Short answer
- Ionic bond: one atom transfers electrons to another, forming a positive cation and a negative anion. The bond is the electrostatic attraction between them, and the ions pack into a crystal lattice.
- Covalent bond: two atoms share pairs of electrons attracted to both nuclei. The atoms stay neutral, and the result is a molecule.
- Typical elements: ionic bonding usually pairs a metal with a nonmetal; covalent bonding usually pairs nonmetals with each other.
- Melting points: ionic compounds generally melt and boil high, because melting means breaking apart a lattice of strong attractions. Covalent molecular compounds generally melt and boil lower, because melting usually loosens the weak forces between molecules, not the bonds inside them.
- Conductivity: solid ionic compounds do not conduct, because their ions are locked in place. Melted or dissolved ionic compounds conduct well, because their ions can move. Covalent compounds are generally poor conductors in any state, because they have no mobile charged particles.
- Solubility: many ionic compounds dissolve readily in water. Covalent compounds vary: polar ones often mix with water, nonpolar ones usually do not.
- The bridge case: polar covalent bonding sits between the extremes. Electrons are shared, but unequally, so one end carries a partial negative charge and the other a partial positive charge.
- The honest rule: bonding is a continuum, not two sealed boxes. Electronegativity differences guide the call, but they do not decide every case alone.
Introduction
Table salt is hard, crystalline, and melts only at a very high temperature. Water is a liquid at room temperature that does not conduct electricity on its own. Both are held together by chemical bonds, but the bonds work in opposite ways.
This article does not re-teach either bond type from the start. Its job is comparison: for each property, what does each bond type do, and why? Companion articles cover each side in depth.
In short: ionic bonding transfers electrons, so atoms become charged ions that attract each other, while covalent bonding shares electrons, so atoms stay neutral and form molecules. Almost everything else follows from that contrast.
The Core Difference in One Picture
Picture two atoms meeting. In the ionic picture, one atom gives up electrons and the other takes them. The giver is left positive, because it now has more protons than electrons; the taker is left negative, because it now has more electrons than protons. Opposite charges attract, and that attraction is the bond. Sodium and chlorine give the classic example: sodium forms a positive ion, chlorine a negative ion, and the result is sodium chloride, ordinary table salt.
In the covalent picture, neither atom gives up its electrons completely. Instead, the two atoms pool electrons between them. The shared pair sits between the two nuclei, pulled on by both, and that shared pull holds the atoms together. Two hydrogen atoms do this to make hydrogen gas; oxygen shares with hydrogen to make water; carbon shares with hydrogen to make methane, the main component of natural gas.
Two consequences shape the whole comparison. First, ionic bonding produces charged particles, while covalent bonding produces neutral ones. Second, ionic bonding produces an extended structure, a lattice in which each ion is surrounded by ions of the opposite charge, while covalent bonding produces a discrete unit, a molecule with a definite number of atoms. When chemists write NaCl, they mean the simplest ratio of sodium ions to chloride ions in the crystal, not a single NaCl molecule. When they write H2O, they mean one actual molecule: two hydrogen atoms and one oxygen atom, bonded together.
That distinction clears up a common confusion: there is no single link to point to as “the” ionic bond in salt. Each ion is attracted to all of its neighbors of opposite charge at once, so the strength that matters is the whole lattice, which is why so much heat is needed to melt it.
The Comparison at a Glance
| Feature | Ionic bonding | Covalent bonding |
|---|---|---|
| What happens to electrons | Transferred from one atom to another | Shared between atoms |
| Particles formed | Cations and anions | Neutral molecules (or, in network solids, an extended covalent structure) |
| Typical elements | Metal plus nonmetal | Nonmetal plus nonmetal |
| Structure | Three dimensional crystal lattice of ions | Discrete molecules, except for network solids such as diamond |
| State at room temperature | Solid, usually crystalline | Gas, liquid, or solid |
| Melting and boiling points | Generally high | Generally lower for molecular compounds; very high for network solids |
| Hardness and toughness | Hard and rigid, but brittle | Molecular solids are usually soft; network solids are very hard |
| Electrical conductivity, solid | Poor: ions cannot move | Poor: no mobile charged particles |
| Electrical conductivity, melted or dissolved | Good for ionic compounds, once ions are free to move | Generally poor, with exceptions such as acids in water |
| Solubility in water | Many dissolve readily | Variable: polar covalent compounds often dissolve, nonpolar ones usually do not |
| Formula meaning | Simplest ratio of ions in the crystal | Actual number of atoms in one molecule |
| Classic examples | Sodium chloride, magnesium oxide, calcium carbonate | Water, carbon dioxide, methane, oxygen gas |
Most rows say “generally” or “usually” for a reason; exceptions are covered below.
What Actually Happens to Electrons
Transfer: full charges, strong attraction
Ionic bonding starts with a mismatch in how strongly atoms hold their electrons. Metals, on the left and lower parts of the periodic table, hold their outer electrons loosely and lose them easily. Nonmetals, in the upper right, attract electrons strongly and gain them readily to complete their outer shells. When a metal meets a nonmetal, electrons move from metal to nonmetal, and both end up with a more stable arrangement, usually that of a nearby noble gas.
The resulting ions are not a polite fiction. A sodium atom and a sodium ion behave entirely differently: sodium metal reacts vigorously with water, while sodium chloride simply dissolves in it. Likewise, chlorine gas is poisonous and corrosive, while chloride ions in table salt are a normal part of human diets. Bonding changes what a substance is.
Pulling the lattice apart also takes real energy. OpenStax reports that it takes 769 kilojoules to separate one mole of solid sodium chloride into separate gaseous ions. That number measures lattice strength, not one bond, and it is why salt does not melt until 801 degrees Celsius and does not boil until 1,413 degrees Celsius.
Sharing: equal, then unequal
Covalent bonding starts from a different situation. Two nonmetals both attract electrons strongly, so neither gives electrons up entirely. Instead, they share. In a pure covalent bond, between two identical atoms, sharing is equal: the electrons spend the same time near each nucleus, and neither end carries a charge. The bonds in hydrogen gas and oxygen gas are of this kind.
Sharing is not always equal, and that inequality is the single most important idea in this comparison. The property that decides how equally electrons are shared is electronegativity: the ability of an atom to attract a pair of electrons in a chemical bond. When two atoms have the same electronegativity, sharing is equal. When they differ, the more electronegative atom pulls the shared electrons closer, and the bond becomes polar. The atom that pulls harder gains a partial negative charge, and the other gains a partial positive charge. These are partial charges, not full ionic charges. A polar covalent bond is still a shared bond: no electron has fully changed hands.
The continuum, not two boxes
Put the three cases in order and the picture becomes a sliding scale: equal sharing (nonpolar covalent), then unequal sharing (polar covalent), then transfer (ionic). As the difference grows, the bond becomes more ionic in character, but no single value flips a bond from one category to the other, so cutoff guides are guides, not laws.
Polar Covalent Bonding: The Bridge Case
Polar covalent bonding deserves its own section, because it is where the simple two box story breaks down most usefully. Consider hydrogen fluoride, HF. Hydrogen and fluorine have an electronegativity difference of 1.9 on the scale OpenStax uses. By a common cutoff rule, a difference above about 1.7 is supposed to mean ionic. Yet hydrogen fluoride is polar covalent, not ionic. Fluorine pulls the shared electrons strongly toward itself, so the bond has a strongly negative fluorine end and a strongly positive hydrogen end, but the electrons are still shared, and the substance behaves as a molecular compound, not as an ionic lattice.
The same lesson appears at the other end of the scale. Manganese and iodine have a difference of only 1.0, which a cutoff rule would call polar covalent, yet manganese iodide forms an ionic compound. Sodium and chlorine, with a difference of 2.1, do form an ionic compound, as the rule would predict, but the two exceptions show why OpenStax gives better advice than any cutoff: the best guide to bond character is the types of atoms involved and their positions in the periodic table. Bonds between two nonmetals are generally covalent, and bonding between a metal and a nonmetal is often ionic.
Water is the polar covalent example that matters most in daily life. Oxygen pulls the shared electrons in each oxygen hydrogen bond toward itself, leaving the oxygen end partially negative and the hydrogen ends partially positive. The molecule as a whole is neutral, but its ends are not. That uneven charge distribution lets water attract other water molecules, dissolve many ionic and polar substances, and behave very differently from a nonpolar molecule such as methane. Ammonia, NH3, is another polar covalent case, with a nitrogen hydrogen difference of 0.9, again in the polar range and again molecular in behavior.
Polarity also explains why “covalent compounds do not dissolve in water” is a statement to avoid. Some do, readily. What dissolves well in water is generally either charged, like an ionic compound, or polar, like water itself. What dissolves poorly is generally nonpolar, like oil, wax, and methane. The rule is about charge distribution, not about the label ionic or covalent alone.
Why the Properties Differ: Structure First
An ionic solid is a lattice. Ions of opposite charge alternate in a regular three dimensional pattern, each ion surrounded by neighbors of the opposite charge. Nothing in the solid is a separate molecule that could melt away on its own. To melt it, heat must overcome attractions throughout the lattice at once, which is why melting points are high and why the solid is rigid.
A covalent molecular compound is a collection of separate molecules. The strong bonds are inside each molecule; between molecules, only much weaker forces act. To melt such a compound, heat mainly needs to loosen those weak forces, while the bonds inside the molecules stay intact. That is why water boils at 100 degrees Celsius while sodium chloride boils at 1,413 degrees Celsius, and why many covalent compounds are gases or liquids at room temperature even though the bonds inside their molecules are strong. Melting a molecular compound tests how strongly molecules cling to each other, not bond strength inside the molecule.
Covalent network solids break this pattern. In diamond, every carbon atom is bonded covalently to four other carbon atoms in a continuous network. There are no separate molecules to loosen apart. Melting diamond means breaking covalent bonds throughout the network, which takes extreme temperatures: LibreTexts notes that diamond does not melt under ordinary conditions, but vaporizes above 3,500 degrees Celsius. Quartz, the main mineral in sand, is a similar network solid. So the careful statement is not “covalent compounds have low melting points.” It is “covalent molecular compounds generally have low melting points, while covalent network solids have very high ones.”
Property by Property
Melting and boiling points
Ionic compounds: generally high, for the lattice reason above. Sodium chloride melts at 801 degrees Celsius, and many other ionic compounds need similar or higher temperatures.
Covalent molecular compounds: generally low, because only forces between molecules must be overcome. Water boils at 100 degrees Celsius, while carbon dioxide and methane are gases at room temperature.
Covalent network solids: very high, because covalent bonds throughout the network must be broken. Diamond and quartz are the standard examples.
Melting point alone never proves bond type. A high melting point could mean an ionic lattice or a covalent network. A low melting point points more strongly toward a molecular covalent compound, but it does not prove it alone, which is why the identification section below uses several tests together.
Electrical conductivity
This is the cleanest practical test, because it follows directly from what carries charge.
In a solid ionic compound, ions are present but locked in the lattice, so the solid does not conduct electricity. Melt it or dissolve it in water, and the ions are free to move, so the liquid or solution conducts well. That change with state is itself diagnostic: a substance that does not conduct as a solid but conducts when melted or dissolved is very likely ionic.
In a covalent molecular compound, there are no ions to move, so the substance is generally a poor conductor as a solid, as a liquid, and in solution. Two exceptions trap students. Hydrogen chloride gas is covalent, but dissolved in water it forms hydrochloric acid, and the solution conducts because the dissolved molecules produce ions. Graphite, a form of carbon, conducts within its layers even though it is covalent, because its structure lets some electrons move freely. Neither exception makes the original compound ionic: conductivity here comes from dissolving, or from a special structure, not from an ionic lattice.
Solubility
Ionic compounds: many dissolve readily in water, because polar water molecules can surround and separate ions. Not all do: calcium carbonate in limestone dissolves only slightly, which is why limestone landscapes and seashells persist.
Covalent compounds: solubility depends on polarity. Polar covalent compounds, such as sugar and ethanol, mix well with water. Nonpolar ones, such as oil and wax, do not. This is the familiar “like dissolves like” pattern: polar solvents dissolve polar and ionic substances best, and nonpolar solvents dissolve nonpolar substances best. Solubility therefore tells you about polarity as much as about bond type.
Hardness, brittleness, and state
Ionic solids are hard and rigid, but brittle. If a blow shifts one layer slightly, ions of the same charge can end up next to each other, and their repulsion cracks the crystal along a clean plane. That brittleness follows directly from the charge pattern in the lattice.
Covalent molecular solids are usually soft, because only weak forces between molecules must be overcome to deform them, and they are often waxy or crumbly rather than brittle in the ionic sense. Covalent network solids reverse the softness: diamond is famously hard, precisely because deforming it means breaking strong covalent bonds.
At room temperature, ionic compounds are essentially always solids, while covalent compounds can be gases, liquids, or solids. That variety is itself a clue: if a substance is a gas or liquid at room temperature, it is almost certainly not ionic.
Compounds That Contain Both Bond Types
Some compounds contain both ionic and covalent bonds in different places.
Calcium carbonate, CaCO3, is one. The calcium ion and the carbonate ion are held together ionically, but inside the carbonate ion, carbon and oxygen are held together covalently. Potassium nitrate, KNO3, works the same way: potassium and nitrate ions attract ionically, while nitrogen and oxygen inside the nitrate ion share electrons covalently. Ammonium chloride, NH4Cl, is another: nitrogen and hydrogen inside the ammonium ion are covalently bonded, while the ammonium ion and the chloride ion are ionically bonded to each other.
These examples matter for two reasons. First, chemists describe bonding per bond, not per compound, when precision matters: a compound can be ionic overall while containing covalent bonds inside its polyatomic ions. Second, they explain why the identification method below looks for polyatomic ions: recognizing carbonate, nitrate, sulfate, or ammonium immediately tells you the compound contains covalent bonds inside the ion.
How to Tell Which Type a Compound Has
Use the tests together, in this order.
Step 1: Look at the elements
Find each element on the periodic table. A metal plus a nonmetal points to ionic; two nonmetals point to covalent. This is the fastest and most reliable starting test, and it handles the famous cutoff exceptions: hydrogen fluoride is two nonmetals, so it is covalent despite its large difference, and manganese iodide is a metal plus a nonmetal, so it is ionic despite its small difference.
Step 2: Check for polyatomic ions
If the formula contains a familiar group such as carbonate, nitrate, sulfate, phosphate, hydroxide, or ammonium, the compound contains covalent bonds inside that group, even if the compound as a whole is ionic, as in calcium carbonate. This step prevents the common error of calling every bond in such a compound ionic.
Step 3: Use electronegativity difference as a guide, not a verdict
As a rough guide, differences below about 0.4 suggest nonpolar covalent, differences from about 0.4 to about 1.7 suggest polar covalent, and larger differences suggest ionic character. Treat these as approximate bands with many exceptions, exactly as OpenStax presents them. If the guide disagrees with Step 1, trust the element test and the observed properties over the number.
Step 4: Test the physical properties
Check state at room temperature, melting point if known, brittleness, and conductivity in different states. A hard, brittle, crystalline solid that does not conduct when solid but conducts when melted or dissolved is ionic. A gas, liquid, or soft solid that does not conduct in any state is likely covalent and molecular. A very hard solid with a very high melting point could be either ionic or a covalent network solid, so this test alone cannot separate them.
Step 5: Check solubility with care
Ready solubility in water supports an ionic or polar covalent identification, but does not settle the question alone, because polar covalent compounds also dissolve well. Poor solubility in water with good solubility in a nonpolar solvent points toward nonpolar covalent. Insolubility in both may mean a low solubility ionic compound, such as calcium carbonate, or a network solid, so combine this test with the others.
Practice examples
| Compound | Elements | Call | Why |
|---|---|---|---|
| Sodium chloride, NaCl | Metal plus nonmetal | Ionic | Lattice; melts at 801 degrees Celsius; conducts when melted or dissolved |
| Magnesium oxide, MgO | Metal plus nonmetal | Ionic | Metal cation plus nonmetal anion in a lattice |
| Water, H2O | Nonmetal plus nonmetal | Polar covalent | Oxygen pulls shared electrons more strongly; boils at 100 degrees Celsius |
| Carbon dioxide, CO2 | Nonmetal plus nonmetal | Covalent | Gas at room temperature; discrete molecules |
| Methane, CH4 | Nonmetal plus nonmetal | Covalent | Gas at room temperature; discrete molecules |
| Oxygen gas, O2 | Same nonmetal twice | Nonpolar covalent | Identical atoms share equally |
| Hydrogen chloride, HCl | Nonmetal plus nonmetal | Polar covalent | Unequal sharing; conducts only after dissolving in water |
| Calcium carbonate, CaCO3 | Metal plus polyatomic ion | Both | Ionic between ions; covalent inside the carbonate ion |
Elements first, polyatomic ions second, numbers as a guide, properties as confirmation.
Common Mistakes to Avoid
“Ionic bonds are always stronger than covalent bonds.” This compares two different kinds of strength. Lattice energy measures pulling an ionic crystal apart; bond energy measures breaking one covalent bond. Covalent bonds inside molecules can be extremely strong: the triple bond in nitrogen gas is one of the strongest bonds known, and the covalent network in diamond is among the hardest structures known. Compare specific substances and processes, not labels.
“Sharing means equal sharing.” Equal sharing is only the nonpolar case. Most covalent bonds between different atoms are polar to some degree, because different atoms attract electrons differently. Polar covalent is the normal case for bonds between different nonmetals.
“A large electronegativity difference proves a bond is ionic.” Hydrogen fluoride, with a difference of 1.9, is the standard counterexample, and manganese iodide, with a difference of only 1.0, is the reverse. Cutoffs are teaching guides; element types, structure, and observed properties decide close cases.
“Covalent compounds never conduct electricity.” Pure covalent compounds generally do not, but solutions of covalent acids in water conduct, and graphite conducts within its layers. The precise statement is that covalent bonding itself provides no mobile ions, so conductivity, when it appears, needs a separate explanation.
“Melting a covalent compound breaks its bonds.” For molecular compounds, melting breaks the weak forces between molecules, not the strong bonds within them. That is why water can boil at 100 degrees Celsius while its oxygen hydrogen bonds remain intact in every steam molecule. For network solids, the opposite is true.
“The formula tells you the molecule.” For a covalent compound, yes: H2O is one molecule. For an ionic compound, no: NaCl is a ratio in a crystal, not a molecule.
Frequently Asked Questions
What is the main difference between ionic and covalent bonds?
What happens to electrons. In ionic bonding, electrons are transferred, creating charged ions that attract each other. In covalent bonding, electrons are shared, creating neutral molecules. That difference leads to differences in structure, melting point, conductivity, and solubility.
Which is stronger, ionic or covalent?
There is no single winner, because the comparison mixes lattice energy for separating an ionic crystal with bond energy for breaking one covalent bond. Some ionic lattices are very strong, and some covalent bonds, such as the triple bond in nitrogen gas and the network bonds in diamond, are also very strong. Compare specific substances, not labels.
What is a polar covalent bond, and how is it different from an ionic bond?
A polar covalent bond is a shared bond in which electrons are shared unequally, because one atom is more electronegative than the other. The more electronegative end carries a partial negative charge and the other end a partial positive charge, but no electron has fully transferred and no full ion has formed. An ionic bond, by contrast, involves full transfer and full charges, with ions arranged in a lattice.
Can a compound have both ionic and covalent bonds?
Yes. Calcium carbonate, potassium nitrate, and ammonium chloride all do. In each case, a polyatomic ion is held together by covalent bonds inside the ion, while the ion as a whole is held to an oppositely charged ion by ionic attraction. Describe bonding per bond when precision matters, not just per compound.
How can you tell if a compound is ionic or covalent?
Use the five steps above in order: elements first, polyatomic ions second, electronegativity difference as a rough guide only, then properties and solubility as confirmation. Together, these tests settle nearly every ordinary case.
Why do ionic compounds conduct electricity when dissolved but not as solids?
Because conductivity needs charged particles that can move. In a solid ionic compound, ions exist but are locked in the lattice; dissolving or melting frees them, and moving ions carry charge. Solid covalent compounds contain no mobile charged particles.
Why do covalent compounds usually have lower melting points than ionic compounds?
Because melting a molecular covalent compound mainly loosens the weak forces between molecules, not the strong bonds inside them, while melting an ionic compound means overcoming strong attractions throughout a lattice. Covalent network solids are the exception: melting them breaks covalent bonds throughout the network, so their melting points are very high.
Is water ionic or covalent?
Water is polar covalent. Oxygen and hydrogen are both nonmetals, and they share electrons, with oxygen pulling the shared electrons more strongly. The molecule is neutral overall, with a partially negative oxygen end and partially positive hydrogen ends. That polarity makes water a good solvent.
Are all bonds either purely ionic or purely covalent?
No. Bonding forms a continuum from equal sharing through unequal sharing to transfer, and many real bonds sit between the textbook extremes. Polar covalent bonds are the large middle territory. Calling a bond ionic or covalent summarizes its dominant character.
Sources & references
- OpenStax Chemistry 2e (Rice University): Chapter 7.1, Ionic Bonding, for cation and anion formation, the ionic lattice, brittleness, high melting and boiling points, poor solid conductivity with good conductivity when melted or dissolved, and the 769 kilojoules per mole lattice separation energy for sodium chloride; Chapter 7.2, Covalent Bonding, for shared electrons attracted by both nuclei, pure and polar covalent bonding, electronegativity, the Figure 7.8 guide bands with the hydrogen fluoride, ammonia, sodium chloride, and manganese iodide exceptions, and polyatomic ions containing both bond types.
- LibreTexts Chemistry: Section 7.2, Contrasting Ionic Compounds and Covalent Compounds, for melting of ionic compounds breaking ionic bonds compared with melting of covalent compounds disrupting forces between molecules, and for conductivity and solubility contrasts; Section 4.2, Covalent Compounds, Formulas and Names, for water boiling at 100 degrees Celsius compared with sodium chloride boiling at 1,413 degrees Celsius, for covalent compounds as liquids or gases at room temperature, and for compounds containing both bond types; Section 2.1.01, Comparing Ionic and Molecular Substances, for the side by side property table and for diamond vaporizing above 3,500 degrees Celsius; Section 11.8, Bonding in Solids, for graphite conductivity and diamond thermal conductivity exceptions among covalent solids.
