Periodic Table Trends

Question 1

Where are the alkali metals located on the periodic table?

Answer 1

the first column

(Group IA)

Question 2

What is the valence shell configuration of all alkali metals, and to which oxidation state do they ionize?

Answer 2

s¹

Alkali metals are relatively electropositive, so they will lose 1 valence electron and form a +1 oxidation state.

Question 3

Where are the alkaline earth metals located on the periodic table?

Answer 3

the second column

(Group IIA)

Question 4

What is the valence shell configuration of all alkaline earth metals, and to which oxidation state do they ionize?

Answer 4

s²

They are relatively electropositive, so they will lose 2 valence electrons and form a +2 oxidation state.

Question 5

Where are the halogens located on the periodic table?

Answer 5

the fifth column of the p block

(Group VIIA)

Question 6

What is the valence shell configuration of all halogens? What oxidation state do they ionize to?

Answer 6

s²p⁵

They are quite electronegative, so they will accept one additional valence electron to take on a -1 oxidation state.

Question 7

Where are the noble gases located on the periodic table?

Answer 7

the sixth column of the p block

(Group VIIIA)

Question 8

What is the valence shell configuration of all noble gases, and to which oxidation state do they ionize?

Answer 8

s²p⁶

Trick question! Since they already have a completely filled octet, noble gases do not ionize, and they typically exist in the 0 oxidation state as free particles.

Exceptions are Kr and Xe, since they are below the 3rd row, can exceed their octet and make coordinate covalently bonded compounds such as XeF₆.

Question 9

What is the oxygen group, and where is it located on the periodic table?

Answer 9

It is the group (column) below oxygen.

It includes elements such as S and Se that are chemically similar to oxygen. Elements generally share similar properties with the other elements in their group.

Question 10

Where are the transition metals located on the periodic table?

Answer 10

the entire d block

Question 11

Why do transition metals have high conductivity?

Answer 11

Due to their unfilled d subshells

d electrons, by their nature, are loosely bound to the atom. As such, elements with partially-filled d subshells can be thought of as nuclei floating in a sea of unattached electrons, prime conditions for electrical conductivity.

Question 12

What are the representative elements, and where are they located on the periodic table?

Answer 12

They are the most common elements in the solar system and the universe.

They are found in the s block and the p block of the Periodic Table.

By standard nomenclature, these are groups IA, IIA, IIIA, IVA, VA, VIA, VIIA, and VIIIA.

Question 13

What is the valence subshell for the elements in the first two columns of the periodic table?

Answer 13

s

Group IA has an s¹ valence configuration, while IIA is s².

Note that helium also has a valence s subshell, but is typically listed on the farthest column with the noble gases, as it is chemically more similar to them than the alkaline earth metals.

Question 14

What is the valence subshell for the elements in the last six columns of the periodic table?

Answer 14

p

For example, Group IIIA has an s²p¹ valence configuration, while VIIIA is s²p⁶.

Note that although helium is typically listed on the farthest column with the noble gases in VIIIA, it actually has a valence s subshell.

Question 15

Describe the properties of metals in terms of their:

• position in the periodic table
• electronegativity
• preferred oxidation state

Answer 15

Metals are generally:

• found in the lower-left areas of the periodic table.
• low in electronegativity, losing electron density when bonded to nonmetals.
• found in positive oxidation states when in compounds.

Question 16

What are the main physical properties of metals?

Answer 16

Metals generally are/have:

• good conductors of heat and electricity.
• malleable, ductile, lustrous, and dense solids at room temp.
• fairly high melting and boiling points.

Question 17

Describe the properties of nonmetals in terms of:

• position in the periodic table
• electronegativity
• preferred oxidation state

Answer 17

Nonmetals are generally:

• found in the upper-right areas of the periodic table.
• high in electronegativity, gaining electron density when bonded to metals.
• found in negative oxidation states when in compounds.

Question 18

What are the main physical properties of nonmetals?

Answer 18

Nonmetals are/have:

• poor conductors of heat and electricity.
• dull and brittle if they form solids at room temperature.
• significantly lower melting and boiling points than metals (carbon is the primary exception).

Question 19

How many valence electrons does oxygen, element 8, have?

Answer 19

6

Oxygen is the 6th element in its row. Its valence shell configuration is 2s²2p⁴, for a total of 6 valence electrons.

Question 20

How many valence electrons does iron, element 26, have?

Answer 20

8

Iron is the 8th element in its row. Its valence shell configuration is 4s²3d⁶, for a total of 8 valence electrons.

Question 21

Define:

first ionization energy

Answer 21

It is the energy required to remove one valence electron from an atom in the gas phase.

The generic ionization energy equation is:

X(g) ⇒ X⁺(g) + e^-

Question 22

Describe the general trend of ionization energy across a row of the periodic table.

Answer 22

It increases from left to right across a row of the periodic table.

Other notes about ionization energy:

• Atoms with fully-filled subshell will have high ionization energies.
• Atoms with half-filled subshells will have higher ionization energies than their neighbors.
• The alkali and alkaline earth metals have very low ionization energies.

Question 23

Which has a higher first ionization energy, Cl or Br?

Answer 23

Cl

Remember that ionization energy decreases going down a column. Br is below Cl in the halogen column.

Question 24

Describe the general trend of ionization energy heading down a column of the periodic table.

Answer 24

Ionization energy decreases heading down a column of the periodic table.

The further down a column an element lies, the easier to remove. These atoms have higher n values for their valence electrons. Higher n electrons sit further from the atomic nucleus, and are therefore less strongly bound to the nucleus.

Question 25

Which has a higher first ionization energy, Si or P?

Answer 25

P ## Footnote Remember that ionization energy increases from left-to-right across a column. P is to the right of Si in period 3.

Question 26

What is an atom's **second ionization energy**?

Answer 26

It is the energy required to remove a subsequent (second) valence electron from a singly-charged ion in the gas phase. ## Footnote The generic second ionization energy equation is: X⁺(g) ⇒ X²⁺(g) + e^-

Question 27

What are the relative magnitudes of any atom's first and second ionization energies?

Answer 27

The second ionization energy is always larger in magnitude than the first ionization energy for every atom. ## Footnote The removal of the first electron reduces the electron-electron repulsion energy of the molecule, allowing the positive nucleus to attract the remaining electrons more strongly and increasing the energy needed to remove subsequent electrons.

Question 28

How does atomic radius vary as atomic shell increases down a column of the periodic table?

Answer 28

Atomic radius increases down a column of the periodic table. ## Footnote Each increasing shell can be thought of as another "layer" of electrons, outside the previous layer, increasing the atom's size. Below. the highlighted elements represent one column of the Periodic Table.

Question 29

Which has a larger radius, Cl or Br?

Answer 29

Br ## Footnote Remember that atomic radius increases going down a column, and Br is below Cl in the halogens column. The valence electrons from Br are in the *n*=4 shell, those from Cl are in the *n*=3 shell.

Question 30

How does **atomic radius** vary as atomic number increases across a row of the Periodic Table?

Answer 30

It decreases across a row from left to right on the periodic table. ## Footnote This can be explained by effective nuclear charge. As Z_eff increases, the nucleus binds the electrons more tightly, pulling them in closer. Below, the highlighted elements represent one full row of the Periodic Table.

Question 31

Which has a larger radius, Si or P?

Answer 31

Si ## Footnote Remember that atomic radius decreases from left to right across a column, and P is to the right of Si in the third period.

Question 32

# Define: electron affinity

Answer 32

It is the energy released when one valence electron is added to an atom in the gas phase. ## Footnote The generic electron affinity equation is: X(g) + e^- ⇒ X^-(g)

Question 33

Describe the general trend of electron affinity across a row of the periodic table.

Answer 33

Electron affinity increases from left to right across a row. ## Footnote The smaller an atom is, the closer a newly-added valence electron gets to the positively-charged nucleus, and the more energy released when that electron is added.

Question 34

Describe the general trend of electron affinity down a column of the periodic table.

Answer 34

Electron affinity decreases down a column. ## Footnote The further down a column an element lies, the higher the value of *n* for its valence electrons. Higher-*n* electrons sit further from the atomic nucleus, and so are less bound to the nucleus and release less energy when added.

Question 35

# Define: electronegativity

Answer 35

It describes that atom's tendency to attract electron density towards itself through a chemical bond. ## Footnote Note that electronegativity only applies to atoms in a bond. There is no such concept as the electronegativity of a bare atomic species, although certain elements are said to be "more electronegative" due to their behavior when bonded.

Question 36

Describe the general trend of electronegativity across a row of the periodic table.

Answer 36

Electronegativity increases from left to right across a row. ## Footnote The smaller an atom is, the closer the electrons in its bonds get to the positively-charged nucleus, and the more strongly the electrons are attracted to the nucleus.

Question 37

Describe the general trend of electronegativity heading down a column of the periodic table.

Answer 37

Electronegativity decreases heading down a column of the periodic table. ## Footnote The further down a column an element lies, the larger its radius. This puts more space between the positively-charged nucleus and the electrons in any bonds it makes, reducing the attraction between the nucleus and the electrons.

Question 38

Which element has the highest electronegativity?

Answer 38

Fluorine ## Footnote Electronegativity increases toward the top right of the periodic table. Fluorine, which is the rightmost and uppermost element that isn't a noble gas, has the greatest tendency to be electronegative. On the Pauling scale, the most commonly-used scale for determining electronegativity values, fluorine has the highest value possible: 4.0.

Question 39

What class of elements have the highest values of electronegativity?

Answer 39

Nonmetals ## Footnote Fluorine is the element with the highest electronegativity, and as a general rule, the closer an element is to fluorine, the higher its electronegativity. Some other notes about electronegativity: * The halogens are the most electronegative group. * Noble gases capable of making bonds (Xe and Kr) are relatively electronegative. * Metals, particularly the alkali and alkali earth metals, are generally electropositive (very low electronegativity).

Question 40

Which has a higher electronegativity, Cl or Br?

Answer 40

Cl ## Footnote Remember that electronegativity decreases going down a column, and Br is below Cl in the halogens column.

Question 41

Which has a higher electronegativity, Si or P?

Answer 41

P ## Footnote Remember that electronegativity increases from left to right across a column, and P is to the right of Si in period 3.