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Chemistry 131 > Quantum mechanics > Flashcards

Quantum mechanics Flashcards

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1
Q

What is provided by “ψ2”?

A

Information about tan electron’s location in an allowed energy state; it gives us the probability that an electron is at that given position.

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2
Q

What is an orbital?

A

A specific distribution of electron density around an atom.

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3
Q

How many quantum numbers we use to describe orbitals?

A

Principal quantum number = n (size)
Azimuthal quantum number = l (shape)
Magnetic quantum number = ml (orientation)
mS = Orientation of electron in an orbital

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4
Q

As n increases the orbital becomes_____________, and the electron is __________ from the __________. Also, the energy associated with an orbital __________ as ___ increases.

A

Bigger, further away, nucleus, increases, n.

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5
Q

If orbitals have the same value of n, they are considered to be in the same…

A

Electronic shell

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6
Q

If orbitals have the same value of n and l, they are considered to be in the same…

A

Electronic subshell

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7
Q

An individual orbital is defined by a unique set which numbers?

A

n, l, and m1

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8
Q

How is the azimuthal number defined?

A

Values from 0 to n-1 and letters (s, p, d, f)

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9
Q

How is the m1 number defined?

A

-l to +l

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10
Q

How is the ms number defined?

A

2l + 1

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11
Q

What are the numbers related to the letters of the azimuthal number?

A

s(0), p(1), d(2), f(3)

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12
Q

When describing orbitals we are talking about the probability that an electron…

A

Will be in a given position.

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13
Q

The p orbital has 3 subshells

A

True

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14
Q

How many subshells does a d orbital have?

A

Five

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15
Q

How many subshells does a f orbital have?

A

Seven

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16
Q

What is the possible number for ms?

A

-1/2 or +1/2, referred to as spin up or spin down

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17
Q

Two electrons can have the same four quantum numbers.

A

False. Each electron has a specific set of quantum numbers. This is called the Pauli Exclusion Principle

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18
Q

Electrons will start filling up the ______ energy orbitals. This is referred to as the _________.

A

Lowest; ground state

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19
Q

What are called two electrons or opposite spin?

A

Paired

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20
Q

What is called an electro by itself?

A

Unpaired

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21
Q

What is the Hund’s rule?

A

“For degenerate orbitals, the lowest energy state (most stable) is obtained when the number of electrons with the same spin in maximized.”

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22
Q

What is the primary force behind Hund’s rule?

A

Electrostatic repulsion between electrons

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23
Q

What are the most stable elements?

A

The ones with full shells. There is no room for electrons to move, thus elements with full shells are particularly stable and nonreactive.

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24
Q

What are the electrons that drive reactivity and atom behavior?

A

Valence electrons

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25
Q

What are the other electrons called?

A

Core electrons

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26
Q

Elements where d orbitals are being filled are called…

A

Transition Metals

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27
Q

What is the similarity between Na, K, and Rb?

A

All have noble gas core and one s electron; also have similar chemical properties

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28
Q

The elements in which the ____ orbitals are being filled are called ____________; the elements in which the ____ orbitals are being filled are called ______________.

A

4f; lanthanides; 5f; actinides

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29
Q

What is the electron configuration of Chromium?

A

[Ar]4s13D5; this is an exception and it happens because subshells that are either completely full or exactly half full, are especially stable.

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30
Q

In a periodic table we can see that __________ come after a __________.

A

Soft reactive metal comes after a nonreactive gas

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31
Q

What do we use to describe the average environment created by the nucleus and other electrons in the atom?

A

Effective nuclear charge, Zeff; Zeff = Z-S, with S being a constant related to the core electrons

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32
Q

As we move across the periodic table from left to right, Zeff ________. As we go down a column, Zeff stays _________ .

A

Increases, constant.

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33
Q

What is the atomic size?

A

The average radius of the electron distribution in an atom.

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34
Q

We have similar radius numbers in bonding and nonbonding molecules.

A

False.

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35
Q

According to chemical trends, the atomic radius:

A

Increases if we go down a column, and decreases if we move across the periodic table.

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36
Q

What does ionic radii depend on?

A

Nuclear charge, the number of electrons, and the orbitals of the valence electrons.

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37
Q

For cations the ionic radius is…

A

Smaller than the atomic radius.

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38
Q

For ions of the same charge the size…

A

Increases as we go down the periodic table.

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39
Q

The higher the ionization energy, the easier it is to remove an electron.

A

False, the harder it is.

40
Q

Ionization energy increases as we move from left to right and up a group.

A

True.

41
Q

What are the most important periodic trends?

A

Electronegativity, ionization energy, electron affinity, atomic radius and metallic character.

42
Q

What is one exception in the electron configuration of ions?

A

Fe. Its regular configuration is [Ar]4s23d6. Upon the first ionization it loses one electron in the s orbital; upon the second one it loses the last electron on the s orbital. Only in the third ionization it actually loses electrons in the d orbital.

43
Q

Why are atoms in group 6A different than the rest of the periodic table?

A

Because they have a full n shell and an exactly half full p shell, which is specially stable.

44
Q

Electron affinity ____________ as we go across the periodic table from left to right.

A

Becomes more negative

45
Q

If EA is greater than zero, then…

A

An electron will not attach itself to an atom.

46
Q

We describe an element as metal, nonmetal, or metalloid depending on the…

A

Configuration of the element.

47
Q

What are chemical bonds?

A

The force that holds the atoms in position in a substance

48
Q

What are the three types of chemical bonds?

A

Ionic bonds - Electrostatic interaction between ions; generally found between a metal and a nonmetal.
Covalent bonds - Sharing of electrons; generally found between nonmetals.
Metallic bonds - Electrons are shared through the lattice; generally found between metals.

49
Q

The number of valence electrons as well as the electron configuration determine…

A

how two atoms are bonded together.

50
Q

What are the conditions for the formation of ionic bonds?

A

Ionization of the metal (should be low), and electron affinity of the nonmetal (should be high)

51
Q

Lattice Energy

A

It is the energy required to separate a mole of an ionic solid into its gaseous ions.

52
Q

Formation of ionic solids from their constituent elements is a highly endothermic process.

A

False, it is highly exothermic

53
Q

As ions get larger, the spacing between ions in a lattice decreases and as a result:

A

The lattice energy decreases.

54
Q

Changes in lattice energy based on charge are much greater than changes based on…

A

size

55
Q

Bonds where electron are shared equally are called…

A

Nonpolar covalent bond

56
Q

If one atom exerts a greater pull on the electrons than the other, we refer to it as a…

A

polar covalent bond

57
Q

What values do we use to predict bond polarity?

A

EA<0.5 - nonpolar covalent
EA>2.0 - Ionic
0.5

58
Q

Electronegativity ______ as we go from left to right, and it ________ if we go down a group.

A

Increases, decreases

59
Q

What is Electronegativity?

A

The ability of an atom to attract electrons to itself in a molecule. It is similar to electron affinity but it refers to atoms in a molecule, instead of single atoms.

60
Q

What is the most electronegative atom? And the least electronegative?

A

Fluorine (F; 4.0); Cesium (Cs; 0.7)

61
Q

Effective nuclear charge

A

Increases when we go across a period. Number of protons increase but number of shielding e is constant, More or less constant down a group

62
Q

Atomic size

A

More protons pulling electrons in tighter; from left to right there is a decrease, from top to bottom it increases (shells get larger)

63
Q

Ionization energy

A

Left to right ionization energy increases (elements get closer to noble gases, it takes more energy to remove electrons) It decreases from top to bottom.

64
Q

What are the three exceptions for the Octet rule?

A

Molecules or ions with an odd number of electrons; molecules or ions in which an atom has fewer than 8 valence electrons; molecules or ions in which an atom has more than 8 valence electrons.

65
Q

What are radicals?

A

Molecules or atoms with an odd number of valence electrons. They are highly reactive. One example is NO.

66
Q

What is the strength of the covalent bond (or bond enthalpies)?

A

The amount of energy required to break that bond.

67
Q

The more we increase the number of bonds,

A

bond strength increases

68
Q

The more we increase the number of bonds,

A

the shorter they get

69
Q

What does VESPR stand for and how does it define electron groups?

A

Valence Electron Shell Pair Repulsion; lone pairs, single bonds, double bonds, and triple bonds.

70
Q

What are the five geometric arrangements we can have for electron groups?

A

Linear (2 groups)
Trigonal planar (3 groups)
Tetrahedral (4 groups)
Trygonal bipyremidal (5 groups)
Octahedral (6 groups)

71
Q

Electron geometry is…

A

The placement of electrons in a molecule

72
Q

Molecular geometry is…

A

The placement of atoms in a molecule

73
Q

In any molecule with four electron groups, the electrons will rearrange themselves in a _________

A

Tetrahedron

74
Q

An atom with linear geometry has ___ atoms groups and will present an angle of:

A

Two

75
Q

An atom with trigonal planar geometry has ___ atoms groups and will present an angle of:

A

Three

76
Q

What kind of structure is this and how many bonds does it have?

A

Tetrahedral geometry; it is seen in molecules with 4 bonds. with four sp3 orbitals)

77
Q

What are the positions we can have in an atom?

A

Axial positions and equatorial positions.

78
Q

Trigonal Bipyramidal Geometry has how many bonds? And which angles?

A

5 bonds; 90 and 120 angles

79
Q

Which geometrical structure is this, how many bonds does it have, and what is its angle?

A

Octahedral geometry, 6 bonds, angle 90°

80
Q

In molecular geometry we desconsider lone pairs when drawing structures.

A

True. A tetrahedral structure for electron geometry is not necessarily going to be the same for molecular geometry, like in ammonia (NH3), for example

81
Q

If all electron groups are bonds then the electron geometry and molecular geometry will be similar

A

True

82
Q

In moleculr geometry, if you have 3 electron groups but one of them is a lone pais, then the structure will be

A

Linear, bent

83
Q

As the number of electron domain increases, the number of possible molecular shapes increases

A

True

84
Q

What do methane, ammonia, and water have in common?

A

They all have four electron domains, and a tetrahedral electron structure.

85
Q

Non-bonding pairs are slightly more ___________ than bonding pairs and have a slighty __________ density than bonding pairs.

A

Repulsive, higher charge; this causes them to push the electrons a little bit more, distorting the angles.

86
Q

Placing lone pairs in equatorial positions does what towards electron pair repulsion?

A

Minimizes it

87
Q

Dipole moments are __________ therefore they have ___________ and _____________.

A

Vectors, magnitude, momentum

88
Q

If the vectors cancel each other out, then the dipole moment is 1

A

False. It is zero, they cancel each other, therefore there is no displacement.

89
Q

In general, if the central atom of a molecule is symetrically surrounded by atoms of the same element, the it is:

A

Non-polar

90
Q

Orbitals do not have to be the same type to do what?

A

Overlap and form a bond

91
Q

What is the optimal distance?

A

It is the distance where we will have bonds connecting.

92
Q

What are Sigma bonds?

A

Bonds where the electron density is located on the axis

93
Q

What are pi bonds?

A

Bonds where the electron density is located off the axis

94
Q
A
95
Q

The lobes of the p-orbitals overlpa to form the pi bond, located off axis. On-axis electrons are from the overlap of sp2 electrons in a theta bond.

A

Chemistry 131 (14 decks)

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