Ionic, Covalent, Dative & Metallic Bonding

Last updated: 15 August 2026

Diagnose → Learn → Master: Every study-note section includes 15 free MCQs — start with 5 Diagnostic MCQs to identify gaps in your understanding, then complete 10 Mastery MCQs near the end to check what you have learned.

Sodium chloride forms a rigid crystal, oxygen exists as O₂ molecules, aluminium chloride can form Al₂Cl₆, and copper forms a solid that conducts electricity and can be bent into shape. These substances behave differently because their particles are held together in different ways.

At first, ionic, covalent, co-ordinate and metallic bonding can feel like four separate definitions to memorise. A better approach is to look beneath the terminology and ask one question: what positive and negative charges are attracting each other?

Once you can identify those particles, the definitions, electron diagrams and many physical properties become much easier to reason out instead of memorising them as disconnected facts.

Diagnostic Quiz

How Well Do You Understand Chemical Bonding?

Try these five questions before reading the chapter. They test the basic ideas behind ionic, covalent, co-ordinate and metallic bonding.

5 quick questions
Spot common bonding misconceptions
Every answer option is explained

What is the underlying nature of chemical bonding?

All bonds form because electrons must always be transferred between atoms.

Electron transfer is associated with the formation of ions in many ionic compounds, but it is not how covalent or metallic bonding is described.

Chemical bonding involves electrostatic attraction between positive and negative charges.

Correct. The particles involved differ between bonding models, but electrostatic attraction is the common underlying idea.

All bonds form because atoms must contain exactly eight electrons.

The octet rule is a useful pattern for many main-group species, but it is not the fundamental physical origin of bonding and has important exceptions.

All bonds involve attraction between complete positive and negative ions.

That describes ionic bonding. Covalent bonding instead involves positively charged nuclei attracting shared electrons.

Which statement correctly defines an ionic bond?

The transfer of electrons from a metal atom to a non-metal atom.

Electron transfer can form the ions, but the transfer itself is not the ionic bond.

The sharing of electrons between positive and negative ions.

Ionic bonding does not require a shared electron pair. Sharing is associated with covalent bonding.

The electrostatic attraction between oppositely charged ions.

Correct. Positive and negative ions attract because they carry opposite charges.

The attraction between two neutral atoms after electron transfer.

Electron transfer produces charged ions rather than leaving both particles neutral.

Which statement best describes a covalent bond?

Electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

Correct. The negatively charged shared electrons are attracted to the positively charged nuclei of both bonded atoms.

Electrostatic attraction between two negatively charged electron clouds.

Electron clouds repel one another because they carry the same negative charge.

Transfer of a shared electron pair from one atom to another.

Covalent bonding involves shared electrons rather than complete transfer of the bonding pair.

Attraction between two neutral nuclei.

Both nuclei are positively charged and therefore repel rather than attract each other.

What makes a co-ordinate bond different in the way it is formed?

It contains only one electron.

A co-ordinate bond contains a shared pair of two electrons, just like an ordinary covalent bond.

The electrons are transferred completely to form ions.

Complete electron transfer describes ion formation rather than co-ordinate covalent bonding.

One bonding electron is supplied by each atom.

That is the usual electron-accounting picture for an ordinary covalent bond.

Both electrons in the shared pair are supplied by the same atom.

Correct. This is the defining feature of a co-ordinate, or dative covalent, bond.

Which description best represents metallic bonding?

Neutral metal atoms joined by localised pairs of electrons.

Metallic bonding is not described using localised electron pairs between particular pairs of atoms.

Positive metal ions attracted to delocalised electrons.

Correct. The electrostatic attraction between the positive ions and the delocalised electrons holds the metallic lattice together.

Positive metal ions attracted to negative metal ions.

A metallic lattice is not made from alternating positive and negative metal ions.

Metal atoms held together only by intermolecular forces.

Metallic bonding is a strong chemical bonding interaction, not an intermolecular attraction between separate molecules.

One electrostatic idea links all four bonding types

The word electrostatic simply refers to interactions between electrical charges. Opposite charges attract, while like charges repel.

This gives us a useful way to organise chemical bonding:

Bonding type	Positive component	Negative component
Ionic	Positive ions	Negative ions
Covalent	Positively charged nuclei	Shared electrons
Co-ordinate	Positively charged nuclei	Shared electrons supplied initially by one atom
Metallic	Positive metal ions	Delocalised electrons

The definition must identify the attraction, not merely what happens to the electrons.

For example, saying that ionic bonding is “electron transfer” describes how ions may be produced. It does not define the force that holds the ions together.

Why does bonding make a system more stable?

Atoms do not form bonds because they consciously “want” complete outer shells. The deeper reason is energetic: a bond can form when the bonded arrangement is lower in energy than the separated particles.

Consider two atoms approaching one another. Several interactions occur at the same time:

each positively charged nucleus attracts the negatively charged electrons of the other atom;
the two positive nuclei repel one another; and
the negatively charged electron clouds repel one another.

At a suitable internuclear distance, the attractive interactions outweigh the repulsive interactions sufficiently for the total potential energy to decrease. A stable bond corresponds to an arrangement near a minimum in potential energy.

If the atoms are pushed too close together, nucleus–nucleus and electron–electron repulsions rise sharply, so there is an optimum bond distance rather than the nuclei simply collapsing together.

Useful introductory idea: atoms often form bonds that give them noble-gas-like electron arrangements.

Deeper A-Level idea: bonding occurs because the overall bonded arrangement can be energetically more stable than the separated particles.

Ionic bonding: attraction between oppositely charged ions

Ionic compounds contain positively and negatively charged ions. These ions are held together by electrostatic attraction.

Definition of an ionic bond

An ionic bond is the electrostatic attraction between oppositely charged ions.

Electron transfer forms the ions

Ionic compounds are often introduced by showing electrons being transferred from one atom to another.

For sodium chloride, sodium loses one electron:

Na → Na⁺ + e⁻

Chlorine gains one electron:

Cl + e⁻ → Cl⁻

Electron transfer therefore explains how Na⁺ and Cl⁻ ions can be formed. Once these oppositely charged ions exist, they attract one another electrostatically.

Formation of ions: electron transfer.

Ionic bond: electrostatic attraction between the resulting oppositely charged ions.

There are no separate NaCl molecules in solid sodium chloride

It is tempting to imagine each Na⁺ ion paired permanently with one Cl⁻ ion. That is not how a solid ionic compound is arranged.

In sodium chloride, every Na⁺ ion is attracted to several nearby Cl⁻ ions, while every Cl⁻ ion is attracted to several nearby Na⁺ ions. These attractions extend throughout the solid to form a giant ionic lattice.

The formula NaCl therefore represents the simplest whole-number ratio of Na⁺ ions to Cl⁻ ions in the lattice. It does not represent a discrete NaCl molecule.

Magnesium oxide involves ions with larger charges

Magnesium loses two electrons:

Mg → Mg²⁺ + 2e⁻

Oxygen gains two electrons:

O + 2e⁻ → O²⁻

The Mg²⁺ and O²⁻ ions then attract one another throughout a giant ionic lattice.

What affects the strength of attraction between ions?

Electrostatic attraction becomes stronger when the charges are larger and when the charged particles are closer together.

For ionic compounds, this means attraction tends to be stronger when:

the ions have larger charge magnitudes; and
the ions have smaller ionic radii, so their centres can approach more closely.

Mg²⁺ and O²⁻ therefore experience particularly strong electrostatic attraction because both ions have charges of magnitude 2 and are relatively small.

Do not compare ionic bonding using charge alone.

The electrostatic attraction depends on both the magnitudes of the charges and the distance separating their centres.

Drawing ionic dot-and-cross diagrams

Dot-and-cross diagrams are electron-accounting diagrams. Dots can represent the outer electrons originally belonging to one atom, while crosses represent those originally belonging to another.

For an ionic dot-and-cross diagram:

show the relevant outer-shell electrons;
show the transferred electron using the symbol of the atom from which it came;
place each ion inside square brackets;
write the ionic charge outside the brackets;
show the completed outer shell where appropriate; and
do not draw a shared pair between the oppositely charged ions.

Dots and crosses do not represent different kinds of electrons.

They simply allow you to keep track of where the electrons originated. Electrons themselves are indistinguishable.

Covalent bonding: attraction involving shared electrons

Covalent bonding commonly occurs between non-metal atoms. Instead of transferring electrons completely, the atoms share one or more pairs of electrons.

The shared electrons are negatively charged and are attracted to the positively charged nuclei of both bonded atoms.

Definition of a covalent bond

A covalent bond is the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

Why “a shared pair of electrons” is not the full definition

A shared pair tells you where the bonding electrons are located, but it does not explicitly state the attractive force holding the nuclei together.

Incomplete: A covalent bond is a shared pair of electrons.

Better: A covalent bond is the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

Formation of H₂

Each hydrogen atom contains one electron. When two hydrogen atoms form H₂, each atom contributes one electron to a shared pair.

This negatively charged pair is attracted to both hydrogen nuclei. The attraction holds the two atoms together as an H–H covalent bond.

Single, double and triple covalent bonds

Two atoms may share more than one pair of electrons.

Bond type	Shared electron pairs	Shared electrons	Examples
Single	1	2	H₂, Cl₂, HCl, CH₄
Double	2	4	O₂, CO₂, C₂H₄
Triple	3	6	N₂, HCN

Between the same two atoms, increasing bond order generally gives a shorter and stronger bond. A triple bond is therefore generally shorter and stronger than a double bond between the same atoms, while a double bond is generally shorter and stronger than the corresponding single bond.

However, a double bond is not simply “twice as strong” as a single bond, and a triple bond is not simply “three times as strong”.

Examples worth being able to draw

In Cl₂, the chlorine atoms share one pair of electrons, giving a Cl–Cl single bond. Each chlorine atom also has three lone pairs.

In O₂, the oxygen atoms share two electron pairs, giving an O=O double bond. Each oxygen atom retains two lone pairs.

In N₂, the nitrogen atoms share three electron pairs, giving an N≡N triple bond. Each nitrogen atom retains one lone pair.

In CH₄, carbon forms four C–H single bonds.

In CO₂, carbon forms one C=O double bond to each oxygen atom:

O=C=O

Drawing covalent dot-and-cross diagrams

In a covalent dot-and-cross diagram, bonding electrons are placed between the two bonded atoms because they form shared pairs.

Remember to:

use different symbols to show the original source of electrons;
show one shared pair for a single bond;
show two shared pairs for a double bond;
show three shared pairs for a triple bond;
include non-bonding electrons as lone pairs; and
use brackets and an overall charge when drawing a polyatomic ion.

Co-ordinate (dative covalent) bonding

A co-ordinate bond is not a completely different type of attractive force. It is a particular way in which a covalent bond is formed.

In the simplest electron-accounting picture of an ordinary covalent bond, each atom contributes one electron to the shared pair. In a co-ordinate bond, both electrons in that shared pair originate from the same atom.

Definition of a co-ordinate bond

A co-ordinate, or dative covalent, bond is a covalent bond in which both electrons in the shared pair are supplied by the same atom.

The donor needs a lone pair

For a co-ordinate bond to form, one species must possess an available lone pair. This species acts as the electron-pair donor.

The other species must be able to accept that pair, for example because it has an empty orbital or is electron-deficient. It acts as the electron-pair acceptor.

A co-ordinate bond is often shown using an arrow:

electron-pair donor → electron-pair acceptor

The arrow records where the bonding pair came from. It does not mean that electrons continuously move in that direction after the bond has formed.

Formation of NH₄⁺

An ammonia molecule, NH₃, contains three N–H covalent bonds and one lone pair on the nitrogen atom.

A hydrogen ion, H⁺, has no electrons of its own. Nitrogen can donate its lone pair to H⁺:

NH₃ + H⁺ → NH₄⁺

Both electrons in the newly formed N–H bond came from nitrogen, so this bond is formed by co-ordinate bonding.

The arrow points from nitrogen towards H⁺.

Nitrogen is the electron-pair donor, while H⁺ is the electron-pair acceptor.

Once NH₄⁺ has formed, all four N–H bonds are equivalent

It would be incorrect to imagine that one N–H bond in NH₄⁺ permanently remains special or weaker because it originated as a co-ordinate bond.

Co-ordinate bonding describes how the bond was formed. Once the NH₄⁺ ion exists, all four N–H bonds are equivalent in length and strength.

Formation of H₃O⁺

Water also contains lone pairs. An oxygen atom in H₂O can donate one lone pair to H⁺:

H₂O + H⁺ → H₃O⁺

The new O–H bond is formed using a pair of electrons originally supplied entirely by oxygen.

Why AlCl₃ can form Al₂Cl₆

In a simple AlCl₃ molecule, aluminium forms three covalent bonds. This places only six electrons around the aluminium atom.

Aluminium is therefore electron-deficient and can accept a lone pair from a chlorine atom belonging to another AlCl₃ molecule.

Two AlCl₃ units can therefore combine to form Al₂Cl₆. The structure contains two bridging chlorine atoms between the aluminium atoms.

Each bridging chlorine donates a lone pair towards an electron-deficient aluminium centre. When co-ordinate arrows are used, they therefore point:

Cl → Al

NH₄⁺: N donates a lone pair to H⁺.

Al₂Cl₆: bridging Cl atoms donate lone pairs to electron-deficient Al atoms.

Metallic bonding

Metals require a different bonding model. Their outer electrons are not confined to one particular atom or to one bond between a particular pair of atoms.

Instead, these electrons are delocalised: they can move throughout the metallic structure.

Definition of a metallic bond

A metallic bond is the electrostatic attraction between positive metal ions and delocalised electrons.

The metallic lattice

A metal can be pictured as a regular lattice of positive metal ions surrounded by delocalised electrons.

These electrons do not belong permanently to any one particular metal ion. Each positive ion is attracted to the surrounding delocalised electrons, while the electrons are attracted to many positive ions.

Metallic bonding therefore extends throughout a giant metallic lattice.

Why metals conduct electricity

The delocalised electrons are mobile. When a potential difference is applied, they undergo a net movement through the metal and carry electric charge.

Metals can therefore conduct electricity in the solid state.

In a solid ionic compound: the ions are held in fixed lattice positions, so there are no mobile charge carriers.

In a solid metal: the delocalised electrons are already mobile, so the metal can conduct.

Why metals are malleable and ductile

Metals are malleable, meaning that they can be hammered or rolled into sheets. They are also ductile, meaning that they can be drawn into wires.

When layers of positive metal ions shift relative to one another, the delocalised electrons can redistribute around the new arrangement. Electrostatic attraction is therefore maintained rather than requiring a fixed directional bond between each particular pair of ions.

What affects the strength of metallic bonding?

Metallic bonding tends to become stronger when the attraction between the positive metal ions and the delocalised electrons becomes stronger.

Important factors include:

the charge on the positive metal ion;
the number of delocalised electrons contributed per atom; and
the size of the positive metal ion.

A higher positive charge produces stronger attraction to the electron cloud, while a smaller ion allows the positive charge and surrounding electrons to be closer together.

Comparing sodium and magnesium

Sodium contributes one outer electron per atom to the delocalised electron system and is represented by Na⁺ ions within the metallic lattice.

Magnesium contributes two outer electrons per atom and is represented by Mg²⁺ ions. Mg²⁺ also has a higher charge and is smaller than Na⁺.

The electrostatic attraction between Mg²⁺ ions and the delocalised electrons is therefore stronger than the corresponding attraction in sodium.

Stronger metallic bonding requires more energy to overcome.

This helps explain why magnesium has a higher melting point than sodium, although detailed metallic structures can also influence melting-point trends.

Comparing ionic, covalent, co-ordinate and metallic bonding

Once the four models are placed side by side, the common electrostatic foundation becomes much easier to see.

Bonding	What happens to electrons?	What is attracted?	Typical structural setting
Ionic	Electrons may be transferred to form ions	Positive ions and negative ions	Giant ionic lattice
Covalent	Electrons are shared	Shared electron pairs and bonded nuclei	Simple molecules or giant covalent structures
Co-ordinate	Both bonding electrons are supplied initially by one atom	Shared electron pair and bonded nuclei	Molecules, ions and electron-deficient species
Metallic	Outer electrons are delocalised	Positive metal ions and delocalised electrons	Giant metallic lattice

Composition gives a clue, but it is not the definition

At an introductory level, metal + non-metal often suggests ionic bonding, while non-metal + non-metal often suggests covalent bonding.

These are useful patterns, not definitions. For example, aluminium chloride has substantial covalent character despite containing aluminium and chlorine.

Define bonding using the particles and electrostatic attraction involved.

Do not rely only on whether the elements are labelled as metals or non-metals.

A substance can contain more than one type of bonding

Ammonium chloride, NH₄Cl, is a useful example.

Within each NH₄⁺ ion, nitrogen and hydrogen are joined by covalent bonds. Between NH₄⁺ ions and Cl⁻ ions, there is ionic attraction.

One of the N–H bonds can also be described as having been formed initially by co-ordinate bonding when NH₃ donated its lone pair to H⁺.

Within NH₄⁺: covalent N–H bonding.

Between NH₄⁺ and Cl⁻: ionic bonding.

Common chemical-bonding mistakes

Common mistake	Why it is wrong	Better approach
Defining an ionic bond as electron transfer	Electron transfer forms ions but does not describe the attractive force between them	State the electrostatic attraction between oppositely charged ions
Drawing solid NaCl as separate NaCl molecules	NaCl forms a giant ionic lattice rather than discrete molecules	Treat NaCl as the simplest ratio of Na⁺ to Cl⁻ ions
Defining a covalent bond only as a shared pair of electrons	The statement does not identify the attraction holding the nuclei together	Refer to attraction between the shared pair and the bonded nuclei
Forgetting lone pairs in dot-and-cross diagrams	Not all outer-shell electrons are accounted for	Count all valence electrons before completing the diagram
Pointing a co-ordinate arrow towards the donor	The arrow should show the direction in which the electron pair is donated	Draw the arrow from lone-pair donor to electron-pair acceptor
Thinking the co-ordinate bond in NH₄⁺ remains permanently different	All four N–H bonds are equivalent after the ion has formed	Use co-ordinate bonding to describe bond formation, not a permanent label
Describing a metal as neutral atoms surrounded by free electrons	The bonding model requires attraction between positive metal ions and delocalised electrons	Use the precise metallic-bond definition
Saying delocalised electrons are completely free	They are mobile through the lattice but remain electrostatically attracted to the positive ions	Describe them as mobile or delocalised within the metallic structure
Assuming every metal–non-metal compound must be completely ionic	Bonding is not classified solely by the periodic-table labels of the elements	Consider the actual electron distribution and structure
Assuming a triple bond must be exactly three times as strong as a single bond	Bond strength does not scale directly with the number of shared pairs	State that higher bond order generally gives shorter and stronger bonds between the same atoms

Chemical-bonding checklist

Ionic bonding

Can I define an ionic bond precisely?
Can I distinguish electron transfer from ionic attraction?
Can I draw brackets, charges and transferred electrons correctly?
Do I recognise a giant ionic lattice rather than separate molecules?

Covalent bonding

Can I give the full electrostatic definition?
Can I identify single, double and triple bonds?
Can I account for shared pairs and lone pairs?
Can I draw accurate dot-and-cross diagrams?

Co-ordinate bonding

Can I identify the lone-pair donor?
Can I identify the electron-pair acceptor?
Does my arrow point from donor to acceptor?
Can I explain NH₄⁺ and Al₂Cl₆?

Metallic bonding

Can I identify positive metal ions and delocalised electrons?
Can I explain electrical conductivity?
Can I explain malleability and ductility?
Can I compare metallic bond strength using charge, size and electron number?

Mastery Quiz

Chemical Bonding Mastery Quiz

Now apply the bonding models to structures and unfamiliar examples. These questions focus on distinguishing the particles involved and explaining the Chemistry rather than recalling definitions alone.

10 application-focused questions
Ionic, covalent, co-ordinate and metallic bonding
Every answer option is explained

Which description of solid sodium chloride is most accurate?

It contains separate NaCl molecules held together by intermolecular forces.

NaCl does not form discrete molecules in the solid. It forms a giant ionic lattice.

Each Na⁺ ion is permanently bonded to only one Cl⁻ ion.

Each ion experiences electrostatic attractions to several nearby oppositely charged ions.

Na⁺ and Cl⁻ ions form a giant lattice held together by electrostatic attractions in all directions.

Correct. Ionic attraction extends throughout the lattice rather than occurring only within isolated NaCl pairs.

Neutral sodium and chlorine atoms alternate throughout the crystal.

Sodium chloride contains Na⁺ and Cl⁻ ions, not neutral Na and Cl atoms.

Why are the ionic attractions in MgO generally stronger than those in NaCl?

MgO contains covalent bonds as well as ionic bonds, whereas NaCl does not.

The comparison here is between electrostatic attractions in ionic lattices, not the presence of an additional covalent bond.

Mg²⁺ and O²⁻ have larger charge magnitudes, and their ionic sizes also allow strong electrostatic attraction.

Correct. Ionic attraction depends on both charge magnitude and the distance between the centres of the ions.

Oxygen has more electrons than chlorine.

The total number of electrons in the isolated atoms does not by itself determine the strength of ionic attraction in the lattice.

Magnesium transfers electrons more quickly than sodium.

The rate of electron transfer is not the factor determining the strength of the electrostatic attraction in the final lattice.

A student draws Na⁺ and Cl⁻ correctly but places one shared pair of electrons between the two ions. What is the main error?

Sodium should have a 2⁺ charge.

Sodium commonly loses one outer electron and forms Na⁺, so the 1⁺ charge is correct.

Chloride should be drawn without brackets.

Ions in dot-and-cross diagrams should normally be placed in brackets with the charge shown outside.

Chloride should have a positive charge.

Chlorine gains one electron to form Cl⁻, so its charge is negative.

A shared electron pair suggests covalent bonding rather than ionic attraction between Na⁺ and Cl⁻.

Correct. In an ionic diagram the transferred electron belongs to the chloride ion; a shared pair should not be placed between Na⁺ and Cl⁻.

In the formation of NH₄⁺ from NH₃ and H⁺, which description is correct?

NH₃ + H⁺ → NH₄⁺

Nitrogen donates its lone pair to H⁺, so a co-ordinate arrow is drawn from N towards H.

Correct. Nitrogen supplies both electrons in the newly formed bonding pair, while H⁺ accepts the pair.

H⁺ donates two electrons to nitrogen.

H⁺ has no electrons available to donate.

Nitrogen transfers its lone pair completely to H⁺ to form H⁻.

The pair becomes a shared bonding pair. H⁻ is not formed in this process.

The arrow should point from H⁺ towards nitrogen because H⁺ attracts electrons.

The co-ordinate arrow specifically records the direction of electron-pair donation, so it starts at nitrogen's lone pair.

Why can two AlCl₃ molecules combine to form Al₂Cl₆?

2AlCl₃ → Al₂Cl₆

Aluminium transfers all its electrons to chlorine to form Al³⁺ and Cl⁻ ions.

The Al₂Cl₆ structure is explained using covalent and co-ordinate bonding, not complete electron transfer into isolated ions.

Aluminium atoms donate lone pairs to chlorine atoms.

Aluminium is electron-deficient in AlCl₃. It acts as the electron-pair acceptor rather than the lone-pair donor.

Lone pairs on chlorine can be donated to electron-deficient aluminium atoms, producing chlorine bridges.

Correct. Bridging chlorine atoms act as electron-pair donors, while the aluminium centres act as acceptors.

Aluminium atoms form an Al=Al double bond between the two AlCl₃ units.

The dimer is held together through bridging chlorine atoms rather than an Al=Al double bond.

Which statement about the N–H bonds in a completed NH₄⁺ ion is correct?

One bond is permanently weaker because it was formed by co-ordinate bonding.

The method of formation does not make one N–H bond permanently weaker.

All four N–H bonds are equivalent after NH₄⁺ has formed.

Correct. Co-ordinate bonding describes how one bond originated; in the completed ion the four N–H bonds are equivalent.

Three bonds are covalent while the fourth is ionic.

All four N–H connections in NH₄⁺ are covalent bonds.

The donated electron pair remains located entirely on nitrogen.

Once the bond forms, the pair is a shared bonding pair between nitrogen and hydrogen.

Why can a piece of solid copper conduct electricity?

Copper ions move freely through the solid lattice.

The positive metal ions remain around fixed lattice positions in the solid. They are not the mobile charge carriers.

Covalent bonds break repeatedly to release electrons.

Electrical conduction in a metal does not require covalent bonds to be continuously broken.

Copper contains alternating positive and negative ions.

The metallic model contains positive metal ions and delocalised electrons, not alternating positive and negative ions.

Delocalised electrons can move through the metallic lattice and carry charge.

Correct. The mobility of delocalised electrons allows a solid metal to conduct.

Which explanation best accounts for stronger metallic bonding in magnesium than in sodium?

Mg²⁺ has a higher charge and smaller radius than Na⁺, and each Mg atom contributes more delocalised electrons.

Correct. These factors increase the electrostatic attraction between the positive ions and the delocalised electron cloud.

Sodium atoms have no outer-shell electrons.

Sodium has one outer-shell electron, which becomes part of the delocalised electron system in the metallic model.

Magnesium contains negative Mg²⁻ ions that attract positive electrons.

Mg²⁺ ions are positively charged, while electrons carry negative charge.

Magnesium forms discrete Mg₂ molecules in the solid.

Metallic magnesium forms a giant metallic lattice rather than a molecular solid.

Which description of bonding in solid NH₄Cl is correct?

Only covalent bonding is present because NH₄⁺ contains covalent N–H bonds.

This ignores the electrostatic attraction between NH₄⁺ and Cl⁻ ions in the solid.

Only ionic bonding is present because NH₄Cl consists of ions.

Ionic attraction occurs between the ions, but covalent N–H bonds exist within the NH₄⁺ ion.

Covalent bonding occurs within NH₄⁺, while ionic attraction occurs between NH₄⁺ and Cl⁻.

Correct. Bonding must be described at the appropriate structural level: within the polyatomic ion and between the ions.

Metallic bonding occurs between nitrogen and chlorine atoms.

Metallic bonding requires positive metal ions and delocalised electrons; that is not the bonding present in NH₄Cl.

BF₃ can accept a lone pair from NH₃ to form an adduct. Which description of the newly formed bond is correct?

BF₃ + NH₃ → F₃B←NH₃

Boron donates a lone pair to nitrogen.

Boron in BF₃ is electron-deficient and acts as the electron-pair acceptor, not the donor.

A fluorine atom transfers one electron completely to nitrogen.

The new bond is formed by donation of an electron pair rather than complete transfer of a single electron.

One electron is supplied by nitrogen and one by boron.

In a co-ordinate bond both electrons in the newly formed shared pair are supplied by the donor.

Nitrogen donates its lone pair to electron-deficient boron, forming a co-ordinate bond.

Correct. NH₃ is the electron-pair donor and BF₃ is the electron-pair acceptor, so the co-ordinate arrow points from N towards B.

Frequently asked questions

Questions students often ask about chemical bonding

Once the four definitions are clear, the important next step is understanding how they connect to electron diagrams, structures and the language used in A-Level Chemistry.

Do atoms bond because they “want” a full outer shell?

A complete outer shell is a useful pattern for many main-group species, but it is not the fundamental physical reason for bonding. A bond forms when electrostatic interactions allow the bonded arrangement to be lower in energy and therefore more stable than the separated particles.

Why is electron transfer not the definition of an ionic bond?

Electron transfer explains how positive and negative ions may be produced. Once those ions exist, the ionic bond is the electrostatic attraction between the oppositely charged ions. Separating these two ideas makes the definition much more precise.

Why does NaCl not contain separate NaCl molecules?

In the solid, each Na⁺ ion is attracted to several surrounding Cl⁻ ions, and each Cl⁻ ion is attracted to several surrounding Na⁺ ions. These attractions continue throughout a giant ionic lattice. NaCl therefore gives the simplest ratio of ions rather than identifying a discrete molecule.

Why is a covalent bond still electrostatic if there are no ions?

Electrostatic attraction does not require complete ions. The nuclei carry positive charge because of their protons, while electrons carry negative charge. In a covalent bond, the shared electrons are attracted to both positively charged nuclei.

Are dots and crosses physically different electrons?

No. The symbols are only an electron-accounting device used to show where electrons originated. Electrons are indistinguishable, so an electron represented by a dot is not physically different from one represented by a cross.

Is a co-ordinate bond weaker than an ordinary covalent bond?

Not simply because both electrons originally came from the same atom. Co-ordinate bonding describes how the shared pair was supplied. Once formed, the bond is a covalent bond, and its actual strength depends on the atoms and bonding environment involved.

Why are all four N–H bonds equivalent in NH₄⁺?

One N–H bond can be described as being formed when nitrogen donates a lone pair to H⁺. After NH₄⁺ has formed, however, there is no permanent label on that electron pair. The four N–H bonds are equivalent in the finished ion.

What exactly does the arrow in a co-ordinate bond mean?

The arrow points from the electron-pair donor towards the electron-pair acceptor. It shows where both bonding electrons originated. It does not mean that electrons continually travel along the arrow after the bond is formed.

Are delocalised electrons in a metal completely free?

They are mobile throughout the metallic lattice and are not confined to one particular atom or bond. However, they remain part of the metal and remain electrostatically attracted to the positive metal ions. “Delocalised” is therefore more precise than simply saying “free”.

Can one substance contain more than one type of bonding?

Yes. NH₄Cl is a good example. Covalent N–H bonds occur inside NH₄⁺ ions, while ionic attraction occurs between NH₄⁺ and Cl⁻ ions. One of the N–H bonds can also be described as having formed initially by co-ordinate bonding when NH₃ donated its lone pair to H⁺.

Ionic, covalent, co-ordinate and metallic bonding become much easier to understand when you stop treating them as four unrelated definitions. In every case, begin by identifying the positive component, the negative component and the electrostatic attraction between them. Then ask how the electrons are arranged: transferred to form ions, shared between nuclei, donated as a lone pair or delocalised throughout a metal. Once that framework is secure, dot-and-cross diagrams, bonding definitions and explanations of structure become much more systematic.