Acids, Bases and pH

ILFORD publishes the pH of almost everything it sells. MULTIGRADE paper developer at 1+9 reads 10.45 to 10.55; ID-11 film developer stock reads 8.60 to 8.70; RAPID FIXER at 1+4 reads 5.0 to 5.5; ILFOSTOP concentrate reads 2.1. Those numbers span more than eight units, which sounds like a modest range and is in fact a range of more than a hundred million to one in the quantity that actually matters. This page is about what that quantity is, why the scale hides its own size, and what each bath is using it for.

Acids and bases, twice over

The first useful definition, due to Arrhenius, is the one most people half-remember: an acid releases hydrogen ions, H⁺, into water, and a base releases hydroxide ions, OH⁻. Hydrochloric acid gives H⁺ and Cl⁻; sodium hydroxide gives Na⁺ and OH⁻. It works, and for the darkroom it is nearly always enough.

It fails on the substances photography actually uses most. Sodium carbonate contains no hydroxide at all, and its solution is strongly alkaline. Borax contains no hydroxide, and its solution is alkaline. Sodium sulfite contains no hydroxide, and its solution is alkaline. The Arrhenius picture has nothing to say about any of them.

The Brønsted–Lowry definition fixes it by moving attention from what a substance contains to what it does. An acid is a proton donor; a base is a proton acceptor. A hydrogen ion is a bare proton, so donating H⁺ and accepting H⁺ are the whole story. Carbonate is a base because it accepts a proton:

CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻

Carbonate acts as a base by taking a proton from water

It has taken a proton from a water molecule and left hydroxide behind. That is where the alkalinity of a carbonate developer comes from, and no hydroxide had to be in the tin.

A conjugate pair is one idea, not two

Notice what carbonate became: hydrogencarbonate, HCO₃⁻. That species can give the proton back. The two are a conjugate acid–base pair — the same chemical entity with and without one proton, and the only difference between them.

Acid	Conjugate base	Where you meet the pair
CH₃COOH, acetic acid	CH₃COO⁻, acetate	Stop baths, and the acid reserve of a fixer
H₂CO₃ / HCO₃⁻	HCO₃⁻ / CO₃²⁻	Carbonate developers
H₃BO₃, boric acid	B(OH)₄⁻, borate	Borax and metaborate developers; hardening fixers
HSO₃⁻, hydrogensulfite	SO₃²⁻, sulfite	Every bath containing sulfite or metabisulfite
H₂O	OH⁻	Everywhere, including water itself

Water appears twice in that table, as an acid in one row and a base in the reaction above. A substance that can do both is amphiprotic, and water’s ability to be both is the reason the next section exists.

Water reacts with itself

Two water molecules can trade a proton. One donates, one accepts, and you are left with a hydronium ion, H₃O⁺, and a hydroxide ion.

2 H₂O ⇌ H₃O⁺ + OH⁻

The autoionisation of water

It happens very little. OpenStax puts it at about two water molecules in a billion ionised at 25 °C. But it never stops, and it sets a floor under every aqueous solution: there is no such thing as water with no hydrogen ion in it.

The equilibrium constant for it is the ion product of water, Kw, and because water itself is the solvent it does not appear in the expression:

K_w = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C

The ion product of water

That single equation controls more darkroom chemistry than any other. The two ion concentrations multiply to a fixed number, so you cannot raise one without lowering the other. Push hydrogen ion up with an acid and hydroxide falls in exact proportion. Push hydroxide up with an alkali and hydrogen ion falls. There is only one dial.

In pure water the two must be equal, so each is the square root of 10⁻¹⁴, that is 1.0 × 10⁻⁷ mol/L. That is what neutral means.

!Neutral is 7 only at 25 °C, and your darkroom is not at 25 °C

Neutrality is defined by the hydrogen and hydroxide concentrations being equal, not by the number 7. The number 7 follows from Kw at 25 °C and from nothing else.

Kw changes with temperature. OpenStax states that the autoionisation is endothermic, so both ion concentrations rise as the water gets hotter, and gives Kw at 100 °C as about 5.6 × 10⁻¹³ — some fifty times its value at 25 °C — at which temperature pure water is neutral at pH 6.31.

Below 25 °C the effect runs the other way: Kw is smaller, both ions are scarcer, and the neutral point sits a little above 7.00. A darkroom at 20 °C is therefore not a place where 7.00 means neutral, though it is close. This course has not verified a published value of Kw at 20 °C and does not quote one; what it states is the direction, which is sourced, and the practical consequence, which is that a pH reading only means what it says when its temperature is recorded beside it. That is why the pH laboratory session records a temperature with every reading.

pH, and the logarithm you actually need

Hydrogen ion concentrations in photography run from about 10⁻² mol/L in a stop bath to about 10⁻¹¹ in a paper developer. Writing those out is unbearable and comparing them by eye is worse, so we take the logarithm.

This is the first logarithm in this course, so here it is in plain language. The base-ten logarithm of a number is the power of ten that makes it. log 10 = 1, because 10 is 10¹. log 100 = 2. log 1000 = 3. log 0.001 = −3, because 0.001 is 10⁻³. That is all a logarithm is: it answers “ten to the what?”

pH is the negative of that, applied to the hydrogen ion concentration in moles per litre:

pH = −log₁₀[H₃O⁺]

Definition of pH

The minus sign is there for convenience: hydrogen ion concentrations are small numbers, their logarithms are therefore negative, and flipping the sign gives a scale that runs positive over the useful range. It also, unavoidably, makes the scale run backwards: a higher pH means less hydrogen ion.

Going the other way is the same operation in reverse:

[H₃O⁺] = 10^−pH

From pH back to concentration

And because Kw fixes the product, the two ions have a matching relation:

pH + pOH = 14.00 at 25 °C, where pOH = −log₁₀[OH⁻]

pH and pOH

One unit is a factor of ten, and it never stops being one

What one pH unit means

1. 1One rung — one pH unit: ten times less hydrogen ion, ten times more hydroxide
2. 2The product never moves — [H₃O⁺] × [OH⁻] = 1.0 × 10⁻¹⁴ on every rung, at 25 °C
3. 3Two real baths — ID-11 film developer at 8.6 and MULTIGRADE paper developer at 10.5: 1.9 units, a factor of 79

The dot columns show the ratio between rungs, not an absolute number of ions.

Work one conversion in each direction, because you will need both.

🧮Two conversions, worked

Forward. A print developer reads pH 10.5. What is the hydrogen ion concentration? [H₃O⁺] = 10⁻¹⁰·⁵ = 3.2 × 10⁻¹¹ mol/L. And the hydroxide: pOH = 14.00 − 10.5 = 3.5, so [OH⁻] = 10⁻³·⁵ = 3.2 × 10⁻⁴ mol/L.

Backward. A solution is found to contain 2.5 × 10⁻⁹ mol/L of hydrogen ion. Its pH is −log(2.5 × 10⁻⁹) = 8.6 — which is ID-11 stock developer.

And the comparison that matters. From pH 8.6 to pH 10.5 is 1.9 units, so the paper developer has 10¹·⁹ = 79 times as much hydroxide ion and one seventy-ninth as much hydrogen ion. “A bit more alkaline” is a description of the number; a factor of eighty is a description of the chemistry.

Part XIII takes this same operation and applies it to light and to density, where log exposure and optical density are the same trick played on a different quantity. It will not teach the logarithm again; it will assume this page.

Strong against weak, and why per cent is not a strength

An acid’s strength is not how much of it there is. It is what fraction of it lets go of its proton in water.

A strong acid ionises essentially completely: put hydrochloric acid in water and there is effectively no HCl left, only H₃O⁺ and Cl⁻. A weak acid ionises only partly, and sits at an equilibrium between the intact molecule and its ions:

CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺

Acetic acid, most of which stays intact

The equilibrium constant for that is the acid ionisation constant Ka, and because Ka values span many orders of magnitude they are usually quoted as pKa = −log Ka, on the same logarithmic principle as pH. A small Ka, and so a large pKa, means a weak acid.

Acid	Ka at 25 °C	pKa	Source
Sulfurous acid, first proton	1.6 × 10⁻²	1.80	OpenStax Appendix H
Sulfurous acid, second proton	6.4 × 10⁻⁸	7.19	OpenStax Appendix H
Acetic acid	1.8 × 10⁻⁵	4.74	OpenStax Appendix H
Carbonic acid, first proton	4.3 × 10⁻⁷	6.37	OpenStax Appendix H
Carbonic acid, second proton	4.7 × 10⁻¹¹	10.33	OpenStax Appendix H
Boric acid	5.4 × 10⁻¹⁰	9.27	OpenStax Appendix H
Citric acid, three protons	—	2.87, 4.35, 5.68	IUPAC dataset, one determination at 20 °C in 0.1 mol/L NaClO₄

The citric acid row carries its conditions because it has to. The IUPAC compilation lists several determinations of the same three constants that differ by up to a few tenths of a unit depending on ionic strength and method, and quoting one without its conditions would be quoting a number that does not exist on its own.

Two per cent of what?

Take 2 per cent w/v solutions of acetic acid and of hydrochloric acid — the same mass of acid in the same volume of water — and compare them.

🧮Same percentage, different universe

Acetic acid. 20 g/L ÷ 60.05 g/mol = 0.333 mol/L. Because it is weak, only a little ionises, and for a weak acid the hydrogen ion concentration is close to the square root of Ka times the concentration: √(1.8 × 10⁻⁵ × 0.333) = √(6.0 × 10⁻⁶) = 2.4 × 10⁻³ mol/L. So pH 2.6, and only 2.4 × 10⁻³ ÷ 0.333 = 0.7 per cent of the acetic acid molecules have ionised at all.

Hydrochloric acid. 20 g/L ÷ 36.46 g/mol = 0.549 mol/L, ionised completely, so [H₃O⁺] = 0.549 mol/L and pH 0.26.

The gap is 2.35 pH units, which is a factor of 224 in hydrogen ion concentration, between two solutions a label would describe identically.

And now the fact that makes weak acids useful rather than merely feeble. Kodak’s 1928 primer states it with complete clarity and without using the word pH at all: the strength of an acid depends on the proportion of its hydrogen that has dissociated, “however, the quantity of alkali which the acid can neutralize depends upon the total quantity of the hydrogen present, and not on the dissociated portions only.”

Read that twice. Strength and capacity are different properties. At equal molar concentration, acetic acid and hydrochloric acid neutralise exactly the same amount of alkali; the acetic acid simply does it from a much less acidic solution. That is why Kodak concludes that “the best acid for the purpose” of an acid fixing bath is acetic acid: a fixer needs a large reserve of acid to absorb alkaline developer carried in on wet film, and yet must not be strongly acid, because strong acid decomposes thiosulfate and precipitates sulfur. A weak acid in quantity is the only way to have both.

Strong and weak acid at the same concentration

1. 1Strong acid — essentially every molecule has given up its proton: low pH
2. 2Weak acid at the same molarity — most molecules are still intact: much higher pH
3. 3Neutralising capacity — equal for both, because it counts total hydrogen and not only the dissociated part

The proportions in the weak-acid box are drawn to make the point visible; for 0.33 mol/L acetic acid the real dissociated fraction is about 0.7 per cent.

The acids and alkalis of the darkroom, and why a formula picks one

Acetic acid is the standard stop bath and the standard acid of a fixer, for the reason above: weak enough to be used in quantity, cheap, and its salt, acetate, does nothing to the process it is left behind in. Its disadvantage is its smell, and its concentrate is corrosive — the aggregated ECHA notifications on PubChem give glacial acetic acid H226, flammable, and H314, causes severe skin burns and eye damage, which is a property of the concentrate and not of a working stop bath.

Citric acid does the same job as a solid with three ionisable protons and no smell. ILFORD’s own ILFOSTOP is described on its technical sheet as “a low odour citric acid stop bath”, which is the reason it appears on the scale above at pH 2.1 rather than smelling of vinegar.

Sulfuric acid is the strong acid photography does use, and Kodak Ltd’s 1949 handbook shows where: in the hardening solution of the chrome-alum fixing bath F-16, in the permanganate reducer R-2, and in the quinone–thiosulfate intensifier IN-6. It is never the acid of a stop bath, because it is strong, and a strong acid in the quantity a stop bath needs would take the pH far below where thiosulfate is stable. Note what the handbook prints beside two of those formulas: “add the sulphuric acid very gradually to the water with constant stirring.” That instruction is the subject of the next section.

Sodium hydroxide and potassium hydroxide are the caustic alkalis. Kodak’s 1924 primer notes the problem with them in a developer: a caustic alkali is present as a fixed dose, and once it has been consumed there is no more. It also warns that too alkaline a developer over-swells the gelatin and gives trouble with frilling and blisters in warm weather.

Sodium and potassium carbonate are the standard developer alkalis, and Kodak’s 1924 explanation of why is the best short account of buffering in the course’s sources: carbonic acid is a very weak acid, so a carbonate solution is alkaline; and the carbonate acts as “a sort of reservoir of alkali, only a small amount of alkali being present at any time, but more being generated by dissociation of the carbonate as it is used up,” so that a small alkali concentration is kept “nearly constant during use.” That mechanism is the whole subject of the next page.

Borax and sodium metaborate are the mild alkalis. A borax solution is alkaline — HSDB puts a saturated one at about pH 9.5 — and it holds a much lower pH than a carbonate bath. Sodium metaborate sits between the two, and Kodak Ltd’s 1949 handbook introduces it in exactly those terms: Kodalk is “a new alkali, introduced by Kodak Ltd., intermediate in activity between sodium carbonate and borax”, and films developed in a Kodalk developer “will not blister when placed in an acid fixing bath, even at high temperatures” — because, unlike a carbonate, it releases no carbon dioxide when the acid reaches it.

Sodium sulfite, bisulfite and metabisulfite are three points on one acid–base scale. Sulfite is the base, hydrogensulfite the conjugate acid, and disulfite (metabisulfite) becomes hydrogensulfite on dissolving. A formula chooses among them partly for the sulfite it wants and partly for the pH that choice brings with it. Kodak’s 1928 primer names sodium bisulfite as the substitute acid for a fixing bath when acetic acid cannot be had, describing it as “intermediate between sodium sulphite and sulphurous acid” and therefore equal in acidity to a mixture of the two.

Neutralisation, and the order you mix in

An acid and a base react to give a salt and water. Stripped to the ions that actually do something, the neutralisation of any strong acid by any strong base is one reaction:

H₃O⁺ + OH⁻ → 2 H₂O

Neutralisation, net ionic form

The arithmetic is moles, not grams and not millilitres: one mole of hydrogen ion is neutralised by one mole of hydroxide ion. The molarity conversions are on the concentration page.

Neutralisation releases heat, and so does dissolving a concentrated acid or a solid alkali in water. The two together are why the mixing order is not a preference. Princeton’s guidance for photographic work states the rule plainly — always add acid to the water when diluting — and Kodak Ltd’s 1949 handbook prints it beside the individual formulas that need it: “add the sulphuric acid very gradually to the water with constant stirring.”

Poured that way, a small amount of acid is diluted into a large mass of water, which absorbs the heat and barely warms. Poured the other way, water lands on concentrated acid, the heat is released in a small volume right at the surface, and the mixture can boil and spit acid back out of the vessel. The same argument applies to solid sodium hydroxide, whose dissolution CAMEO records as liberating enough heat to make a solution steam and spatter.

pH, bath by bath

Where the baths of this course sit on the pH scale, with their published values

1. 1Stop bath — ILFOSTOP concentrate, pH 2.1; diluted 1+19 for use, and the working pH is not published on the sheet
2. 2Fixers — RAPID FIXER 1+4 and HYPAM 1+4 and 1+9, all pH 5.0 to 5.5
3. 3Neutral region — ILFOTOL 7.00 to 7.02 and WASHAID 1+4 at 7.00 to 7.20
4. 4Film developers — PERCEPTOL 7.68 to 7.82, DD-X 8.45 to 8.55, ID-11 8.60 to 8.70, MICROPHEN 8.67 to 8.93, LC29 8.9 to 9.00
5. 5Paper developers — BROMOPHEN 10.30 to 10.50, MULTIGRADE 10.45 to 10.55, PQ UNIVERSAL 10.48 to 10.58

Every value is the manufacturer's published figure for a fresh solution at the stated dilution; ILFORD notes that its figures were obtained under controlled laboratory conditions and that users should measure their own.

Developer activity. Nearly every developing agent works in its deprotonated, anionic form, so the more alkaline the bath the more of the agent is in the active form and the faster development runs. That is the mechanism, and it is why paper developers sit two units above film developers: a print is developed to completion in a minute or two and a negative is developed part-way over ten. The direction is certain. The shape of the relationship is not something this course has a measured curve for, and the plot below says so on its face.

How development rate responds to pH: the shape, not a measurement

• Relative rate, drawn to show the shape

Show the numbers behind this plot

A rising S-shaped curve. Development is very slow below pH 8, rises steeply through pH 9 and 10, and begins to flatten above pH 11. Vertical guides mark three manufacturer-published developer pH values: PERCEPTOL stock at 7.75, ID-11 stock at 8.65 and MULTIGRADE paper developer at 1+9 at 10.5. The curve places the film developers on the steep lower part of the rise and the paper developer high on it, which is the point: a small pH error matters more to a film developer than to a paper one because the film developers sit where the curve is steepest.

Series	pH of the developer	Relative rate of development, arbitrary units
Relative rate, drawn to show the shape	7.00	0.01
Relative rate, drawn to show the shape	7.50	0.03
Relative rate, drawn to show the shape	8.00	0.06
Relative rate, drawn to show the shape	8.50	0.12
Relative rate, drawn to show the shape	9.00	0.25
Relative rate, drawn to show the shape	9.50	0.45
Relative rate, drawn to show the shape	10.00	0.68
Relative rate, drawn to show the shape	10.50	0.85
Relative rate, drawn to show the shape	11.00	0.94
Relative rate, drawn to show the shape	11.50	0.98
Relative rate, drawn to show the shape	12.00	1.00

The three vertical guides are manufacturer-published pH values and are real. The curve is not: no measured rate-against-pH curve is cited anywhere in this course's sources, and this one shows only the direction and the fact that the response is steep in the middle of the range and flattens at the top. The curve is drawn to show the shape, not measured from a real material. Your own materials will differ, and measuring them is what the sensitometry part of the course is for.

The stop bath. Its job is to halt development immediately by removing the alkalinity the developing agent needs. It works in seconds because it changes the pH of the swollen gelatin layer, not because it destroys anything.

Fixer stability and hardening. A fixer is held mildly acid — ILFORD publishes 5.0 to 5.5 for RAPID FIXER at 1+4 — for two reasons that pull in opposite directions. It must be acid enough to neutralise alkaline developer carried over on wet film, and not so acid that thiosulfate decomposes and drops sulfur, which is exactly the reaction Kodak’s 1928 primer describes when a mineral acid meets hypo. Hardening is affected by the same carryover: the 1928 primer records that a chrome alum fixing bath keeps its hardening properties well until developer is carried over on the plates and films, when they “fall off owing to the presence of sodium sulphite in the developer”. ILFORD’s own sheet gives the correction for the pH, which is the part you can measure: if a stop bath is not used and the fixer reads too high, a few drops of 50 per cent acetic acid, added gradually and with thorough stirring, will bring it back — but not past the published limits.

Toning and emulsion making both depend on pH, and this course does not put numbers on either here. Part XX owns toning chemistry and Part V owns emulsion making, and each will give its figures with its sources.

Your tap water. Every bath you mix starts as tap water, and tap water is not neutral, not consistent between regions and not constant through the year. Thames Water, for example, publishes hardness bands in milligrams per litre as calcium carbonate — soft up to 100, hard 200 to 300, very hard above 300 — and states that all the water in its region is hard, because it passes through chalk. It does not publish a pH on that page. Find your own supplier’s water-quality report for your postcode, and then measure it yourself: that is the first sample in the pH laboratory session, and the difference between what the report says and what your tap gives is itself worth knowing.

What pH does not tell you

This is the most misused number in photography, so be precise about its limits. Concentration, capacity and activity are three different properties, and pH measures none of them directly.

It does not tell you concentration. A 0.33 mol/L solution of acetic acid reads pH 2.6; so does a much more dilute solution of a stronger acid. Two solutions at the same pH may contain wildly different amounts of acid.

It does not tell you capacity. This is the important one, and the buffers page is built on it. A stop bath that has neutralised most of its acid can still read close to its nominal pH right up to the point where it fails, because the pH is set by the ratio of acid to conjugate base and the capacity by the total. A pH reading of a working bath is a snapshot of the ratio, not a fuel gauge.

It does not tell you activity, and a higher-pH developer is not automatically a faster one. Rate depends on which developing agents are present and in what concentration, on the bromide already released, on the sulfite, on temperature, and on how fast fresh solution reaches the grain — every one of which is a later page in this part. Kodak Ltd’s 1949 handbook makes the point from the other side: increasing or decreasing the alkali lets you either change the contrast reached in a given time or change the time needed to reach a given contrast. Which of the two you get depends on what else you hold fixed. pH is one input among several.

Reading a pH statement in a formula or on a sheet

Three habits.

A published range is the tolerance. ILFORD gives ID-11 stock as 8.60 to 8.70 and MICROPHEN stock as 8.67 to 8.93 — a band of 0.10 for one and 0.26 for the other. Those widths are information: they say how tightly the manufacturer expects a correctly mixed fresh batch to fall, and a reading outside the band is a reason to look for a mixing error. A single number with no band, by contrast, tells you nothing about tolerance and should not be treated as though it did.

A published figure is a laboratory figure. ILFORD says so on every sheet: the values “were obtained under carefully controlled laboratory conditions and may differ slightly from measurements made by users in their own working areas”, and it recommends that users take their own control measurements from accurately mixed fresh solutions for later comparison. The published number is a target; your number, measured on the day you mixed it, is the reference you will actually use.

When no pH has ever been published, say so and make your own. Most historical formulas have no published pH at all, and inventing one is exactly what Rule 1 of this course forbids. What you can do is measure the pH of a correctly mixed fresh batch, record it with the temperature and the meter’s calibration, and treat that as your own reference for that formula. It is not the manufacturer’s figure, it is not a literature value, and your notebook should say which it is.

Recap

An acid donates protons and a base accepts them, which is what lets carbonate and borax be alkalis without containing any hydroxide; each acid and its deprotonated form are a conjugate pair. Water ionises slightly and its two ions multiply to a fixed constant, so there is only one dial: raising one lowers the other. pH is minus the base-ten logarithm of the hydrogen ion concentration, so one unit is a factor of ten and the scale runs backwards, and neutrality is 7.00 only at 25 °C. Strength — the fraction ionised — and capacity — the total acid present — are different properties, which is why a fixer uses a weak acid in quantity. The baths of this course run from about pH 2 to about pH 10.6, each value published by its manufacturer as a band rather than a point. And pH tells you neither concentration, nor capacity, nor activity: it tells you a ratio.

Next: how a weak acid and its conjugate base together hold that ratio steady, how much abuse the pair can take, and what it means when the capacity runs out.

Check your understanding

Question 1. Convert pH 10.5 to a hydrogen ion concentration, and say how it compares with pH 9.5.

• 3.2 × 10⁻¹¹ mol/L, and pH 10.5 has ten times less hydrogen ion than pH 9.5
• 10.5 mol/L, and the two differ by one unit of concentration
• 3.2 × 10⁻¹¹ mol/L, and pH 10.5 has ten times more hydrogen ion than pH 9.5
• 1.05 × 10⁻¹⁰ mol/L, and the two are almost the same

Show the answer and why

Answer: 3.2 × 10⁻¹¹ mol/L, and pH 10.5 has ten times less hydrogen ion than pH 9.5

[H₃O⁺] = 10⁻pH = 10⁻¹⁰·⁵ = 3.2 × 10⁻¹¹ mol/L. Because the scale carries a minus sign, higher pH means less hydrogen ion, and one unit is a factor of ten in both directions: pH 9.5 holds 3.2 × 10⁻¹⁰ mol/L, ten times more. The hydroxide concentrations run the other way, so pH 10.5 is ten times the more alkaline of the two.

Question 2. Why do a 2 per cent w/v acetic acid solution and a 2 per cent w/v hydrochloric acid solution have such different pH values?

• Because hydrochloric acid is more concentrated by mass
• Because acetic acid is a weak acid and only a small fraction of its molecules give up a proton, while hydrochloric acid is strong and ionises essentially completely: about 0.7 per cent against 100 per cent, giving pH 2.6 against pH 0.26
• Because acetic acid is organic and organic compounds do not ionise
• Because the two have different densities

Show the answer and why

Answer: Because acetic acid is a weak acid and only a small fraction of its molecules give up a proton, while hydrochloric acid is strong and ionises essentially completely: about 0.7 per cent against 100 per cent, giving pH 2.6 against pH 0.26

Percentage is a statement about how much acid is in the bottle; pH is a statement about how much of it has let go of its proton. For 0.33 mol/L acetic acid, [H₃O⁺] is about √(1.8 × 10⁻⁵ × 0.333) = 2.4 × 10⁻³ mol/L, which is 0.7 per cent of the acid present. The 2.35 pH units between the two solutions is a factor of 224 in hydrogen ion concentration. At equal molarity they would neutralise the same amount of alkali, which is the separate property of capacity.

Question 3. A tray of carbonate-based print developer stands open all evening. What happens to its pH, by what mechanism, and what would the prints show?

• The pH rises as water evaporates and concentrates the alkali; prints would be contrastier
• The pH falls, because carbon dioxide from the air dissolves and forms carbonic acid, which converts carbonate to hydrogencarbonate; prints would be slower to come up and would struggle to reach a full black
• Nothing changes, because a carbonate developer is buffered and buffers do not change
• The pH falls because the developing agents evaporate

Show the answer and why

Answer: The pH falls, because carbon dioxide from the air dissolves and forms carbonic acid, which converts carbonate to hydrogencarbonate; prints would be slower to come up and would struggle to reach a full black

Two things happen to an open tray and they are worth keeping apart. Atmospheric carbon dioxide dissolves and acts as an acid, consuming carbonate and pushing the pH down; and oxygen from the air attacks the developing agents directly, which is the aerial oxidation page later in this part. Both slow the developer. A buffer resists the pH change but has a finite capacity to do so, which the next page quantifies. This course has not verified a published figure for how far an open tray drifts in an evening, so treat the direction as established and the size as something to measure yourself.

Question 4. Which of these does a pH reading of a working bath actually tell you?

• How much acid or alkali is left in it
• How much longer it will work
• The ratio of acid to conjugate base at the moment of measurement, and nothing directly about the total quantity of either
• The concentration of the developing agents

Show the answer and why

Answer: The ratio of acid to conjugate base at the moment of measurement, and nothing directly about the total quantity of either

A buffered solution holds a pH set by the ratio of the pair, while its capacity depends on the total concentration of the pair. Two baths at the same pH can therefore have very different amounts of work left in them, and a bath can read its nominal pH shortly before it fails. This is exactly why capacity and exhaustion are separate terms in this course, and why the stop bath question on the next page is worth working through.

Question 5. Why is concentrated acid added to water, and never water to concentrated acid?

• To avoid diluting the acid too much
• Because neutralisation and dilution both release heat, and adding a little acid to a large volume of water spreads that heat through the water, while adding water to concentrated acid releases it in a small volume at the surface, where the mixture can boil and spit acid out
• Because acid is denser than water and would sink
• It is a convention with no physical basis

Show the answer and why

Answer: Because neutralisation and dilution both release heat, and adding a little acid to a large volume of water spreads that heat through the water, while adding water to concentrated acid releases it in a small volume at the surface, where the mixture can boil and spit acid out

The heat is real: CAMEO records that dissolving sodium hydroxide can liberate enough to make a solution steam and spatter, and concentrated acids behave comparably. Princeton’s guidance for photographic work states the rule as: always add acid to the water when diluting. The same argument covers solid alkalis, which is why they go into water too and never the reverse.

Question 6. ILFORD publishes ID-11 stock at pH 8.60 to 8.70 and MICROPHEN stock at 8.67 to 8.93. What does the difference in the widths of those two bands tell you?

• That MICROPHEN is less accurately made
• That the bands are the manufacturer’s stated tolerance for a correctly mixed fresh solution, so a MICROPHEN batch has more room before a reading counts as suspect than an ID-11 batch does
• That MICROPHEN is a stronger developer
• Nothing; the bands are rounding

Show the answer and why

Answer: That the bands are the manufacturer’s stated tolerance for a correctly mixed fresh solution, so a MICROPHEN batch has more room before a reading counts as suspect than an ID-11 batch does

A published band is information about tolerance, not decoration. ID-11 is expected within 0.10 of a unit and MICROPHEN within 0.26, so the same 0.15-unit deviation means something different for each. ILFORD also states on the same sheets that its figures came from controlled laboratory conditions and that users should measure their own fresh solutions for comparison, which is why your own dated reading of a fresh batch is the number you should actually work against.

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