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Organic Chemistry 1 > Chapter 1 > Flashcards

Chapter 1 Flashcards

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1
Q

The chemistry of compounds that contain the element carbon

A

Organic chemistry

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2
Q

A compound that doesn’t contain carbon

A

Inorganic

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3
Q

Organic compounds that came from living organisms, and only living things could synthesize organic compounds through intervention of a vital force

A

Vitalism

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4
Q

Elements combined in different proportions

A

Compounds

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5
Q

Made up of atoms

A

Elements

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6
Q

Consists of a nucleus containing protons and neutrons surrounded by a cloud of electrons

A

Atom

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7
Q

A number equal to the number of protons in its nucleus

A

Atomic number (Z)

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8
Q

Atoms of the same element with different masses because of different amounts of neutrons

A

Isotopes

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9
Q

A hydrogen atom with one neutron

A

Deuterium (2H)

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10
Q

A hydrogen atom with 2 neutrons

A

Tritium (3H)

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11
Q

The outermost shell of electrons that are used in bonding

A

Valence shell

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12
Q

Electrons in the valence shell

A

Valence electrons

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13
Q

Bonds formed by the transfer of one or more electrons from one atom to another to create ions; attractive force between oppositely charged ions

A

Ionic bonds

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14
Q

Bonds formed when atoms share electrons; sharing of electrons between atoms of similar electronegativities to achieve the configuration of a noble gas

A

Covalent bonds

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15
Q

The tendency for an atom to achieve a configuration where it’s valence shell contains 8 electrons

A

Octet rule

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16
Q

A measure of the ability of an atom to attract electrons

A

Electronegativity

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17
Q

Compounds that form only when atoms of very different electronegativities transfer electrons to become ions

A

Salts (ionic compounds)

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18
Q

Composed of atoms joined exclusively or predominantly by covalent bonds

A

Molecules

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19
Q

A formula that has lines that show bonding electrons pairs and includes elemental symbols for the atoms in a molecule

A

Dash structural formula

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20
Q

A bond between two carbon atoms comprised of two electrons shared in a sigma (σ) bond

A

Carbon-carbon single bond

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21
Q

A representation of a molecule showing electron pairs as a pair of dots or as a dash

A

Lewis dot structure (electron dot structure)

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22
Q

Bonds comprised of one sigma (σ) bond and two pi (π) bonds

A

Triple bond

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23
Q

A bond between two carbon atoms comprised of four electrons; two of the electrons are in a sigma bond and two of the electrons in a pi bond

A

Carbon-carbon double bond

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24
Q

A bond between two carbon atoms comprised of six electrons; two of the electrons are in a sigma bond and four of the electrons are as pairs in each of two pi bonds

A

Carbon-carbon triple bond

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25
Q

A negative ion

A

Anion

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26
Q

A positive ion

A

Cation

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27
Q

A comparison between the number of valence electrons an atom “owns” in a molecule compared with the number of valence electrons it would have as a neutral isolated ion

A

Formal charge

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28
Q

Different molecules that have the same molecular formula

A

Isomers

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29
Q

A formula that gives the total number of each kind of atom in a molecule; is a whole number multiple of the empirical formula

A

Molecular formula

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30
Q

Compounds that have the same molecular formula but that differ in their connectivity (i.e., same molecular formula but have their atoms connected in different ways)

A

Constitutional isomers

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31
Q

The sequence, or order, in which the atoms of a molecule are attached to each other

A

Connectivity

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32
Q

A formula that shows how the atoms of a molecule are attached to each other

A

Structural formula

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33
Q

A chemical formula written using letters of the elemental symbols for the atoms involved, listed in sequence for the connections of the central chain of atoms and without showing the bonds between them

A

Condensed structural formula

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34
Q

A formula that shows the carbon skeleton of a molecule with lines. The number of hydrogen atoms necessary to fulfill each carbon’s valence is assumed to be present but not written in. Other atoms (e.g., O, Cl, N) are written in

A

Bond-line formula (skeletal formula)

35
Q

Structural formulas in which atom symbols are drawn and a line or “dash” represents each pair of electrons (a covalent bond). These formulas show connectivities but don’t represent the true geometries of the species

A

Dash structural formulas

36
Q

The angle between two bonds originating at the same atom

A

Bond angle

37
Q

A molecular geometry where a central atom is located at the center with four substituents that are located at the corners of a tetrahedron; bond angles are approximately 109

A

Tetrahedral geometry

38
Q

A molecular geometry with one atom at the center and three atoms at the corners of an equilateral triangle, called peripheral atoms, all in one place; bond angles are 120

A

Trigonal planar geometry

39
Q

A molecular geometry where a central atom is bonded to two other atoms; bond angles are 180

A

Linear geometry

40
Q

Lewis structures that differ from one another only in the position of their electrons. A single resonance structure will not adequately represent a molecule. The molecule is better represented as a hybrid of all the resonance structures

A

Resonance structures (resonance contributors)

41
Q

These show the direction of electron flow in a reaction mechanism. They point from the source of an electron or electron pair to the atom receiving the electron or electron pair. Double-barbed curved arrows are used to indicate the movement of a pair of electrons; single-barbed curved arrows are used to indicate the movement of a single electron. They are not used to show the movement of atoms

A

Curved arrows

42
Q

A calculated model in which regions of higher electron density trend toward red and regions of lower electron density trend toward blue

A

Electrostatic potential map

43
Q

A mathematical expression derived from quantum mechanics corresponding to an energy state for an electron, i.e., for an orbital; corresponds to a different energy state for an electron

A

Wave function (ψ)

44
Q

Signs, either + or -, that are characteristic of all equations that describe the amplitudes of waves; indicates whether the solution is positive or negative when calculated for a different point in space relative to the nucleus

A

Phase sign

45
Q

When wave functions with the same phase sign interact. There is a reinforcing effect and the amplitude of the wave function increases

A

Constructive interference

46
Q

When wave functions with opposite phase signs interact. There is a subtractive effect and the amplitude of the wave function goes to zero or changes sign

A

Destructive interference

47
Q

A volume of space in which there is a high probability of finding an electron

A

Orbital

48
Q

A volume of space about the nucleus of an atom where there is a high probability of finding an electron; plots of ψ2 in three dimensions consisting of s, p, and d orbital shapes

A

Atomic orbital

49
Q

A spherical atomic orbital; s orbitals have the quantum number l = 0

A

s orbital

50
Q

A set of three degenerate (equal energy) atomic orbitals shaped like two tangent spheres with a nodal plane at the nucleus; p orbitals have the quantum numbers n = 2, l = 1, m = +1, 0, -1

A

p orbital

51
Q

Orbitals of equal energy (ex. the three 2p orbitals)

A

Degenerate orbitals

52
Q

The principle that states orbitals are filled so that those of lowest energy are filled first

A

Aufbau principle

53
Q

The principle that states a max of two electrons may be placed in each orbital but only when the spins of the electrons are paired

A

Pauli exclusion principle

54
Q

The rule stating when filling degenerate orbitals (ex. p orbitals) we add one electron to each with their spins unpaired until each of the degenerate orbitals are filled

A

Hund’s rule

55
Q

The equilibrium distance between two bonded atoms or groups

A

Bond length

56
Q

A fundamental principle that states that both the position and momentum of an electron (or of any object) can’t be exactly measured simultaneously

A

Heisenberg uncertainty principle

57
Q

Orbitals that encompass more than one atom of a molecule. When atomic orbitals combine to form molecular orbitals, the number of molecular orbitals that results always equals the number of atomic orbitals that combine

A

Molecular orbital

58
Q

Results when two orbitals of the same phase overlap; the energy is lower than the separate orbitals used to create the bonding orbital

A

Bonding molecular orbital (ψmolec)

59
Q

Results when two orbitals of opposite phase overlap; the energy is higher than the separate orbitals used to create the anti bonding molecular orbital

A

Antibonding molecular orbital (ψ*molec)

60
Q

A mathematical method for arriving at wave functions for molecular orbitals that involves adding or subtracting wave functions for atomic orbitals

A

LCAO (linear combination of atomic orbitals)

61
Q

A mathematical (and theoretical) mixing of two or more atomic orbitals to give the same number of new orbitals, called hybrid orbitals, each of which has some of the character of the original atomic orbitals

A

Orbital hybridization

62
Q

An orbital that results from the mathematical combination of pure atomic orbitals, such as the combination of pure s and p orbitals in varying proportions to form hybrids such as sp3, sp2, and sp orbitals

A

Hybrid atomic orbitals

63
Q

The lowest electronic energy state of an atom or molecule

A

Ground state

64
Q

A single bond formed through head-on overlap of atomic orbitals

A

Sigma (σ) bond

65
Q

Shows points in space that happen to have the same electron density

A

Electron density surface

66
Q

Hydrocarbons whose molecules contain a carbon-carbon double bond

A

Alkenes

67
Q

A molecular orbital formed when parallel p orbitals on adjacent atoms overlap

A

Pi (π) molecular orbital

68
Q

A molecular orbital whose energy is higher than that of the isolated atomic orbitals from which it’s constructed

A

Antibonding molecular orbital

69
Q

A molecular orbital formed by end-on overlap of orbitals on adjacent atoms

A

Sigma (σ) orbital

70
Q

Bonds comprised of four electrons: two electrons in a sigma (σ) bond and two electrons in a pi (π) bond

A

Double bonds

71
Q

Two objects are superposable if, when one object is placed on top of the other, all parts of each coincide

A

Superposable

72
Q

Compounds with the same molecular formula that differ only in the arrangement of their atoms in space

A

Stereoisomers

73
Q

Diastereomers that differ in their stereochemistry at adjacent atoms of a double bond or on different atoms of a ring

A

Cis-trans isomers

74
Q

Hydrocarbons in which two carbon atoms share three pairs of electrons between them, and are bonded by a triple bond

A

Alkynes

75
Q

The equilibrium distance between two bonded atoms or groups

A

Bond lengths

76
Q

A method of predicting the geometry at a covalently bonded atom by considering the optimum geometric separation between groups of bonding and non bonding electrons around the atom

A

Valence shell electron pair repulsion (VSEPR) model

77
Q

Electrons shared in covalent bonds

A

Bonding pairs

78
Q

Electrons that are unshared

A

Non bonding pairs (unshared pairs or lone pairs)

79
Q

Consists of 3 bonding pairs and 1 non bonding pair

A

Trigonal pyramidal geometry

80
Q

Consists of 2 or 4 bonding pairs

A

Linear geometry

81
Q

Consists of 3 bonding pairs

A

Trigonal planar geometry

82
Q

Consists of 4 bonding pairs

A

Tetrahedral geometry

83
Q

Consists of 2 bonding and 2 non bonding pairs

A

Angular/bent geometry

Organic Chemistry 1 (10 decks)

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