Phosphorus allotropes (white, red, black); P₄O₆, P₄O₁₀; oxoacids of P (H₃PO₃ vs H₃PO₄ basicity): Foundations

Before you touch the parent note, you must be fluent in the small alphabet of symbols it throws at you. We build each one from nothing, anchor it to a picture, and say why the topic can't live without it.

1. What a chemical symbol and formula actually mean

Why the topic needs this: the whole chapter is a game of counting atoms — 6 bridges, 4 caps, 3 OH groups. If you misread a subscript, every later count collapses.

2. Where an atom keeps its electrons — shells and orbitals

Before we can say why phosphorus makes bonds, we need to know where its outermost electrons live. Electrons around a nucleus are not scattered randomly; they occupy shells and, inside each shell, small "rooms" called orbitals.

Figure — Phosphorus allotropes (white, red, black); P₄O₆, P₄O₁₀; oxoacids of P (H₃PO₃ vs H₃PO₄ basicity)

Figure s01 (above) shows phosphorus' outer shell ( $n = 3$ ): the round $s$ room holds a full pair, and the three $p$ rooms each hold one lonely electron. Read the picture left-to-right and count the single dots — there are three unpaired electrons. That "three" is the number we chase for the rest of the page.

Why the topic needs this: orbital labels ( $s$ , $p$ ) and the shell number $n$ are the vocabulary that lets us predict bonding instead of memorising it.

3. The covalent bond — a shared pair of electrons

Figure s02 (above) contrasts a single line and a double line. Picture two LEGO men each lending one hand to hold the same rope: a single line = one rope (one shared pair), a double line ( $=$ ) = two ropes. Look at the little dots drawn on each line — one pair on the single bond, two pairs on the double bond.

Intuition Why P can form that extra bond (the honest picture)

A double bond needs a second shared pair (the $π$ -bond of the definition above). The old textbook story was that phosphorus promotes an electron into an empty $3 d$ orbital and uses it — you will still see " $s p^{3} d$ " written in many books. The modern, more accurate story: careful quantum-chemical calculations show that the $3 d$ orbitals of phosphorus lie too high in energy to take part meaningfully — promoting an electron there costs far more energy than any bond could pay back. What actually stabilises the extra P–O bonding is (1) strong polarisation: greedy oxygen (see §4) pulls electron density so hard that P's ordinary $3 s$ / $3 p$ orbitals do the extra bonding, and (2) multi-centre bonding, where a bonding pair is spread over several atoms rather than squeezed between just two. What to remember: P can form the fourth (P=O) bond because it is a large third-row atom next to greedy oxygen — not because a real $3 d$ orbital does the work. So every $P = O$ you see (capping each oxoacid and each corner of $P_{4} O_{10}$ ) is genuine, but the " $3 d$ promotion" rationale is a simplification to be treated with caution.

Why the topic needs this: without that fourth bond there would be no $+ 5$ state, no $P_{4} O_{10}$ , no $H_{3} P O_{4}$ — half the chapter.

4. Electronegativity — the tug-of-war number

Values you will use (Pauling): $E_{O} \approx 3.44$ (very greedy), $E_{H} \approx 2.20$ , $E_{P} \approx 2.19$ .

Figure s03 (above) draws the same bond twice. On the left (O–H) the shared pair is dragged far toward oxygen — see the red arrow and the dots crowding O. On the right (P–H) the pair sits dead centre because the two pull numbers are nearly equal. That picture is the acidity rule.

This single contrast is the whole basicity rule of the parent note: only O–H hydrogens leave; P–H hydrogens stay. It is also the same strong pull that (from §3) lets oxygen coax that extra P=O bond out of phosphorus.

5. Oxidation state — a bookkeeping charge

Three rules cover everything here:

Oxygen is usually $- 2$ — but note the exceptions: in peroxides (an O–O bond, e.g. $H_{2} O_{2}$ ) it is $- 1$ , in superoxides $- \frac{1}{2}$ , and bonded to fluorine it can be positive. None of these occur in P oxides/oxoacids, so we safely use $- 2$ throughout this chapter.
Hydrogen bonded to a non-metal is $+ 1$ .
The oxidation states of all atoms must add up to the overall charge ( $0$ for a neutral molecule).

Why the topic needs this: $+ 3$ vs $+ 5$ is the label distinguishing $P_{4} O_{6}$ from $P_{4} O_{10}$ , and $+ 1/ + 3/ + 5$ ranks the three oxoacids.

6. Geometry words — tetrahedron and bond angle

Figure s04 (above) projects the $P_{4}$ tetrahedron flat. Count the orange lines — there are 6 edges joining the 4 corner P atoms; the violet arrow marks the pinched $60°$ angle inside one triangular face.

Notice the "6 edges, 4 corners" — memorise it, because oxide-building is literally "one O in each of the 6 edges, then one O capping each of the 4 corners."

7. Reading the bonds inside an oxoacid

The parent's golden rule needs you to see which H sits where. Here is the recipe for any $H_{3} P O_{n}$ :

Central $P$ , always with one $P = O$ (terminal double-bond oxygen — one $σ$ + one $π$ , made possible by the §3 story of large-atom P next to greedy O).
The remaining connections split between $- O H$ groups (acidic) and $- H$ (P–H, not acidic).

Acid	Structure around P	O–H groups	P–H bonds	Basicity
$H_{3} P O_{4}$	one $= O$ , three $- O H$	3	0	3
$H_{3} P O_{3}$	one $= O$ , two $- O H$ , one $- H$	2	1	2
$H_{3} P O_{2}$	one $= O$ , one $- O H$ , two $- H$	1	2	1

Why the topic needs this: without §3 (bond lines) and §4 (polar O–H vs non-polar P–H) this table is memorisation; with them it is understanding.

Prerequisite map

The box below is a Mermaid diagram — a simple flow-chart. Read it top to bottom: each rounded box is a concept from this page, and an arrow $\to$ means "you need the box behind the arrow before the one it points to." Follow the arrows and they all funnel into the phosphorus topic at the bottom.

Equipment checklist

Test yourself — cover the right side and answer out loud.

What does the subscript in

P_{4} O_{10}

tell you?

There are 4 P atoms and 10 O atoms in one unit.

What does the shorthand

# (\dots)

mean?

"The number of" — a plain counting instruction, e.g.

# (P-OH groups)

= how many P–OH groups there are.

What do the labels

s

and

p

describe, and what does

n

mean?

They are shapes of orbitals (rooms for electrons) inside a shell;

n

is the principal quantum number, telling you which shell (how far out) the electrons live.

How many unpaired electrons does phosphorus' outer shell (

3 s^{2} 3 p^{3}

) have, and so how many bonds does it naturally form?

3 unpaired

p

-electrons, so 3 covalent bonds by default.

What is the difference between a

σ

-bond and a

π

-bond?

σ

-bond overlaps head-on along the atom–atom line (the plain single bond); a

π

-bond overlaps sideways above and below it, forming the second bond of a double bond.

Why can phosphorus form a fourth bond (

P = O

) when nitrogen cannot — and what is the modern reason?

P is a large third-row atom next to greedy oxygen; strong polarisation and multi-centre bonding (using its

3 s

3 p

orbitals) allow the extra bond. The old "

3 d

orbital promotion" story is now considered an oversimplification, since

3 d

orbitals lie too high in energy to contribute much.

What is the difference between a single line

-

and double line

=

in a bond diagram?

A single line is one shared electron pair; a double line is two shared pairs (

σ + π

Which bond is polar, O–H or P–H, and what makes it so?

O–H is polar because oxygen (

E \approx 3.44

) pulls electrons far harder than hydrogen (

E \approx 2.20

); P–H is non-polar because

E_{P} \approx E_{H} \approx 2.2

(Pauling scale).

Why is a P–H hydrogen NOT acidic?

The P–H bond is essentially non-polar, so H is held firmly and cannot ionise as

H^{+}

What oxidation state is oxygen usually assigned, and name an exception?

Usually

- 2

; exceptions include peroxides (

- 1

, as in

H_{2} O_{2}

), superoxides (

- \frac{1}{2}

), and O bonded to fluorine (positive). None occur in P oxides/oxoacids.

Assign the oxidation state of P in

H_{3} P O_{4}

, showing the sum.

(+ 3) + x + (- 8) = 0 \Rightarrow x - 5 = 0 \Rightarrow x = + 5

(add

+ 5

to both sides).

How many corners and edges does a tetrahedron have?

4 corners and 6 edges.

Why does the

60°

P–P–P angle make white P reactive?

It forces strained bonds (angular strain stores energy), so bonds break easily.

Basicity means the number of what?

Ionisable (P–OH) hydrogens the acid can release as

H^{+}

— not the total number of H atoms.

Using the rule, what is the basicity of

H_{3} P O_{2}

3 - 2 (P-H bonds) = 1

Connections

Parent note (Hinglish)
p-Block Group 15 (Nitrogen family) trends
Oxidation states and disproportionation
Oxides — acidic vs basic character
Reducing agents and P–H bonds
Allotropy — Carbon, Sulfur, Oxygen comparison