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Positive ions in a sea of pooled electrons, and an honest account of what all three bonding models fail to explain.
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By the end of this lesson you will be able to describe a metal as a lattice of positive ions in a sea of delocalized electrons, and use that picture to account for conduction in the solid, malleability, a high melting point and a shine. You will also be able to compare the metallic, ionic and simple molecular models across those same properties, say what has to be overcome to melt each kind of solid, and name something each of the three models cannot explain.
You have two models. In the ionic one, a metal hands its outer electrons to a non-metal and the charged particles left behind attract each other. In the covalent one, two non-metals share a pair and both nuclei hold on to it. What neither of them covers is a lump of metal on its own: there is no non-metal to hand the electrons to, and no second kind of atom to share with. So what holds a piece of copper together?
| Term | What it means |
|---|---|
| Delocalized electrons | Electrons that belong to no one atom and move through the whole solid. |
| Sea of electrons | The pool of delocalized electrons in a metal. |
| Metallic bond | The attraction between the lattice of positive ions and the sea. |
| Malleable | Able to be hammered into a sheet without shattering. |
| Ductile | Able to be drawn into a wire. |
| Giant covalent | A solid with covalent bonds all the way through and no molecules. |
A metal atom has few outer electrons and holds them loosely. Put a great many metal atoms together and each of them lets its outer electrons go — but there is nothing to give them to, so the electrons simply stop belonging to any one atom. They are delocalized: free to move anywhere in the solid.
What is left is a regular lattice of positive ions sitting in a sea of negative electrons. Every ion is attracted to the sea and the sea is attracted to every ion, and that mutual attraction is the metallic bond. Notice what it is not: it is not a bond between two named atoms. There are no pairs here, and no partners.
Almost everything a metal does follows from those two features.
Another way: picture
Think of marbles in molasses. The marbles are the positive ions, sitting in a regular pattern; the molasses is the electron sea, filling every gap and sticking to every marble. Tip the tray and the marbles slide past one another without the molasses letting go of any of them — which is a metal being hammered flat, and is the whole reason a metal changes shape where a salt crystal splits.
Another way: steps
To decide which model a solid uses, and what to expect of it:
Classify the solid. Metal atoms only: metallic. Metal with non-metal: ionic. Non-metals only: covalent, then decide whether the bonds run right through (giant) or stop at small molecules (simple).
Ask what is charged and free to move. Electrons in a metal, always. Ions in an ionic compound, only once melted or dissolved. Nothing in a simple molecular solid.
Ask what melting must overcome. A lattice held by an electron sea, a lattice of full charges, attractions between whole molecules, or covalent bonds throughout.
Predict the properties. Conduction from the free charges; melting point from what must be overcome; malleability from whether a slid layer faces like charges.
Check the work. Have you said what moves when a current flows, and not just that it conducts? Have you avoided assuming that every metal melts high? Sodium melts below the boiling point of water. And if a prediction disagrees with a measured value, is the substance one of those between the boxes, such as aluminum chloride?
Classifying by the kinds of atom is allowed because whether atoms give up, take or share electrons depends on how hard they hold them. Two metals both let go; a metal and a non-metal hand over; two non-metals share.
Linking conduction to free charges is allowed because a current is moving charge, by definition. If nothing charged can move, nothing can carry a current, however polar or charged the particles are.
Linking melting point to what must be overcome is allowed because melting supplies enough energy to let the particles move past one another. The stronger the attractions holding them in place, the more energy, and so the higher temperature, it takes.
Predicting more strength from more outer electrons is allowed because each ion then carries a larger positive charge and the sea holds more electrons. Both make the attraction between ions and sea stronger, and the measured melting points of sodium, magnesium and aluminum rise in that order.
Diamond is carbon, a non-metal, bonded covalently — and it melts at about $3550\,^\circ\mathrm{C}$, which is nothing like iodine's $114\,^\circ\mathrm{C}$. Both are covalent. The model as stated so far cannot tell them apart.
The difference is not the bond, it is what the bonds build. In iodine the covalent bonds stop after two atoms, and the solid is a heap of separate $\mathrm{I_2}$ molecules held to each other by something much weaker. Melting it separates molecules and breaks no bonds at all. In diamond every carbon is bonded to four others and those to four more, on and on through the whole crystal: there are no molecules, and melting it means breaking covalent bonds.
| Structure | What melting has to overcome | Melting point |
|---|---|---|
| simple molecular (iodine) | attractions between whole molecules | $114\,^\circ\mathrm{C}$ |
| metallic (sodium) | a lattice held by one electron per atom | $98\,^\circ\mathrm{C}$ |
| metallic (aluminum) | a lattice held by three electrons per atom | $660\,^\circ\mathrm{C}$ |
| ionic (sodium chloride) | a lattice of full opposite charges | $801\,^\circ\mathrm{C}$ |
| giant covalent (diamond) | covalent bonds, throughout | $\approx 3550\,^\circ\mathrm{C}$ |
Sodium melting lower than iodine is the row worth staring at. Metallic does not mean strong, and the melting point of a solid is a fact about what has to be overcome, not about which model has the more impressive name.
The ionic model assumes the electrons went across completely and the ions are hard spheres. In many real compounds they did not quite. A small, highly charged cation next to a large, soft anion pulls the anion's electron cloud toward itself, and the bond acquires some covalent character — aluminum chloride is written as an ionic formula and behaves in many ways like a molecular substance, melting at $192\,^\circ\mathrm{C}$ rather than in the hundreds.
The covalent model assumes the pair is shared equally. Between two different atoms it is not, and the more unequal the sharing, the more the bond resembles an ionic one. There is no line to be drawn: hydrogen chloride, water and sodium chloride are three points on one scale.
The metallic model is the vaguest of the three. A sea of electrons explains conduction and malleability well. It says nothing at all about why tungsten is hard and lead is soft, why some metals are magnetic, or why a transition metal has the colors it has. Those need a picture of the electrons that is far more detailed than a sea.
So the honest summary is that bonding is a spectrum, not three boxes. The three models mark the corners of it, and a great many useful substances live somewhere in the middle. A model is judged by what it lets you predict, and the moment you can say what a model fails to predict, you are using it properly rather than reciting it.
Nearly every house in the United States is wired with copper, and the metallic model explains why. Copper's outer electrons are delocalized, so a current flows through solid wire with nothing but electrons moving; the copper ions stay put for decades, which is why a wire installed in the 1950s is still the same wire.
For a time in the 1960s and 1970s, when copper was expensive, many American homes were wired with aluminum instead. Aluminum is also metallic and conducts well, though only about 61 percent as well as copper for the same thickness, so the wires had to be larger. The trouble came at the connections. Aluminum's surface reacts with air to form aluminum oxide, an ionic solid in which nothing is free to move until it melts, so a thin oxide skin at a screw terminal resists the current and heats up.
That heating, repeated over years, caused fires, and the Consumer Product Safety Commission still warns owners of older homes to have aluminum connections inspected. The lesson is the one in this lesson's table: the metal conducts because its electrons are free, and its oxide does not, because in an ionic solid nothing charged can move. Utilities still use aluminum for long-distance power lines, where its lightness matters and the connections are designed for it.
Foundries in Pittsburgh and Detroit size their furnaces to the metal: zinc melts at 420 degrees Celsius, aluminum at 660, copper at 1085 and iron at 1538. Each step up is a stronger metallic lattice to loosen.
A metal is said to conduct because its ions move. This is the big one. In a solid metal the ions are the lattice and stay where they are; the electrons move. Moving ions are the right answer for a molten or dissolved ionic compound, and getting the two the wrong way round loses both.
A metallic bond is described as between two atoms. It is not a pair and it has no partners. Every ion is attracted to the whole pool, which is exactly why the lattice survives being hammered out of shape.
Metals are assumed to melt high and non-metals low. Sodium melts below the boiling point of water and below solid iodine; tungsten melts above $3400\,^\circ\mathrm{C}$. The question is always what has to be overcome.
Melting a simple molecular solid is said to break its bonds. It breaks none. It separates whole molecules from each other, and every covalent bond inside them survives — which is why melted iodine is still iodine.
Every solid is forced into one of three boxes. Bonding is a spectrum; aluminum chloride, written ionically, behaves much like a molecular substance.
Count sodium's contribution to the sea.
$1 \text{ electron per atom}; \ \mathrm{Na^+}$
The group number is what each atom contributes.
Count magnesium's contribution.
$2 \text{ electrons per atom}; \ \mathrm{Mg^{2+}}$
Twice the charge on each ion.
Compare the seas.
$\text{magnesium's is twice as dense}$
More electrons given up per atom.
Predict the melting points.
$\text{magnesium higher}$
A bigger charge pulling on more electrons is a stronger hold.
Check against measurement.
$650 - 98 = 552 \ ^\circ\mathrm{C} \text{ higher}$
The model predicted the direction.
Set up the same blow for both.
$\text{one layer slides by one particle}$
The starting situation is the same.
Look at the ionic layers.
$+ - + - \text{ alternating}$
Positive and negative ions take turns.
Slide the ionic layer.
$\text{positive now faces positive}$
Like charges come face to face.
Read the ionic result.
$\text{the crystal splits}$
The lattice pushes itself apart.
Slide the metal layer.
$\text{ions still face the electron sea}$
Every layer is positive ions in the same sea.
Read the metal result.
$\text{the lattice holds; the shape changes}$
Nothing repels.
Take a solid that conducts when solid.
$\text{copper: metallic}$
Only a metal has charges free in the solid.
Take one that conducts only when molten.
$\text{sodium chloride: ionic}$
Its ions are freed by melting.
Take one that never conducts and melts low.
$\text{iodine, } 114\,^\circ\mathrm{C}: \text{ simple molecular}$
Melting separates molecules only.
Take one that never conducts and melts very high.
$\text{diamond}, \approx 3550\,^\circ\mathrm{C}: \text{ giant covalent}$
Melting breaks covalent bonds.
Check the melting points rise as expected.
$114 < 801 < 3550$
Simple molecular, then ionic, then giant covalent.
Place copper among them.
$1085\,^\circ\mathrm{C}$
A metallic lattice with a dense sea.
Note the exception to watch for.
$\text{sodium}, 98\,^\circ\mathrm{C}$
Metallic does not always mean high-melting.
Name the model.
$\text{metallic}: \mathrm{Al^{3+}} \text{ in a sea}$
Metal atoms only means a metallic lattice.
Account for conduction and melting.
Account for rolling into foil.
Match each property of a solid to the part of its bonding model that accounts for it.
| the delocalised electrons move through the lattice while the ions stay put | melting or dissolving lets the ions move, and a moving charged ion is a current | the free electrons at the surface take up light and give it straight back out | melting separates whole molecules, and what holds those together is not a bond | |
|---|---|---|---|---|
| conducts electricity while it is still solid | ||||
| conducts once it is melted or dissolved | ||||
| is shiny where it has just been cut | ||||
| melts below the temperature of a cup of tea |
Complete the worked solution: sodium melts at $98$ degrees Celsius, magnesium at $650$ and aluminum at $660$. Each atom gives one, two and three electrons to the sea. Find how much higher magnesium melts than sodium, aluminum than magnesium, and aluminum than sodium.
Compare magnesium with sodium.
$(\text{magnesium}) - (\text{sodium}) =$ a
Two electrons per atom against one: a far stronger hold.
Compare aluminum with magnesium.
$(\text{aluminum}) - (\text{magnesium}) =$ b
A third electron adds much less here.
Compare aluminum with sodium.
$(\text{aluminum}) - (\text{sodium}) =$ c
The sum of the two steps.
A copper wire has been carrying a current for years and has not changed in any way you could weigh. What is moving through it while the current flows?
Put these four solids in order of melting point, lowest first. Each is a different kind of structure, and what has to be overcome to melt it is different in each case.
Number the steps in order (write the number in the box):
A tiny cluster of $8$ magnesium atoms forms a metallic lattice. Each atom has $2$ outer electron(s). How many delocalized electrons does the sea hold?
Answer:
A foundry in Pittsburgh casts both copper, which melts at $1085$ degrees Celsius, and zinc, which melts at $420$. How many degrees hotter must the furnace run to melt the copper?
The answer: a.
A notice pinned up in a workshop explains why the wiring is copper and why the furnace liner is not. Exactly one sentence gets the chemistry wrong. Mark it.
This task has no paper form; do it on a device.
Lesson test: one question per skill, one attempt each, no hints. Your answers are checked when you submit.
Fill in the comparison. For each of the three kinds of solid, say whether it conducts electricity while solid, whether it conducts once melted, whether its melting point is high or low, and what does the moving when a current flows.
| conducts while solid? | conducts when molten? | melting point | what moves when a current flows | |
|---|---|---|---|---|
| a metal, such as copper | ||||
| an ionic solid, such as sodium chloride | ||||
| a simple molecular solid, such as iodine |
You can say what is moving when a metal conducts, and what is moving when a molten salt does — and they are not the same thing. Say out loud why sodium melts lower than iodine, and name one property of metals that a sea of electrons does not account for. Next: the forces between whole molecules, which are what melting a simple molecular solid actually overcomes.
15. Your turn: aluminum conducts electricity, melts at 660 degrees Celsius, and can be rolled into kitchen foil. Which model, and which feature of it accounts for each?, step 2
$\text{free electrons}; \ 3+ \text{ ions, dense sea}$
Free charges conduct; a strong hold melts high.
15. Your turn: aluminum conducts electricity, melts at 660 degrees Celsius, and can be rolled into kitchen foil. Which model, and which feature of it accounts for each?, step 3
$\text{layers slide; the sea holds them}$
No like charges meet when a layer moves.