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General Chemistry Study Guide

Chapter 8. The Periodic Table

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Yu Wang

OpenStax 6.5 Periodic Variations in Element Properties. Brown 7 Periodic Properties of the Elements.

1. Electron Configuration and Periodic Classification

1.1. Electron Configuration of Elements

According to the type of subshell being filled, the elements can be divided into categories—the representative elements, the noble gases, the transition elements (or transition metals), the lanthanides, and the actinides, as shown in the following Figure.

The representative elements (also called main group elements) are the elements in Groups 1A through 7A, all of which have incompletely filled $s$ or $p$ subshells of the highest principal quantum number.

With the exception of helium, the noble gases (the Group 8A elements) all have a completely filled $p$ subshell. (The electron configurations are $1s^2$ for helium and $ns^2np^6$ for the other noble gases, in which $n$ is the principal quantum number for the outermost shell.)

The transition metals are the elements in Groups 1B and 3B through 8B, which have incompletely filled $d$ subshells or readily produce cations with incompletely filled $d$ subshells.

The Group 2B elements are Zn, Cd, and Hg, which are neither representative elements nor transition metals.

The lanthanides and actinides are sometimes called f-block transition elements because they have incompletely filled $f$ subshells.

The ground-state electron configurations of the elements are shown in the following Figure.

The outer electrons of an atom, which are those involved in chemical bonding, are often called the valence electrons. Having the same number of valence electrons accounts for similarities in chemical behavior among the elements within each of these groups.

1.2. Electron Configuration of Ions

Ions Derived from Representative Elements

In the formation of a cation from the neutral atom of a representative element, one or more electrons are removed from the highest occupied $n$ shell so that cation has a noble-gas outer electron configuration.

For example,
$\ce{Na}$: $[\ce{Ne}]3s^1$
$\ce{Na^+}$: $[\ce{Ne}]$

In the formation of an anion, one or more electrons are added to the highest partially filled $n$ shell so that cation has a noble-gas outer electron configuration.

For example,
$\ce{F}$: $[\ce{He}]2p^5$
$\ce{F^-}$: $[\ce{Ne}]$

Isoelectronic: have the same number of electrons, and hence the same ground-state electron configuration

Cations Derived from Transition Metals

When a cation is formed from an atom of a transition metal, electrons are always removed first from the $ns$ orbital and then from the $(n − 1)d$ orbitals.

For example,
$\ce{Fe}$: $[\ce{Ar}]4s^23d^6$
$\ce{Fe^{2+}}$: $[\ce{Ar}]4s^03d^6$
$\ce{Fe^{3+}}$: $[\ce{Ar}]4s^03d^5$

Requirements

1. Understand the classification of elements.
2. Understand the valence electrons.
3. Understand the electron configuration chance upon ion formation.

2. Effective Nuclear Charge and Size

2.1. Effective Nuclear Charge

The effective nuclear charge ($Z_{eff}$) is the nuclear charge felt by an electron when both the actual nuclear charge ($Z$) and the repulsive effects (shielding) of the other electrons are taken into account.

The increase in effective nuclear charge from left to right across a period (thus the nuclear holds the out most layer electrons stronger) and from top to bottom in a group for representative elements (but in this case the nuclear does not hold the out most layer electrons stronger because the distance between the nuclear and the electrons is larger).

2.2. Atomic Radius

We define the size of an atom in terms of its atomic radius, which is one-half the distance between the two nuclei in two adjacent metal atoms. For elements that exist as diatomic molecules, such as iodine, the radius of the atom is defined as one-half the distance between the centers of the atoms in the molecules.

Trend for representative elements:

2.3. Ionic Radius

Ionic radius is the radius of a cation or an anion.

For isoelectronics, i.e. ions and atoms have the same electron configuration, the greater the nuclear charge is, the smaller the ionic or atomic radius would be, because the nuclius holds the electrons stronger.

For example, $\ce{N^{3-}}$, $\ce{O^{2-}}$, $\ce{F^-}$, $\ce{Ne}$, $\ce{Na^+}$, $\ce{Mg^{2+}}$, and $\ce{Al^{3+}}$ have the same electron configuration. The order of the sizes is $\ce{N^{3-}}$>$\ce{O^{2-}}$>$\ce{F^-}$>$\ce{Ne}$>$\ce{Na^+}$>$\ce{Mg^{2+}}$> $\ce{Al^{3+}}$

For the same nuclius, i.e. the same number of protons, the greater the electron number is, the larger the ionic or atomic radius would be, because the effect that electrons repulse each other is more significant.

For example, $\ce{Fe^{3+}}$<$\ce{Fe^{2+}}$<$\ce{Fe}$.

The radii (in picometers) of ions of familiar elements arranged according to the elements' positions in the periodic table:

Requirements

1. Understand effective nuclear charge.
2. Learn how to compare the sizes of atoms and ions.

3. Ionization Energy and Electron Affinity

3.1. Ionization Energy

Ionization energy (IE) is the minimum energy (in kJ/mol) required to remove an electron from a gaseous atom in its ground state.

To remove the first electron -> first ionization energy ($IE_1$);
To remove the second electron -> second ionization energy ($IE_2$);
ect.

$$IE_1 < IE_2 < IE_3 < ...$$

The increase in first ionization energy from left to right across a period and from bottom to top in a group for representative elements.

3.2. Electron Affinity

Electron affinity (EA) is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion.

Except noble gases, which have close to zero or negtive electron affinity, the increase in electron affinity has the same trend as that of first ionization energy.

Requirements

1. Remember the trend of increasing first ionization energy and increasing electron affinity for representative elements.

4. A Survy of Chemical Properties of Elements

4.1. General Trends

1. The same group of elements have the same valence electron configuration; thus, there chemical properties must have something in common.
2. Another trend in the chemical behavior of the representative elements is the diagonal relationship. Diagonal relationships are similarities between pairs of elements in different groups and periods of the periodic table. Specifically, the first three members of the second period (Li, Be, and B) exhibit many similarities to those elements located diagonally below them in the periodic table. The reason for this phenomenon is the closeness of the charge densities of their cations.

Bear in mind that a comparison of the properties of elements in the same group is most valid if we are dealing with elements of the same type with respect to their metallic character. This guideline applies to the elements in Groups 1A and 2A, which are all metals, and to the elements in Groups 7A and 8A, which are all nonmetals. In Groups 3A through 6A, where the elements change either from nonmetals to metals or from nonmetals to metalloids, it is natural to expect greater variation in chemical properties even though the members of the same group have similar outer electron configurations.

4.2. Survy of the Representative Elements

Hydrogen

Hydrogen has only one electron. It can lose one electron to form a cation $\ce{H+}$. For example, $\ce{H}$ in $\ce{HF}$, $\ce{HCl}$, $\ce{HBr}$, etc.

It can also gain one extra electron to fill the $1s$ orbital and become $\ce{H-}$. For example, $\ce{H}$ in $\ce{NaH}$, $\ce{KH}$, $\ce{CaH2}$, etc.

Group 1A Elements

All of these elements, the alkali metals, have low ionization energies and therefore a great tendency to lose the single valence electron. These metals are so reactive that they are never found in the pure state in nature. They react with water to produce hydrogen gas and the corresponding metal hydroxide. When exposed to air, they gradually lose their shiny appearance as they combine with oxygen gas to form oxides.

Group 2A Elements

The alkaline earth metals are somewhat less reactive than the alkali metals. They can form $\ce{M^{2+}}$ ions (where $\ce{M}$ denotes an alkaline earth metal atom). The metallic character and activity increases from top to bottom.

Beryllium does not react with water; magnesium reacts slowly with steam; calcium, strontium, and barium are reactive enough to attack cold water.

Group 3A Elements

The first member of Group 3A, boron, is a metalloid; the rest are metals. Boron does not form binary ionic compounds and is unreactive toward oxygen gas and water. The next element, aluminum, readily forms aluminum oxide when exposed to air.

Aluminum that has a protective coating of aluminum oxide is less reactive than elemental aluminum. Aluminum forms only tripositive ions.

The other Group 3A metallic elements form both unipositive and tripositive ions. Moving down the group, we find that the unipositive ion becomes more stable than the tripositive ion.

From left to right across the periodic table, we are seeing a gradual shift from metallic to nonmetallic character in the representative elements.

Group 4A Elements

The first member of Group 4A, carbon, is a nonmetal, and the next two members, silicon and germanium, are metalloids. The metallic elements of this group, tin and lead, do not react with water, but they do react with acids (hydrochloric acid, for example) to liberate hydrogen gas.

The Group 4A elements form compounds in both the +2 and +4 oxidation states. For carbon and silicon, the +4 oxidation state is the more stable one.

In tin compounds, the +4 oxidation state is only slightly more stable than the +2 oxidation state. In lead compounds, the +2 oxidation state is unquestionably the more stable one.

Group 5A Elements

Nitrogen and phosphorus are nonmetals, arsenic and antimony are metalloids, and bismuth is a metal.

Elemental nitrogen is a diatomic gas ($\ce{N2}$). It forms a number of oxides ($\ce{NO}$, $\ce{N2O}$, $\ce{NO2}$, $\ce{N2O4}$, and $\ce{N2O5}$). Nitrogen also has a tendency to accept three electrons to form the nitride ion, $\ce{N^{3−}}$.

Phosphorus exists as $\ce{P4}$ molecules. It forms two solid oxides with the formulas $\ce{P4O6}$ and $\ce{P4O10}$. The important oxoacids $\ce{HNO3}$ and $\ce{H3PO4}$ are formed when the following oxides react with water.

Group 6A Elements

The first three members of Group 6A (oxygen, sulfur, and selenium) are nonmetals, and the last two (tellurium and polonium) are metalloids.

Oxygen is a diatomic gas; elemental sulfur and selenium have the molecular formulas $\ce{S8}$ and $\ce{Se8}$, respectively.

Oxygen has a tendency to accept two electrons to form the oxide ion ($\ce{O^{2−}}$) in many ionic compounds. Sulfur, selenium, and tellurium also form dinegative anions ($\ce{S^{2−}}$, $\ce{Se^{2−}}$, and $\ce{Te^{2−}}$).

The elements in this group (especially oxygen) form a large number of molecular compounds with nonmetals. The important compounds of sulfur are $\ce{SO2}$, $\ce{SO3}$, and $\ce{H2S}$. Sulfuric acid is formed when sulfur trioxide reacts with water.

Group 7A Elements

All the halogens are nonmetals with the general formula $\ce{X2}$, where $\ce{X}$ denotes a halogen element. Because of their great reactivity, the halogens are never found in the elemental form in nature. Fluorine is so reactive that it attacks water to generate oxygen.

The halogens have high ionization energies and large positive electron affinities. Anions derived from the halogens ($\ce{F^{−}}$, $\ce{Cl^{−}}$, $\ce{Br^{−}}$, and $\ce{I^{−}}$) are called halides. They are isoelectronic with the noble gases immediately to their right in the periodic table.

The vast majority of the alkali metal halides and alkaline earth metal halides are ionic compounds. The halogens also form many molecular compounds among themselves (such as $\ce{ICl}$ and $\ce{BrF3}$) and with nonmetallic elements in other groups (such as $\ce{NF3}$, $\ce{PCl5}$, and $\ce{SF6}$). The halogens react with hydrogen to form hydrogen halides.

The hydrogen halides dissolve in water to form hydrohalic acids. Hydrofluoric acid ($\ce{HF}$) is a weak acid (that is, it is a weak electrolyte), but the other hydrohalic acids ($\ce{HCl}$, $\ce{HBr}$, and $\ce{HI}$) are all strong acids (strong electrolytes).

Group 8A Elements

All noble gases exist as monatomic species. Their atoms have completely filled outer $ns$ and $np$ subshells, which give them great stability. The Group 8A ionization energies are among the highest of all elements, and these gases have no tendency to accept extra electrons. For years these elements were called inert gases, and rightly so.

4.3. Properties of Oxides Across a Period

We observed earlier that oxygen has a tendency to form the oxide ion. This tendency is greatly favored when oxygen combines with metals that have low ionization energies, namely, those in Groups 1A and 2A, plus aluminum. Ionic compounds have high melting points and boiling points.

As the ionization energies of the elements increase from left to right, so does the molecular nature of the oxides that are formed. Silicon is a metalloid; its oxide ($\ce{SiO2}$) also has a huge three-dimensional network, although no ions are present.

The oxides of phosphorus, sulfur, and chlorine are molecular compounds composed of small discrete units. The weak attractions among these molecules result in relatively low melting points and boiling points.

Most oxides can be classified as acidic or basic depending on whether they produce acids or bases when dissolved in water or react as acids or bases in certain processes. Some oxides are amphoteric, e.g. $\ce{Al2O3}$, which means that they display both acidic and basic properties.

Metal oxides are generally basic because they can react with acid to form salt and water. For example, $\ce{Na2O}$ and $\ce{MgO}$

Non-metal oxides are generally non-basic. Some of them are acidic bacause they can react with base to form salt and water, e.g. $\ce{SO3}$, $\ce{Cl2O7}$, $\ce{SiO2}$. Some of them are neutral, e.g. $\ce{CO}$, $\ce{NO}$.

An example of amphoteric oxide is $\ce{Al2O3}$ which can react with both acids and bases.

Requirements

1. Be familiar with the general properties of each group.
2. Understand basic and acidic oxides. Remember $\ce{Al2O3}$ is amphoteric.

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