

266
Q

Naught

A

indicates standard state conditions

267
Q

Endothermic reaction

A

a reaction with a positive enthalpy change

268
Q

Exothermic reaction

A

a reaction with a negative enthalpy change

269
Q

Entropy

A

nature’s tendency to create the most probably arrangement that can occur within a system; state function; extensive property

270
Q

Gibbs free energy

A

extensive property; state function; represents the maximum work available/free for a reaction

271
Q

Endergonic reaction

A

positive delta G; non-spontaneous

272
Q

Exergonic reaction

A

negative delta G; spontaneous

273
Q

Hess’s law of heat summation

A

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

274
Q

Transition state

A

peak of reaction coordinate where old bonds are breaking and new bonds are forming

275
Q

Intermediate location on reaction coordinate

A

would show multiple peaks (transition states) and the intermediates would be the troughs lying between these peaks

276
Q

Chemical equilibrium

A

when the forward reaction rate equals the reverse reaction rate and there is no net change in the concentration of the products or reactants

277
Q

Law of Mass Action

A

the mathematical relationship between a chemical equation and the associated equilibrium constant

278
Q

Le Chatelier’s principle

A

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
Q

Relationship between K and G

A

if K = 1, G = 0; if K >1, G < 0; if K <1, G > 0

280
Q

Ideal gas model/kinetic molecular theory

A
  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
281
Q

Simple mercury barometer

A

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

282
Q

Boyle’s law

A

pressure and volume inversely related

283
Q

Charles’ law

A

volume of a gas is directly proportional to temperature

284
Q

Avogadro’s law

A

volume of a gas is directly proportional to the number of moles of that gas

285
Q

Adiabatic process

A

no heat transfer

286
Q

Isothermal process

A

no change in internal energy

287
Q

Isovolumetric process

A

no change in volume, therefore no work

288
Q

Partial pressure

A

the total pressure of the gaseous mixture multiplied by the mole fraction of the particular gas

289
Q

Real gases

A

volume is larger, pressure is smaller (due to intermolecular attractive forces)

290
Q

Deviations from ideal gas law

A

if pv/rt > 1, deviation due to molecular volume; of pv/rt < 1, deviation due to attractive intermolecular forces

291
Q

Heat capacity

A

the added energy required to increase the temperature of a given substance by one K (or degree C)

292
Q

Constant volume heat capacity

A

system cannot due PV work and thus all input energy must be in the form of heat

293
Q

Constant pressure heat capacity

A

added energy can be lost by work done by the system by expanding its volume therefore it is typically greater than Cv

294
Q

Calorimeter

A

measures heat change; holds a liquid with a thermometer placed inside to measure any changes in temperature

295
Q

Coffee cup calorimeter

A

constant pressure calorimeter that measures energy change at atmospheric pressure (open top); used to measure heats of reactions

296
Q

Bomb calorimeter

A

measures energy change at constant volume and thus indicates the internal energy change in a reaction; reaction takes place in a rigid container

297
Q

Heat of fusion

A

the enthalpy change associated with melting

298
Q

Heat of vaporization

A

enthalpy change associated with boiling

299
Q

Sublimation

A

solid to gas

300
Q

Deposition

A

gas to solid

301
Q

Condensation

A

gas to liquid

302
Q

Phase diagram

A

indicates the phases of a substance at different pressures and temperatures

303
Q

Critical temperature

A

the temperature above which a substance cannot be liquefied regardless of the pressure applied

304
Q

Critical pressure

A

the pressure required to produce the liquid phase when the substance is at the critical temperature

305
Q

Triple point

A

the only point where a substance can exist in equilibrium between the solid, liquid, and gas phases

306
Q

Solution

A

homogeneous mixture of 2 or more compounds in a single phase

307
Q

Solvent

A

the compound of which there is more of in a solution

308
Q

Solute

A

the compound of which there is less of in a solution

309
Q

Solvation

A

when ionic compounds dissolve, the cations and anions break apart and are surrounded by the oppositely charged ends of the polar solvent

310
Q

Hydration

A

solvation in water; something that is hydrated is said to be in an aqueous phase

311
Q

Hydration number

A

the number of water molecules that must surround an ion for hydration to occur; commonly 4-6

312
Q

Electrolyte

A

a compound that forms ions in aqueous solutions; conduct electricity in solution

313
Q

Parts per million

A

10^6 multiplied by the ratio of the mass of solute to the total mass of the solution

314
Q

Normality

A

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
Q

Solution formation

A
  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
Q

Vapor pressure

A

the pressure created by the molecules in the open space at equilibrium

317
Q

Evaporation

A

a compound evaporates when the vapor pressure of its liquid phase is greater than the partial pressure of its gaseous phase

318
Q

Condensation

A

a compound condenses when the vapor pressure of its gas phase is greater than the partial pressure of its liquid phase

319
Q

Raoult’s law for nonvolatile solutes

A

if 97% of the solution is solvent, the vapor pressure will be 97% of the vapor pressure of the pure solvent

320
Q

Raoult’s law for volatile solutes

A

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
Q

Solubility

A

quantifies a solute’s tendency to dissolve in a solvent

322
Q

Saturated solution

A

the concentration of dissolved salt has reached a maximum

323
Q

Spectator ions

A

ions that have no effect on a reaction

324
Q

Henry’s law

A

the solubility of a gas is proportional to its vapor partial pressure

325
Q

Redox titration

A

used to find the molarity of a reducing agent, an oxidizing agent of known molarity is slowly added

326
Q

Half reaction

A

no half reaction can occur by itself; any reduction half reaction must be accompanied by an oxidation half reaction

327
Q

Electric potential

A

the ability of a molecule to be reduced, more positive = more readily reduced

328
Q

Galvanic/voltaic cell

A

generates a current from one phase to another in a conversion of chemical to electrical energy

329
Q

Terminals

A

conductors such as metal wires

330
Q

Electrodes

A

conductors

331
Q

Electromotive force (emf)

A

the potential difference between the terminals when they are not connected

332
Q

Salt bridge

A

a type of liquid junction that minimizes the potential difference caused by ions moving across a liquid junction

333
Q

Concentration cell

A

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
Q

Electrolytic cell

A

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

MCAT (5 decks)

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