2.1 — Matter

This section covers the fundamental nature of matter — what everything in the physical world is made of. Understanding atomic structure, chemical bonding, and the states of matter is essential for aircraft maintenance engineers, since the materials used in aviation (metals, composites, fluids, gases) all behave according to these principles.

Nature of Matter: Elements, Atoms, and Molecules

What Counts as Matter

Matter is anything that has mass and occupies space. Both halves of that test have to be met, and it is the definition examiners lean on: a forged undercarriage leg, a litre of hydraulic fluid and the nitrogen charge in a tyre are all matter, because each has mass and each takes up volume. Heat, light, sound and electrical energy are not matter. They have no mass and occupy no space, although they routinely change the temperature, motion or state of matter that does — which is exactly what the rest of this module is about.

Two quantities are easy to confuse at this point. Mass is the quantity of matter in a body. It is the same on the ramp, at 40,000 ft and in orbit, and it is what resists a change of motion. Weight is the force gravity exerts on that mass, so it varies with the strength of the gravitational field. An engineer reading a spring balance is measuring weight and inferring mass from it; the mechanics section of this module develops the relationship properly. Volume is the space the matter occupies, and mass divided by volume gives density — the property that decides whether a material earns its place in an airframe, and the reason two metals of comparable strength are not equally attractive to a designer.

Matter is organised in a hierarchy, and a surprising number of exam questions are really asking which level of that hierarchy a term belongs to:

• Sub-atomic particles — protons, neutrons and electrons. They exist and they determine everything above them, but on their own they are not a chemical substance you could weigh out or react.
• Atom — the smallest particle of an element that can exist on its own and still be that element. This is the answer the exam wants when it asks for the smallest particle of matter that can exist in nature. Physicists can name smaller objects — quarks inside the proton, for instance — but a quark is never found in isolation, so it is not something that can exist by itself as matter.
• Molecule — two or more atoms chemically bonded together. A molecule built from identical atoms is still an element; one built from different atoms is a compound.
• Element, compound and mixture — the bulk substances an engineer actually handles, stores and signs for.

One rule underpins the whole of this section: in any process an engineer meets, matter is neither created nor destroyed, only rearranged. Melt a kilogram of ice and you have a kilogram of water. Burn a kilogram of kerosene and the exhaust products weigh more than the fuel did, because oxygen drawn from the air has combined with it — nothing has vanished, and the extra mass came in through the intake. Whenever a question seems to hinge on something appearing or disappearing, the answer is almost always an audit of where the particles went.

Chemical Elements

An element is a pure substance that cannot be broken down into simpler substances by chemical means. There are 118 known elements, each identified by an atomic number (the number of protons in its nucleus). Elements are organised in the Periodic Table.

Elements important in aviation include:

Element	Symbol	Atomic No.	Aviation Use
Aluminium	Al	13	Primary airframe material (lightweight, corrosion-resistant)
Titanium	Ti	22	Engine components, high-strength fasteners
Iron	Fe	26	Steel alloys for landing gear, engine parts
Chromium	Cr	24	Plating, stainless steel alloys
Copper	Cu	29	Electrical wiring and connectors
Nitrogen	N	7	Tyre inflation, fuel tank inerting
Oxygen	O	8	Crew and passenger breathing systems
Hydrogen	H	1	Lightest element; component of fuels and water

Reading the Periodic Table

The Periodic Table is not simply a list of the elements — it is arranged so that an element's position predicts its behaviour, and that is what makes it worth learning rather than memorising. The elements are laid out in order of increasing atomic number, one proton at a time, reading left to right along each row and then down to the next. (Early versions ordered them by atomic weight, which put a handful of pairs in the wrong sequence; ordering by proton count fixed that and is the arrangement in use today.)

• A period is a horizontal row. Each new period begins a new outer electron shell, so an element's period number tells you how many shells are in use.
• A group is a vertical column. Every member of a group has the same number of electrons in its outer shell, which is why they behave chemically alike: sodium and potassium sit in the same group and both form chlorides of the form NaCl and KCl.

The horizontal position separates the two great families. Metals occupy the left and centre. They hold few outer electrons, part with them readily to form positive ions, and in bulk are lustrous, malleable, ductile and good conductors of both heat and electricity. Non-metals occupy the upper right. Their outer shells are nearly full, so they gain or share electrons instead of losing them; as solids they tend to be dull and brittle, several are gases at room temperature, and as a family they are poor conductors of heat and electricity. Between the two runs a diagonal staircase of metalloids — boron, silicon, germanium, arsenic — whose conductivity sits between the extremes and which are the raw material of the semiconductor industry.

The extreme right-hand column is the special case: the noble gases (helium, neon, argon, krypton, xenon and radon), whose outer shells are already full. They are non-metals, so the general statement that the elements at the far right of the table are non-metals holds all the way to the last column.

Region of the table	Outer-shell electrons	What the atom does with them	Typical members	Where the engineer meets them
Groups 1–2 (reactive metals)	1 or 2	Loses them, forming a positive ion	Lithium, sodium, magnesium, calcium	Lithium cells; magnesium alloy castings
Transition metals (centre block)	Variable	Loses a variable number; alloys readily	Titanium, chromium, iron, nickel, copper	Engine and airframe alloys, plating, wiring
Group 13	3	Loses three, forming \( \text{Al}^{3+} \)	Aluminium	The principal airframe metal
Group 14	4	Shares four rather than losing them	Carbon, silicon	Composite fibres, semiconductors, carbides
Groups 15–17 (non-metals)	5 to 7	Gains or shares to complete the octet	Nitrogen, oxygen, sulphur, chlorine	Inerting gas, breathing oxygen, corrosive salts
Group 18 (noble gases)	8 (2 for helium)	Nothing — the shell is already full	Helium, neon, argon	Argon shielding for welding; helium leak testing

Because every member of a group carries the same outer-electron count, a group also shares a combining power, which is what lets you predict a compound's formula from the table alone. That connection is worked through under valency below.

Structure of the Atom

Every atom consists of three sub-atomic particles:

Particle	Location	Charge	Relative Mass
Proton	Nucleus	Positive (+1)	1
Neutron	Nucleus	Neutral (0)	1
Electron	Shells (orbits)	Negative (−1)	≈ 1/1836

Key definitions:

• Atomic number (Z) — the number of protons in the nucleus. This defines the element.
• Mass number (A) — the total number of protons + neutrons in the nucleus.
• Isotopes — atoms of the same element with different numbers of neutrons (same Z, different A).
• Ion — an atom that has gained or lost electrons, giving it a net electrical charge.

The Scale of an Atom and Its Particles

Relative mass and charge are what an exam question normally asks for, but the measured values behind them explain several answers that otherwise look arbitrary:

Particle	Rest mass (kg)	Mass in atomic mass units (u)	Electric charge (C)
Proton	\( 1.673 \times 10^{-27} \)	1.0073	\( +1.602 \times 10^{-19} \)
Neutron	\( 1.675 \times 10^{-27} \)	1.0087	0
Electron	\( 9.109 \times 10^{-31} \)	0.000549	\( -1.602 \times 10^{-19} \)

Three consequences follow directly from those numbers.

A proton and a neutron weigh essentially the same. The neutron is heavier by about one part in a thousand, which is why both are given a relative mass of 1 and why the mass of a proton is correctly described as equal to the mass of a neutron. That approximation is what makes the mass number a simple count of nucleons rather than a sum that has to be looked up.

The electron is negligible by comparison. Dividing the proton mass by the electron mass gives a ratio of about 1,836, which is where the fraction 1/1836 comes from; rounded to one significant figure that ratio becomes "about 2,000 times", so an answer option built on 2,000 and a textbook figure of 1,836 are describing the same ratio. The practical consequence is that virtually all of an atom's mass sits in its nucleus. Take an aluminium-27 atom: its 13 electrons together weigh about \( 1.2 \times 10^{-29} \) kg, against a nucleus of roughly \( 4.5 \times 10^{-26} \) kg, so the electrons contribute less than 0.03 % of the atom's mass. Strip them all off and the atom's weight is essentially unchanged — but its chemistry is destroyed.

The proton and electron charges are exactly equal and opposite. An atom is therefore electrically neutral only when it holds exactly as many electrons as protons, which is why the number of electrons in a stable, un-ionised atom always equals the number of protons in its nucleus. An atom with two protons carries two electrons; one with 13 protons carries 13. Break that equality by adding or removing electrons and you no longer have a neutral atom but an ion.

The dimensions are as lopsided as the masses. A whole atom is of the order of \( 10^{-10} \) m across — a tenth of a nanometre — while its nucleus measures only about \( 10^{-15} \) to \( 10^{-14} \) m. The nucleus is therefore some ten thousand to a hundred thousand times smaller in diameter than the atom that contains it, which means it occupies less than a million-millionth of the atom's volume. An atom is overwhelmingly empty space, with more than 99.9 % of its mass concentrated in a speck at the centre and the electrons sweeping out everything else.

That raises an obvious question: if protons all carry the same positive charge, and like charges repel, why does a nucleus not fly apart? The answer is the strong nuclear force, an attraction between nucleons that is far stronger than the electrostatic repulsion but which acts only over distances of about \( 10^{-15} \) m — roughly the width of the nucleus itself. Neutrons contribute that attraction without adding any repulsion, which is why they act as the nuclear "glue" and why larger nuclei need proportionally more neutrons than protons to stay together. Where the balance is wrong, the nucleus is unstable.

Counting Protons, Neutrons and Electrons

Almost every atomic-structure question you will meet reduces to the same three-step count, so it is worth having it as a reflex:

• The atomic number Z is the number of protons, and nothing else. It is what identifies the element — change it and you have a different element altogether.
• The mass number A is the number of protons plus neutrons, i.e. the total number of nucleons. Electrons are never included; they weigh too little to matter.
• The neutron count is therefore \( A - Z \), and in a neutral atom the electron count equals \( Z \).

Two traps recur. The first is offering "protons plus electrons" for the mass number. In a neutral atom the electron count equals the proton count, so what that option actually adds up is \( Z + Z \), twice the atomic number, when the mass number \( A \) is protons plus neutrons. The two agree only for an atom that happens to hold exactly as many neutrons as protons, and neither worked example above is such an atom: aluminium-27 would come out at \( 13 + 13 = 26 \) rather than 27, and the shell-diagram nucleus at \( 3 + 3 = 6 \) rather than 5. Electrons are not nucleons and never enter a mass number at all, so the option is wrong whether the atom is neutral or ionised, and whichever isotope it is. The second is confusing mass number with relative atomic mass. The mass number is always a whole number, because it is a count of particles; relative atomic mass usually is not, for the reason set out under isotopes below.

The atomic-structure diagram above is drawn on exactly this principle. Its nucleus is labelled only "p+ n" rather than with a count, but the electrons can be counted: two occupy the innermost ring and five the next, and the third ring is drawn empty. Read as a real atom, that arrangement is nitrogen — seven electrons in a neutral atom imply seven protons, so \( Z = 7 \). The empty outer ring is there to make the point that an energy level exists whether or not anything occupies it. The rings themselves are schematic: neither their radii nor the spacing between them is drawn to scale, and the electrons are spaced evenly around each ring for legibility rather than being pinned to any real position.

Models of the Atom

Every picture of an atom is a model, chosen because it predicts something useful, and it is worth knowing which model is being used and what it is good for.

• The earliest useful picture treated the atom as a solid indivisible sphere. It explains why elements combine in fixed ratios, and nothing else.
• The discovery of the electron produced an atom of positive material with electrons embedded in it. That model was killed by experiment: firing alpha particles at thin gold foil sent a few of them straight back, which is impossible unless the positive charge and nearly all the mass are concentrated in a very small central body. That experiment is where the nucleus comes from.
• The shell model that followed puts electrons in fixed circular orbits at discrete energy levels, and this is the model used throughout this section, throughout the syllabus and in the diagram above. Its virtue is that it gets the chemistry right: the electron count per shell, the octet rule, valency, ionisation and the emission of light at definite wavelengths all fall straight out of it.
• The modern picture replaces the orbit with an orbital — a region in which an electron of a given energy is likely to be found, with no defined path at all. The s orbital is spherical, the three p orbitals are dumbbell-shaped and mutually perpendicular, and the electron's position within any of them is a matter of probability rather than position.

The two later models do not contradict each other on anything an engineer needs. Energy levels, capacities and the drive towards a full outer shell are identical in both; only the picture of where the electron is at a given instant differs. That is why the shell model remains the working tool, and why a shell diagram should be read as a bookkeeping device for electron counts rather than as a photograph of an orbit.

Isotopes and Relative Atomic Mass

Because chemical behaviour is decided by the electrons, and the electron count follows the proton count, adding or removing a neutron changes an atom's mass without changing its chemistry at all. Atoms of one element that differ only in neutron count are isotopes of that element: same atomic number, different mass number, different number of nucleons in the nucleus. Carbon-14 and carbon-12 have the same number of protons and the same number of electrons; only the neutrons differ, and both are unambiguously carbon.

Hydrogen is the textbook case because it has so few particles that the effect is obvious:

Isotope	Protons	Neutrons	Mass number	Abundance and notes
Protium (ordinary hydrogen)	1	0	1	Over 99.9 % of natural hydrogen; the only common nuclide with no neutron at all
Deuterium	1	1	2	About 0.02 %; combined with oxygen it gives heavy water, \( \text{D}_2\text{O} \)
Tritium	1	2	3	Trace amounts only; radioactive, and forms tritiated water rather than heavy water

Hydrogen is also the lightest and structurally simplest element there is — a single proton with a single electron — and the most abundant in the universe, which is why it is the element questions reach for when they ask which material is the smallest or simplest.

Carbon behaves the same way: carbon-12 accounts for about 98.9 % of natural carbon and carbon-13 for most of the rest, while carbon-14 exists only in traces, is radioactive, and decays at a known rate — the basis of radiocarbon dating. The half-life of a radioactive isotope is simply the time taken for half of a given quantity of it to decay.

Isotopes are also why relative atomic masses are rarely whole numbers. The value quoted for an element is the weighted average over its natural isotopes. Chlorine is quoted as 35.5 not because any chlorine atom weighs 35.5 u, but because about 75 % of chlorine atoms are chlorine-35 and about 25 % are chlorine-37:

The unit those masses are quoted in is fixed by international agreement against one particular nuclide: the atomic mass unit is defined as exactly one twelfth of the mass of a carbon-12 atom, which works out at about \( 1.661 \times 10^{-27} \) kg. Every atomic weight on the Periodic Table is therefore referred to carbon. Hydrogen was the original reference and oxygen was used for a time afterwards, but carbon-12 is the standard in force today, which is the point behind questions asking which element atomic weights are related to.

Two practical consequences. Radioactive isotopes are used deliberately in maintenance: iridium and cobalt gamma sources are standard for radiographic inspection, and they work precisely because an unstable nucleus is indifferent to the chemistry going on around it. And an isotope is not an ion — the distinction is worth nailing down now, because the wrong answers on isotope questions are often built out of electron counts, which is the ion's defining property and not the isotope's. An isotope differs in neutrons and is electrically neutral; an ion differs in electrons and carries a net charge.

Unstable Nuclei and Ionising Radiation

A nucleus is stable only when the strong nuclear force and the electrostatic repulsion between its protons are in balance, and that balance depends on the neutron-to-proton ratio. Where the ratio is wrong — too few neutrons, too many, or simply too many nucleons in total — the nucleus sheds particles or energy until it reaches a stable arrangement. That is radioactive decay, and it is a property of the nucleus alone: heating, cooling, dissolving or chemically combining a radioactive isotope changes its decay rate not at all.

Emission	What it is	Charge	Stopped by
Alpha	A helium nucleus — two protons and two neutrons	Positive	A sheet of paper or the outer layer of skin; dangerous chiefly if inhaled or ingested
Beta	A high-speed electron ejected as a neutron converts to a proton	Negative	A few millimetres of aluminium
Gamma	Electromagnetic radiation of very short wavelength — no particle at all	None	Only substantial thicknesses of lead or concrete; never absorbed completely

All three are described as ionising radiation because each carries enough energy to knock electrons out of the atoms it passes through, leaving a trail of ions behind. That is what makes it useful and what makes it hazardous, and the two are the same mechanism: the ionisation that exposes a radiographic film is the ionisation that damages living tissue.

The practical consequence in maintenance is gamma radiography. A sealed isotope source is used to inspect welds, castings and structure that cannot be reached with an X-ray set, and because gamma radiation cannot be switched off, control is by distance, shielding and time: a controlled area is established and cleared, warning signs and audible or visual indicators are posted, the source is kept in its shielded container until the moment of exposure, and personnel carry dose monitoring. Anyone working near an active radiography operation must respect the barriers even if their own task is unrelated — the source has no way of knowing who is in the beam.

Ions and Ionisation

The direction of the charge is the part that catches people out, so reason it through rather than memorising it. Electrons are negative. Remove one and the atom is left with more protons than electrons, so it becomes a positive ion — a positive ion is one that has lost electrons, not gained them. Add an electron and the atom has more negative charge than positive, so it becomes a negative ion. A positive ion is also called a cation and a negative ion an anion, and the charge is written as a superscript: \( \text{Na}^{+} \), \( \text{Mg}^{2+} \), \( \text{Al}^{3+} \), \( \text{Cl}^{-} \), \( \text{O}^{2-} \).

Which way an atom goes is decided by its outer shell. Metals hold one to three outer electrons loosely and shed them to reach the full shell underneath, so metals form positive ions. Non-metals are within one to three electrons of a full shell and capture what they need, so they form negative ions. Noble gases do neither.

The energy needed to pull the most loosely held electron away from a neutral atom is its ionisation energy. It is lowest for the outermost electron of a large metal atom, because that electron is furthest from the nucleus and partly screened from it by the shells in between, and highest for a small atom with a full shell. In practice atoms are ionised by any of the following:

• Friction — charge separation as one material rubs or flows over another. Fuel moving through a hose and filter, and air flowing over an airframe in flight, both do this.
• Heat — a flame or an electric arc is hot enough to strip electrons from gas atoms.
• Radiation — ultraviolet and X-rays carry enough energy per photon to eject an electron. Solar UV is what maintains the ionosphere.
• Dissolving — an ionic compound such as sodium chloride splits into free ions when it dissolves, which is what makes the solution an electrolyte.

Ions matter to an engineer for several reasons at once. An aircraft accumulates static charge in flight through friction with the air and precipitation, and that charge is bled back off through static dischargers so it cannot build to a level that interferes with radio reception or produces a discharge. Refuelling is done with the aircraft, the bowser and the nozzle bonded together, because charge separated in flowing fuel is otherwise free to jump as a spark. Batteries pass current through their electrolyte as moving ions rather than as free electrons. And corrosion is fundamentally an ionic process: it needs a metal that gives up ions and an electrolyte to carry them, which is why trapped moisture and salt are so much more damaging than dry air. A gas in which a substantial fraction of the atoms has been ionised stops behaving like an ordinary gas altogether; that state is dealt with under states of matter below.

Electron Shells

Electrons orbit the nucleus in energy levels (shells), labelled K, L, M, N, etc. (or 1, 2, 3, 4...). Each shell holds a maximum number of electrons:

Shell	Number	Max Electrons	Formula
K	1	2	\( 2n^2 \)
L	2	8
M	3	18
N	4	32

The outermost occupied shell contains the valence electrons, which determine the chemical behaviour of the element. Elements with a full outer shell (noble gases) are chemically stable and unreactive.

Sub-shells and Where 2n² Comes From

Each shell is itself divided into sub-shells, labelled s, p, d and f, and each type of sub-shell has a fixed capacity no matter which shell it belongs to: an s sub-shell holds 2 electrons, a p holds 6, a d holds 10 and an f holds 14. The shell maxima are simply those capacities added up, which is why the \( 2n^2 \) formula works:

Shell	n	Sub-shells present	Capacities added	Shell maximum
K	1	1s	2	2
L	2	2s, 2p	2 + 6	8
M	3	3s, 3p, 3d	2 + 6 + 10	18
N	4	4s, 4p, 4d, 4f	2 + 6 + 10 + 14	32

Two exam answers fall straight out of that table. The innermost shell holds a maximum of two electrons, because it contains only a 1s sub-shell — and any s sub-shell, in any shell, holds exactly two. The second shell holds a maximum of eight, from one s and one p sub-shell.

Electrons fill the lowest available energy first, working outward. The shell closest to the nucleus is the lowest energy level, because that is where the electrostatic attraction to the positive nucleus is strongest; each shell further out is a higher energy level and its electrons are held more weakly. Filling is not strictly shell by shell, however: the 4s sub-shell lies slightly below 3d in energy, so it fills before the M shell has taken its full 18. That is why an outer shell in the elements an engineer actually handles never carries more than eight electrons, even though the M shell's theoretical maximum is 18 — once a shell reaches eight in its outer position, the next electron starts the shell beyond it. The 18 and 32 figures are reached only when those shells are buried underneath an occupied outer shell.

The spacing of the shells is worth stating carefully, because two things move in opposite directions. In the shell model the radius grows roughly as the square of the shell number — in proportion to 1, 4, 9 and 16 for the first four shells — so the gaps between successive shells are about 3, 5 and 7 in the same units. The further from the nucleus the shells lie, the further apart from each other they get. Their energies, however, converge: for hydrogen the four lowest levels sit at about \( -13.6 \), \( -3.4 \), \( -1.5 \) and \( -0.85 \) electron-volts, so the steps between them shrink from 10.2 to 1.9 to 0.7 eV as you move outward. An outer electron is therefore both physically further from its neighbours and energetically much easier to dislodge, which is precisely why chemistry and electrical conduction are governed by the outermost shell and never by the inner ones.

Give an electron a modest amount of energy — from heat, from a collision, or from absorbing light — and it can jump to a higher shell, leaving the atom in an excited state. It does not stay there; when it falls back it re-emits that energy as light of one definite wavelength, set by the size of the gap it fell across. That is why a sodium lamp is yellow, why a welding arc has its characteristic colour, and why every element produces its own line spectrum. Give the electron enough energy instead and it leaves the atom altogether, which is ionisation.

Electron Configurations of Some Familiar Elements

Writing out the electrons shell by shell turns the rules above into something you can use. Fill from the innermost shell outward, respect each shell's capacity, and stop when you have used up as many electrons as the element has protons:

Element	Z	Electrons per shell (K, L, M, N)	Outer shell	Consequence
Hydrogen	1	1	1	One electron short of a full K shell; forms one bond
Helium	2	2	2 (full)	Inert — the K shell needs only two
Nitrogen	7	2, 5	5	Completes its octet by sharing three
Oxygen	8	2, 6	6	Needs two; combining power 2
Neon	10	2, 8	8 (full)	Inert
Sodium	11	2, 8, 1	1	Loses one readily; a very reactive metal
Magnesium	12	2, 8, 2	2	Loses two; combining power 2
Aluminium	13	2, 8, 3	3	Three outer electrons; the principal airframe metal
Silicon	14	2, 8, 4	4	Shares four; the semiconductor case
Chlorine	17	2, 8, 7	7	One electron short of a full octet
Argon	18	2, 8, 8	8 (full)	Inert; the arrangement every other atom is working towards
Titanium	22	2, 8, 10, 2	2	M shell still filling beneath a settled outer shell
Chromium	24	2, 8, 13, 1	1	Variable valency — commonly 3 or 6
Iron	26	2, 8, 14, 2	2	Variable valency — commonly 2 or 3
Copper	29	2, 8, 18, 1	1	One loose outer electron above a filled M shell

Three things in that table repay a second look. First, nitrogen's 2, 5 is exactly the arrangement drawn in the atomic-structure diagram earlier in this section. Second, the noble gases are the only entries whose outer shell is full, and they are the only ones that do nothing chemically — the pattern is not a coincidence but the whole explanation. Third, the transition metals break the neat progression: titanium, chromium, iron and copper all carry only one or two electrons in their outermost shell while a deeper shell — the M shell, capacity 18 — is the one whose filling is in progress across the row. Titanium, chromium and iron are caught part-way through it, holding 10, 13 and 14; copper is the element that has just completed it, at the full 18, leaving a single electron outside. Because those inner electrons are close in energy to the outer ones, they can also take part in bonding, which is why the transition metals show more than one valency — iron as ferrous or ferric, copper as cuprous or cupric, chromium as trivalent or hexavalent — and why they form so many alloys and coloured compounds.

Valency and the Octet Rule

A full outer shell is the arrangement every atom behaves as though it is trying to reach. For all the elements an engineer meets that means eight electrons in the outer shell — the octet rule — with helium the exception, needing only two because its shell is the K shell. An atom's valency is its combining power: the number of electrons it must gain, lose or share to get there.

The route it takes depends on how far it has to go. Up to three outer electrons and it is quicker to lose them, exposing the full shell underneath; five or more and it is quicker to gain or share the few that are missing. Four sits on the boundary and shares. So the combining power is the outer-electron count for atoms with three or fewer, and eight minus that count for atoms with five or more.

Outer-shell electrons	Term used for that count	Usual combining power	How the octet is completed	Examples
1	Univalent (monovalent)	1	Loses one	Hydrogen, lithium, sodium, potassium
2	Bivalent (divalent)	2	Loses two	Magnesium, calcium, zinc
3	Trivalent	3	Loses three	Aluminium, boron
4	Tetravalent	4	Shares four	Carbon, silicon
5	Pentavalent	3 (5 in some compounds)	Gains or shares three	Nitrogen, phosphorus
6	Hexavalent	2	Gains or shares two	Oxygen, sulphur
7	Heptavalent	1	Gains or shares one	Fluorine, chlorine
8 (2 for helium)	Inert, zero valency	0	Nothing to do — already full	Helium, neon, argon

Atoms whose outer shell is already full have nothing to gain, lose or share, so their valency is zero and they are practically inert — that is the whole of the chemistry of the noble gases, and the reason helium, neon and argon are the standard examples. Inertness is not the same as being a good conductor: an atom with a full outer shell has no loose electrons at all, so a full outer shell means unreactive, not conductive.

Valency is also the fastest way to predict a formula. Combine two elements so that the total combining power on each side balances — in practice, swap the valencies and use them as subscripts:

A shortage of a valence electron is always a shortage in the outer shell, never an inner one, because the inner shells are full and their electrons are far too tightly held to take part. That is the single idea behind every question that asks where a missing or extra bonding electron sits.

Valence Electrons and Electrical Behaviour

The same outer-shell count that predicts chemistry also predicts whether a material will carry current, which is why this paragraph of the syllabus underpins the electrical modules that follow.

• One to three outer electrons — conductor. The outer electrons are so loosely held that they detach and drift through the lattice, and an applied voltage sets them moving as a current. Copper is the default wiring material on the strength of its single loose outer electron; silver is marginally better and gold is used where a contact must not corrode.
• Four outer electrons — semiconductor. Silicon and germanium sit exactly on the boundary. Their electrons are locked into shared bonds at low temperature but can be freed by heat, light or the deliberate addition of impurities.
• Five to eight outer electrons — insulator. The outer electrons are tightly bound into completed or nearly completed shells and cannot move, so no current flows. PTFE, polyimide, rubber, glass and ceramics all work on this principle: an insulator is defined by having its electrons tightly bound, not by having none.

Temperature moves the two classes in opposite directions, and the direction is a favourite question. In a metal the free electrons are already there, so heating only makes the lattice vibrate more violently and scatter them — resistance rises with temperature. In a semiconductor heating frees additional carriers that were not previously available, and that effect dominates — resistance falls with temperature. Aluminium is worth a note of its own here: with three valence electrons it conducts about 60 % as well as copper for a given cross-section, but it is roughly a third of the density, so for a given weight it carries more current. That is why heavy feeder cables and busbars are frequently aluminium while general airframe wiring is copper.

Molecules

A molecule is the smallest particle of a substance that retains the chemical properties of that substance. It consists of two or more atoms bonded together. Molecules of elements contain atoms of the same type (e.g. \( \text{O}_2 \), \( \text{N}_2 \)), while molecules of compounds contain different types of atoms (e.g. \( \text{H}_2\text{O} \), \( \text{CO}_2 \)).

Monatomic, Diatomic and Polyatomic Molecules

How many atoms a molecule contains is set by how those atoms complete their outer shells, and the count has a name:

• Monatomic — a single atom is the whole molecule. Only the noble gases behave this way, precisely because they have no bonding to do: a cylinder of argon holds single argon atoms, never \( \text{Ar}_2 \).
• Diatomic — two atoms. Hydrogen, nitrogen, oxygen, fluorine and chlorine all exist as \( \text{H}_2 \), \( \text{N}_2 \), \( \text{O}_2 \), \( \text{F}_2 \) and \( \text{Cl}_2 \), which is why free oxygen is written \( \text{O}_2 \) and not O.
• Polyatomic — three or more. Ozone is \( \text{O}_3 \); water and carbon dioxide are polyatomic compounds; the hydrocarbon chains in kerosene run to a dozen or more carbons apiece.

A molecule of a single kind of atom is still an element, however many atoms it holds — \( \text{O}_2 \) is elemental oxygen, not a compound. Only when unlike atoms are bonded together does the substance become a compound. Nitrogen makes the practical case: its two atoms are held by a triple bond, one of the strongest bonds between like atoms, and that is why nitrogen is so nearly unreactive at normal temperatures. It is that inertness — not simply the fact that it is a gas — that makes dry nitrogen the choice for tyre inflation, oleo charging and fuel-tank inerting, since it cannot support combustion and carries no moisture to freeze.

Reading a formula is a matter of distinguishing two kinds of number. The subscript counts atoms within one molecule; the coefficient in front counts whole molecules. So \( \text{H}_2 \) is one molecule containing two hydrogen atoms, while \( 2\text{H}_2 \) is two such molecules containing four hydrogen atoms in total. Putting hydrogen and oxygen together to make water is written:

Count both sides: four hydrogen atoms and two oxygen atoms before, four hydrogen atoms and two oxygen atoms after. Nothing has been created or destroyed, only rearranged — which is the conservation rule from the start of this section written in chemical shorthand. Adding the relative atomic masses the same way gives the relative molecular mass: water is \( 2 \times 1 + 16 = 18 \), and carbon dioxide is \( 12 + 2 \times 16 = 44 \).

Brownian Motion

Everything above treats molecules as real objects in constant motion. The direct evidence for that is Brownian motion: view smoke particles in air, or a fine powder suspended in a liquid, under a microscope and each visible speck jitters about along a random, endlessly changing path. Nothing is stirring the sample, and the specks are far too large to be moving of their own accord.

The explanation is that the visible particle is being struck from every side by molecules that are themselves invisible. At any instant the bombardment does not balance exactly, and the particle is knocked in whichever direction happens to be receiving fewer or slower impacts; a moment later the imbalance is elsewhere. Brownian motion therefore describes the motion of molecules, not their bonding or their breakdown, and it establishes two things at once: that matter is made of discrete particles small enough to be individually unnoticeable, and that those particles are in permanent random motion.

The temperature of a substance is a measure of the average kinetic energy of exactly that motion. Raise the temperature and the molecules move faster, so the jitter becomes more vigorous; cool the sample and it slows. That single link between temperature and molecular speed is what the rest of this note rests on — it explains why gases exert pressure (molecules striking the container walls), why smells and gas leaks spread through still air without any draught to carry them (diffusion), why a liquid evaporates faster when warm, and ultimately why heating a solid far enough will melt it.

Chemical Compounds

A compound is a substance formed when two or more elements chemically combine in a fixed ratio. The properties of a compound are usually very different from those of its constituent elements.

Compounds, Mixtures, Solutions and Alloys

The word that does the work in the definition above is chemically. Put two elements in the same container and you have a mixture; make them bond and you have a compound. The distinction is not academic — it decides how the substance behaves, how you separate it and what its properties will be.

Feature	Compound	Mixture
How the components are held	Chemically bonded — electrons transferred or shared	Not bonded — simply intermingled
Proportions	Fixed and definite (water is always two hydrogens to one oxygen)	Any proportion at all
Separation	Only by chemical means (electrolysis, heating, reaction)	By physical means (filtering, distilling, settling, magnetism)
Properties	Quite unlike those of the constituent elements	Each component keeps its own properties
Energy change on forming	Heat is given out or taken in	Essentially none
Melting and boiling	A single definite melting and boiling point	A range, set by the components present

Sodium chloride is the standard demonstration of how completely properties change. Sodium alone is a soft metal that reacts violently with water; chlorine alone is a poisonous green gas. Bonded, they make ordinary table salt. Nothing of the parent elements' behaviour survives the bond, because the outer-shell arrangement that produced that behaviour no longer exists.

Water is a compound and not a mixture, however casually the word "mixture" gets used, because its hydrogen and oxygen are chemically bonded in a fixed two-to-one ratio and can be separated only by electrolysis or another chemical process. Air, by contrast, is a mixture: its nitrogen, oxygen, argon and carbon dioxide are simply intermingled, each keeping its own properties, and they can be separated physically by cooling the air until the components liquefy at their different temperatures. That is exactly how commercial oxygen and nitrogen are produced.

Two intermediate cases matter in a hangar. A solution is a mixture in which one substance is dispersed through another right down to molecular scale — salt in water, or the small quantity of water that dissolves in kerosene. It is uniform throughout and it is still a mixture, because no chemical bond has formed between solute and solvent. An alloy is a solid solution of one or more added elements in a host metal: the atoms of the added element occupy sites in the host lattice, distorting it, which is why alloying makes a metal stronger and harder than the pure element while lowering its electrical conductivity. Aluminium alloyed with copper, magnesium and zinc gives the airframe alloys; iron alloyed with carbon and chromium gives the steels. In neither case has a compound been formed, which is why an alloy's composition can be varied continuously to tune its properties instead of being locked to one fixed ratio.

Solutions, Solubility and Saturation

In any solution the substance that dissolves is the solute and the substance it dissolves in is the solvent. A solution that will accept no more solute at a given temperature is saturated, and how much it can hold before it reaches that point is the solute's solubility. Two temperature effects run in opposite directions, and both matter in a hangar.

• A solid dissolved in a liquid is generally more soluble hot. Heat helps break the solid apart and helps the solvent molecules surround its particles, so warming a solution lets it hold more — and cooling a saturated solution forces the excess back out as crystals.
• A gas dissolved in a liquid is less soluble hot. A dissolved gas molecule only needs enough kinetic energy to escape the surface, and heating supplies it. Warm a liquid and its dissolved gases come out of solution; this is why bubbles appear in water long before it boils.

Pressure works on gases the other way again: the more gas is pressed against the surface, the more dissolves, so raising the pressure increases the amount held in solution and dropping the pressure releases it. Between them these two rules explain a set of otherwise unconnected observations. Air dissolved in hydraulic fluid or engine oil under system pressure comes out as bubbles wherever the pressure falls — at a pump inlet, across a restrictor, or in a return line — producing foaming, a spongy response and, where the bubbles collapse again, cavitation erosion. Oxygen dissolved in water accelerates corrosion. And fuel always carries a small quantity of dissolved water, which comes out of solution as free water as the fuel cools.

A solution stays a mixture throughout: nothing is chemically bonded, the proportions are variable up to the saturation limit, and the solute can be recovered unchanged by physical means — evaporating the solvent, distilling it off, or simply letting the temperature change and collecting what separates out.

Types of Chemical Bonds

Bond Type	Mechanism	Example	Properties
Ionic	Electrons transferred from one atom to another, creating oppositely charged ions that attract	NaCl (sodium chloride)	High melting point, conducts electricity when dissolved, crystalline solid
Covalent	Atoms share one or more pairs of electrons	\( \text{H}_2\text{O} \), \( \text{CO}_2 \)	Can be gas, liquid, or solid; generally lower melting points
Metallic	Positive metal ions surrounded by a "sea" of delocalised electrons	Aluminium, copper	Good conductors of heat and electricity, malleable, ductile

How a Bond Forms: the Drive to a Full Outer Shell

All three bonding mechanisms above are the same idea carried out three different ways: each atom ends up with the electron arrangement of a noble gas. What differs is who ends up owning the electrons.

Ionic bonding — electrons are transferred. Sodium carries one electron in its outer shell and chlorine has seven. Sodium hands its single outer electron over, exposing the full shell beneath it and leaving a \( \text{Na}^{+} \) ion; chlorine accepts it and completes its own octet as a \( \text{Cl}^{-} \) ion. The two are now oppositely charged, and the bond is the electrostatic attraction between them — there is no sharing involved. Because that attraction acts equally in all directions, ionic solids build up into a regular crystal lattice with a very large number of such attractions to break, which is why their melting points are high. In the solid the ions are locked in place and cannot carry current; dissolve or melt the compound and the ions become free to move, and it conducts. Ionic bonding is characteristically a metal joining a non-metal.

Covalent bonding — electrons are shared. Two non-metals both need to gain electrons, and neither can, so they pool a pair instead: the shared pair counts towards the outer shell of both atoms at once. Two hydrogen atoms share one pair to give a single bond in \( \text{H}_2 \). Two oxygen atoms, each two electrons short, share two pairs — a double bond. Two nitrogen atoms, each three short, share three pairs — the triple bond that makes nitrogen so unreactive. Oxygen forms two single bonds to two separate hydrogens in \( \text{H}_2\text{O} \); carbon, with four outer electrons, forms four bonds and gives the whole of organic chemistry, including the hydrocarbon chains of kerosene. Non-metals bonding with other non-metals form covalent bonds, and that is the pairing to look for in a question.

Metallic bonding — electrons are pooled across the whole solid. Metal atoms have few outer electrons and nothing to gain them from, so they release them into a shared cloud that belongs to the lattice as a whole rather than to any one atom. The positive ions left behind are held together by their attraction to that cloud.

The difference in what happens under load follows directly. Push on an ionic crystal hard enough to slide one plane of ions over the next and you bring like charges into contact; the crystal splits, which is why ionic solids are brittle. Push on a metal and the planes of positive ions slide while the electron cloud simply flows with them, so the metal deforms instead of fracturing. That is the mechanism behind every forming, bending and riveting operation carried out on an aircraft.

Polar Molecules and the Forces Between Molecules

A bond inside a molecule and a force between whole molecules are two different things, and mixing them up is the usual reason a boiling point looks inexplicable. Breaking the covalent bonds inside water would give hydrogen and oxygen gas; boiling water only has to separate whole \( \text{H}_2\text{O} \) molecules from each other, and that takes far less energy. It is the weaker, between-molecule forces that set melting and boiling points, surface tension and viscosity.

In a covalent bond between unlike atoms the shared pair is not held equally. Oxygen attracts electrons more strongly than hydrogen does, so in a water molecule the shared electrons sit closer to the oxygen, leaving that end slightly negative and the hydrogen ends slightly positive. A molecule with such a permanent charge separation is polar. Carbon dioxide is not polar even though its individual bonds are, because the molecule is symmetrical and the two pulls cancel; hydrocarbons such as kerosene are essentially non-polar for the same reason.

Three consequences follow, and all three are met daily:

• Hydrogen bonding. The positive hydrogen end of one water molecule is strongly attracted to the negative oxygen end of the next. This is much stronger than an ordinary intermolecular attraction, and it is why water boils at 100 °C when comparable molecules of similar size boil far below room temperature. It is also why water has an unusually high surface tension and an unusually high capacity to absorb heat.
• Van der Waals forces. Even non-polar molecules attract each other weakly, through the momentary uneven distribution of their own electrons. The forces are feeble but they add up along a long molecule, which is why heavier hydrocarbons are liquids and the lightest ones are gases.
• Like dissolves like. Polar solvents dissolve polar and ionic substances; non-polar solvents dissolve non-polar ones. Water dissolves salt and glycol; it does not dissolve kerosene, which is why water separates out and collects at the bottom of a fuel tank instead of staying mixed.

Two named cases of intermolecular attraction come up repeatedly in questions, and the only thing to keep straight is whether the molecules concerned are alike:

Force	Acts between	What it produces	Everyday example
Cohesion	Molecules of the same substance	Holds the substance together; surface tension; a droplet pulling itself into a bead	A drop of water holding its shape on a waxed surface
Adhesion	Molecules of different substances, at the surface where they touch	Makes one substance cling to or wet another	Water wetting glass; paint, adhesive and penetrant clinging to metal

Which of the two wins decides what a liquid does on a surface. Where adhesion to the solid beats cohesion within the liquid, the liquid spreads and wets the surface; where cohesion wins, it pulls away and beads instead. Every cleaning, painting and bonding operation depends on getting adhesion to win, which is why degreasing beforehand is not optional — a contaminated surface is one the liquid is being asked to adhere to and cannot.

Simple Molecular and Giant Covalent Structures

Covalent substances split into two families whose bulk properties could hardly be more different, and the difference is structural rather than chemical. In a simple molecular substance the covalent bonds tie a small fixed number of atoms into a discrete molecule, and only weak intermolecular forces hold one molecule to the next. Melting or boiling need only overcome those weak forces, so the melting points are low and many such substances are gases or liquids at room temperature. Carbon dioxide, water and the hydrocarbons in fuel all belong here, and that is the family the general statement about lower melting points describes.

In a giant covalent structure there are no separate molecules at all. The covalent bonding continues in three dimensions throughout the whole crystal, so melting it means breaking strong covalent bonds by the million. These substances have some of the highest melting points known:

Structure	Arrangement	Behaviour	Aviation use
Diamond (carbon)	Each carbon covalently bonded to four others in a rigid three-dimensional network	Hardest natural material; an electrical insulator, since every electron is committed to a bond	Cutting, grinding and honing tools; hardness testing
Graphite (carbon)	Flat sheets of carbon bonded to three neighbours each, the sheets held together only weakly	Sheets slide over one another, so it is soft and slippery; the fourth electron is free, so it conducts	Dry lubricant; brushes and brush-type contacts; carbon fibre precursor
Silicon carbide, alumina, quartz	Giant networks of silicon, aluminium, carbon and oxygen	Extremely hard, high melting point, chemically stable, brittle	Abrasives and blast media; ceramic coatings; insulators

Carbon is the instructive case because diamond and graphite are the same element, differently arranged. One is the hardest substance in common use and an insulator; the other is soft enough to write with and conducts electricity. Nothing about the atoms differs — only the way the bonds are laid out — which is as clear a demonstration as there is that structure, not just composition, determines properties.

Chemical Reactions: Synthesis, Analysis and Catalysts

A chemical change rearranges the bonds between atoms and produces a substance with new properties; a physical change does not. Melting, dissolving and evaporating are physical — the substance is recoverable unchanged. Burning, corroding and curing are chemical.

Two words name the two directions, and each is the exact reverse of the other, so the pair is best learned together:

• Synthesis — building a compound up by chemically bonding two or more elements together, which is also how a compound is produced from commonly available raw materials. Sodium plus chlorine to give salt, or hydrogen plus oxygen to give water, are syntheses.
• Analysis — the reverse: splitting a chemical compound into simpler substances. Electrolysing water back to hydrogen and oxygen is an analysis.

A catalyst is a substance that speeds a reaction up without being consumed by it and without appearing in the final products. It does not start a reaction that could not otherwise happen, it does not stop one, and it does not reverse the direction — it lowers the energy barrier so that more of the collisions between reacting molecules succeed. Because it is not used up, a very small quantity keeps working indefinitely until it is poisoned or physically degraded. The everyday example is a catalytic converter, in which a thin layer of platinum-group metal converts exhaust gases without itself being consumed; the same principle appears in two-part adhesives and sealants, where an accelerator brings the cure time down without becoming part of the cured polymer.

Energy always moves in a reaction, in one direction or the other. An exothermic reaction releases heat, because the bonds formed are stronger than those broken: combustion, curing resin and corroding metal all warm their surroundings. An endothermic reaction absorbs heat from its surroundings. Combustion is the one an engineer must be able to write down:

Burn kerosene with sufficient oxygen and the products are carbon dioxide and water vapour. Starve it of oxygen and combustion is incomplete, producing carbon monoxide and soot instead — which is why carbon monoxide is the hazard associated with a leaking exhaust or a faulty combustion heater, and why visible smoke indicates a fuelling or combustion problem rather than merely a dirty engine. Combustion is a specific case of oxidation, which means any reaction in which a substance combines with oxygen or, more generally, loses electrons; reduction is the opposite. Corrosion is oxidation running slowly at ambient temperature, and it is the same chemistry as a fire, only spread over years.

Common Aviation-Related Compounds

Compound	Formula	Aviation Relevance
Water	\( \text{H}_2\text{O} \)	Cooling, cleaning, weather effects
Carbon dioxide	\( \text{CO}_2 \)	Fire extinguishing agent, combustion product
Aluminium oxide	\( \text{Al}_2\text{O}_3 \)	Protective corrosion layer on aluminium
Iron oxide (rust)	\( \text{Fe}_2\text{O}_3 \)	Corrosion product on steel — must be prevented
Kerosene (Jet A-1)	Hydrocarbon mix	Primary aviation turbine fuel

Oxides, Corrosion and Protection

The compounds table above puts aluminium oxide and iron oxide side by side, and the contrast between them is one of the most useful pieces of chemistry an airframe engineer has. Both are oxides. Both form spontaneously when the metal meets oxygen. One protects the metal underneath and the other destroys it, and the reason is structural.

Aluminium oxide forms almost instantly on bare aluminium and builds a film only a few molecules thick that is dense, hard, tightly adherent and impervious to oxygen. Once it has formed, it seals the surface and the reaction stops itself. That self-limiting film is the whole reason a metal as reactive as aluminium can be used as an airframe material at all — and it is why a freshly filed or abraded aluminium surface must be protected promptly, because the film that reforms on it is not the same as the controlled one that was removed.

Iron oxide forms a film that is porous, flaky and of greater volume than the metal it came from. It neither adheres nor seals; it lifts, exposing fresh metal underneath, and the process continues until the section is consumed. That is rust, and it is the reason steel components must be protected by a coating rather than allowed to build their own.

The processes that follow from this are ordinary maintenance work:

• Anodising thickens the natural aluminium oxide film electrolytically into a much deeper, harder layer, which also holds paint well. Chromic and sulphuric acid processes are the usual routes.
• Chemical conversion coatings replace the natural film with a chemically formed one that resists corrosion and provides a paint key. Traditional chromate coatings used chromium in its hexavalent form, which is why their use is now tightly controlled on health grounds and trivalent chromium alternatives have replaced them in many applications.
• Cladding bonds a thin layer of nearly pure aluminium to each face of a high-strength alloy sheet. The pure surface builds a better oxide film than the alloy would, so the strong core is protected by the weaker skin.
• Titanium forms an oxide film of the same self-sealing kind as aluminium's, but tougher and stable to much higher temperatures, which is much of the reason it is chosen for hot engine and exhaust-area structure.

Corrosion accelerates sharply when two dissimilar metals are in electrical contact in the presence of an electrolyte such as salt-laden moisture. The more reactive of the two gives up ions preferentially and corrodes, protecting the less reactive one — the reason dissimilar-metal joints are separated by a barrier, sealant or a compatible primer rather than being bolted directly together. The principle is deliberately exploited in a sacrificial anode, where a deliberately reactive metal is fitted to be consumed first.

Carbon dioxide appears in the same table for two unrelated reasons, and both are worth being clear about. It is a product of complete combustion, and it is also an extinguishing agent: it is denser than air and chemically stable, so it displaces the oxygen a fire needs without itself burning. Stored in a cylinder it is a liquid under its own vapour pressure; when it is discharged, the sudden expansion cools it so violently that some of it freezes into solid carbon dioxide, which is why a discharge produces visible white snow and why a bare-hand grip on the horn risks a cold burn.

Hydrocarbons and Aviation Fuel

A hydrocarbon is a compound of carbon and hydrogen only, and the whole family exists because carbon is tetravalent: with four outer electrons to share, a carbon atom can bond to four neighbours, including other carbon atoms, so the atoms can be strung into chains and rings of any length. Each carbon in the middle of a simple chain uses two of its bonds to link to its neighbours and the other two to carry hydrogen, while the two carbons at the ends of the chain carry three hydrogens each.

Chain length is what sets the physical properties, and it does so through the intermolecular forces rather than the bonds inside the molecule. A long molecule offers more surface for its neighbours to attract, so the attraction between molecules grows with chain length — and with it the boiling point, the density and the viscosity. That single relationship separates the products of a refinery: the shortest chains are gases at room temperature, the next are petrol and then kerosene, and the longest are lubricating oils, greases and waxes.

Aviation turbine fuel is described in the table above as a hydrocarbon mix, and the word mix is exact. It is not a compound with a formula but a mixture of many different hydrocarbons, blended to a specification. Because it is a mixture it has no single boiling point: it distils over a range, the lighter fractions evaporating first. Two properties of that mixture are worth carrying:

• The flash point is the lowest temperature at which a liquid gives off enough vapour to form a mixture with air that will ignite if a flame is applied. It is a property of the vapour above the liquid, not of the liquid itself, and it is the number that governs how a fuel must be handled. Jet A-1 has a specification minimum around 38 °C, so at ordinary ambient temperatures its tank vapour is normally too lean to ignite. Aviation gasoline has a flash point far below any temperature it will meet in service, so its vapour is ignitable essentially all the time — which is why the two fuels demand quite different handling precautions even though both burn identically once alight.
• The freezing point of a kerosene fuel is the temperature at which wax crystals that have formed as it cools disappear again on warming. It is a specification limit of the mixture, and it is a fuel-management concern on long high-altitude sectors.

Because it is a mixture and not a compound, fuel can also carry things it has not bonded with: dissolved water, dissolved air, and additives put in deliberately — anti-oxidants, static-dissipator additives to reduce the charge separated by fuel flowing through pipes and filters, and icing inhibitors. None of these is chemically part of the fuel, which is precisely why they can be blended in at controlled concentrations and why their effect can be lost by dilution or settling.

Polymers: Very Large Molecules

A polymer is a compound whose molecules are built by linking a small repeating unit, the monomer, into a chain thousands of units long. Nothing new happens chemically — the bonds along the chain are ordinary covalent bonds — but the sheer size of the molecule changes the bulk behaviour completely, because the weak attractions between neighbouring chains, feeble at any one point, add up over an enormous length.

How the chains are connected to each other divides polymers into three families that an engineer handles quite differently:

Family	Chain arrangement	Behaviour	Typical aircraft use
Thermoplastic	Separate chains held to each other only by weak intermolecular attraction	Softens when heated and hardens again on cooling, repeatedly; can be formed, welded and recycled	Acrylic and polycarbonate transparencies, cable insulation, ducting, PTFE hose liners
Thermoset	Chains chemically cross-linked into one continuous network during cure	Cannot be re-softened — heating it enough degrades it instead; cure is irreversible	Epoxy and phenolic matrices in composites, structural adhesives, potting compounds
Elastomer	Long coiled chains with a few widely spaced cross-links	Stretches greatly under load and returns to shape when released	Seals, O-rings, hoses, vibration mounts, tyre compounds

The cross-linking is the whole distinction, and it explains the storage rules as well as the mechanical ones. A thermoset resin is supplied as separate components, or frozen as a pre-impregnated fabric, precisely because the cross-linking reaction begins as soon as the components meet and warm up; once it has run, the network is a single molecule spanning the part and there is no way back. That is also why a cured composite is repaired by bonding on new material rather than by melting the old, and why cure temperature, cure time and out-time are controlled so carefully.

Polymers age in service in ways metals do not. Ultraviolet light, ozone, heat and contact with solvents or fluids all attack the chains or the cross-links, so an elastomer hardens, cracks and takes a permanent set, and a transparency crazes. That is the reasoning behind cure-date and shelf-life limits on rubber items, fluid-compatibility restrictions on seals, and the protection of transparencies from ultraviolet.

States of Matter

Matter exists in three principal states. The state depends on the balance between the kinetic energy of the particles (which tends to separate them) and the intermolecular forces (which tend to hold them together).

Property	Solid	Liquid	Gas
Shape	Fixed	Takes shape of container	Fills entire container
Volume	Fixed	Fixed	Variable (fills container)
Particle arrangement	Regular, closely packed	Close but disordered	Widely spaced, random
Particle motion	Vibrate in fixed positions	Slide past each other	Move freely at high speed
Intermolecular forces	Very strong	Moderate	Very weak
Compressibility	Almost incompressible	Almost incompressible	Easily compressed
Density	High	High	Low

Why a Substance Is Solid, Liquid or Gas at a Given Temperature

The balance described above is not a metaphor — it is a straight competition between two measurable quantities, and knowing which one is winning tells you the state without having to memorise it.

One side of that competition can be changed and the other cannot. The particles' kinetic energy rises and falls with temperature. The intermolecular forces are fixed by what the substance is: how polar its molecules are, whether they can hydrogen-bond, how large they are, and whether the bonding continues through the whole solid rather than stopping at the edge of a molecule. Raising the temperature does not weaken those forces at all; it simply gives the particles more energy to fight them with. So:

• Forces winning comfortably — particles are locked into fixed positions and can only vibrate about them. The substance is a solid and holds its own shape.
• Roughly matched — particles have enough energy to break out of fixed positions and slide past each other, but not enough to escape one another's attraction altogether. The substance is a liquid: it keeps its volume but takes the shape of whatever holds it.
• Kinetic energy winning — particles escape the attraction entirely and move independently at high speed. The substance is a gas, with neither a shape nor a volume of its own; it expands until it fills the container.

This is why a substance's melting point is a direct readout of the strength of the forces holding its particles together. Compare, at atmospheric pressure:

Substance	What holds the particles	Melts at	Boils at
Nitrogen	Very weak forces between small non-polar molecules	About −210 °C	About −196 °C
Water	Hydrogen bonding between polar molecules	0 °C	100 °C
Aluminium	Metallic bonding	About 660 °C	Around 2,500 °C
Copper	Stronger metallic bonding	About 1,085 °C	Around 2,560 °C
Iron	Stronger metallic bonding again	About 1,538 °C	Around 2,860 °C
Titanium	Strong metallic bonding	About 1,668 °C	Around 3,290 °C
Tungsten	Exceptionally strong metallic bonding	About 3,420 °C	Around 5,550 °C

Nothing in that table is arbitrary: the order of the melting points is the order of the bond strengths. It also explains a fact worth carrying into engine work — tungsten is chosen for filaments and electrodes because it stays solid where every other common metal has vaporised, and aluminium alloys are restricted to cool structure because their melting point is low and their strength falls away well before it.

One qualification matters in a workshop. A pure substance melts at a single sharp temperature, because every bond in it is the same. An alloy or any other mixture does not: it begins to melt at one temperature and is not fully liquid until a higher one, so it passes through a pasty range in between. Welding and brazing depend on that range, and heat-treatment temperatures are chosen with reference to it.

Whether a solid is crystalline or amorphous matters for the same reason. In a crystalline solid — every metal, and most minerals — the particles occupy a regular repeating lattice, and there is one temperature at which the whole lattice comes apart. In an amorphous solid such as glass or a cured thermoset resin the particles are frozen in a disordered arrangement, and there is no single melting point at all: the material softens progressively over a range. That is why glass can be worked over a wide range of temperature rather than at one, and why a composite structure has a maximum service temperature quoted rather than a melting point.

Inside a Solid: Lattice, Grains and Why Metals Bend

Calling a solid's arrangement "regular and closely packed" is accurate but hides the detail that decides how a metal behaves under a hammer. The atoms of a metal sit in a repeating three-dimensional lattice, and which lattice it is has real consequences. Aluminium, copper, nickel and austenitic stainless steel adopt a face-centred cubic arrangement, which offers many planes along which atoms can slide, so those metals are notably ductile and stay ductile when cold. Ordinary steel at room temperature is body-centred cubic, with fewer such planes, and titanium and magnesium are close-packed hexagonal, with fewer still — which is a large part of why magnesium is difficult to form cold and why titanium is usually worked hot.

A real component is not one crystal. As metal solidifies, crystals start growing at many points at once and meet each other, so the finished part is a mosaic of small crystals called grains, each with its lattice pointing in a different direction, joined at grain boundaries. Deformation happens by planes of atoms slipping over one another within a grain, and a grain boundary obstructs that slip. Finer grains therefore mean more boundaries and a stronger, harder metal; this is why grain size is controlled during heat treatment and why overheating a component, which lets the grains grow, is so damaging.

Slip does not require a whole plane of atoms to move at once — it happens by the movement of line defects in the lattice, so the force needed is far less than a perfect crystal would demand. Everything that makes a metal harder works by making those defects harder to move: alloying (foreign atoms distort the lattice), work hardening (the defects tangle with one another), and heat treatment (fine particles are precipitated in their path). Annealing reverses it, letting the lattice re-order and restoring ductility. This is the mechanism behind the statement earlier in this note that metals are malleable and ductile, and it explains why a metal that has been repeatedly flexed becomes hard and eventually cracks.

Solids are described as "almost incompressible" for a straightforward reason: the atoms are already as close as their mutual repulsion allows, so squeezing them further needs enormous pressure. What a solid does do under load is deform elastically — strain in proportion to stress, recovering when the load is removed — and that behaviour is the subject of the statics section rather than this one.

Inside a Liquid: Surface Tension, Capillarity and Viscosity

A liquid keeps its volume because its molecules are still within each other's attraction, and takes its container's shape because they are free to slide. Three consequences of that half-way condition are directly useful in maintenance.

Surface tension. A molecule in the body of a liquid is pulled equally in every direction by its neighbours. A molecule at the surface has neighbours only below and to the sides, so the net pull is inward. The surface behaves as though it were under tension, pulling itself to the smallest possible area — which is why a free droplet is spherical, why a small quantity of mercury beads up, and why a clean needle can be floated on water. Surface tension is a cohesive effect, and it falls as temperature rises, because the molecules are moving faster and holding each other less effectively.

Capillarity. Put a narrow tube into a liquid that wets it and the liquid climbs, because adhesion to the tube wall beats cohesion within the liquid and surface tension drags the bulk up behind the wetted edge. The narrower the gap, the higher it climbs. Put the same tube into a liquid that does not wet it — mercury in glass — and the level is pushed down instead. Capillary action is what draws dye penetrant into a surface-breaking crack during a penetrant inspection, and what draws it back out into the developer afterwards, so the whole method depends on the crack being clean enough for the penetrant to wet it. It is also why a hairline crack or an unsealed lap joint will pull water in and hold it there, and why capillary gaps are a classic corrosion site.

Viscosity. Viscosity is a liquid's internal resistance to flow, and it comes from the same intermolecular attraction plus the mechanical interlocking of molecules sliding past one another. The temperature dependence goes in opposite directions for the two fluid states, and the reason is worth understanding rather than memorising:

• In a liquid, heating gives the molecules enough energy to slip past their neighbours more easily, so viscosity falls as temperature rises. Cold oil is thick and hot oil is thin, which is why oil pressure is high and flow poor on a cold start, and why a viscosity index describing how little an oil thins with heat is a specified property of an aircraft lubricant.
• In a gas, there is almost no attraction to overcome; resistance comes from molecules crossing between fast and slow layers and transferring momentum. Heating makes them cross more often, so gas viscosity rises as temperature rises — the opposite of a liquid.

Liquids are described as almost incompressible on the same grounds as solids: the molecules are already nearly touching. "Almost" is the honest word, though. At the pressures a hydraulic system works at — commonly 3,000 psi, about 207 bar — the fluid does compress measurably, by of the order of one per cent of its volume. That is small enough for the hydraulic system to work at all, but large enough to matter: it is part of the sponginess felt in a system with a poorly matched accumulator, and it is one reason a system must be bled properly, since trapped air compresses by orders of magnitude more than the fluid does.

Inside a Gas: Spacing, Diffusion and the Air Itself

In a gas the molecules are perhaps ten diameters apart on average, moving in straight lines at hundreds of metres per second until they collide. Almost all of the volume is empty space, and that single fact accounts for the properties listed for a gas at the head of this section: it has no shape and no fixed volume because nothing holds its molecules together, it exerts pressure because those molecules hammer continuously on the container walls, its density is low because so few molecules occupy a given volume, and it compresses easily because there is space to compress into.

Diffusion follows from the same random motion. Release a gas at one point in a still room and it spreads throughout, without any draught to carry it, simply because its molecules travel in random directions and collisions gradually distribute them evenly. This is why a fuel or hydraulic-fluid leak can be smelled some distance from its source, and why a leaking oxygen system enriches a whole compartment rather than a local pocket — a serious fire risk, because an oxygen-enriched atmosphere makes materials that are normally difficult to ignite burn fiercely.

It is worth separating two words that everyday usage runs together — while keeping hold of the fact that the separation is between two kinds of gas, not between a gas and something else: a vapour is a gas, and the three principal states of matter remain solid, liquid and gas. A gas is above its critical temperature and cannot be liquefied by pressure alone. A vapour is the gaseous form of something that is a liquid at ordinary temperature, and it will condense if compressed enough. Water vapour in the air is a vapour; the nitrogen alongside it is a gas.

Air itself is a mixture of gases, and its composition is examined directly:

Component	Approximate proportion by volume	Notes
Nitrogen	78 % (about four fifths)	Chemically inert at normal temperatures; the bulk of the atmosphere
Oxygen	21 % (about one fifth)	Supports combustion and respiration; the component that matters at altitude
Argon	About 0.9 %	A noble gas; the third most abundant component
Carbon dioxide and trace gases	Under 0.1 %	Includes water vapour, which varies widely with conditions

The working approximation is one fifth oxygen to four fifths nitrogen. Each gas in a mixture contributes its own share of the total pressure in proportion to how much of the mixture it forms, so at sea level oxygen accounts for roughly a fifth of the atmospheric pressure. That proportion does not change with altitude — air remains about 21 % oxygen far above the cruising levels of any transport aircraft — but the total pressure falls, so the oxygen's share of it falls with it. That, and not any change in composition, is why supplemental oxygen and cabin pressurisation are needed.

One practical point follows directly from a gas being compressible: it stores a great deal of energy in its own expansion, so a failed pneumatic component, a burst tyre or a dropped high-pressure cylinder releases that energy explosively — which is why a nitrogen charging trolley is fitted with a regulator and why a wheel assembly is deflated before it is split. A hydraulic system, holding a nearly incompressible fluid, stores far less energy for the same pressure and volume, so a hydraulic leak is a hazard chiefly through its jet and its fluid, not through stored expansion.

How State Changes Density

The density entries in this section's opening table — high for solids and liquids, low for gases — follow directly from particle spacing, and the size of the difference is worth having a feel for. In a solid or a liquid the particles are touching, while at the gas spacing described above the same number of particles occupies of the order of a thousand times the volume. Liquid water at 1,000 kg/m³ becomes steam at atmospheric pressure with a density of only about 0.6 kg/m³, and air at sea level is about 1.225 kg/m³ — roughly eight hundred times lighter than the water it sits above.

Between solid and liquid the change is slight, because the particles do not move appreciably further apart — only their arrangement changes. That is why a solid usually sinks in its own melt, and why ice floating on liquid water stands out as the exception.

Material	Approximate density (kg/m³)	Note
Air at sea level	1.2	Falls with altitude as pressure falls
Aviation kerosene	About 800	Varies with temperature and between batches
Ice at 0 °C	917	Less dense than its own liquid — the anomaly
Water	1,000	The reference against which relative density is quoted
Carbon fibre composite	About 1,600	Roughly 40 % lighter than aluminium alloy
Magnesium alloy	About 1,800	The lightest structural metal in aircraft use
Aluminium alloy	About 2,800	The principal airframe metal
Titanium alloy	About 4,500	Denser than aluminium but far stronger and heat resistant
Steel	About 7,800	Landing gear, fasteners, engine components

Those figures are the reason composite and aluminium structure dominate an airframe while steel and titanium are confined to the places where their strength or temperature capability is genuinely needed. Density — mass per unit volume — is developed further, together with the way materials are compared against water and its role in fluid pressure and buoyancy, in the fluids sections of this module.

Plasma: the Fourth State

Solid, liquid and gas are the three states a substance passes through as it is heated, and they are the three the syllabus concentrates on. Keep heating a gas, though, and something further happens: the collisions become violent enough to knock electrons off the atoms themselves. What remains is a mixture of free electrons and positive ions, still gaseous in the sense that it has no fixed shape or volume, but electrically quite unlike a gas. That is plasma, and it is normally counted as the fourth state of matter — which is why a question offering a choice of four states is not a trick.

A plasma differs from a gas in three ways that matter. Because it contains free charge carriers it conducts electricity, where an ordinary gas is an insulator. Because those carriers are charged it responds to electric and magnetic fields. And because recombining ions release their energy as light, a plasma glows.

Engineers meet it more often than the exotic name suggests: in a lightning channel and in the corona around a highly charged surface; in the arc of an arc welder and the much hotter arc of a plasma cutter; in the ionised gas used in plasma-spray coating to deposit ceramic thermal barriers on turbine components; and in every fluorescent tube, neon sign and old-style discharge lamp. On the largest scale, plasma is the most common state of matter in the universe, since stars are made of it — but on Earth it needs to be produced deliberately.

Changes Between States

When heat energy is added to or removed from a substance, it may change state. These changes are physical (not chemical) — the substance itself is unchanged, only the arrangement and energy of its particles differs.

Change	Direction	Energy
Melting (fusion)	Solid → Liquid	Heat absorbed (endothermic)
Boiling (vaporisation)	Liquid → Gas	Heat absorbed (endothermic)
Freezing (solidification)	Liquid → Solid	Heat released (exothermic)
Condensation	Gas → Liquid	Heat released (exothermic)
Sublimation	Solid → Gas (directly)	Heat absorbed
Deposition	Gas → Solid (directly)	Heat released

What Happens to the Particles at Each Change

Each of the six changes is one event described at particle level, and picturing that event is what makes the energy direction obvious rather than something to be memorised.

• Melting. The particles are already touching; heating them only has to shake them free of their fixed lattice positions so they can slide. They stay in contact, so the volume barely changes — for most substances it increases by a few per cent, because the disordered liquid packs slightly less efficiently than the ordered crystal.
• Boiling. Far more is being asked: every particle has to be pulled clear of its neighbours' attraction altogether and given room to move independently. That is why the volume change is enormous. One kilogram of water occupies about a litre; as steam at 100 °C and atmospheric pressure that same kilogram occupies roughly 1,670 litres — an expansion of more than 1,600 times — and it is why a small quantity of trapped water flashing to steam can do structural damage, and why an autoclave or a steam line is treated with such respect.
• Sublimation. A particle at the surface of a solid escapes straight into the gas phase without ever being part of a liquid. It happens because the particles in a solid do not all carry the same energy: a few in the tail of the distribution have enough to break away, even far below the melting point. A material capable of doing this is a sublimate.
• Deposition. The exact reverse — gas molecules losing energy at a cold surface and locking straight into a solid lattice without a liquid stage. Frost forming on a cold wing overnight is deposition, not freezing: the water was vapour in the air, never liquid on the metal.
• Freezing and condensation simply run melting and boiling backwards, and release exactly the energy those changes absorbed.

The energy direction follows from that picture with no memory work at all. Any change that pulls particles further apart — melting, boiling, sublimation — must supply the energy to overcome the attraction, so heat is absorbed and the change is endothermic. Any change that lets particles fall back together — freezing, condensation, deposition — releases that same energy, so heat is given out and the change is exothermic. If you can decide which way the particles are moving, you can always decide the sign.

Because these are physical rather than chemical changes, they are all reversible and none of them alters the substance itself. Water is water whether it is ice, liquid or steam, and a compound like water can therefore exist in all three states — which is why boiling a kettle produces steam and not hydrogen and oxygen. Only a chemical change would break the molecule apart.

Evaporation Is Not Boiling

Both turn a liquid into a gas, but they are not the same process and the differences are worth holding separately.

	Evaporation	Boiling
Where it happens	Only at the free surface	Throughout the bulk of the liquid, with bubbles forming inside it
Temperature	Any temperature at all	Only at the boiling point for the prevailing pressure
Rate	Slow; increases with temperature, surface area and air movement	Rapid, and limited only by how fast heat can be supplied
Effect on the liquid left behind	Cools it	Temperature stays constant while it boils

The mechanism of evaporation is worth stating properly, because the cooling effect and the humidity behaviour that follow both fall straight out of it. Molecules in a liquid do not all move at the same speed; there is a spread. A molecule that happens to be at the surface, moving upward, and carrying more than the average energy can break free of its neighbours' attraction and escape as gas. It does not need the liquid to be hot — a puddle dries at 5 °C.

Two consequences follow. First, evaporation is selective: it removes the fastest molecules, so the average energy of those left behind falls and the liquid cools itself. That is why solvent or fuel on the skin feels cold, why a wetted surface reads below air temperature, and why an aircraft standing in rain or with fuel spilled on it can chill locally below the surrounding air. Second, evaporation continues until the air in contact with the liquid can hold no more — the air is then saturated, and molecules are returning to the liquid as fast as they leave.

How much vapour air can hold rises steeply with its temperature. Relative humidity is the amount of water vapour present expressed as a percentage of the maximum the air could hold at that temperature. Cool a sample of air without changing its water content and the maximum falls until it equals what is actually there: at that temperature, the dew point, the air is saturated and any further cooling forces water out as condensation. If the dew point is below freezing, the water is deposited directly as frost instead, and that temperature is called the frost point.

This is the single most useful idea in this section for practical work. Condensation and frost do not appear because the air is wet; they appear because a surface is at or below the dew point of the air touching it. Warm humid air meeting cold aircraft structure — a fuselage descending from cruise, a fuel tank cold-soaked after a long flight, a windscreen on a humid morning — produces water on the metal whatever the weather is doing. That is one of the ways water accumulates in a fuel tank sump, and it is why the sump is drained before flight rather than after.

Pressure Changes the Melting and Boiling Points

The melting and boiling points quoted for a substance are only valid at a stated pressure, and boiling is the more sensitive of the two by far. A liquid boils when its molecules can form vapour bubbles inside it, and they can only do that when the vapour's own pressure has risen to match the pressure pushing down on the surface. Raise the surrounding pressure and the liquid must be hotter before that happens; lower it and boiling starts sooner.

• At the standard sea-level pressure of 1,013 hPa (about 14.7 psi) water boils at 100 °C. That figure is a consequence of the pressure, not a property of water on its own.
• At a cabin altitude of 8,000 ft the pressure is around 75 kPa and water boils at roughly 92 °C. On an unpressurised aircraft at high altitude it is lower still.
• Raise the pressure instead and the boiling point rises. That is the principle of a pressurised cooling system, which allows a coolant to run above 100 °C without boiling, and of an autoclave, which cures composite structure at a temperature that would boil off the volatiles at atmospheric pressure.

Melting points shift far less with pressure, and water is the notorious exception to the usual direction. For almost every substance the solid is denser than the liquid, so squeezing it favours the solid and raising the pressure raises the melting point. Water's solid is less dense than its liquid, so pressure favours the liquid and raising the pressure lowers the melting point of ice. It is a small effect — it takes a great deal of pressure to move the melting point by a degree — but it is the reason ice under high local contact pressure can develop a film of water, and it is a direct consequence of the expansion-on-freezing dealt with below.

The Triple Point and the Critical Point

Plot pressure against temperature for a substance and the three states occupy three regions, separated by lines along which two states can coexist. The melting point line, for instance, is the set of conditions at which solid and liquid are in equilibrium — which is the precise definition of a melting point: not the temperature at which a solid vanishes, but the temperature at which its solid and liquid phases can exist together indefinitely, neither one growing at the other's expense.

Two points on that diagram have names. The triple point is asked about directly; the critical point earns its place for a different reason, because it is what divides a gas from a vapour:

• The triple point is the single combination of temperature and pressure at which all three states — solid, liquid and vapour — coexist in equilibrium. For water it is 0.01 °C, so for exam purposes approximately 0 °C, at a pressure of only about 611 Pa (roughly 6 mbar, well under one hundredth of atmospheric). Do not confuse it with the dew point, which is a property of a particular sample of air and its moisture content, not a fixed property of the substance.
• The critical point is the temperature above which a substance cannot be liquefied by pressure alone, however hard it is squeezed. For water it is about 374 °C at roughly 22 MPa. Above its critical temperature a substance is a gas in the strict sense; below it, the gaseous form is properly a vapour.

The diagram also explains why some substances sublime in ordinary conditions and others never do. If a substance's triple-point pressure is below atmospheric pressure, then at atmospheric pressure heating the solid takes it into the liquid region first, and it melts — which is what water does, its triple point being far below one atmosphere. If the triple-point pressure is above atmospheric, there is no liquid region to pass through at atmospheric pressure at all, and the solid goes straight to gas.

Carbon dioxide is the standard demonstration: its triple point is at about 5.1 atmospheres, so at ordinary atmospheric pressure liquid carbon dioxide cannot exist and solid carbon dioxide sublimes directly to gas at about −78.5 °C. That is why it is called dry ice, why it leaves no puddle, and why it is used to keep cargo cold without wetting it. Inside a pressurised extinguisher cylinder, where the pressure is far above 5.1 atmospheres, the same substance sits quite happily as a liquid.

Why Water Expands When It Freezes

Almost every substance contracts as it solidifies, because an ordered lattice packs its particles more tightly than a disordered liquid. Water does the opposite, and the reason is the hydrogen bonding described earlier. As water cools towards freezing, its molecules line up so that each hydrogen bond can point directly at a neighbouring oxygen, and that arrangement is an open hexagonal framework with a hole at the centre of every ring. The ordered structure is less densely packed than the jumbled liquid.

The numbers make the consequences concrete. Liquid water has a density close to 1,000 kg/m³, while ice at 0 °C is about 917 kg/m³. Freezing a given mass of water therefore increases its volume by about 9 %. Run it the other way and the same arithmetic says that melting a block of ice leaves the water occupying about 8 % less space than the ice did: melted ice takes up less room, not more.

Being less dense, ice floats: it displaces its own weight of water, and in fresh water roughly 8 % of its bulk stands proud of the surface. Water is also unusual in reaching its maximum density at about 4 °C rather than at its freezing point, so the coldest water in a still pond rises to the top and freezes there, leaving the depths liquid.

In maintenance the expansion is a hazard rather than a curiosity, because it is powerful and it happens in confined spaces. Water trapped in a blind hole, a lap joint, a honeycomb core or a control-surface hinge line expands as it freezes with enough force to spread a joint, unseat a fastener, split a water line or crush a honeycomb cell — and the damage is only discovered later, when the ice has melted and left no evidence. It is also why potable-water and waste systems in unheated areas are drained or heated, and why water found in a fuel sump matters as much for what it can do to structure as for what it does to the fuel.

Supercooling and the Effect of a Dissolved Solute

A liquid does not always freeze the moment it reaches its freezing point. To start freezing it needs a nucleus — a speck of dust, a scratch on the container, an existing ice crystal — on which the first ordered lattice can build. Clean water in a smooth container, or a small airborne droplet with nothing to nucleate on, can be cooled well below 0 °C and remain liquid. That condition is called supercooling, and it is unstable: disturb the liquid or give it a nucleus and it freezes almost instantly, its temperature jumping back up to the freezing point as the latent heat it had been holding is released.

This is not a laboratory curiosity for an aircraft engineer — it is the whole basis of in-flight icing. Cloud droplets are small, clean and suspended, so they routinely exist as liquid water at temperatures far below freezing. The aircraft arriving through them is the nucleus: on impact the droplet freezes onto the structure at once. No supercooling, no airframe icing: ice crystals that are already frozen largely bounce off a cold surface rather than adhering to it.

Dissolving anything in water moves both of its fixed points, and the directions are opposite:

• The freezing point falls. Solute particles get in the way of water molecules trying to lock into the ice lattice, so a lower temperature is needed before freezing can proceed. Salt water freezes below 0 °C, and a glycol and water mixture stays liquid far below it — the more concentrated the mixture, the lower the freezing point.
• The boiling point rises. Solute particles at the surface reduce the rate at which water molecules escape, so a higher temperature is needed to reach boiling. This is why an engine coolant mixture both resists freezing in winter and boils at a higher temperature than plain water.

Both effects depend on how much is dissolved rather than on what it is, so they are undone by dilution: add more water to the mixture and its freezing point climbs back towards 0 °C. That is the physics a ground de-icing operation is working with, and it is taken up again at the end of this section.

Latent Heat

During a change of state, the temperature remains constant even though heat is being added or removed. The energy is used to break (or form) intermolecular bonds rather than to change the temperature. This energy is called latent heat.

The two latent heats have names and standard values. The latent heat of fusion is the energy needed to melt one kilogram of a solid already at its melting point; for water it is about 334 kJ/kg. The latent heat of vaporisation is the energy needed to boil one kilogram of a liquid already at its boiling point; for water it is about 2,260 kJ/kg. Freezing and condensing release exactly those same quantities back again — the direction reverses, the magnitude does not.

Setting them beside the energy needed simply to warm the water shows how large they are. Raising 1 kg of water through the whole 100 °C from freezing to boiling takes about 419 kJ. Boiling that same kilogram away, at constant temperature, takes about 2,260 kJ — more than five times as much energy for no rise in temperature whatever.

The reverse direction is the reason steam is so dangerous. Steam at 100 °C condensing on skin delivers the whole 2,260 kJ/kg of latent heat at the point of contact before the resulting water has cooled by a single degree, which is why a steam burn is far worse than a scald from boiling water at the same temperature. The same effect works usefully elsewhere: a vapour-cycle refrigeration system moves heat by evaporating a refrigerant where heat is to be absorbed and condensing it where heat is to be rejected, and the latent heat is what does the carrying. An air-cycle machine, which is what most turbine-powered aircraft use instead, is the instructive contrast: its working fluid is the air itself, passing straight through the machine and into the cabin rather than round a sealed loop, and it is cooled by the shaft work taken out of it as it expands through a turbine. Nothing changes state to carry the load there, so the heat is carried as sensible heat — a temperature change rather than a change of state — and a much larger mass flow is needed for the same cooling duty. Any moisture that condenses in such a system is water being wrung out of the air stream for removal, not a refrigerant condensing to reject the load.

The relative size of the two latent heats also sets the shape of the temperature-against-time graph. Supply heat at a steady rate to a block of ice and the trace rises, flattens at 0 °C while the ice melts, rises again through the liquid range, then flattens a second time at 100 °C while the water boils. The second plateau is far the longer of the two — about seven times as long for water, since 2,260 kJ/kg has to be delivered at the boiling point against 334 kJ/kg at the melting point. Cooling the same substance reproduces the identical trace in reverse, with heat coming out at each plateau instead of going in.

Changes of State in Aircraft Operation

The mechanisms behind that are worth separating, because they call for different remedies. Airframe icing in flight is supercooling put into practice: small droplets freeze on impact almost immediately and give opaque rime ice, while larger ones spread across the surface before freezing and give clear ice, which is heavier, harder and adheres more strongly. Ground frost is a different process altogether: it is deposition, water vapour turning straight to solid on structure that is below the frost point, and it forms on a clear night with no precipitation at all. A wing that has been cold-soaked by cold fuel can grow frost on its upper surface in the morning even when the air temperature is above freezing.

The counter-measures divide the same way. An anti-icing system is run before ice can form and prevents accretion; a de-icing system is allowed to accrete a layer and then removes it. Thermal systems, whether they bleed hot air or use electrical heating elements, supply the latent heat needed to melt the ice at the surface. Ground de-icing and anti-icing fluids work on a different principle again: they are glycol-based, so they depress the freezing point of the water film rather than adding heat to melt it, and the holdover time they provide is the period before dilution undoes that protection.

Sublimation and deposition appear in the same operation. Frost on a cold-soaked upper surface can sublime away in dry air at altitude without ever melting; conversely, moisture in an oxygen or pneumatic system can deposit as ice at a restriction and block it, which is why aviator's breathing oxygen is specified dry and why moisture separators are fitted where compressed air is used.

Fuel deserves its own note because two separate things happen to it in the cold. Jet A-1 is specified to a maximum freezing point of about −47 °C, below which wax crystals are present in the fuel and can block filters; that is a property of the fuel, and it is managed by fuel specification and by monitoring fuel temperature. Quite separately, the small quantity of water that fuel always carries in solution comes out as the fuel cools and freezes into ice crystals at around 0 °C, long before the fuel is anywhere near its own freezing point. Ice crystals are the reason for water drains, filter icing indications and inhibitor additives. Confusing the two leads to the wrong troubleshooting entirely.

Finally, the same physics runs in the useful direction throughout an aircraft. Water drains work because water is denser than fuel and separates to the lowest point. Carbon dioxide and halogenated extinguishing agents are stored as liquids and discharge as gases, the change of state cooling the discharge sharply. Air-conditioning water separators condense moisture out of cabin air by cooling it below its dew point. Cabin condensation forms on cold structure behind trim panels during descent into warm humid air, and it drips back out later — a nuisance, and a corrosion risk in exactly the places that are hardest to inspect.