Chemical Bonding

Chemical Bonding

Why atoms stick together, and why the way they stick shapes so much of what we see

Table salt and water are both made of ordinary atoms, yet salt is a hard crystal that melts at 801 °C and water is a liquid you can drink.[9] The difference comes down to how their atoms hold on to each other. That hold is a chemical bond: an attraction between atoms that comes from their electrons, especially the outermost ones, called valence electrons.

On this page we'll meet the three main kinds of bond, use electronegativity to predict which kind two atoms will form, learn to draw Lewis structures, and then use them to predict a molecule's 3D shape and whether it is polar. You can try the bond-type predictor, flip through the molecule explorer, and finish with the quiz. Keep the periodic table handy; you'll use it a lot.

Three Kinds of Bond

Atoms bond because they end up lower in energy, and more stable, together than apart. How they get there depends on how hard each one pulls on electrons.[1]

Ionic: give and take

A metal atom loses one or more electrons and a nonmetal atom gains them. Sodium gives up its single valence electron to become Na+; chlorine takes it to become Cl−. The opposite charges attract. There are no separate “NaCl molecules”: each ion is surrounded by ions of the other charge in a repeating crystal.[1][6]

Typical traits: hard, brittle, high melting points, conduct electricity when melted or dissolved.

Covalent: sharing

Two nonmetal atoms share a pair of electrons between them. One shared pair is a single bond (H–H), two pairs a double bond (O=O) and three pairs a triple bond (N≡N). More shared pairs pull the atoms closer and hold them more tightly.[2]

If both atoms pull equally the bond is nonpolar. If one pulls harder, the shared electrons lean its way and the bond is polar, with a partial negative end (δ−) and a partial positive end (δ+).[2]

Metallic: a shared sea

In a metal, each atom lets its valence electrons roam through the whole piece. You can picture positive metal ions sitting in a “sea” of mobile electrons. Those free-moving electrons explain why metals conduct electricity and heat, and why you can bend or hammer them without shattering them.[12]

Carbon–carbon bonds get shorter as more electron pairs are shared[8]
BondExampleShared pairsLength
C–C singleethane1154 pm
C=C doubleethylene2134 pm
C≡C tripleacetylene3120 pm

1 pm (picometre) is a trillionth of a metre.

Electronegativity Difference as a Guide

Electronegativity is how strongly an atom pulls on the electrons in a bond. Linus Pauling built the scale most textbooks use, from cesium at 0.79 up to fluorine at 3.98.[3][10] Subtract the smaller value from the larger to get the difference, written ΔEN. The bigger the gap, the more lopsided the sharing.[2]

ΔENUsual bond typeExample
less than 0.4Nonpolar covalentH–H 0; C–H 0.35
0.4 to 1.8Polar covalentH–Cl 0.96; O–H 1.24
more than 1.8Mostly ionicNa–Cl 2.23

Differences use the Pauling values on our periodic table page. OpenStax, using rounded values, gets 0.9 for H–Cl and 2.1 for Na–Cl.

Treat these cutoffs as a guide, not a law. Textbooks draw the lines in slightly different places (some books use 1.7 or 2.0 instead of 1.8). Hydrogen fluoride is a good test: with the Pauling values on this page its ΔEN is 1.78, just under 1.8, so the cutoffs correctly call it polar covalent. It only looks like an exception if you use OpenStax's rounded 1.9. The real exception is MnI2: its ΔEN is only about 1.0, yet it is ionic. Chemists also look at the kinds of atoms involved: metal plus nonmetal usually means ionic, and two nonmetals usually means covalent.[2]

Bond-Type Predictor

Pick two elements. The predictor finds the electronegativity difference and makes a first guess at the bond.

Sodium (0.93) and chlorine (3.16): ΔEN = 2.23. A metal and a nonmetal with a big gap: ionic.

The scale runs to 3.0 to keep the bar readable. Two metals get “metallic”; electronegativity difference doesn't apply to them.

Lewis Structures

In 1916 the American chemist Gilbert N. Lewis proposed that a covalent bond is a pair of electrons shared between two atoms, and drew valence electrons as dots around the element symbol.[4] His drawings, now called Lewis structures, are still the quickest way to see where a molecule's electrons are.

The octet rule

Many atoms bond until they are surrounded by eight valence electrons, the same as a noble gas. This is the octet rule. Hydrogen is the exception from the start: it needs only two, like helium.[5]

How to draw one

  1. Add up the valence electrons of every atom (use the group number: O has 6, N has 5, C has 4, halogens have 7, H has 1). Add one for each negative charge on an ion, subtract one for each positive charge.
  2. Pick the central atom, usually the one that can make the most bonds and is least electronegative. Hydrogen is never central.
  3. Connect each outer atom to the center with a single bond (2 electrons each).
  4. Give the outer atoms lone pairs until each has an octet (H gets none).
  5. Put any leftover electrons on the central atom as lone pairs.
  6. If the central atom still has fewer than 8, turn an outer lone pair into a shared pair to make a double or triple bond.[5]

Worked example, CO2: 4 + 6 + 6 = 16 valence electrons. Two single bonds use 4. Giving each O three lone pairs uses the other 12, but then carbon only has 4 electrons. Move one lone pair from each O into a bond, and you get O=C=O with every atom at 8. Find it in the explorer below.

Common exceptions

Molecular Shapes and Polarity

A Lewis structure is flat, but molecules are not. VSEPR theory (valence shell electron pair repulsion) says the electron groups around a central atom push each other as far apart as possible. Count the electron domains: each bond (single, double or triple) counts as one, and so does each lone pair. Lone pairs take up a little more room than bonds, so they squeeze bond angles slightly.[7][11]

DomainsLone pairsShapeAngleExample
20Linear180°CO2
30Trigonal planar120°BF3
40Tetrahedral109.5°CH4
41Trigonal pyramidalabout 107°NH3
42Bent104.5°H2O

Measured angles for NH3 and H2O are from NIST.[8]

Is it polar? First check the bonds: are any of them polar? Then check the shape. If polar bonds point symmetrically in opposite directions, their pulls cancel and the molecule is nonpolar overall (CO2, BF3, CH4). If they don't cancel, usually because lone pairs make the shape lopsided, the molecule is polar (H2O, NH3).[7]

Molecule Explorer

Choose a molecule to see its Lewis structure, its shape, its bond angle and whether it is polar. In the shape drawings, shaded lobes are lone pairs and a dashed stick points away from you. A yellow arrow points toward the negative end of a polar molecule.

Water (H2O)

Lewis structureOHH3D shapeOHH104.5°δ−δ+
Shape
Bent
Bond angle
104.5°
Electron domains
4 electron domains: 2 bonds + 2 lone pairs
Polarity
Polar (dipole moment 1.85 D)
The two lone pairs push the O–H bonds closer together than 109.5°. Both O–H bond dipoles point up toward oxygen, so they add up: water is polar.

Carbon dioxide (CO2)

Lewis structureOCO3D shapeCOO180°
Shape
Linear
Bond angle
180°
Electron domains
2 electron domains: 2 double bonds, no lone pairs on C
Polarity
Nonpolar (dipole moment 0 D)
Each C=O bond is polar, but the two point in exactly opposite directions and cancel. CO2 is nonpolar overall.

Ammonia (NH3)

Lewis structureNHHH3D shapeNHHH107°δ−δ+
Shape
Trigonal pyramidal
Bond angle
about 107°
Electron domains
4 electron domains: 3 bonds + 1 lone pair
Polarity
Polar (dipole moment 1.47 D)
The lone pair sits on top and squeezes the three N–H bonds down into a pyramid. The bond dipoles all lean toward N, so ammonia is polar.

Methane (CH4)

Lewis structureCHHHH3D shapeHCHHH109.5°
Shape
Tetrahedral
Bond angle
109.5°
Electron domains
4 electron domains: 4 bonds, no lone pairs
Polarity
Nonpolar (dipole moment 0 D)
The four C–H bonds point to the corners of a tetrahedron, as far apart as they can get. They are only slightly polar and cancel by symmetry, so methane is nonpolar.

Boron trifluoride (BF3)

Lewis structureBFFF3D shapeBFFF120°
Shape
Trigonal planar
Bond angle
120°
Electron domains
3 electron domains: 3 bonds, no lone pairs on B
Polarity
Nonpolar (dipole moment 0 D)
Boron has only 6 valence electrons around it here, an octet exception. The three very polar B–F bonds sit 120° apart in a flat triangle and cancel, so BF3 is nonpolar.

Hydrogen chloride (HCl)

Lewis structureHCl3D shapeHClδ+δ−
Shape
Linear (two atoms)
Bond angle
none (only one bond)
Electron domains
Chlorine has 1 bond + 3 lone pairs
Polarity
Polar (dipole moment 1.11 D)
Any two-atom molecule is linear. Chlorine (3.16) pulls harder than hydrogen (2.20), so the shared pair leans toward Cl and HCl is polar.

Nitrogen (N2)

Lewis structureNN3D shapeNN
Shape
Linear (two atoms)
Bond angle
none (only one bond)
Electron domains
Each N has a triple bond + 1 lone pair
Polarity
Nonpolar (dipole moment 0 D)
A triple bond: three shared pairs. It is one of the strongest bonds known, which is why the N2 in air is so unreactive. Identical atoms share equally, so it is nonpolar.

Oxygen (O2)

Lewis structureOO3D shapeOO
Shape
Linear (two atoms)
Bond angle
none (only one bond)
Electron domains
Each O has a double bond + 2 lone pairs
Polarity
Nonpolar (dipole moment 0 D)
The Lewis structure shows a double bond with every electron paired. That is a useful model, but it misses something: see “How do we know?” below.

How Do We Know?

For most of chemistry's history nobody could see a bond at all, and even today's images need special microscopes. So where do these pictures come from?

Seeing bonds directly

In 2012 a team at IBM Research in Zurich, led by Leo Gross, used an atomic force microscope with a single carbon monoxide molecule on its tip to image individual molecules. Their images showed the bonds between carbon atoms, and they could even tell single bonds from double-bond-like ones by their length and brightness.[14] A year later, Dimas de Oteyza and colleagues used the same method to watch the bonds of a molecule rearrange during a chemical reaction.[15] The pictures matched the structures chemists had worked out decades earlier.

X-rays map where atoms sit

When X-rays pass through a crystal, the regularly spaced atoms scatter them into a pattern of spots. In 1913 William Lawrence Bragg and his father William Henry Bragg used those patterns to work out crystal structures, starting with simple salts. Sodium chloride turned out to be a grid in which every sodium ion is surrounded by six chloride ions, with no separate NaCl pairs at all, just as the ionic model predicts. The Braggs shared the 1915 Nobel Prize in Physics.[6] X-ray crystallography is still how chemists find the exact positions of atoms in crystals.

Bond lengths and angles are measured

Using X-ray and electron diffraction and the way molecules absorb microwave and infrared light, chemists measure bond lengths and angles very precisely. That is how we know the H–O–H angle in water is 104.5° and that a C≡C triple bond (120 pm) is shorter than a C–C single bond (154 pm). These measurements are collected in reference databases such as NIST's.[8]

Dipole moments reveal polarity

A polar molecule lines up in an electric field, like a tiny compass needle. Measuring how strongly it does that gives its dipole moment, measured in debye (D). Water's is 1.85 D and ammonia's is 1.47 D, while CO2, BF3 and CH4 measure zero.[8] That zero is strong evidence for their symmetric shapes: a bent CO2 would be polar.

From Lewis to Pauling

Lewis's 1916 paper introduced the shared electron pair.[4] In 1932 Linus Pauling noticed that bonds between unlike atoms are stronger than expected, and used that extra energy to build his electronegativity scale.[3] His 1939 book The Nature of the Chemical Bond pulled the ideas together, and he won the 1954 Nobel Prize in Chemistry for that work.[10] In 1957 Ronald Gillespie and Ronald Nyholm set out the electron-pair repulsion rules that became VSEPR.[11]

When the model breaks: oxygen is magnetic

The Lewis structure of O2 has every electron paired. But pour liquid oxygen between the poles of a strong magnet and it sticks there, which only happens when a substance has unpaired electrons. A more advanced model, molecular orbital theory, predicts the two unpaired electrons correctly.[13] Lewis structures still work well for most molecules; O2 is a reminder that every model has limits.

Test Yourself

Ten questions. Click an answer to check it and read why. The explorer and predictor above are allowed!

The Bond Builder Challenge

Think About It

Questions to puzzle over in class or at home.

  1. Both CO2 and H2O have two polar bonds to a central atom. Why is only one of them polar, and how could that affect which one dissolves salt?
  2. Metals bend and ionic crystals shatter. Using the “sea of electrons” and the grid of charged ions, explain the difference.
  3. HF's ΔEN is 1.78 with Pauling's values but 1.9 with rounded ones, which lands it on opposite sides of the 1.8 line. What does that tell you about using any single number to sort the world into categories?
  4. Lewis structures say O2 should not be magnetic, but it is. Should chemists stop teaching them? Why or why not?

Next unit: Chemical Reactions →

Back to all chemistry lessons

Sources

  1. OpenStax, Chemistry 2e, “7.1 Ionic Bonding.” https://openstax.org/books/chemistry-2e/pages/7-1-ionic-bonding
  2. OpenStax, Chemistry 2e, “7.2 Covalent Bonding” (electronegativity and bond type, HF and MnI2 exceptions). https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding
  3. Linus Pauling, “The Nature of the Chemical Bond. IV. The Energy of Single Bonds and the Relative Electronegativity of Atoms,” Journal of the American Chemical Society 54 (1932): 3570–3582. https://doi.org/10.1021/ja01348a011
  4. Gilbert N. Lewis, “The Atom and the Molecule,” Journal of the American Chemical Society 38 (1916): 762–785. https://doi.org/10.1021/ja02261a002
  5. OpenStax, Chemistry 2e, “7.3 Lewis Symbols and Structures” (octet rule, drawing steps, exceptions). https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures
  6. Nobel Prize, “The Nobel Prize in Physics 1915” (William Henry Bragg and William Lawrence Bragg, crystal structure by X-rays). https://www.nobelprize.org/prizes/physics/1915/summary/
  7. OpenStax, Chemistry 2e, “7.6 Molecular Structure and Polarity” (VSEPR shapes, angles, polarity). https://openstax.org/books/chemistry-2e/pages/7-6-molecular-structure-and-polarity
  8. NIST Computational Chemistry Comparison and Benchmark Database (experimental bond lengths, angles and dipole moments). https://cccbdb.nist.gov/
  9. PubChem, “Sodium Chloride” (melting point 801 °C). https://pubchem.ncbi.nlm.nih.gov/compound/Sodium-Chloride
  10. Nobel Prize, “Linus Pauling: Facts” (Chemistry 1954, the nature of the chemical bond). https://www.nobelprize.org/prizes/chemistry/1954/pauling/facts/
  11. R. J. Gillespie and R. S. Nyholm, “Inorganic Stereochemistry,” Quarterly Reviews, Chemical Society 11 (1957): 339–380. https://doi.org/10.1039/QR9571100339
  12. OpenStax, Chemistry 2e, “10.5 The Solid State of Matter” (metallic solids and bonding). https://openstax.org/books/chemistry-2e/pages/10-5-the-solid-state-of-matter
  13. OpenStax, Chemistry 2e, “8.4 Molecular Orbital Theory” (paramagnetism of O2). https://openstax.org/books/chemistry-2e/pages/8-4-molecular-orbital-theory
  14. Leo Gross et al., “Bond-Order Discrimination by Atomic Force Microscopy,” Science 337 (2012): 1326–1329. https://doi.org/10.1126/science.1225621
  15. Dimas G. de Oteyza et al., “Direct Imaging of Covalent Bond Structure in Single-Molecule Chemical Reactions,” Science 340 (2013): 1434–1437. https://doi.org/10.1126/science.1238187

Electronegativity values are Pauling values from the PubChem periodic table, the same data used on our periodic table page.