Chapter 3: Periodic Properties of the Elements

Question 1

How is the density of elements related to the periodic table, and why?

Answer 1

The density of elements tends to increase as we move down a column in the table. The mass-volume ratio of an atom is an important characteristic in determining its density. As you move down the column, the density of the elements increases even though the radius increases as well; this happens because the mass of each successive atom increases even more than its volume does.

Question 2

Define periodic property.

Answer 2

A property that is generally predictable based on an element’s position within the periodic table.

Question 3

State the periodic law.

Answer 3

When the elements are arranged in order of increasing mass, certain sets of properties recur periodically.

Question 4

What is the difference between main-group elements and transition elements/metals?

Answer 4

The properties of main-group elements are largely predictable based on their position in the periodic table. Transition elements and inner transition elements are less predictable based simply on their position within the table.

Question 5

What is the name of a column of main-group elements in the periodic table?

Answer 5

A family or group.

Question 6

What are the two possible values for an electron’s spin? (ms)

Answer 6

Spin up (+ 1/2) or spin down (- 1/2).

Question 7

State the Pauli exclusion principle. Why is this true?

Answer 7

No two electrons in an atom can have the same four quantum numbers. There are only two possible values for an electron’s spin, so there can only be two electrons/orbital, each with an opposing spin.

Question 8

In multi-electron atoms, what does the energy level of the orbitals depend on?

Answer 8

The orbitals within a principal level of a multi-electron atom are not degenerate (all having the same energy). Instead, their energy depends on l; we say that the energies of the sublevels are split. In general, the lower the value of l within a principal level, the lower the energy of the corresponding orbital.

Question 9

What is Coulomb’s law?

Answer 9

This law states that the potential energy of two charged particles depends on their charges and on their separation.

Question 10

What is the relationship between energy and the distance between charged particles?

Answer 10

The potential energy associated with the interaction of like charges is positive but decreases as the particles get further apart (as r increases). Like charges that are close together have high potential energy and tend to move away from each other. The opposite is true of opposite charges.

Question 11

How is the magnitude of the interaction between charged particles related to the charge of those particles?

Answer 11

The magnitude increases as the charges of the particles increase. An electron with a charge of 1- is more attracted to a nucleus with a charge of 2+ than to a nucleus with a charge of 1+.

Question 12

Explain how electron shielding works.

Answer 12

The inner electrons in effect are repulsed by outer electrons. This screens the outer electrons from the full effect of the nuclear charge.

Question 13

Explain how penetration works.

Answer 13

As outer electrons enter the innermost orbital of an atom, it undergoes penetration into the region occupied by the inner electrons. It experiences a greater nuclear charge and therefore (according to Coulomb’s law) a lower energy.

Question 14

Even though an electron in a 2p orbital is more likely to be found closer to the nucleus than in a 2s orbital, why is the energy higher in the 2p orbital? What prerequisite must be met for this statement to be true?

Answer 14

First of all, the 1s orbital must be filled (When it is empty, the 2s and 2p orbitals are degenerate). Almost all the 2p orbital is shielded from nuclear charge by the 1s orbital. The 2s orbital, because it experiences more of the nuclear charge due to its greater penetration, is lower in energy than the 2p orbital.

Question 15

What is the Aufbau principle?

Answer 15

The pattern of orbital filling is based on the Pauli exclusion principle.

Question 16

State Hund’s rule and explain why it is so.

Answer 16

Hund’s rule says that when filling degenerate orbitals, electrons fill them singly first, with parallel spins. This rule is a result of an atom’s tendency to find the lowest energy state possible. When two electrons occupy separate orbitals of equal energy, the repulsive interaction between them is lower than when they occupy the same orbital because the electrons are spread out over a larger region of space.

Question 17

State the order in which electrons fill orbitals.

Answer 17

1s2s2p3s3p4s3d4p5s4d5p6s

Question 18

What does the number of columns in a block tell you?

Answer 18

It corresponds to the maximum number of electrons that can occupy the particular sublevel of that block. For example, the s block has two columns, the p block has six columns, the d block has 10 columns, and the f block has 14 columns.

Question 19

What does the lettered group number tell you?

Answer 19

The lettered group number tells you the number of valence electrons for any main-group element. For example, we know that chlorine has seven valence electrons because it is in group number 7A.

Question 20

What does the row number of main-group elements tell you?

Answer 20

The row number of main-group elements is equal to the n value of the highest principal level. For example, because chlorine is in row 3, its highest principal level is the n = 3 level.

Question 21

What does the row number of transition series elements tell you?

Answer 21

The row number minus one tells you the principal quantum number of the d orbitals.

Question 22

What are the two exceptions for electron configurations within the transition series elements? Why are they irregular?

Answer 22

Cr and Cu. Respectively, their configurations are 4s13d5 and 4s13d10. They are irregular because 4s orbitals are close in energy to 3d orbitals. This involves factors of shielding and penetration, and the relative energies depend heavily on the specific sample and conditions under which it is studied.

Question 23

What does the row number of inner transition series elements tell us?

Answer 23

The row number minus two tells us the principal quantum number of the f orbitals that fill across each row.

Question 24

What are some basic properties of metals?

Answer 24

1. Good conductors of heat and electricity.
2. They can be pounded into flat sheets.
3. They can be drawn into wires.
4. They are often shiny.
5. They tend to lose electrons when they undergo chemical changes (most important).

Question 25

What are some basic properties of nonmetals?

Answer 25

1. Poor conductors of heat and electricity.
2. They tend to gain electrons when they undergo chemical changes.

Question 26

Define alkali metals and describe some characteristics.

Answer 26

Alkali metals are the group 1A elements. They all have an outer electron configuration of ns1. Alkali metals readily and violently react in chemical reactions.

Question 27

Define alkaline earth metals and describe some characteristics.

Answer 27

Alkaline earth metals are the group 2A elements. They all have an outer electron configuration of ns2.

Question 28

Define halogens and describe some characteristics.

Answer 28

Halogens are the group 7A elements. They all have an outer electron configuration of ns2np5. They tend to gain one electron to have an overall charge of 1-.

Question 29

Define the nonbonding atomic radius (also known as the van der Waals radius).

Answer 29

The van der Waals radius represents the radius of an atom when it is not bonded to another atom.

Question 30

Define the bonding atomic radius (also known as the covalent radius) in terms of nonmetals.

Answer 30

This radius equals one-half the distance between two of the atoms bonded together.

Question 31

Define the bonding atomic radius (also known as the covalent radius) in terms of metals.

Answer 31

This radius equals one-half the distance between two of the atoms next to each other in a crystal of the metal.

Question 32

Define atomic radius.

Answer 32

Atomic radius refers to a set of average bonding radii determined from measurements on a large number of elements and compounds. The atomic radius represents the radius of an atom when it is bonded to another atom and is always smaller than the van der Waals radius.

Question 33

What are the general trends in the atomic radii of main-group elements?

Answer 33

1. As we move down a column (or family) in the periodic table, the atomic radius increases.
2. As we move to the right across a period (or row) in the periodic table, the atomic radius decreases.

Question 34

Why does the atomic radius increase as we move down a column?

Answer 34

It increases because as we move down a column, the valence electrons move farther from the nucleus. As we move down, the highest principal quantum number (n) of the valence electrons increases. Consequently, the valence electrons occupy larger orbitals, resulting in larger atoms.

Question 35

What is the effective nuclear charge of an electron?

Answer 35

The average or net charge experienced by an electron. This equals the actual nuclear charge minus the charge shielded by other electrons.

Question 36

How does the shielding of the outermost electrons by the core electrons differ from the shielding of the outermost electrons by each other?

Answer 36

Core electrons efficiently shield electrons in the outermost principal energy level from nuclear charge, but outermost electrons do not efficiently shield one another from nuclear charge.

Question 37

Why do the atomic radii decrease as we move to the right across a period?

Answer 37

The effective nuclear charge experienced by the electrons in the outermost principal energy level increases, resulting in a stronger attraction between the outermost electrons and the nucleus, and smaller atomic radii.

Question 38

Why don’t the atomic radii of the transition elements follow the same trend as the main group elements as we move to the right across a row?

Answer 38

Instead of decreasing in size, the radii of transition elements stay roughly constant. Across a row of transition elements, the number of electrons in the outermost principal energy level (highest n value) is nearly constant (because the 4s orbital fills before the 3d). Therefore, the electrons experience a roughly constant effective nuclear charge, keeping the radius approximately constant.

Question 39

What were the results of the Stern-Gerlach experiment?

Answer 39

The experiment sent a beam of gaseous Ag atoms through a strong magnetic field. The researchers chose silver as the element because it is electrically neutral but has 47 electrons; therefore, one of them must be unpaired. The beam of silver gas was deflected to two different spots. This proved that an electron’s angular momentum was quantized and complementary.

Question 40

What happens if you try to solve a wave function with two electrons having the same spin?

Answer 40

The equation collapses to zero because two electrons in the same orbital cannot have the same spin.

Question 41

Which elements will “promote” an electron to a d orbital?

Answer 41

Cr, Mo, Cu, & Ag. Found in columns 4 & 9.

Question 42

What happens to d orbitals when they become full?

Answer 42

Once the d orbitals are full, they become core electrons; they are no longer valence electrons and will not participate in chemical reactions.

Question 43

What does it mean when two atoms are isoelectronic?

Answer 43

It means they have the same electron configuration.

Question 44

Describe Group 16 of the periodic table.

Answer 44

These are the chalcogens. They form 2- ions during reactions.

Question 45

The metalloids are found along a zig-zag line on the periodic table. Which elements along that line are the exception (are not metalloids)?

Answer 45

Al & Po are not metalloids.

Question 46

Compare the energy levels of ns and nd orbitals in the transition metals.

Answer 46

For transition metals, ns is the higher energy orbital and the first to lose electrons.

Question 47

How do you write the electron configuration of a transition metal cation?

Answer 47

For transition metal cations, you must remove the electrons in the highest n-value first, even if this does not correspond to the reverse order of filling. During filling, the 4s orbital normally fills before the 3d orbital. When a fourth-period transition metal ionizes, however, it normally loses its 4s electrons before its 3d electrons.

Question 48

What two factors contribute to the phenomenon that electrons are not filled and removed in reverse for transition metal cations?

Answer 48

1. ns and (n-1)d orbitals are extremely close in energy, and depending on the exact configuration, can vary in relative energy ordering.
2. As the (n-1)d orbitals begin to fill, the increasing nuclear charge stabilizes the d orbitals relative to the ns orbitals. This happens because the d orbitals are not the outermost (or highest n) orbitals and are therefore not effectively shielded from the increasing nuclear charge by the ns orbitals.

Question 49

What does it mean for an atom or ion to be paramagnetic?

Answer 49

When an atom or ion that contains unpaired electrons is attracted to an external magnetic field, we say it is paramagnetic.

Question 50

What does it mean for an atom or ion to be diamagnetic?

Answer 50

When an atom or ion has all of its electrons paired and is not attracted to an external magnetic field (it is instead slightly repulsed), it is diamagnetic.

Question 51

How does the radius of an atom compare to its corresponding cation?

Answer 51

Cations are much smaller than their corresponding neutral atoms. This is because the outermost electrons that the cation lost used to be shielded by the core electrons (increasing the radius dramatically).

Question 52

How does the radius of an atom compare to its corresponding anion?

Answer 52

Anions are much larger than their corresponding neutral atoms. Extra electrons in the outermost shell increase the repulsions among the valence electrons, causing the atomic radius to be large.

Question 53

Define ionization energy.

Answer 53

The ionization energy of an atom or ion is the energy required to remove an electron from the atom or ion in the gaseous state. IE is always positive because removing an electron always requires energy.

Question 54

How does ionization energy change as you move down a column, and why?

Answer 54

IE decreases as you move down a column in the periodic table. n increases as you move down a column. Orbitals with higher principal quantum numbers are larger than orbitals with smaller principal quantum numbers. Therefore, electrons in the outermost principal level are farther away from the positively charged nucleus and are therefore held less tightly as we move down a column.

Question 55

How does ionization energy change as you move to the right across a row, and why?

Answer 55

IE increases as you move to the right across a period. The outermost electrons experience a higher effective nuclear charge as you move to the right. Therefore, they also have a higher IE.

Question 56

Define first ionization energy.

Answer 56

The energy required to remove the first electron from an atom.

Question 57

Name the exceptions to the trends in first ionization energies.

Answer 57

Boron has a smaller ionization energy than beryllium.
Oxygen has a smaller ionization energy than nitrogen.

Question 58

Define electron affinity.

Answer 58

The electron affinity of an atom or ion is the energy change associated with the gaining of an electron by the atom in the gaseous state. Electron affinity is usually negative because an atom or ion usually releases energy when it gains an electron. (the process is exothermic)

Question 59

In general, how does electron affinity change as you move right across the periodic table?

Answer 59

Electron affinity generally becomes more negative (more exothermic) as you move to the right across a row in the periodic table. Therefore, the halogens (Group 7A) have the most negative electron affinities.

Question 60

Is there any trend in electron affinity as you move down a column?

Answer 60

Most groups do not exhibit a trend. However, among the Group 1A metals, electron affinity becomes more positive as you move down the column (because adding an electron becomes less exothermic).

Question 61

How does metallic character change as you move along the periodic table?

Answer 61

Metallic character decreases as you move to the right and increases as you move down a column.

Question 62

How are penetration and energy related?

Answer 62

Penetration and energy are inversely related.

Question 63

Name the electron configuration of Cr.

Answer 63

Cr = [Ar]4s1 3d5

Question 64

Name the electron configuration of Mo.

Answer 64

Mo = [Kr] 5s1 4d5

Question 65

Name the electron configuration of Cu.

Answer 65

Cu = [Ar]4s1 3d10

Question 66

Name the electron configuration of Ag.

Answer 66

Ag = [Kr]5s1 4d10

Question 67

Why is hydrogen unique on the periodic table (ie it doesn’t really fit in anywhere)?

Answer 67

It shares the chemical properties of the alkali metals in that it forms a 1+ ion, but none of the physical properties. It also shares the characteristic of the halogens that it could form a 1- ion to have a full valence, but it generally won’t.

Question 68

When does ionization energy encounter a barrier, and why?

Answer 68

Once a cation achieves a noble gas configuration, it does not want to change from that. Therefore, the energy required to remove an electron is immense–much larger than the previous ionization energies.

Question 69

What is electron affinity?

Answer 69

Electron affinity refers to the energy required to add an electron to a neutral atom in the gas phase.