Atoms joined together, four different ways
, HCl, NaCl, and a piece of copper all contain atoms joined to other atoms, yet their properties could hardly be more different: is a gas, NaCl a hard, brittle crystal, and copper conducts electricity and can be drawn into wire. Chapter 2.4 studied covalent, ionic, and metallic bonds one at a time. AP topic 2.1 goes a step further: predict the type of bond from the properties of the elements involved, and explain why.
Bond type depends on how the electrons are shared
The valence electrons decide everything. How strongly each atom attracts bonding electrons (its electronegativity) determines how they end up distributed:
- Shared equally: when both atoms have the same electronegativity (H–H, Cl–Cl), the shared pair sits evenly between them: a nonpolar covalent bond非极性共价键.
- Shared unequally: when the electronegativities differ, the shared pair is pulled toward the more electronegative atom: a polar covalent bond极性共价键. The more electronegative end carries a partial negative charge部分负电荷, δ−, and the other end a partial positive charge, δ+. The larger the difference, the larger the partial charges and the more polar the bond.
- Essentially transferred: metals have low electronegativity and low ionization energy, nonmetals high electronegativity, so electrons move almost completely from the metal to the nonmetal. The resulting ions, with whole-number charges, are held together by Coulombic attraction: an ionic bond.
- Shared by everyone: metal atoms all give up their valence electrons easily. Those electrons don’t belong to any one pair of atoms but move freely through the whole piece of metal: a metallic bond.
Two rules of thumb
Rule 1: the kinds of elements. A metal with a nonmetal usually forms an ionic bond; two nonmetals form a covalent bond; two metals form a metallic bond.
Rule 2: the electronegativity difference, ΔEN. The cutoffs below are rules of thumb, and textbooks vary slightly:
| ΔEN | below about 0.4 | about 0.4 to 1.7 | above about 1.7 |
|---|---|---|---|
| Bond type | nonpolar (or very slightly polar) covalent | polar covalent | ionic |
| Bond | H–H | C–H | N–H | H–Cl | O–H | H–F | Na–Cl | Cs–F |
|---|---|---|---|---|---|---|---|---|
| ΔEN | 0 | 0.35 | 0.84 | 0.96 | 1.24 | 1.78 | 2.23 | 3.19 |
(Electronegativities on the revised Pauling scale, for example H 2.20, C 2.55, O 3.44, F 3.98, Na 0.93, Cl 3.16.)
Polarity depends on the electronegativity difference, not on whether the two atoms are the same element. C (2.55) and S (2.58), or P (2.19) and H (2.20), are different elements with almost equal electronegativities, so the bonds between them are essentially nonpolar covalent. To compare the ionic character of a set of compounds, compare the ΔEN of their bonds: Na–Cl (2.23) is more ionic than Mg–Cl (1.85), which is more ionic than Al–Cl (1.55).
Bond type is a continuum
Bonds don’t jump from one type to another; they change continuously. So the two rules occasionally disagree:
- Al and Cl: a metal and a nonmetal, so rule 1 says ionic, but ΔEN is only 1.55 and the bond has considerable covalent character. In fact, exists as molecules in the gas phase.
- H and F: ΔEN is 1.78, which rule 2 would call ionic, but both are nonmetals and no electrons are fully transferred. HF is a very polar covalent compound made of molecules.
In cases like these, AP expects you to address both sides: what kinds of elements are bonding, and what the electronegativity difference says about how polar (how ionic) the bond is. Even NaCl has a little covalent character, and even a very polar covalent bond has only partial charges, not ions.
One substance, two kinds of bond
Compounds made of polyatomic ions contain both ionic and covalent bonds. In , the N atom and the 4 H atoms inside are joined by covalent bonds, while and are held together by ionic bonding. contains no metal at all, yet it has ionic bonds: treat a polyatomic ion as a single ion. NaOH (the O–H bond in ), (the C–O bonds in ), and work the same way.
Pick two elements to see where the electronegativity difference puts them on the continuum, then open the enrichment section for the bond-type triangle:
Enrichment: the bond-type triangle (not on the AP exam)
The horizontal axis is the average electronegativity of the two elements, and the vertical axis is their difference. The lower left is like cesium (metallic), the lower right like fluorine gas (covalent), and the top like cesium fluoride (ionic). The closer a point is to a corner, the more the bond resembles that type; the three types blend continuously.
Electronegativity: revised Pauling scale (trends.ts). 0.4 and 1.7 are rules of thumb, and textbooks vary slightly.
Example 1: Ranking bond polarity
Rank C–H, N–H, O–H, and F–H from least to most polar, and identify the atom with the partial negative charge in each.
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Identify: what decides bond polarity. The larger the electronegativity difference (ΔEN), the more unequally the pair is shared and the more polar the bond. H is 2.20; C, N, O, and F are 2.55, 3.04, 3.44, and 3.98.
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Find ΔEN for each bond.
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Compare and rank. The larger ΔEN, the more polar:
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Find δ−. The shared pair is pulled toward the more electronegative atom, which carries δ−. Carbon (2.55) is also more electronegative than hydrogen, so in C–H the carbon carries a small δ−; in the other three, N, O, and F carry δ−, and H is δ+ in all four.
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Note. F–H has a ΔEN above 1.7, but it is still covalent: both atoms are nonmetals, there is no metal to give up electrons, and the electrons are shared, not transferred.
Example 2: Predicting bond type from the elements
Predict the type of bonding in each of the following and justify your answer from the properties of the elements involved: (1) MgO; (2) ; (3) between copper and zinc in brass.
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MgO: identify the elements. Mg is a metal with low electronegativity (1.31) that easily loses its 2 valence electrons; O is a nonmetal with high electronegativity (3.44).
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MgO: compare electronegativities.
The difference is large, so electrons are essentially transferred from Mg to O: ionic bonding between and .
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: identify and compare. Si is a metalloid and O a nonmetal; neither gives up electrons easily, so they share pairs:
The shared pairs are pulled toward O: polar covalent bonds.
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Copper and zinc: identify and compare. Both are metals with low electronegativities (1.90 and 1.65) that give up valence electrons easily, and those electrons are delocalized through the solid: metallic bonding.
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State the reasons. Each judgment rests on whether the elements are metals or nonmetals and on the electronegativity difference, not on atom size or melting point.
- Cl–Cl
Identical atoms with equal electronegativity, so the shared pair isn't pulled either way.
- C–H (ΔEN 0.35)
The difference is below about 0.4, so it is usually treated as nonpolar covalent; strictly, it is very slightly polar.
- C–S (ΔEN 0.03)
C and S are different elements with almost equal electronegativities: nonpolar covalent.
- N–H (ΔEN 0.84)
Two nonmetals with a clear difference: polar covalent, with δ− on N.
- Between K and Br in KBr (ΔEN 2.14)
A reactive metal and a nonmetal with a large difference: ionic.
- Between and in
The cation and the anion are held by ionic bonding; only the N–H bonds inside are covalent.
- S–O (ΔEN 0.86)
Two nonmetals: polar covalent, with δ− on O.
- Between Na atoms in a piece of sodium
Identical metal atoms: metallic bonding, with valence electrons moving through the whole piece of metal.
- Between Ca and O in CaO (ΔEN 2.44)
A metal and a nonmetal with a large difference: ionic.
- H–F (ΔEN 1.78)
The difference is above 1.7, but both are nonmetals: a very polar covalent bond, and HF is made of molecules.
- P–Cl (ΔEN 0.97)
Two nonmetals: polar covalent, with δ− on Cl.
- Between Ni and Cu in a nickel–copper alloy
Two metals: metallic bonding.
Common mistakes
Mistake 1: “ΔEN above 1.7 always means ionic, and below 1.7 always covalent.” 1.7 is only a rule of thumb. HF has ΔEN = 1.78 but is a covalent molecule; Al–Cl has ΔEN = 1.55, and that metal–nonmetal bond has clear covalent character. Look at the kinds of elements too.
Mistake 2: “Atoms in a polar covalent bond carry whole-number charges.” They carry only partial charges, written δ+ and δ−, smaller than the charge of one electron. Only the ions in an ionic compound carry whole-number charges.
Mistake 3: “A metallic bond is a pair of electrons shared between two metal atoms.” A metal’s valence electrons don’t belong to any particular pair of atoms; they move freely through the whole piece of metal (the sea of electrons). That is also why metals conduct electricity.
Quick check
Which bond is the most polar?
Electronegativity differences: C–O 0.89, N–O 0.40, O–H 1.24, C–Cl 0.61. The larger the difference, the more polar the bond.
In a C–Cl bond, which atom carries the partial negative charge, and why?
Partial charges follow electronegativity: Cl 3.16, C 2.55. Choice A states a true fact (C is smaller) but reaches the wrong claim, because what matters is attraction for the shared electrons. D ignores polar covalent bonds.
Rb and Br react to form a compound. Based only on where the two elements are in the periodic table, which claim is best supported?
Electrons are essentially transferred completely between an alkali metal and a halogen (ΔEN = 2.14), giving ionic bonds. A: bond type alone doesn't establish solubility (many ionic compounds are insoluble in water, such as and AgCl). B: Rb forms and Br forms , so the formula is RbBr. C describes covalent bonding.
In which compound do the bonds have the greatest ionic character?
Every bond is to O (3.44), so compare the other element: Mg 1.31, Al 1.61, Si 1.90, P 2.19. The differences are 2.13, 1.83, 1.54, and 1.25; Mg–O has the largest difference and the most ionic character.
Which substance contains both ionic and covalent bonds?
and are held by ionic bonding, and N and O inside are joined by covalent bonds. KCl has only ionic bonding; and have only covalent bonds.
The electronegativity difference between H and F is 1.78, above 1.7. Is HF an ionic compound?
1.7 is only a rule of thumb, and two nonmetals don't form an ionic bond. B is a wrong fact. In C, aqueous HF conducts because HF reacts with water to make ions, which doesn't make HF itself ionic.
AP-style free response
Electronegativity data: H 2.20, C 2.55, O 3.44, Na 0.93, Mg 1.31, Al 1.61, Cl 3.16.
(a) Identify the type of bonding: Na–Cl, C–O, C–H, and Mg–Mg in a piece of magnesium.
(b) In a C=O bond, which atom carries the partial negative charge? Justify your answer.
(c) A student claims, “Al is a metal and Cl is a nonmetal, so the bonds in must be typical ionic bonds.” Evaluate this claim using electronegativity.
(d) Rank C–H, C–O, and C–Cl from most to least polar.
Model answer and scoring
(a) Na–Cl: ionic (a metal and a nonmetal, ΔEN = 2.23). C–O: polar covalent (two nonmetals, ΔEN = 0.89). C–H: covalent and only very slightly polar (ΔEN = 0.35), usually treated as nonpolar. Mg–Mg: metallic.
(b) O carries the partial negative charge. Oxygen is more electronegative (3.44) than carbon (2.55), so it attracts the shared electrons more strongly and they are pulled toward it.
(c) Not entirely correct. The electronegativity difference is only 3.16 − 1.61 = 1.55, below 1.7, so electrons are not fully transferred and the bonds have considerable covalent character ( exists as molecules in the gas phase). “Metal + nonmetal = ionic” is only a rule of thumb.
(d) C–O (0.89) > C–Cl (0.61) > C–H (0.35).
Scoring: (a) all four correct; (b) names O and says it is more electronegative and attracts the shared electrons more strongly; (c) calculates ΔEN and explains the significant covalent character; (d) correct order.
Summary
- Bond type depends on how the bonding electrons are distributed: shared equally (nonpolar covalent), pulled toward one atom (polar covalent, δ+/δ−), essentially transferred (ionic), or shared among many metal atoms (metallic).
- Two rules of thumb: whether the elements are metals or nonmetals, and the electronegativity difference (cutoffs near 0.4 and 1.7).
- Bond type is a continuum: a larger electronegativity difference means a more polar (more ionic) bond. When the rules disagree (HF, ), address both.
- Compounds containing polyatomic ions (such as and ) have both ionic and covalent bonds.
AP tip: Questions often give electronegativity data and ask you to identify a bond type, compare polarity, label δ+ and δ−, or explain why two elements form a certain kind of bond. Base the reasons on the properties of the elements involved: electronegativity (attraction for bonding electrons) and metal versus nonmetal (how easily valence electrons are lost). How bond type shapes macroscopic properties such as melting point and conductivity continues in U3.1 and U3.2.