Chemistry · Ch 4 — Chemical Bonding and Molecular Structure

Kössel-Lewis Approach to Chemical Bonding

4.1

Kössel-Lewis Approach to Chemical Bonding

The Kössel-Lewis Approach: A Foundation for Chemical Bonding

By 1916, chemists had a puzzle: what drives atoms to form bonds? Two scientists, Walther Kössel and Gilbert N. Lewis, independently cracked the code. Their shared insight was brilliant in its simplicity — the key lay in the electron arrangements of the noble gases.

Lewis pictured an atom as a positively charged "kernel" (the nucleus plus all inner-shell electrons) surrounded by an outer shell. This outer shell, he proposed, could hold a maximum of eight electrons, which he imagined sitting at the eight corners of a cube. A sodium atom, with its single outer electron, would have that electron occupying just one corner. A noble gas like neon or argon would have all eight corners filled. That complete octet, Lewis argued, represented a state of exceptional stability.

The central postulate was this: atoms achieve the stable octet when they form chemical bonds. For sodium and chlorine, this happens by electron transfer — sodium gives up its outer electron to chlorine, producing Na⁺ and Cl⁻ ions. For molecules like Cl₂, H₂, or F₂, the bond forms by sharing a pair of electrons between atoms. In both cases, each atom ends up with a stable outer octet.

Kössel, working independently, focused on the electrostatic consequences. He noted three key facts:

In the periodic table, highly electronegative halogens and highly electropositive alkali metals are separated by the noble gases.
A halogen atom gains an electron to form a negative ion; an alkali metal loses an electron to form a positive ion.
The resulting ions both attain stable noble gas configurations (for helium, a duplet of 2 electrons; for others, an octet of $ns^2 np^6$ electrons).
These oppositely charged ions are then stabilized by electrostatic attraction.

The bond formed by this electrostatic attraction was termed the electrovalent bond. The electrovalence of an ion equals the number of unit charges on it — calcium has a positive electrovalence of two, chlorine a negative electrovalence of one.

Note

Kössel's work laid the foundation for understanding ionic compounds. However, he recognized that many compounds — like Cl₂ or H₂ — simply did not fit the electron-transfer model. This limitation set the stage for Lewis's complementary theory of shared-electron bonds.

Lewis Symbols: Representing Valence Electrons

Only the outermost electrons — the valence electrons — participate in chemical bonding. The inner-shell electrons remain protected and uninvolved. G.N. Lewis introduced a simple notation to represent these valence electrons: Lewis symbols (also called electron-dot symbols).

For elements of the second period, the Lewis symbols are:

Element	Lewis Symbol	Number of Valence Electrons
Li	$\text{Li} \cdot$	1
Be	$\cdot \text{Be} \cdot$	2
B	$\cdot \text{B} \cdot$	3
C	$\cdot \dot{\text{C}} \cdot$	4
N	$\cdot \dot{\text{N}} \cdot$	5
O	$\cdot \ddot{\text{O}} \cdot$	6
F	$\cdot \ddot{\text{F}} \cdot$	7
Ne	$\cdot \ddot{\text{Ne}} \cdot$	8

The number of dots around the symbol directly represents the number of valence electrons. This number helps determine the common or group valence of the element. The group valence is generally either equal to the number of dots in the Lewis symbol, or equal to $8$ minus the number of dots (or valence electrons).

Tip

For main-group elements, the group number (1, 2, 13–18) tells you the number of valence electrons directly. Group 1 has 1 valence electron, Group 2 has 2, Group 13 has 3, and so on up to Group 18 (noble gases) which have 8 (except helium with 2).

The Octet Rule and Bond Formation

The driving force behind chemical bonding, according to Lewis, is the tendency of atoms to achieve a stable outer shell of eight electrons — the octet. For hydrogen and helium, the stable configuration is a duplet of two electrons (the $1s^2$ configuration of helium).

This octet can be achieved in two ways:

1. Electron Transfer (Ionic Bonding): One atom loses one or more electrons, and another atom gains them. The resulting ions are held together by electrostatic attraction.

2. Electron Sharing (Covalent Bonding): Two atoms share one or more pairs of electrons. Each shared pair counts toward the octet of both atoms.

Formation of Ionic Compounds: The Kössel Scheme

Kössel's scheme explains the formation of ionic compounds step by step. Let's work through the examples the textbook provides.

Example 1: Sodium Chloride (NaCl)

Step 1: Sodium loses one electron to form a sodium ion.

\text{Na} \rightarrow \text{Na}^+ + e^-

Sodium's electron configuration: $[\text{Ne}]\,3s^1$

Sodium ion's configuration: $[\text{Ne}]$ (octet achieved)

Step 2: Chlorine gains one electron to form a chloride ion.

\text{Cl} + e^- \rightarrow \text{Cl}^-

Chlorine's electron configuration: $[\text{Ne}]\,3s^2 3p^5$

Chloride ion's configuration: $[\text{Ne}]\,3s^2 3p^6$ or $[\text{Ar}]$ (octet achieved)

Step 3: The oppositely charged ions are held together by electrostatic attraction.

\text{Na}^+ + \text{Cl}^- \rightarrow \text{Na}^+\text{Cl}^- \quad \text{or simply} \quad \text{NaCl}

Watch out

The formula NaCl does not represent a discrete molecule. It represents the simplest ratio of ions in the entire crystal lattice. Solid sodium chloride is a three-dimensional array of Na⁺ and Cl⁻ ions, not a collection of NaCl "molecules."

Example 2: Calcium Fluoride (CaF₂)

Step 1: Calcium loses two electrons to form a calcium ion.

\text{Ca} \rightarrow \text{Ca}^{2+} + 2e^-

Calcium's electron configuration: $[\text{Ar}]\,4s^2$

Calcium ion's configuration: $[\text{Ar}]$ (octet achieved)

Step 2: Each fluorine atom gains one electron to form a fluoride ion.

\text{F} + e^- \rightarrow \text{F}^-

Fluorine's electron configuration: $[\text{He}]\,2s^2 2p^5$

Fluoride ion's configuration: $[\text{He}]\,2s^2 2p^6$ or $[\text{Ne}]$ (octet achieved) …

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