Difficulty: Intermediate | Prerequisites: Chapters 2, 7, and 8 (periodic table, electron configurations, periodic trends)
Chapter 9 ties everything together: electron configurations (Chapter 7), periodic trends (Chapter 8), and the ions and naming rules from Chapter 2. This is where you learn why atoms bond, how to represent their bonding with Lewis symbols, and what makes a bond ionic, covalent, or polar covalent. The bonding concepts here are the foundation for molecular geometry and polarity in later chapters.
Atoms bond to achieve stable (noble-gas) electron configurations. Ionic bonds involve electron transfer between metals and nonmetals. Covalent bonds involve electron sharing between nonmetals. Electronegativity differences determine whether a covalent bond is polar or nonpolar. Lewis dot symbols are the visual tool for tracking valence electrons through all of this.
Lewis dot symbol
A representation of an atom showing its chemical symbol surrounded by dots for its valence electrons. Each side of the symbol (top, bottom, left, right) can hold up to two dots. Think of it as a quick visual inventory of the electrons available for bonding.
Ionic bond
The electrostatic attraction between a cation and an anion, formed when one atom transfers electrons to another. Typically occurs between a metal and a nonmetal. In simple terms, one atom gives electrons and the other takes them, and the opposite charges hold them together.
Covalent bond
A bond formed when two atoms share one or more pairs of electrons. Typically occurs between two nonmetals. Think of it as two atoms holding onto the same electrons rather than giving them away.
Octet rule
Atoms tend to form bonds so that each atom has eight electrons in its valence shell (like a noble gas). Hydrogen is an exception; it is satisfied with two electrons (a "duet"). The octet rule is a guideline with exceptions, but at this level it covers the vast majority of exam questions.
Electronegativity
A measure of how strongly an atom attracts shared electrons in a chemical bond. Fluorine is the most electronegative element (4.0 on the Pauling scale). Electronegativity generally increases across a period and decreases down a group.
Polar covalent bond
A covalent bond in which electrons are shared unequally because the two atoms have different electronegativities. The more electronegative atom carries a partial negative charge (δ-) and the less electronegative atom carries a partial positive charge (δ+).
Nonpolar covalent bond
A covalent bond in which electrons are shared equally (or nearly so), because the atoms have the same or very similar electronegativities. Examples: H-H, Cl-Cl.
Electrolyte
A substance that dissociates into ions when dissolved in water and conducts electricity. Strong electrolytes dissociate completely (e.g. NaCl, HCl). Weak electrolytes dissociate only partially (e.g. acetic acid). Nonelectrolytes do not dissociate (e.g. sugar).
Nonelectrolyte
A substance that dissolves in water but does not produce ions, so the solution does not conduct electricity. Most molecular compounds are nonelectrolytes.
Draw the element's symbol and place dots around it representing valence electrons.
Place one dot on each of the four sides before pairing any. This reflects the filling of p orbitals.
For ions: cations have fewer dots (or none for Group 1/2 metals); anions have more, often a full octet shown in brackets with the charge.
Example: Carbon (Group 14, 4 valence electrons) has one dot on each of the four sides.
Formed by the transfer of electrons from a metal to a nonmetal.
The metal loses electrons to become a cation; the nonmetal gains electrons to become an anion.
The electrostatic attraction between oppositely charged ions holds the compound together.
Ionic compounds form crystal lattices, not individual molecules. The formula (e.g. NaCl) represents the simplest whole-number ratio of ions.
Formed by the sharing of electron pairs between two nonmetal atoms.
The octet rule drives covalent bonding: each atom shares enough electrons to fill its valence shell to eight (two for hydrogen).
From a Lewis symbol, predict how many bonds an atom forms:
Carbon: 4 valence electrons, needs 4 more → forms 4 bonds.
Nitrogen: 5 valence electrons, needs 3 more → forms 3 bonds (plus one lone pair).
Oxygen: 6 valence electrons, needs 2 more → forms 2 bonds (plus two lone pairs).
Halogens: 7 valence electrons, need 1 more → form 1 bond (plus three lone pairs).
Ionic vs covalent comes down to what the atoms are:
Metal + nonmetal (or polyatomic ion) → ionic.
Nonmetal + nonmetal → covalent (molecular).
This distinction connects directly to Chapter 2 naming: ionic compounds follow one set of rules, molecular compounds follow another.
Electrolytes and nonelectrolytes:
Strong electrolytes dissociate completely in water (ionic compounds, strong acids like HCl, strong bases like NaOH).
Weak electrolytes dissociate partially (weak acids like acetic acid, weak bases like NH₃).
Nonelectrolytes do not dissociate (most molecular compounds like sugar, ethanol).
Electronegativity trend: increases across a period (left to right) and decreases down a group. Fluorine (4.0) is the highest; francium is the lowest.
Bond type from electronegativity difference (ΔEN):
ΔEN = 0: nonpolar covalent (e.g. H₂, Cl₂).
0 < ΔEN < ~1.7: polar covalent (e.g. H-Cl, where Cl is δ- and H is δ+).
ΔEN > ~1.7: ionic (e.g. NaCl).
These thresholds are guidelines, not hard cutoffs.
In a polar bond, the more electronegative atom is δ- (partial negative) and the less electronegative atom is δ+ (partial positive).
To compare the polarity of two bonds, compare their ΔEN values. Larger ΔEN = more polar.
Predicting number of covalent bonds from Lewis symbols
Element | Group | Valence e⁻ | Bonds typically formed | Lone pairs |
|---|---|---|---|---|
C | 14 | 4 | 4 | 0 |
N | 15 | 5 | 3 | 1 |
O | 16 | 6 | 2 | 2 |
F, Cl, Br, I | 17 | 7 | 1 | 3 |
H | 1 | 1 | 1 | 0 |
Bond type from electronegativity difference
ΔEN range | Bond type | Example |
|---|---|---|
0 | Nonpolar covalent | Cl₂ |
0.1 to ~1.7 | Polar covalent | H-Cl (ΔEN ≈ 0.9) |
Above ~1.7 | Ionic | NaCl (ΔEN ≈ 2.1) |
The ionic vs covalent distinction is why table salt (NaCl, ionic) dissolves in water and conducts electricity, while sugar (C₁₂H₂₂O₁₁, molecular) dissolves without conducting. Electronegativity differences explain why water is such a good solvent: the O-H bonds are polar enough to interact with both ionic and other polar substances. Electrolyte chemistry is the basis of how sports drinks replenish ions (Na⁺, K⁺, Cl⁻) lost through sweat.
Students often think ionic compounds form molecules. They do not: they form extended crystal lattices. The formula unit (e.g. NaCl) is a ratio, not a molecule.
Students often assume all bonds between different atoms are polar. Two atoms with the same electronegativity (like C-S, ΔEN ≈ 0) form essentially nonpolar bonds.
Students often mix up which atom is δ+ and which is δ-. The more electronegative atom always carries the partial negative charge. Electronegativity tells you who "wins" the tug-of-war for shared electrons.
Students often forget that hydrogen only needs two electrons (a duet) to be stable, not eight.
⚠️ Drawing Lewis dot symbols for atoms and monatomic ions is a fundamental skill tested on its own and as a building block for Lewis structures.
⚠️ Classifying a compound as ionic or covalent from its formula (metal + nonmetal = ionic, nonmetal + nonmetal = covalent) is tested directly and underpins the naming questions from Chapter 2.
⚠️ Using electronegativity differences to classify bonds as nonpolar covalent, polar covalent, or ionic, and to assign δ+ and δ- to atoms, is a high-frequency exam topic.
⚠️ Comparing the polarity of two bonds by comparing their ΔEN values is a common question format.
⚠️ Electrolyte classification (strong, weak, nonelectrolyte) based on whether a compound is ionic, a strong acid/base, or molecular often appears.
True or False: A Lewis dot symbol for neon would show 8 dots arranged around the symbol. (True)
Fill in the blank: In the bond H-F, the fluorine atom carries a partial ______ charge. (negative, δ-)
True or False: CO₂ is an ionic compound. (False; both C and O are nonmetals, so it is molecular/covalent)
Fill in the blank: Oxygen typically forms ______ covalent bonds. (2)
True or False: A strong electrolyte only partially dissociates in water. (False; strong electrolytes dissociate completely)
Q: Draw the Lewis dot symbol for nitrogen (N).
A: Nitrogen has 5 valence electrons. Place one dot on each of the four sides (4 dots), then pair one of them on one side: three unpaired dots and one pair. The three unpaired dots represent the three bonds N typically forms.
Q: Classify the bonding in MgBr₂ as ionic or covalent. Explain.
A: Ionic. Mg is a metal (Group 2) and Br is a nonmetal (Group 17). Metal + nonmetal = ionic bonding. Mg transfers two electrons (one to each Br) to form Mg²⁺ and two Br⁻ ions.
Q: Which bond is more polar: H-Cl or H-F? Explain using electronegativity.
A: H-F is more polar. The electronegativity of F (4.0) is higher than Cl (3.0), so ΔEN for H-F (4.0 - 2.1 = 1.9) is larger than for H-Cl (3.0 - 2.1 = 0.9). Greater ΔEN = more polar.
Q: Is KBr a strong electrolyte, weak electrolyte, or nonelectrolyte when dissolved in water?
A: Strong electrolyte. KBr is an ionic compound, so it dissociates completely into K⁺ and Br⁻ ions in water.
Q: From its Lewis symbol, how many covalent bonds does carbon typically form, and why?
A: Four. Carbon has 4 valence electrons and needs 4 more to complete an octet. Each shared pair is one bond, so carbon forms 4 bonds.
Lewis dot symbols are the first step toward drawing full Lewis structures, which come later and are used to predict molecular geometry (VSEPR theory). The ionic vs covalent distinction made here determines which naming rules from Chapter 2 apply. Electronegativity trends follow the same periodic logic as atomic radius and effective nuclear charge from Chapter 8.
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