Bonding decides structure — and structure decides every physical property
The big idea
BONDING
→
STRUCTURE
→
PHYSICAL PROPERTIES
Pure substances have different structures depending on the type of bonding they have, and the
structure of a substance decides what its physical properties will be. That is the whole unit in one line — every
melting point, every conductivity answer and every solubility prediction on this page comes from reading that arrow
left to right.
The exam skill. You will be given a substance and asked to predict its properties, or given properties and asked
to identify the structure. Both directions need the same three questions: what particles are there, what holds
them together, and can anything charged move?
The four structure types
Click any branch to jump to it. Note that covalent splits into simple and
giant — and those two behave almost as differently from each other as ionic does from metallic.
IONIC
Metal + non-metal. A giant lattice of cations and anions held by strong electrostatic attraction.
ex. sodium chloride, copper(II) oxide
COVALENT — SIMPLE
Small discrete molecules. Strong bonds inside each molecule, weak IMFs between them.
ex. carbon dioxide, water
COVALENT — GIANT / NETWORK
One continuous network of atoms joined by strong covalent bonds — no separate molecules at all.
ex. diamond, graphite, quartz — group 14 elements: C, Si
METALLIC
Positive metal ions in a regular lattice, surrounded by a mobile sea of delocalized electrons.
ex. copper, iron, gold
Graphite is treated as a fifth category throughout this unit. It is technically giant covalent, but its layered
structure gives it properties — soft, slippery, and electrically conducting — that break every other giant covalent rule.
See why →
Three questions that decide everything
1What are the particles? Ions, molecules, atoms, or cations-plus-electrons. This is what you are pulling apart when you melt or boil something.
2What holds them together — and how strongly? Strong bonds (ionic, covalent, metallic) mean a high melting point. Weak intermolecular forces mean a low one.
3Is anything charged and free to move? No mobile charge, no current. This single question answers every conductivity part of every question in the unit.
✓Answer those three and the whole summary table reconstructs itself from memory.
Cations + anions locked in a giant crystal lattice
What they are
Ionic substances are compounds of a metal and a non-metal — sodium chloride, copper(II) oxide, calcium oxide.
The metal atoms lose electrons to become cations (+) and the non-metal atoms gain them to become
anions (−). They are all solids at room temperature.
The + ions and − ions strongly attract each other by electrostatic force, and because that attraction
acts in every direction at once, the ions pack into a regular repeating crystal lattice — not into pairs or
molecules. There is no such thing as "one NaCl molecule"; the formula only gives the ratio of ions in the lattice.
Property 1 — high melting and boiling points, hard and brittle
Melting an ionic solid means overcoming the strong electrostatic attraction between every ion and all of its
neighbours. That takes an enormous amount of energy, so ionic compounds have high melting and boiling
points — boiling points typically run from about 800 °C to 3500 °C.
The same strong attractions make ionic crystals hard. But they are also brittle: hit the crystal hard
enough to shift one layer of ions by a single position, and suddenly + sits next to + and − next to −. The layers now
repel, and the crystal splits cleanly along that plane.
Ionic compound
Melting point (°C)
Iron(II) chloride, FeCl2
677
Potassium chloride, KCl
770
Sodium chloride, NaCl
801
Copper(II) oxide, CuO
1446
Calcium oxide, CaO
2707
Notice CaO. Ca2+ and O2− are doubly charged, so the attraction between them is far stronger
than between the singly charged ions in NaCl — and the melting point is more than three times higher.
Property 2 — conductivity: melt it or dissolve it
Ionic compounds are made of charged ions, so they can conduct electricity — but only if the ions are
free to move. In the solid, every ion is locked in the lattice: it can vibrate on the spot, but it cannot travel, so
no current flows. Two things set the ions free:
MELT it (≈ 800 °C). Enough heat breaks the lattice apart. The ions are still charged, and now
they can move — molten NaCl conducts.
DISSOLVE it (20 °C, in water). Polar water molecules surround each ion and pull it out of the
lattice. The ions drift freely in solution — NaCl(aq) conducts.
Melt vs. Dissolve — ion animator
Same compound, two completely different routes to the same result: mobile ions. Watch what happens to the
lattice in each case, and notice that the water route needs no heat at all.
Ions free to move?
No
Conducts electricity?
No
Temperature
20 °C
Property 3 — solubility depends on the compound
Some ionic substances are soluble in water; others are not. Whether water can pull the ions apart depends on how
strongly the lattice is holding them together — its lattice energy — compared with how much energy is released
when water molecules surround the freed ions.
Low lattice energy — water wins, the lattice comes apart.
NaCl is soluble.
High lattice energy — the ions hold on to each other, water cannot separate them.
PbCl2 is insoluble.
There is no trick or trend here — check the solubility chart. "It's ionic, so it dissolves" is one of the most common
errors in this unit. Ionic means it might dissolve; only the chart tells you whether it actually does.
Ionic — the summary row
Property
Ionic substances
Because…
State at RT
All solids
Strong attractions in every direction
MP / BP
High (BP ≈ 800–3500 °C)
Strong electrostatic attraction between ions
Hardness
Hard but brittle
Shifted layers put like charges together, which repel
Conducts — solid
NO
Ions are fixed in the lattice; they cannot move
Conducts — molten
YES
Ions are free to move and carry charge
Conducts — dissolved
YES
Ions are free to move and carry charge
Soluble in water?
Check the chart
Depends on the lattice energy of that compound
Simple Covalent (Molecular) Substances
Strong bonds inside the molecule — weak forces between molecules
What they are
These are substances like carbon dioxide CO2, water H2O and methane CH4, which are always made
of simple, separate molecules — whether they happen to be solids, liquids or gases at room
temperature. Freezing water does not join the molecules together; it just packs the same H2O molecules into a
fixed arrangement.
The two forces — and why students mix them up
VERY STRONG bonds BETWEEN ATOMS. The covalent bonds holding H to O inside a water
molecule are strong, so the molecule itself is very hard to break up.
WEAK attractions BETWEEN MOLECULES. The intermolecular forces (IMFs) holding
one water molecule to the next are weak, so the molecules are easy to separate — i.e. easy to change state.
When you boil water you do not break a single O–H bond. You only pull whole H2O molecules away from
each other. That is exactly why the boiling point is 100 °C and not several thousand — melting and boiling break IMFs,
not covalent bonds. Getting this the wrong way round is the classic error in this unit.
Property 1 — low melting and boiling points, soft when solid
Compound
MP (°C)
BP (°C)
State at RT
Water, H2O
0
100
Liquid
Butane, C4H10
−138
−0.5
Gas
Methane, CH4
−182
−164
Gas
Carbon dioxide, CO2
—
−78
Gas
Oxygen, O2
−218
−183
Gas
Hydrogen, H2
−259
−252
Gas
CO2 has no ordinary melting point at normal pressure — at −78 °C solid CO2 (dry ice)
sublimes, going straight from solid to gas without ever being a liquid.
Because only weak IMFs have to be overcome, simple molecular substances have low melting and boiling points —
boiling points run from about −270 °C to 600 °C, and most of them are gases at room temperature.
For the same reason, the ones that are solid are weak and soft. Ice shatters under a light tap; solid
carbon dioxide crumbles. Compare that with an ionic crystal or a diamond.
Watch the pattern. Bigger molecules have more surface contact, so their IMFs are stronger and their boiling
points are higher: H2 (−252 °C) → CH4 (−164 °C) → C4H10 (−0.5 °C). Water breaks the
pattern because it is polar — that is what makes it a liquid at all.
Property 2 — they never conduct electricity
As the molecules are NOT CHARGED, simple molecular substances do not conduct electricity in
any state — solid, liquid or gas. There are no free ions and no free electrons; every electron is either in a
covalent bond or in a lone pair on an atom. Melting them does not help: you get mobile molecules, not mobile
charges.
The pure liquid is the giveaway. Pure water is a very poor conductor. Tap water conducts because of the
dissolved ionic compounds in it — which is why a hair dryer in the bath is dangerous even though H2O
itself is a molecular substance.
Property 3 — polarity decides solubility
POLAR molecules are soluble in water (water is itself polar), while NON-POLAR molecules are insoluble.
This is the "like dissolves like" rule. To use it you need the molecule's shape (from VSEPR), not just
its bonds — CO2 has polar bonds but is a non-polar molecule, because its linear shape makes the two bond dipoles
cancel exactly.
Polarity & Solubility Predictor
Pick a molecule. The tool shows its VSEPR shape, whether the individual bonds are polar, whether the whole
molecule ends up polar once the shape is taken into account, and therefore whether it dissolves in water.
Molecule
VSEPR shape
Bond type
Molecule polarity
Soluble in water?
Polar bonds ≠ polar molecule. Both CO2 and H2O contain polar bonds. CO2 is
linear, so the two dipoles point in exactly opposite directions and cancel — the molecule is non-polar and insoluble.
H2O is bent, so its dipoles do not cancel — polar, and the best solvent you own.
Simple covalent — the summary row
Property
Simple covalent substances
Because…
State at RT
Solid, liquid or gas — usually gas
Weak IMFs are easily overcome
MP / BP
Low (BP ≈ −270 – 600 °C)
Only weak IMFs between molecules have to break
Hardness
Soft and weak when solid
Weak attractions between the molecules
Conducts — any state
NEVER
Molecules are not charged; no free ions or electrons
Soluble in water?
Check the polarity (VSEPR)
Polar molecules dissolve in polar water; non-polar ones do not
Giant / Network Covalent Substances
One molecule the size of the whole crystal
What they are
In these materials strong covalent bonds join atoms to other atoms of the same type to make giant structures,
rather than little groups. There are no separate molecules at all — a diamond is, in effect, a single molecule
containing every carbon atom in the stone. They are typically built from group 14 elements: carbon and silicon.
DIAMOND. Every carbon atom is joined to 4 others by strong covalent bonds, in
a rigid three-dimensional tetrahedral network extending through the whole crystal.
SILICA, SiO2. A similar structure to diamond — every Si atom is joined to
4 O atoms. Silica is the main substance in rocks; pure silica is called quartz.
Properties — everything follows from "strong bonds everywhere"
1High melting and boiling points (BP ≈ 2200 °C – 4000 °C). To melt one you must break millions of strong covalent bonds — not weak IMFs.
2Usually hard and strong. Every atom is locked to its neighbours in three dimensions, so nothing can shift. Diamond is the hardest natural substance known.
3Solid at room temperature — always.
4Never conduct electricity. All the atoms are uncharged, and every electron is tied up in a covalent bond, so there is nothing free to carry a current.
5Insoluble in water — always. Dissolving would mean breaking those strong covalent bonds, and water cannot do it.
Graphite — the special case
Graphite is the common form of carbon found in soot, charcoal and pencils — and it breaks nearly every giant covalent rule
above. The reason is its layered structure: each carbon atom is joined to only 3 others by strong
covalent bonds, forming flat hexagonal sheets. Those sheets are then held to each other by nothing but
weak IMFs.
3 bonds, not 4 — so each carbon has one electron left over. Those spare
electrons form clouds of free ("delocalized") electrons between the layers.
Weak IMFs between the sheets — so the layers slide over each other very
easily, even though the sheets themselves are extremely strong.
Graphite — layer slider
Drag the force slider to push sideways on the top layer, and switch on the voltage to send the delocalized electrons
drifting. Notice that the sheets themselves never break — only the weak forces between them give way.
Graphite's properties — and what each one comes from
Property
Graphite
Because…
MP / BP
High
Strong covalent bonds between the atoms within each sheet
Hardness
Soft and slippery
Weak IMFs between the layers, so they slide over each other easily — this is why it writes on paper
Conducts electricity
YES (fairly well)
Only 3 of the 4 outer electrons are bonded, so free electrons move between the layers — used for sliding contacts in electric motors
Soluble in water?
Never
Non-polar, and held together by strong covalent bonds
The classic exam question: why does graphite conduct but diamond does not? In diamond every carbon uses all
four outer electrons in covalent bonds, so none is free. In graphite each carbon bonds to only three
others, leaving one delocalized electron per atom free to move between the layers. Same element — completely different
structure, completely different property.
Giant covalent — the summary rows
Property
Giant / network
Graphite
State at RT
All solids
Solid
MP / BP
High (BP ≈ 2200–4000 °C)
High
Hardness
Hard & strong
Soft & slippery
Conducts
NEVER
YES
Soluble in water?
Never
Never
Metallic Substances
Positive ions in a sea of delocalized electrons
The structure
In a metal, the atoms lose several of their outer electrons, which then drift around between the metal ions as
free "delocalized" electrons. What is left is a regular repeating lattice of positive metal ions
immersed in a mobile sea of electrons. The attraction between those cations and the electron sea is the
metallic bond.
The atoms are held together by strong bonds in regular structures (body-centred cubic, face-centred cubic,
hexagonal — you do not need to memorise the names). This means most metals have high melting and boiling
points, with boiling points running from roughly 700 °C to 5600 °C.
Metals are all solids at room temperature except mercury, Hg(l). They are usually hard and strong,
though hardness varies a lot — sodium can be cut with a knife, while tungsten is used for drill bits.
Sea of electrons — interactive
Switch between the three modes to see where each metallic property comes from.
The two properties the electron sea explains
GOOD CONDUCTORS — of electricity AND heat. The large number of free electrons can drift
through the whole structure. They carry charge (electricity) and they also carry kinetic energy from a hot end to a cold
end (heat). No other structure type conducts both this well.
MALLEABLE and DUCTILE. Because the structure is regular and the electron sea is
non-directional, layers of ions can slide over each other without breaking the bonds (though not as easily as
graphite's layers). So metals bend rather than snap (malleable) and can be drawn into wires (ductile).
Compare with ionic. Slide the layers of an ionic crystal and like charges end up facing each other, so it shatters —
brittle. Slide the layers of a metal and every cation is still surrounded by the same electron sea, so nothing repels — it
just changes shape. Same action, opposite outcome, and the reason is the structure.
Alloys
An alloy is a mixture of chemical elements where the primary component is a metal. Alloying can result in
the physical properties being more desirable than those of the pure metal — usually harder and stronger.
SUBSTITUTIONAL. Atoms of one metal are substituted with others of equivalent
size. Examples: brass, bronze. Because the radii are comparable, substitutional alloys show
less variation in properties than interstitial ones.
INTERSTITIAL. Small atoms are trapped within the gaps of the metal crystal
matrix. Example: steel (small carbon atoms in iron). Both kinds can happen at once — stainless steel
has both atom exchange and interstitial placement.
Why alloys are harder: a pure metal has regular layers that slide easily, so it is soft and malleable. Adding
differently sized atoms distorts the layers, so they can no longer slide past each other cleanly — the alloy is
much harder (and correspondingly less malleable).
Alloy Mixer
Load one of the named alloys, or build your own: choose a base metal, an alloying element, and whether it goes in by
substitution or into the interstices. Watch how much the layers get distorted.
Presets:
Base metal
Alloying element
Arrangement
Amount added: 10%
Pure metal
Pure
Hardness
Soft
Malleability
Very high
Layer distortion
None
Reactivity — a chemical property worth knowing here
1Very active metals (alkali metals and some alkaline earth metals) react with water or steam.
2Somewhat active metals will only react with acids, producing hydrogen gas:
Mg + 2 HCl → H2(g) + MgCl2
3Less active metals (ex. Ag, Au) will not react with water, nor with acids — which is exactly why they are used for jewellery and electrical contacts.
✓Which tier a metal falls into is read off the activity series — the same table you used in the Chemical Reactions unit.
Regular structure lets layers slide without breaking bonds
Conducts
YES — heat & electricity
Free delocalized electrons between the ions
Structure Visualizer
Drag to rotate — the structural reason behind every property
Lattice Viewer
Drag the model to rotate it yourself
What to look for in each
Structure
Look for
Property it explains
Ionic NaCl
Alternating + and − ions, each surrounded by 6 of the opposite charge
High MP, hard but brittle, conducts only when the ions are freed
Giant Diamond
Every carbon joined to 4 others in a rigid 3D web
Extremely high MP, hardest natural substance, no free electrons ⇒ insulator
Giant Silica
Every Si joined to 4 O atoms — the same idea, two elements
High MP, hard, insulating — glass and quartz
Graphite
Flat hexagonal sheets, wide gaps between the layers
Soft & slippery (layers slide), but conducts via the free electrons in the gaps
Metallic
Regular cations with electrons moving freely between them
Conducts heat & electricity, malleable and ductile
Conductivity State Simulator
The one question that decides it: is anything charged AND free to move?
Build the circuit
Choose a structure type, then choose the state it is in. The bulb lights only if there are charged particles that
are free to move — mobile ions or delocalized electrons.
Structure:
State:
Charge carriers
—
Free to move?
—
Bulb
Off
All five categories at once
Structure
Solid
Molten / liquid
Dissolved in water
Charge carrier
Ionic
NO
YES
YES
Mobile ions
Simple covalent
NO
NO
NO
None — molecules are neutral
Giant covalent
NO
NO
n/a — insoluble
None — all electrons bonded
Graphite
YES
YES
n/a — insoluble
Delocalized electrons between the layers
Metallic
YES
YES
n/a
Delocalized electrons throughout
The two rows that look alike are the two you must be able to separate. An ionic solid and a metallic solid are both
solid, both made of ions, and both have high melting points — but only the metal conducts, because only the metal has a
separate sea of electrons. In the ionic solid the electrons were handed over completely and are locked onto the
anions.
Melting & Boiling Point Comparator
Four categories are high — one is dramatically lower
Typical boiling point ranges
Every bar below is drawn on the same linear temperature axis. Drag the thermometer to sweep a temperature
across all five categories at once.
25 °C
Why the ranges sit where they do
Structure
BP range
What has to break to boil it
Simple covalent
≈ −270 – 600 °C
Only the weak IMFs between molecules — the covalent bonds inside each molecule stay intact
Metallic
≈ 700 – 5600 °C
Strong metallic bonds between the cations and the electron sea
Ionic
≈ 800 – 3500 °C
Strong electrostatic attraction between every ion and its neighbours
Giant covalent
≈ 2200 – 4000 °C
Millions of strong covalent bonds throughout the network
Graphite
≈ 3600 – 4200 °C
The covalent bonds within the sheets — the weak IMFs between layers do not set the boiling point
The one that stands out is simple covalent, and it stands out by thousands of degrees. That is the single most
useful fact for identifying a structure from data: if the melting point is below a few hundred degrees, it is a
simple molecular substance. Nothing else comes close.
Graphite deserves a second look. It is soft enough to write with, yet it has one of the highest boiling points of
any element. Softness comes from the weak IMFs between layers; the boiling point comes from the strong covalent
bonds within them. Two different forces, two different properties — which is exactly why graphite is a special case.
Structure Classifier
Name or formula in — structure type and full property set out
Classify a substance
Type a name or a formula — sodium chloride, NaCl, CO2, diamond,
copper — and press Classify. Anything not in the built-in list is worked out from the elements in the formula,
the same way you would do it in an exam.
Particles present
What holds them
State at RT
MP / BP
Hardness
Conductivity
Soluble in water?
How we knew
How to classify by hand
1One element only? If it is a metal → metallic. If it is C or Si → check which form: diamond/silicon are giant covalent, graphite is the special case. Any other non-metal (O2, N2, I2, S8) → simple covalent.
2Metal + non-metal? → ionic. A polyatomic ion such as NO3− or SO42− counts as the non-metal part, and NH4+ counts as the metal part.
3Non-metals only? → covalent. Then decide simple vs giant: SiO2 and SiC are giant networks; almost everything else you will meet is simple molecular.
4A mixture of metals? → an alloy, so still metallic, but harder and less malleable than the pure metal.
✓Then read the properties straight off the summary table. You never have to memorise the properties of an individual substance — only its structure.
"Predict the Properties" Practice
Substance in — melting point, conductivity and solubility out
Round 1
Score: 0 / 0 ·
Predict the properties of sodium chloride
1 · Structure type
2 · Melting / boiling point
3 · Conducts electricity?
4 · Soluble in water?
Reasoning practice — pick the correct explanation
Score: 0 / 0 ·
Summary — Properties
The master reference for the whole unit
The master table
Structure
Property
Reason
Ionic ALL solids
· HIGH mp/bp; hard, brittle · CONDUCT: solid NO, molten YES, dissolved YES · Soluble in water? Check the solubility chart
· Strong attractions between IONS · Ions can't move in the solid · Ions can MOVE to carry current · Differing lattice energies
Covalent — simple (s, l or g)
· LOW mp/bp (usually gas at RT); soft when solid · CONDUCT: never · Soluble in water? Check polarity using VSEPR
· Attractions (IMFs) between MOLECULES are very weak · No free ions or electrons · Molecules are NOT charged · Polar molecules dissolve in polar solvents (water)
Covalent — giant/network ALL solids
· HIGH mp/bp; hard & strong · CONDUCT: never · Soluble in water? Never
· Strong bonds between all ATOMS · No free ions or electrons · Structures are non-polar
Covalent — graphite solid
· HIGH mp/bp; soft & slippery · CONDUCT: YES (fairly well) · Soluble in water? Never
· Strong bonds between ATOMS; weak IMFs between LAYERS · Free electrons between the layers · Structure is non-polar
Metallic ALL solids except Hg(l)
· HIGH mp/bp; usually hard & strong but variable · Malleable, ductile · CONDUCT: YES (heat & electricity)
· Strong bonds between CATIONS · Regular structure, so layers slide · Free electrons between the ions
Identify a structure from data
1Low melting point? It is simple covalent. Nothing else melts below a few hundred degrees.
2High MP and conducts as a solid?Metallic — or graphite if it is also soft, slippery and made only of carbon.
3High MP, does not conduct as a solid but does when molten or dissolved?Ionic. That switching behaviour is unique to ionic.
4High MP, hard, and never conducts at all?Giant covalent.
✓Those four questions identify every structure in the unit, in that order.
Common mistakes to avoid
✗Saying that boiling water breaks covalent bonds. It only breaks the weak IMFs between molecules.
✗Saying an ionic compound conducts because it is ionic. It conducts only when the ions are free to move — molten or dissolved, never as a solid.
✗Assuming all ionic compounds dissolve. Some do, some don't — check the solubility chart.
✗Saying diamond and graphite behave the same because both are carbon. Structure, not element, decides the properties.
✗Calling a substance polar because it has polar bonds. Use the VSEPR shape — CO2 has polar bonds but is a non-polar molecule.
✗Saying metals are brittle like ionic solids because both are made of ions. Sliding metal layers keeps every cation in the same electron sea, so metals bend instead of shattering.
✗Describing giant covalent substances as having "strong intermolecular forces". There are no molecules — so there are no intermolecular forces at all.
Where to next
See all of this in use on the Real-World Applications
page — pencils and drill bits, alloys in your kitchen, sports drinks and batteries, glass and silicon chips — then test
yourself with the Kahoot quiz.