Chemical Bonding and Molecular Structure · Lesson 1 of 12
Why Atoms Bond: Lewis Symbols and Covalent Bonds
“Discover how valence electrons help explain electron transfer, electron sharing, and the formation of chemical bonds.”
• Explain chemical bonding in terms of attraction and a decrease in the energy of the system. • Draw Lewis symbols and relate their dots to valence electrons and common valency. • Distinguish electron transfer from electron sharing using simple compounds. • Recognise single, double, and triple covalent bonds and apply the octet or duplet rule.
From separate atoms to a bonded system
A piece of salt, a drop of water, and the oxygen in air contain atoms held together in different ways. The atoms are not joined by tiny physical hooks. Their electrons and positively charged nuclei interact, and some arrangements of particles are more stable than the separated particles. Chemical bonding describes the attractions responsible for those arrangements.
An attractive interaction that holds atoms, ions, or other constituents together in a chemical species.
A stable bonded arrangement has a lower energy than the appropriate separated constituents. Energy is released when a stable bond forms and must be supplied to separate the constituents again. Atoms do not consciously “want” a full shell; the full-shell pattern is a useful observation connected with stability, not an intention or a force by itself.
Noble gases provided an important clue. Their outer electronic arrangements are particularly stable: helium has two electrons in its first shell, while the other noble gases have eight outer-shell electrons. Kössel and Lewis used these patterns to explain many ordinary compounds. Their models were early steps towards the more detailed orbital explanations developed later.
Valence electrons and Lewis symbols
The electrons in the outermost occupied shell are called valence electrons. For the main-group elements considered here, they provide a useful starting point for describing bonding. Inner-shell electrons usually remain associated with the atom and are not shown in a Lewis symbol.
The chemical symbol of an element surrounded by dots representing its valence electrons.
Place one dot on each of four sides of a symbol before pairing dots. This is an electron-counting convention, rather than a picture of electrons resting at four fixed positions. The symbol tells us how many electrons are available; it does not show the full three-dimensional shape of an atom.
| Element | Valence-shell configuration | Dots in its Lewis symbol |
|---|---|---|
| Li | 1 | |
| Be | 2 | |
| B | 3 | |
| C | 4 | |
| N | 5 | |
| O | 6 | |
| F | 7 | |
| Ne | 8 |
For many simple main-group compounds, the common valency is the number of valence electrons or the number needed to complete eight. Thus oxygen often forms two bonds and nitrogen often forms three. This is a useful pattern, but it is not a universal rule for every element or compound.
Problem
What information should a Lewis symbol for oxygen contain?
- 1.Oxygen has atomic number 8 and electronic configuration . Its outer shell is the second shell.
- 2.Count the outer-shell electrons: . Put six dots around O, conventionally as two pairs and two single dots.
- 3.Two more electrons would complete an octet. This helps explain the common valency of two, but does not by itself specify which compound will form.
Electron transfer and the octet rule
One route to a noble-gas arrangement is electron transfer. Sodium can lose its single outer-shell electron and chlorine can gain it. The resulting particles are ions: atoms or groups with a net electrical charge. A positive ion is a cation, and a negative ion is an anion.
The sodium ion has the electron arrangement of neon, and the chloride ion has that of argon. Oppositely charged ions attract. In solid sodium chloride, this attraction acts throughout an ordered lattice containing many ions; an isolated Na–Cl pair is not the complete picture of the solid.
Many main-group atoms form bonds by gaining, losing, or sharing electrons so that eight electrons surround each atom in its valence shell.
Helium is the important starting exception to the “eight” pattern. Its first shell is full with two electrons. Hydrogen commonly forms one covalent bond and obtains a two-electron, or duplet, arrangement when the shared pair is counted.
Problem
Explain the formula of calcium fluoride using electron transfer.
- 1.Calcium has two outer-shell electrons and can form by losing both.
- 2.Each fluorine atom gains one electron to form . Two fluorine atoms are therefore needed to accept the two electrons lost by one calcium atom.
- 3.The charges balance: . The formula is , with one calcium ion for every two fluoride ions.
The signed number of elementary charge units on an ion in the ionic description. Calcium in Ca²⁺ has electrovalence +2; chloride has electrovalence −1.
Lewis’s early historical picture placed an octet at cube corners. Langmuir’s refinement retained electron-pair sharing while abandoning that fixed cubical picture and developing the covalent-bond language. Modern Lewis drawings are electron-accounting diagrams, not claims that electrons remain stationary at cube corners.
Sharing electron pairs
Electron transfer does not describe every compound. Two chlorine atoms have the same tendency to attract electrons, yet they form a stable molecule. In the Lewis model, they share an electron pair. Each chlorine contributes one electron to this pair and counts both shared electrons when checking its octet.
A bond described in the Lewis model by the sharing of an electron pair between two atoms.
The three unshared pairs on each chlorine are lone pairs. A line between two atom symbols represents one shared pair, so the molecule can be written as . A line does not mean one electron: it represents two. In a water molecule, oxygen forms two O–H bonds and retains two lone pairs. In carbon tetrachloride, carbon forms four C–Cl bonds and has no lone pair.
Problem
How many electrons surround oxygen and each hydrogen in the Lewis picture of water?
- 1.Oxygen contributes six valence electrons and two hydrogen atoms contribute one each, giving eight valence electrons in total.
- 2.Two shared pairs form the two O–H bonds. The remaining four electrons make two lone pairs on oxygen.
- 3.Oxygen counts four bonding electrons plus four lone-pair electrons, giving eight. Each hydrogen counts the two electrons in its single shared pair, giving a duplet.
When atoms share more than one pair
A single shared pair is sometimes insufficient to give the atoms their usual stable electron counts. Two atoms can share two pairs or three pairs. These are double and triple covalent bonds. The bond multiplicity counts shared pairs between a particular pair of atoms, not the total number of bonds in the whole molecule.
| Bond | Shared electron pairs | Illustrations |
|---|---|---|
| Single | 1 | and each O–H bond in water |
| Double | 2 | Each C=O bond in ; C=C in ethene |
| Triple | 3 | N≡N in nitrogen; C≡C in ethyne |
In carbon dioxide, the simple Lewis structure is . Each oxygen shares two pairs with carbon, while carbon is surrounded by four shared pairs in total. In ethene, the carbon atoms share two pairs with one another. In ethyne and nitrogen, the linked atoms share three pairs. Later lessons explain the different orbital overlaps hidden inside these bond lines.
A shared pair belongs to the bonding region between two atoms. We count that pair around both atoms when checking octets, but it remains only two actual electrons when counting the total electrons in the molecule.
Check your understanding
Use the electron count and the meaning of a bond line to answer these questions. A full-shell pattern helps organise the reasoning, while energetic stability explains why a bonded arrangement is possible.
Quiz
A neutral oxygen atom has how many dots in its Lewis symbol?
What does one line in a Lewis structure represent?
Why are two fluoride ions required for each calcium ion in calcium fluoride?
What is shared in the bond between the two nitrogen atoms in a nitrogen molecule?
Which statement best explains the formation of a stable bond?
Practice Problems
- Draw Lewis symbols for Mg, Al, S, and Br, and state the number of dots in each.
- Explain the difference between a lone pair and a shared pair using chlorine as your example.
- Use electron transfer and charge balance to explain the formula of magnesium chloride.
- Draw the shared-pair arrangements in water, carbon dioxide, nitrogen, and ethyne. State the bond multiplicity between the central pair of atoms where applicable.
- Explain why hydrogen follows a duplet rule while carbon commonly follows an octet rule.
- A student says that sodium chloride consists of separate molecules, each containing only one attracted ion pair. Explain what this description leaves out.
Key Takeaways
• Bonding is associated with stabilising attractions and a decrease in system energy. • Lewis symbols display valence electrons, not the complete electron cloud. • Electron transfer produces ions; electron sharing describes covalent bonds. • One, two, and three shared pairs represent single, double, and triple bonds. • The octet rule is useful, while hydrogen commonly has a duplet.
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Next · Lesson 2
Drawing and Testing Lewis Structures