AP Chemistry 1.7-1.8: Periodic Trends and Valence Electrons

Why do elements in the same periodic-table group tend to have similar chemical behavior?

They have similar valence-shell electron configurations, including the same number of valence electrons for main-group elements.

How do atomic size and metallic character generally change across periods and down groups?

Atomic size generally decreases from left to right across a period and increases down a group. Metallic character decreases across a period and increases down a group.

Covalent radius

One-half the distance between the nuclei of two identical atoms joined by a covalent bond. It is a practical measure of atomic size.

Effective nuclear charge, $Z_{\mathrm{eff}}$

The net attractive force that the nucleus exerts on a particular electron after accounting for electron shielding and electron–electron repulsions. It is often approximated by $Z_{\mathrm{eff}} = Z - \text{shielding}$.

Which electrons provide the most effective shielding, and why?

Core electrons shield valence electrons effectively because they are located between the nucleus and the valence shell. Electrons in the same valence shell shield one another less effectively.

Why does atomic radius increase down a group?

Each successive period places the valence electrons in a shell with a higher principal quantum number, $n$, farther from the nucleus. Increased distance and shielding make the atom larger.

Why does atomic radius generally decrease from left to right across a period?

The nuclear charge increases as protons are added, while shielding increases only slightly because the added electrons occupy the same principal shell. Therefore, $Z_{\mathrm{eff}}$ increases and pulls the electrons closer to the nucleus.

Rank the following atoms from smallest to largest covalent radius: Kr, Br, Ge, and Fl.

The order is $\mathrm{Kr < Br < Ge < Fl}$. Radius decreases across a period and increases down a group.

How does forming a cation affect an atom's radius?

A cation is smaller than its parent atom because it has lost electrons, often its entire outer shell. The remaining electrons experience a stronger effective nuclear attraction.

How does forming an anion affect an atom's radius?

An anion is larger than its parent atom because added electrons increase electron–electron repulsions and reduce the effective nuclear attraction per electron.

How does ionic radius change as the positive charge of a cation increases?

For cations of the same element, a higher positive charge produces a smaller radius because fewer electrons experience a stronger attraction to the unchanged nucleus. For example, $\mathrm{V^{3+}}$ is smaller than $\mathrm{V^{2+}}$.

How do the radii of same-charge cations generally change down a group?

They generally increase because the valence electrons occupy shells with progressively larger principal quantum numbers.

Isoelectronic species

Atoms or ions that have the same number of electrons and therefore the same electron configuration, such as $\mathrm{N^{3-}}$, $\mathrm{O^{2-}}$, $\mathrm{F^-}$, Ne, and $\mathrm{Na^+}$.

How can the sizes of isoelectronic species be ranked?

The species with more protons is smaller because the same number of electrons is attracted by a greater nuclear charge. For example, $\mathrm{N^{3-} > O^{2-} > F^- > Ne > Na^+ > Mg^{2+} > Al^{3+}}$ in order of decreasing radius.

First ionization energy, $\mathrm{IE_1}$

The energy required to remove the most loosely bound electron from a gaseous atom in its ground state: $\mathrm{X(g) \rightarrow X^+(g) + e^-}$. Ionization is endothermic, so ionization energies are positive.

What is the general periodic trend in first ionization energy?

First ionization energy generally increases from left to right across a period and decreases down a group.

Why does first ionization energy generally increase across a period?

Across a period, increasing $Z_{\mathrm{eff}}$ attracts the valence electrons more strongly and decreases atomic radius, making electron removal more difficult.

Why does first ionization energy generally decrease down a group?

The valence electron is in a higher-energy shell, farther from the nucleus, and more shielded by core electrons, so it is easier to remove.

Why is the first ionization energy of boron lower than that of beryllium?

Beryllium loses a $2s$ electron, whereas boron loses a higher-energy $2p$ electron. The $2p$ electron is easier to remove despite boron's greater nuclear charge.

Why is the first ionization energy of oxygen slightly lower than that of nitrogen?

Nitrogen has a relatively stable half-filled $2p^3$ subshell. Oxygen has one paired $2p$ orbital, so electron–electron repulsion makes one of its electrons easier to remove.

Successive ionization energies

The energy required to remove additional electrons from an increasingly positive ion. They always increase because each remaining electron is attracted more strongly to the nucleus.

What does a large jump between successive ionization energies indicate?

It indicates that all valence electrons have been removed and the next electron comes from a core shell. The position of the jump helps identify the number of valence electrons.

An element has a very large jump between its third and fourth ionization energies. What does this suggest about its valence electrons?

The atom most likely has three valence electrons. The first three electrons are relatively easier to remove, while the fourth is a core electron.

Why is removing an electron from a cation more difficult than removing one from its neutral atom?

The cation has fewer electrons and a greater electrostatic attraction per remaining electron. Removing an electron from a more highly charged cation is even more difficult.

Electron affinity, EA

The energy change when a gaseous atom gains an electron: $\mathrm{X(g) + e^- \rightarrow X^-(g)}$. It may be exothermic or endothermic, so its value may be negative or positive.

What does a negative electron affinity value mean?

Energy is released when the gaseous atom gains an electron; the process is exothermic.

What is the general trend in first electron affinity from left to right across a period?

Electron affinities generally become more negative because increasing $Z_{\mathrm{eff}}$ makes the atom more attractive to an added electron.

Why do group 2 elements show an exception to the general electron-affinity trend?

They have filled $ns$ subshells, so an added electron must enter a higher-energy $np$ subshell, which is less favorable.

Why do group 15 elements show an exception to the general electron-affinity trend?

Their $np^3$ subshell is half-filled and relatively stable. An added electron must pair with another electron in a $p$ orbital, increasing repulsion.

Why do noble gases generally have unfavorable electron affinities?

Their valence shells are filled, so an added electron must enter a higher principal energy level rather than the already filled shell.

Why is chlorine's electron affinity more negative than fluorine's?

The added electron enters chlorine's larger $n = 3$ valence shell, where electron–electron repulsions are lower than in fluorine's compact $n = 2$ shell.

How do atomic radius and metallic character relate to the ease of removing electrons?

Larger atoms and metals hold their valence electrons less tightly, so they lose electrons more readily. Consequently, metallic character increases down a group and decreases across a period.

How does the periodic table help predict common main-group ion charges?

Group 1 metals commonly form $1+$ ions, group 2 metals form $2+$ ions, and group 13 metals often form $3+$ ions. Groups 17, 16, and 15 nonmetals commonly form $1-$, $2-$, and $3-$ ions, respectively, by gaining electrons toward a noble-gas configuration.

Cation versus anion

A cation is a positively charged ion formed by losing electrons; an anion is a negatively charged ion formed by gaining electrons. The element's identity remains unchanged because its number of protons does not change.

What ion is formed by an atom with 13 protons and 10 electrons?

It is $\mathrm{Al^{3+}}$, the aluminum ion. The three fewer electrons than protons give it a $3+$ charge.

What ion is formed by an atom with 34 protons and 36 electrons?

It is $\mathrm{Se^{2-}}$, the selenide ion. The two extra electrons give it a $2-$ charge.

How are transition-metal ion charges different from typical main-group ion charges?

Transition metals commonly form ions with multiple possible charges, so their charge often cannot be predicted from group position alone. For example, iron forms $\mathrm{Fe^{2+}}$ and $\mathrm{Fe^{3+}}$, while copper forms $\mathrm{Cu^+}$ and $\mathrm{Cu^{2+}}$.

Why does fluorine preferentially form an anion rather than a cation?

Fluorine has a favorable tendency to gain an electron, reflected by its favorable electron affinity, while removing an electron requires a large ionization energy. Thus, forming $\mathrm{F^-}$ is much easier than forming a positive fluorine ion.

Monatomic ion versus polyatomic ion

A monatomic ion consists of one atom with a net charge, such as $\mathrm{Na^+}$. A polyatomic ion is a covalently bonded group of atoms that behaves as one charged unit, such as $\mathrm{SO_4^{2-}}$.

What do the suffixes -ate and -ite indicate in oxyanion names?

For related oxyanions of the same central element, -ate generally indicates more oxygen atoms and -ite indicates fewer. The exact number of oxygen atoms depends on the element.

How are perchlorate, chlorate, chlorite, and hypochlorite related?

They form a series with decreasing oxygen content: $\mathrm{ClO_4^-}$ is perchlorate, $\mathrm{ClO_3^-}$ is chlorate, $\mathrm{ClO_2^-}$ is chlorite, and $\mathrm{ClO^-}$ is hypochlorite.

What are the formulas and charges of nitrate, nitrite, sulfate, and sulfite?

Nitrate is $\mathrm{NO_3^-}$, nitrite is $\mathrm{NO_2^-}$, sulfate is $\mathrm{SO_4^{2-}}$, and sulfite is $\mathrm{SO_3^{2-}}$.

What are the formulas and charges of carbonate, bicarbonate, phosphate, and ammonium?

Carbonate is $\mathrm{CO_3^{2-}}$, bicarbonate is $\mathrm{HCO_3^-}$, phosphate is $\mathrm{PO_4^{3-}}$, and ammonium is $\mathrm{NH_4^+}$.

Ionic bond

The electrostatic attraction between oppositely charged ions, usually formed after electron transfer between a metal and a nonmetal.

Covalent bond

The attraction between positively charged nuclei and one or more shared pairs of electrons. Covalent bonding commonly produces molecular compounds between nonmetals.

How can the type of compound be predicted from the elements involved?

A metal combined with a nonmetal generally forms an ionic compound containing cations and anions. Two nonmetals generally form a molecular compound containing covalent bonds.

How is the formula of a simple ionic compound determined?

Choose subscripts so the total positive and negative charges cancel, giving an electrically neutral compound. For example, $\mathrm{Mg^{2+}}$ and $\mathrm{N^{3-}}$ form $\mathrm{Mg_3N_2}$.

What ions and formula result when aluminum reacts with oxygen?

Aluminum forms $\mathrm{Al^{3+}}$ and oxygen forms $\mathrm{O^{2-}}$. Charge balance requires two aluminum ions and three oxide ions, giving $\mathrm{Al_2O_3}$.