In middle school, chemical bonds are often drawn as solid lines or sticks connecting balls. In reality, a bond is a balance of electrical charges. Opposites attract, likes repel, and electrons are caught in a tug-of-war. By the end of this page, you'll be able to use an element's position to predict exactly how it shares its valence electrons — and what type of bond it forms.
A chemical bond is not a physical joint or glue. It is a push or pull between charges — an electrostatic force. Every atom has a positive nucleus (protons) and a cloud of negative electrons. When two atoms get close, they experience both attraction and repulsion:
Opposite charges pull together. Each positive nucleus attracts the negative electrons of the neighboring atom. This attraction pulls the atoms closer.
Like charges push apart. As they get too close, the two positive nuclei repel each other, and the negative electron clouds repel each other.
A stable balance. A bond forms at the exact distance where the attractive pulls are maximized and the repulson pushes are minimized. The atoms settle into this lowest-energy state.
Not all atoms pull on electrons with the same strength. An atom's power to attract shared electrons is called its electronegativity. Small, tight atoms with many protons (like fluorine) have very high electronegativity. Large atoms with shielded outer shells (like sodium) have very low electronegativity.
When two atoms bond, the difference in their pulling power — the electronegativity difference (ΔEN) — decides exactly where the shared electrons spend their time, resulting in one of four bond types:
Equal sharing. Two nonmetals with similar strength (ΔEN < 0.5). Electrons spend equal time between the two nuclei. A nonpolar covalent bond has no charge separation.
Unequal sharing. Two nonmetals with mismatched strength (0.5 ≤ ΔEN < 1.7). The hungrier atom pulls electrons closer, creating a partial negative charge (δ−) on itself and leaving a partial positive charge (δ+) on the other. This is a polar covalent bond.
Electron transfer. A metal (low EN) and a nonmetal (high EN) with a huge strength mismatch (ΔEN ≥ 1.7). The nonmetal rips the electron completely away, forming positive and negative ions held together by attraction. This is an ionic bond.
Delocalized sea. Two metal atoms (both low EN). Neither atom holds its outer electrons tightly, so they let go. The valence electrons float free among a grid of positive metal cores. This shared attraction is a metallic bond, and the electrons form a delocalized sea of electrons.
Those four categories aren't four separate rules — they're checkpoints along one continuous slide. Drag ΔEN below and watch the same shared electron cloud shift from centered, to lopsided, to gone.
Pick Element A and Element B from the list. The explorer reads their electronegativities, calculates the difference (ΔEN), predicts the bond, and renders the electrostatic arrangement at the atomic level.
Choose one element in each dropdown to build a bond. The tug-of-war will calculate in real time.
Carbon (C, EN = 2.55) and hydrogen (H, EN = 2.20) are both nonmetals. What type of bond holds a carbon-hydrogen (C–H) bond together in methane (CH4)?
C-H is nonpolar covalent. The electronegativity difference is ΔEN = 2.55 − 2.20 = 0.35. Since 0.35 is less than the 0.5 threshold, the electrostatic pull from both nuclei is nearly equal. They share the electrons almost perfectly in the middle. This is why organic compounds like oils and fats (made of C-H bonds) are completely nonpolar and do not mix with water.
Hydrogen (H, EN = 2.20) and fluorine (F, EN = 3.98) react to form hydrogen fluoride (HF, ΔEN = 1.78). Both are nonmetals. What type of bond is this?
H-F is polar covalent. This is a key boundary case. The 1.7 threshold is a guideline, not a strict law. Because hydrogen and fluorine are both nonmetals, they both have a strong pull on electrons (high ionization energies), meaning neither wants to let go completely to form separate ions. They still share the electrons (making it covalent), though fluorine pulls them so hard that the bond is extremely polar and holds a massive charge separation.
Copper (Cu, EN = 1.90) and zinc (Zn, EN = 1.65) are melted together to form brass. What type of bond holds the copper and zinc atoms together in the solid alloy?
It is a metallic bond. Both copper and zinc are metal atoms with relatively low electronegativities. Because neither nucleus pulls strongly enough to keep its outer electrons locked in place, their valence shells merge and the electrons flow freely throughout the metal lattice. This shared attraction holds the alloy together, making brass malleable and highly conductive.
Can't remember the details? Tap Reveal. The physical act of retrieval is what makes it stick.
A chemical bond is actually an force of attraction between the positive nuclei and negative electrons. An atom's power to pull shared electrons is its . When two nonmetals share electrons equally, they form a bond. If sharing is unequal, the bond is , which creates partial charges. When a metal and a nonmetal bond, electrons are fully to form ions in an bond. When metals bond together, their loose outer electrons form a delocalized in a bond.
In your own words, explain how the electronegativities of two metal atoms lead to the formation of a "sea of electrons" instead of an ionic or covalent bond.
Writing it out forces your brain to organize the concept. Write your answer first, then peek and compare.
Metals have low electronegativity, meaning their nuclei exert a very weak grip on their valence electrons. When metal atoms get close to each other, neither atom is strong enough to keep its outer electrons locked locally (ruling out covalent sharing) or strong enough to strip them from the neighbor (ruling out ionic transfer). As a result, the valence electrons escape their parent shells and flow freely as a delocalized sea of electrons, attracting all the positive metal cores together like a fluid electrostatic glue.
Give yourself a point for each key concept you wrote down. The flag (⚑) marks the step that separates a complete description from a partial one.
Explain the molecular-level difference in electrical force and electron behavior between a covalent bond in water (H2O) and a metallic bond in aluminum (Al). Contrast where the valence electrons live and how the electrostatic force holds each substance together, linking it to electronegativities. Then name one macroscopic physical property of aluminum that results from this difference.
Self-score: 4 = all four · 3 = missing the EN explanation · 2 = described both bonds but omitted EN and properties · ≤1 = simple definition of one bond type.
Dental crowns and joint replacements utilize custom bond properties. Biocompatible dental ceramics use strong ionic bonds that resist thermal shock and food acids, while metallic titanium implants use delocalized electron seas to absorb stress and blend with bone, merging different chemical bond structures to repair the human body.