1.7 Periodic Trends · 1.8 Valence Electrons and Ionic Compounds

1.7 Periodic Trends · 1.8 Valence Electrons and Ionic Compounds

The periodic table is not merely an index of elements: its position predicts how strongly an atom holds electrons, how large it is, and what ions it tends to form. The central question is: How can an element’s location predict both its microscopic bonding behavior and the macroscopic properties of its compounds?

The shell model behind periodicity

The shell model places electrons in principal energy levels, or shells, around the nucleus. For main-group elements, moving across a period generally adds electrons to the same principal shell while the number of protons increases. Moving down a group adds a new occupied shell, so the valence electrons are farther from the nucleus.

This simple picture requires an important qualification. In the transition series, electrons can enter an inner $(n-1)d$ subshell rather than the newest outer shell; for example, period-4 transition elements fill the $3d$ subshell while the $4s$ subshell is also involved. Therefore, the main-group shell model predicts broad trends especially well, while transition-element bonding may involve both $ns$ and incompletely filled $(n-1)d$ electrons.

Why periodic trends occur

Coulomb’s law states that the electrostatic attraction between opposite charges becomes stronger when charge magnitudes increase and weaker when distance increases:

$$F = k\frac{|q_1q_2|}{r^2}$$

Inner electrons partially block the nucleus from attracting outer electrons. This effect is called shielding. The net positive attraction experienced by an electron is the effective nuclear charge, often represented as $Z_{\text{eff}}$. Across a main-group period, nuclear charge increases while shielding changes less dramatically because electrons are added mainly to the same shell. Thus, $Z_{\text{eff}}$ generally increases.

Down a group, additional shells increase both distance and shielding. Even though the nucleus contains more protons, the outer electron is usually held less tightly because the $r^2$ effect and shielding outweigh the increased nuclear charge.

Property	Across a period, left to right	Down a group	Why
Atomic radius	Generally decreases	Generally increases	Nuclear attraction versus shell distance
Ionization energy	Generally increases	Generally decreases	Energy required to remove an electron
Electronegativity	Generally increases	Generally decreases	Attraction for bonding electrons
Electron affinity	Generally becomes more favorable overall	Generally less favorable overall	Energy change when an electron is added
Ionic radius	Cations are smaller; anions are larger than their atoms	Generally increases	Electron loss or gain plus shell structure

Ionization energy is the energy required to remove an electron from a gaseous atom. A smaller radius and larger $Z_{\text{eff}}$ make removal harder. Electronegativity describes how strongly an atom attracts shared electrons in a bond. Electron affinity describes the energy change when a gaseous atom gains an electron; its numerical sign convention must be read carefully because “more favorable” does not always mean “larger positive number.”

Periodic trends are predictions, not rigid laws. Small exceptions occur when subshell energies and electron pairing alter stability. For example, the first ionization energy of boron is slightly lower than that of beryllium because boron’s removed electron is in a higher-energy $2p$ subshell; oxygen’s is slightly lower than nitrogen’s because oxygen has paired electrons in a $2p$ orbital. The CED does not require writing electron configurations for exceptions to the aufbau principle, but it does require reasoning from models such as shielding and $Z_{\text{eff}}$.

Valence electrons and typical ionic charges

Valence electrons are electrons available for bonding. For main-group elements, these are usually the electrons in the outermost occupied shell. For transition elements, valence electrons can include electrons in an incompletely filled inner $(n-1)d$ subshell as well as $ns$ electrons.

Main-group atoms often form ions that achieve a more stable electron arrangement. Their typical charges follow their periodic-table columns:

Group	Typical ion	Electron change
Group 1	$+1$	Loses one valence electron
Group 2	$+2$	Loses two
Group 13	$+3$	Often loses three
Group 15	$-3$	Gains three
Group 16	$-2$	Gains two
Group 17	$-1$	Gains one
Group 18	$0$	Usually does not form simple monatomic ions

This is Essential Knowledge 1.8.A.3: typical ionic charges are governed by the number of valence electrons and predicted by an element’s location on the periodic table. Essential Knowledge 1.8.A.2 extends the pattern: elements in the same column tend to form analogous compounds because they have related valence-electron arrangements.

Worked example: predicting an ionic compound

Predict the formula and bonding pattern for a compound formed from calcium and chlorine. Calcium is in Group 2, so it commonly forms $Ca^{2+}$ by losing two electrons. Chlorine is in Group 17, so each chlorine atom commonly forms $Cl^-$ by gaining one electron.

Charge balance requires one $Ca^{2+}$ for every two $Cl^-$ ions:

$$Ca^{2+} + 2Cl^- \longrightarrow CaCl_2$$

At the particulate level, $CaCl_2$ is an extended lattice of oppositely charged ions, not a collection of independent calcium-chlorine molecules. Strong Coulombic attractions throughout the lattice help explain its macroscopic properties: a high melting point, brittleness, and electrical conductivity when molten or dissolved because the ions can then move.

Misconception check: “Ionic compounds contain molecules”

A binary ionic compound contains two different elements, but its formula gives the lowest whole-number ratio of ions, not necessarily a discrete molecule. The formula unit $CaCl_2$ communicates the lattice’s $1:2$ ratio. Also, ionic solids do not conduct electricity well as solids because their ions are fixed in place; they conduct when molten or aqueous because the charged particles can move.

Science Practice 4: Model Analysis, skill code 4.A, asks students to predict or explain chemical properties or phenomena using theories, models, and representations. Here, the periodic table, shell model, Coulomb’s law, and ionic-lattice drawing work together: electron arrangement predicts charge; charge and distance predict attraction; attraction and lattice structure predict observable properties. This also develops the required connection between particulate-level structure and macroscopic behavior.

Retrieval check: Which atom should have the larger radius, sodium or chlorine, and why? Which ion is larger, $Na^+$ or $Cl^-$, and why? Finally, predict the formula of the compound formed by aluminum and oxygen. A strong answer should invoke shells, shielding, $Z_{\text{eff}}$, ionic charge, and charge neutrality rather than memorizing isolated formulas.