Chapter 3

Question 1

The energy of atomic orbitals increases as the ____ ____ ____ (____) increases

Answer 1

principal quantum number (n)

Question 2

As the principal quantum number, n, increases, the size of the ____ increases and the electrons spend more time farther from the nucleus

causing

the attraction to the nucleus is ____ and the energy associated with the orbital is ____ (less stabilized).

Answer 2

Orbital

weaker . . . higher

Question 3

What are the four quantum numbers?

Answer 3

principal quantum number n:
shell, the general region for the value of energy for an electron on the orbital

angular momentum or
azimuthal quantum number
l:
subshell, the shape of the orbital

magnetic quantum number ml:
orientation of the orbital

spin quantum number ms:
direction of the intrinsic quantum “spinning” of the electron

Question 4

In any atom with two or more electrons, the repulsion between the electrons makes energies of subshells with different values of ____ differ.

Answer 4

l:
subshell, the shape of the orbital

Question 5

The energy of the orbitals increases within a shell in the order ____<____<____<____

Answer 5

s < p < d < f.

Question 6

Electrons in atoms tend to fill ______________ first.

Answer 6

low-energy orbitals

Question 7

The arrangement of electrons in the orbitals of an atom is called the ____ ____ of the atom.

Answer 7

Electron configuration

Question 8

Electron configuration?

Answer 8

The arrangement of electrons in the orbitals of an atom is called the electron configuration of the atom.

Question 9

An electron configuration consists of symbols that contain three pieces of information: Which three?

Answer 9

1. The principal quantum shell, n.
2. The letter that designates the orbital type (l).
3. A superscript number that designates the number of electrons in that particular subshell.

Question 10

To determine the electron configuration for an atom we add a number of electrons equal to its ____ ____

Answer 10

atomic number

Question 11

Beginning with hydrogen, and continuing across the periods of the periodic table, we add one electron to the subshell of lowest available energy.

This procedure is called the ____ ____, from the German word ____ (“to build up”).

Answer 11

Aufbau principle . . . Aufbau

Question 12

Pauli Exclusion principle?

Answer 12

no two electrons in the same atom can have identical values for all four of their quantum numbers

Question 13

What are orbital diagrams?

Answer 13

Orbital diagrams are pictorial representations of the electron configuration, showing the individual orbitals and the pairing arrangement of electrons.

Question 14

Orbital diagrams:

An upward arrow represents and electron with ms = ____.
A downward arrow represents and electron with ms = ____.

Answer 14

An upward arrow represents and electron with ms = + ½.
A downward arrow represents and electron with ms = - ½.

Question 15

What is Hund’s rule?

Answer 15

Hund’s rule: the lowest-energy configuration for an atom with electrons within a set of degenerate orbitals is that having the maximum number of unpaired electrons.

Question 16

What are valance electrons?

Answer 16

The electrons occupying the orbital(s) in the outermost shell (highest value of n) are called valence electrons.

Question 17

What are core electrons?

Answer 17

The electrons occupying the inner shell orbitals are called core electrons.

Question 18

What is noble gas electron confiuration?

Answer 18

The core electrons represent noble gas electron configurations.

Electron configurations can be expressed in an abbreviated format by writing the noble gas that matches the core electron configuration, along with the valence electrons.

Question 19

Beginning with the ____ ____ scandium (atomic number 21), additional electrons are added successively to the ____ subshell.

Answer 19

transition metal . . . 3d

Question 20

For the two periods of ____ ____ ____, lanthanum (La) through lutetium (Lu) and actinium (Ac) through lawrencium (Lr).
Electrons are added to an ____ ____.

Answer 20

Inner transition metals . . . f subshell

Question 21

The periodic table arranges atoms so that elements with the same chemical and physical properties are in the same ____.

Elements in the same group have similar ____ ____ ____.

Answer 21

Group

valence electron configurations

Question 22

Valence electrons play the most important role in ____ ____.

Answer 22

Chemical reactions

Question 23

____ ____ ____ or ____ ____ are those in which the last electron added enters an s or a p orbital in the outermost shell.

Answer 23

Main group elements or representative elements are those in which the last electron added enters an s or a p orbital in the outermost shell.

Question 24

The valence electrons for main group elements are those with the ____ ____ ____.

Example: Gallium (Ga): [Ar]4s23d104p1
Ga has three valence electrons (4s2 and 4p1).
The completely filled 3d orbitals count as core, not valence electrons.

Answer 24

The valence electrons for main group elements are those with the highest n level.

Example: Gallium (Ga): [Ar]4s23d104p1
Ga has three valence electrons (4s2 and 4p1).
The completely filled 3d orbitals count as core, not valence electrons.

Question 25

____ ____ or ____ ____ are metallic elements in which the last electron added enters a d orbital.

Answer 25

Transition elements or transition metals

Question 26

____ ____ ____ are metallic elements in which the last electron added occupies an f orbital.

Answer 26

Inner transition elements

Question 27

For ____ ____ ____, the electrons that were added last are the first electrons removed. For ____ ____ and ____ ____ ____, the highest ns electrons are lost first, and then the (n – 1)d or (n – 2)f electrons are removed.

Answer 27

main group elements transition metal and inner transition metals

Question 28

How to predict how electrons are added?

Answer 28

The electrons are added in the order predicted by the Aufbau principle.

Question 29

An ____ forms when one or more electrons are added to a parent atom.

Answer 29

anion (negatively charged ion)

Question 30

A ____ forms when one or more electrons are removed from an atom.

Answer 30

cation (positively charged ion)

Question 31

What are periodic properties?

Answer 31

These periodic properties include: 1. Size (radius) of atoms and ions 2. Ionization energies 3. Electron affinities

Question 32

What is covalent radius?

Answer 32

The covalent radius is defined as one-half the distance between the nuclei of two identical atoms when they are joined by a covalent bond

Question 33

Z^eff Example hydrogen compared to all other atoms?

Answer 33

Effective nuclear charge, Zeff, is the pull exerted on a specific electron by the nucleus, taking into account any electron–electron repulsions. For hydrogen, there is only one electron and so the nuclear charge (Z) and the effective nuclear charge (Zeff) are equal. For all other atoms, the electron of interest are partially shielded from the pull of the nucleus by the other electron(s) present.

Question 34

____ electrons are adept at shielding, while electrons in the same valence shell do not block the nuclear attraction experienced by each other as efficiently.

Answer 34

core

Question 35

Each time we move from one element to the next across a period, ____ increases by one, but the shielding increases only slightly due to electrons being added only to the ____ shell.

Answer 35

Z (atomic number) . . . valance

Question 36

____ increases as we move from left to right across a period

Answer 36

Zeff

Question 37

Across a period: A ___ ___ (higher effective nuclear charge) is experienced by the outermost electrons, drawing them closer to the nucleus. The size of the atom (and its covalent radius) ____ across a period.

Answer 37

Stronger pull decreases

Question 38

Ionic radius?

Answer 38

Ionic radius is the measure used to describe the size of an ion.

Question 39

Are cations smaller or larger than the parent atom? Explain?

Answer 39

A cation always has fewer electrons and the same number of protons as the parent atom. As electrons are removed from the outer valence shell, the remaining electrons experience a greater Zeff charge and are drawn even closer to the nucleus. A cation is always smaller than the atom from which it is derived.

Question 40

Are anions smaller or larger than the parent atom? Explain?

Answer 40

An anion always has more electrons and the same number of protons as the parent atom. This results in a greater repulsion among the electrons and a decrease in Zeff. This causes the valence electrons to be farther from the nucleus. An anion is always larger than the atom from which it is derived.

Question 41

What does isoelectric mean?

Answer 41

Atoms and ions that have the same electron configuration are said to be isoelectronic

Question 42

List several and describe isoelectric atoms?

Answer 42

Examples of isoelectronic species: N3–, O2–, F–, Ne, Na+, Mg2+, and Al3+ (1s22s22p6) P3–, S2–, Cl–, Ar, K+, Ca2+, and Sc3+ ([Ne]3s23p6) For atoms or ions that are isoelectronic, the number of protons determines the size. The greater the nuclear charge, the smaller the radius in a series of isoelectronic ions and atoms.

Question 43

What is first ionization energy?

Answer 43

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy (IE1). X(g) ⟶ X+(g) + e− IE1

Question 44

Second ionization energy?

Answer 44

The energy required to remove the second most loosely bound electron is called the second ionization energy (IE2). X+(g) ⟶ X2+(g) + e− IE2

Question 45

What is the general trend of ionization energy?

Answer 45

Energy is always required to remove electrons from atoms or ions, so ionization processes are endothermic and IE values are always positive. As size (atomic radius) increases, the ionization energy decreases. First ionization energies: -Decrease down a group -Increase across a period -There are some deviations from this trend.

Question 46

Deviation in ionization energy oxygen and nitrogen example?

Answer 46

Deviation: Oxygen (O) has a lower IE1 than nitrogen (N). Removing one electron from O will eliminate the electron–electron repulsion caused by pairing the electrons in the 2p orbital. This means that removing an electron from O is more energetically favorable than removing an electron from N.

Question 47

Is it harder to remove an electron from neutral atom or cation?

Answer 47

Removing an electron from a cation is more difficult than removing an electron from a neutral atom because of the greater electrostatic attraction to the cation. Likewise, removing an electron from a cation with a higher positive charge is more difficult than removing an electron from an ion with a lower charge. Thus, successive ionization energies for an element always increase.

Question 48

What is electron affinity? What two forms can this process take? Describe each form?

Answer 48

The electron affinity (EA) is the energy change for the process of adding an electron to a gaseous atom to form an anion. X(g) + e− ⟶ X−(g) EA1 This process can be either endothermic or exothermic, depending on the element. For some elements energy is released when the gaseous atom accepts an electron (negative EA). For other elements energy is required for the gaseous atom to accept an electron (positive EA).

Question 49

It becomes easier to add an electron as the ____ ____ ____ of the atoms increases. Electron affinities becomes more ____ across a period. There are some deviations from this trend.

Answer 49

effective nuclear charge negative

Question 50

_________________, have a completely filled shell and the incoming electron must be added to a higher energy n level. ____ ____ has a filled ns subshell, and so the next electron added goes into the higher energy np subshell. ____ ____ has a half-filled np subshell and the next electron must be paired with an existing np electron. In all of these cases, the initial relative stability of the electron configuration disrupts the trend in EA.

Answer 50

he noble gases, group 18 Group 2 Group 15