Experiment: Precipitating the Silver Halides

Ninety minutes at a bench and half a gram of silver nitrate, and you will have made by hand the three compounds every photograph in this course depends on — and destroyed two of them on purpose, one with light and one with hypo.

Purpose

To make silver chloride, silver bromide and silver iodide by double decomposition, watch each one darken in light against an identical tube that was kept in the dark, dissolve them again in a fixer, and see what changes when the same precipitation is done in gelatin instead of water.

Hypothesis. Three things, each testable in a tube.

1. Mixing equimolar solutions of silver nitrate and a soluble halide gives an insoluble silver halide whose colour runs white, pale yellow, light yellow from chloride to iodide.
2. All three darken in light and none darkens in the dark, so the darkening is caused by light and not by time, air or the glass.
3. All three dissolve in sodium thiosulfate, and the order and speed of dissolution follow their solubility products, chloride fastest and iodide slowest.

The control. Every precipitate is made in duplicate, and one of each pair is wrapped in kitchen foil the moment it forms and stays wrapped until the end. The control tubes go everywhere the exposed tubes go and stand for the same time in the same place. Without them “it went grey” is an observation about a tube; with them it is an observation about light.

The one variable that changes in the main run is light: how much of it a tube receives. Everything else — concentration, volume, temperature, glass, time, the order of addition — is held the same across each pair. Two later stages deliberately change one further variable each: the halide, across the three tubes, and the medium, water against 2 % w/v gelatin.

Learning objectives

By the end you will be able to:

• Prepare a solution of stated molarity from a solid, using the arithmetic of concentration and dilution, and say how accurate your figure really is.
• Write the balanced molecular and net ionic equation for each precipitation you perform.
• Design and use a shielded control, and say what claim it does and does not license.
• Rank the three halides by rate of photolysis and by ease of dissolution, from your own data, and connect each ranking to a published constant.
• Describe what a protective colloid does to a precipitate, having seen it.
• Handle silver nitrate at Level B and account for every drop of silver you used.

Prerequisites

The Silver Halides, which supplies the equations and the solubility products; and from Part II, laboratory safety and PPE, concentration and dilution, silver nitrate handling, chemical storage and incompatibilities and chemical waste and silver waste. The dilution arithmetic is used here and not re-taught.

Safety classification

Level B. The criteria of the rubric that decided it:

• Substances. Silver nitrate is named in the Level B criteria in so many words. Its harmonised European classification carries H272 (may intensify fire, oxidiser), H314 (causes severe skin burns and eye damage) and H400 with H410 (very toxic to aquatic life, with long-lasting effects). A third of notifiers add H318, serious eye damage.
• Procedures whose failure mode is a splash of something corrosive. Pouring one solution into another in a narrow tube, and inverting tubes to mix, are exactly that.
• Optional step. If you take the ammonia route, you handle an irritant vapour and you make, for a few seconds, a solution class that Part II prohibits storing.

It is not Level C. Nothing here is acutely toxic, carcinogenic or a chromium(VI) compound; nothing needs a fume cupboard; nothing is heated above 50 °C except optionally warm water for the gelatin; and the waste is collectable silver-bearing liquid, which is a stream a domestic route already exists for.

What is not a hazard here, and why. The three silver halides you make are almost completely insoluble — 1.93, 0.135 and 0.0028 milligrams per litre — so the precipitate is not a solution hazard: it cannot get into you through skin contact with the liquid, because it is barely in the liquid at all. There is no dust hazard either, because at no point does anything dry out; every solid you make stays wet, in a tube, from the moment it forms until it goes into the waste container. That is a statement about this procedure, not about these compounds: dried and powdered, a silver halide is a fine solid carrying silver that must not be inhaled, and PubChem’s aggregation for silver bromide carries a reproductive-toxicity statement from a minority of notifiers. Keeping it wet is the control, and it is free.

Hazards

Silver nitrate solution, eyes. The irreversible one. Splash goggles that seal against the face, not safety glasses with a gap above the cheek, for every minute of this experiment. Eyewash within reach before the bottle is opened; a personal eyewash bottle is a supplement to a plumbed or portable unit and not a replacement for one.

Silver nitrate solution, skin. Rinse immediately and at length if you are splashed. The brown-black mark that appears later is metallic silver bound to skin protein and will not wash off; the rinse is about limiting absorption, not about the colour, and by the time the colour appears the rinse is too late. Change clothing that gets wet with it.

Silver nitrate as an oxidiser. Keep it away from paper towel you have already used for something else, from alcohols and from every reducing agent on the bench. Spilled solution dried into cloth or card leaves an oxidiser in a combustible material.

Ammonia solution, if you take the optional step. Irritant vapour. Work at an open window with a fan running, keep the bottle capped between pours, and do not put your face over the tube.

☠Ammoniacal silver: why the quench is part of the step and not a tidy-up

Adding ammonia to silver nitrate throws down brown silver oxide, which redissolves in excess ammonia to give the colourless diammine complex — a solution the nineteenth century used constantly. Ware’s cautionary note is that such solutions can, with time, deposit a black precipitate of silver nitride, the “fulminating silver” of the old manuals, which detonates at a touch and sometimes even when wet, and that storing them is not recommended. NIOSH lists ammonia second among the incompatibilities of silver.

The variable in that sentence is time, and drying is how a small volume spends time getting concentrated. So this course’s handling rule, set in Part II, has no exceptions: no ammoniacal silver solution is stored, and none is allowed to dry.

On this page that becomes a procedure, not advice. Every ammoniacal tube is emptied into the quench beaker the moment its reading is taken. The quench beaker holds 50 mL of saturated sodium chloride solution, which does two things at once: chloride at that concentration drives silver back out as silver chloride, and dilution collapses the ammonia concentration the complex needs. The tube is then rinsed twice into the same beaker. Nothing ammoniacal stands on the bench, goes back on a rack, or is left to evaporate — not for the rest of the session, not for five minutes.

This course found no laboratory safety document in its corpus giving a maximum standing time for small ammoniacal silver volumes. It therefore sets the only limit it can defend, which is zero. If you are not willing to run the quench as written, omit the ammonia stage; nothing later in the course depends on it.

Ultraviolet torch, if you use one. A 365 nm source is an eye and skin hazard. Never look into the beam, never point it at anyone, and let it illuminate the rack from the side rather than from where you are standing. ICNIRP’s guidance for workers is written about outdoor sunlight but its personal protective principle transfers: keep the source out of the eye and the skin covered. Daylight alone is enough for this experiment, and a north-facing windowsill on a bright day will separate the three halides perfectly well; the torch shortens the wait and is not required.

Glass. Test tubes crack if they are heated unevenly or knocked. Use a rack, not your hand.

Required PPE

• Chemical splash goggles that seal against the face, worn from before the silver nitrate bottle is opened until after the clean-up.
• Nitrile gloves. HSE’s COSHH essentials sheet P1 takes single-use nitrile of about 0.2 mm as splash protection for manual film and plate development where the safety data sheet gives nothing more specific, and that is the basis used here. Be aware of the honest gap: the glove manufacturer’s own permeation guide this course has read contains no entry for silver nitrate at all, and states that its figures are advisory and that suitability must be determined by testing by the purchaser. So the gloves are splash protection, changed as soon as they are contaminated, and not a permeation barrier you can rely on for immersion.
• An apron or overall, and long sleeves.
• A tray under the whole working area. Silver nitrate stains worktops, grout and sink enamel; the tray is what gets stained.

Ventilation

Nothing in the main procedure produces a vapour, a mist or a dust, so extraction is not among the controls for it: the solutions are cold, dilute and aqueous, and nothing is powdered after the initial weighing. Two exceptions. Weigh the solids in still air, slowly, so nothing becomes airborne. And if you take the optional ammonia step, ventilation becomes a control: work at an open window with a fan drawing air past you and out, keep the bottle capped between pours, and do not do that stage in a sealed room.

Materials

Item	Quantity	Notes
Test tubes, about 15 mL	14	Twelve for the main run, two for the gelatin comparison.
Test tube rack	2	One that can be carried, for the exposure station.
Volumetric flask or graduated bottle, 100 mL	4	One per stock solution. A measuring cylinder will do; your uncertainty will be larger and you will record that.
Graduated pipettes or syringes, 5 mL	4	One per solution, never shared.
Kitchen aluminium foil	a sheet	For the shielded controls.
Beaker, 250 mL	1	The quench beaker, only if you take the ammonia step.
Beaker, 1 L, labelled SILVER WASTE	1	Standing in the tray, open, for the whole session.
Timer	1	A phone is fine.
White card and a pencil	—	Background for judging colour; pencil, because ink runs.
Ultraviolet torch, 365 nm	optional	Daylight works. See the hazard note.

Chemicals

Chemical	Quantity	Form
Silver nitrate	1.70 g	Solid, to make 100 mL of 0.100 mol/L
Sodium chloride	0.58 g	Solid, to make 100 mL of 0.100 mol/L
Potassium bromide	1.19 g	Solid, to make 100 mL of 0.100 mol/L
Potassium iodide	1.66 g	Solid, to make 100 mL of 0.100 mol/L
Sodium thiosulfate pentahydrate	5 g	Solid, to make 50 mL at 10 % w/v
Gelatin	1 g	Any plain culinary gelatin, to make 50 mL at 2 % w/v
Ammonia solution	20 mL	Household strength. Optional stage only, and read the quench note first

🧮Where those masses come from, and why 0.100 mol/L

The arithmetic is Part II’s and is used, not repeated. For 100 mL of 0.100 mol/L you need 0.0100 mol, so the mass is 0.0100 multiplied by the molar mass:

m = 0.0100 mol × M

Mass for 100 mL of 0.100 mol/L

Silver nitrate, M = 169.873, gives 1.699 g. Sodium chloride, M = 58.44, gives 0.584 g. Potassium bromide, M = 119, gives 1.19 g. Potassium iodide, M = 166.0028, gives 1.660 g.

Why 0.100 mol/L? Two reasons, both worth stating. It is concentrated enough to see: 2.0 mL of each solution gives 0.20 mmol of silver halide, which is 38 mg of silver bromide — a solid clearly visible against white card. And it is dilute enough to be cheap in silver: the whole experiment, twelve tubes and the gelatin pair, consumes under 30 mL of the silver nitrate stock, which is about half a gram of the salt and about a third of a gram of silver. The laboratory planner carries dated prices; this page carries the quantity.

The thiosulfate at 10 % w/v is 0.40 mol/L of the pentahydrate. Each tube holds at most 0.20 mmol of silver halide, and Part III’s coupled-equilibrium arithmetic says dissolving x moles of silver takes roughly 2x moles of thiosulfate locked into complexes plus a small free excess. Two millilitres of 0.40 mol/L is 0.80 mmol of thiosulfate, against the 0.40 mmol the stoichiometry demands: a twofold excess. That excess is deliberate rather than generous. Part III’s arithmetic puts the point at which a bath stalls at about 0.45 moles of silver dissolved for every mole of thiosulfate it started with, which here is 0.45 × 0.80 = 0.36 mmol of silver — against the 0.20 mmol actually in the tube. So each tube loads its fixer to a little over half of capacity, comfortably short of the point where clearing times start to lengthen, and a slow tube is therefore telling you about the halide and not about a starved fixer.

Equipment

A balance reading to 0.01 g; a spatula; a wash bottle of distilled or deionised water; labels and a pen that survives water; the tray; and somewhere to put the exposure rack in bright daylight — a windowsill, a garden table, or a bench under a lamp, as long as both racks can be there together.

Estimated cost

Cost band ££. The recurring cost is the silver nitrate, of which this experiment uses under a gram; everything else is a few pence of common salts and a packet of gelatin. Test tubes, a rack and graduated syringes are one-off purchases that the rest of the course reuses. Dated prices live in the laboratory planner, not in this text, so they can be kept current.

Estimated consumables cost

The silver is the whole of the recurring cost and the experiment was designed at 0.100 mol/L partly to keep it small: twelve tubes and the gelatin pair spend under a gram of silver nitrate between them. Test tubes, the rack, the syringes and the flasks are one-off purchases the rest of the course reuses.

Consumed	This session	Sourced price	Cost this session
Silver nitrate	1.70 g weighed; about 0.5 g of it actually spent	£59.95–£112.90 per one jar: 25 g at the lower figure, 10 g at the higher (£2.40–£11.29 a g)	£4.08–£19.19
Potassium bromide	1.19 g	£23.00 per 250 g (£0.09 a g)	£0.11
Sodium chloride	0.58 g	None. sodium-chloride carries a cost band and no dated figure	—
Potassium iodide	1.66 g	None. A named price gap: silver-halide salts other than potassium bromide	—
Sodium thiosulfate pentahydrate	5 g, as 50 mL at 10 % w/v	£14.70 per 1 kg of the raw salt, checked 7 September 2026	£0.07
Gelatin	1 g of plain culinary gelatin	None. culinary-gelatin carries a cost band and no dated figure	—
Ammonia solution, household	20 mL, optional stage only	None. household-ammonia carries a cost band and no dated figure	—
Distilled water	about 500 mL	None. distilled-water carries a cost band and no dated figure	—

The priced rows come to £4.26 to £19.37 for one run of this session, at the retail ranges read on 5 and 7 September 2026 and recorded in the laboratory planner. That is a floor, not a total: 5 of the 8 rows carry no dated price, so they are counted as nothing here and are certainly not free. A priced entry is a dated range to plan against, never a quotation.

The silver row is priced on the mass weighed, not the mass consumed, because a jar is bought whole; the third of a gram of silver metal that ends up in the waste bottle is what the experiment costs in substance. The salt that still dominates the unpriced list is the iodide, which is a named price gap; the raw hypo beside it has a dated price as of 7 September 2026, and both are bought once for the whole of Parts IV, V and XI.

Waste streams

One stream, and it is the important one. Every drop of liquid produced on this page is silver-bearing, including the tubes that look clear: a saturated solution above silver chloride still carries 1.93 mg of silver per litre, and the rinses carry more. Everything — precipitates, supernatants, the thiosulfate tubes, the quench beaker, the rinse water and the paper you wipe with — goes into the labelled silver waste container.

Preparation

1. Label everything before you weigh anything. Four stock bottles, twelve tubes and the waste container. Tubes are labelled with the halide and with E for exposed or C for control.
2. Set up the tray with the waste beaker standing in it and the wash bottle beside it.
3. Make the four stock solutions. Weigh, dissolve in about 60 mL of distilled water, make up to 100 mL, invert to mix, label with the contents, the concentration, the date and your initials. Record the mass you actually weighed, not the mass you meant to.
4. Make the thiosulfate at 10 % w/v: 5 g in water to 50 mL.
5. Make the gelatin. Sprinkle 1 g of gelatin over about 30 mL of cold water, leave it five minutes to swell, then warm it to about 40 °C — a jug of hot tap water is enough — and stir until clear. Make up to 50 mL. Keep it warm; it will set if you let it cool.
6. Silver nitrate last, and into an amber or foil-wrapped bottle, because a solution decomposes in light faster than the solid does.

Bench layout for the session

1. 1Weighing and stock area — balance, four labelled bottles; silver nitrate foil-wrapped
2. 2The tray — everything wet happens inside it
3. 3Working rack, twelve tubes — row E exposed, row C wrapped in foil
4. 4Quench beaker — saturated sodium chloride, 50 mL — only for the ammonia stage
5. 5Silver waste container — open, labelled, standing in the tray from the first pour
6. 6Exposure station — both racks together, white card behind, timer running

Plan view, drawn rather than measured. The one arrangement that matters is number 5: an open, labelled waste container inside the tray, because a container you have to unscrew is a container you will not use.

Procedure

Gloves and goggles on before step 1, and stay on.

Stage 1: the three precipitations, in duplicate

1. Into six tubes, pipette 2.0 mL of a halide solution: two tubes of chloride, two of bromide, two of iodide. The halide goes in first, so that silver is always the ion added to an excess.
2. To each, add 2.0 mL of silver nitrate, a few drops at a time down the side of the tube, as Kodak’s primer directs for an emulsion, so the solid forms evenly rather than in a lump. Watch the first drop land: that instant is the reaction.
3. Stopper or cover and invert twice. Do not shake.
4. Record colour and texture immediately, against white card, for all six tubes.
5. Wrap one of each pair completely in foil — sides, base and top — and mark it C. From now on those three tubes are never unwrapped except to read them.
6. Repeat steps 1 to 4 for a second set of six, which will be used for the dissolution stages. Wrap all six of these in foil too and keep them out of the light; nothing about the dissolution experiment is supposed to depend on how much light a tube has had.

Stage 2: the timed exposure

7. Carry the first rack — three exposed tubes, three wrapped controls — to the exposure station. Put the white card behind them. Start the timer.
8. Read every tube at 1, 2, 5, 10, 20, 40 and 60 minutes. At each reading, unwrap a control just long enough to look at it against the same card, then re-wrap it. Record the time of first visible change for each exposed tube, and the colour at each reading.
9. If you are using a 365 nm torch instead of daylight, hold it to one side of the rack, never pointing towards you or anyone else, and use a shorter series: 15 and 30 seconds, then 1, 2, 5 and 10 minutes.

Stage 3: dissolution in thiosulfate

10. Take the three foil-wrapped tubes from the second set. Let each stand until the solid has settled, then pour off as much clear liquid as you can into the silver waste container, leaving the solid.
11. Add 2.0 mL of the 10 % w/v thiosulfate to each, start the timer, invert twice and watch.
12. Record the time for each tube to go completely clear, and stop at 15 minutes whether it has or not. “Did not clear in 15 minutes” is a result and should be written as one.

Stage 4 (optional): dissolution in ammonia

Take this stage only if you have read the quench note above and are prepared to run it exactly.

13. Have the quench beaker — 50 mL of saturated sodium chloride solution — standing beside the rack before you open the ammonia.
14. To each of the three remaining foil-wrapped tubes, add 2.0 mL of household ammonia solution. Invert twice. Watch for 2 minutes and record what happens in each.
15. Immediately after the reading, pour each tube into the quench beaker and rinse it twice into the same beaker with water from the wash bottle. Do this tube by tube: do not read all three and then quench all three.
16. When all three are quenched, pour the quench beaker into the silver waste container and rinse it twice into the same container.

Stage 5: water against gelatin

17. In tube 13, put 2.0 mL of potassium bromide solution. In tube 14, put 2.0 mL of potassium bromide solution mixed with 2.0 mL of the warm 2 % w/v gelatin.
18. Add 2.0 mL of silver nitrate to each, the same way, at the same rate.
19. Invert twice. Record the appearance of each immediately, and again after standing undisturbed for 5, 15 and 30 minutes. What you are watching for is not colour but whether the solid stays suspended or falls to the bottom.

Expected observations

On colour. Silver chloride is described by its sources as a white solid, silver bromide as yellowish or light yellow, and silver iodide as light yellow. You should be able to arrange the three tubes in that order without hesitating.

On texture. Kodak’s primer predicts the bromide: a thick, curdy precipitate if the solutions are at all concentrated. Expect the chloride to be curdy too, and the iodide to be denser and to settle faster.

On darkening. All three should change under light and none should change under foil. The exposed chloride and bromide are usually described as going grey, then violet-grey, then darker; the iodide darkens more slowly.

=What this page will not tell you

It gives you no expected times. The course has not run this experiment itself and has found no source in its corpus that states how long each halide takes to darken visibly under a stated light source. Anything it printed here would be a guess with a decimal point on it.

What it will do is give you the prediction to test, and the reasoning behind it, so that your data has something to argue with. Argument one, from the absorption edges: the halide that absorbs furthest into the visible has the most of the available light to work with, which puts the yellow halides ahead of the white one. Argument two, from Ware’s halogen-acceptor chemistry: photolysis stalls unless something takes up the liberated halogen, and there is very little in a test tube of water to do that, so all three should slow down and stop well short of black. Those two arguments predict different orderings in the first minute and the same ceiling after an hour. Which one your tubes follow is the experiment.

On dissolution. Expect chloride to clear fastest and iodide slowest, by a wide margin. Whether the iodide clears at all in 15 minutes at this thiosulfate strength is a genuinely open question and one of the more interesting numbers you will take away.

On ammonia, if you run it. Ammonia dissolves silver chloride readily; the formation constant of the diammine ion is 1.7 × 10⁷ against a chloride solubility product of 1.6 × 10⁻¹⁰. Against bromide and especially iodide, whose solubility products are far smaller, expect much less.

On gelatin. The water tube should give a curdy solid that settles to the bottom, leaving clear liquid above. The gelatin tube should give a milky, even suspension that stays suspended.

What is happening chemically

The precipitations, in molecular and net ionic form:

AgNO₃ + NaCl → AgCl + NaNO₃

Silver chloride

AgNO₃ + KBr → AgBr + KNO₃

Silver bromide

AgNO₃ + KI → AgI + KNO₃

Silver iodide

Ag⁺ + X⁻ → AgX

All three, with the spectator ions removed; X is Cl, Br or I

The driving force is the solubility product. Mixing 0.100 mol/L solutions makes an ion product of the order of 10⁻³, against solubility products of 10⁻¹⁰ to 10⁻¹⁶, so the system is supersaturated by six to twelve orders of magnitude and solid appears at once.

The photolysis, in its bookkeeping form:

2 AgX + light → 2 Ag + X₂

Net photolysis, where X is Cl, Br or I

That equation hides the interesting part, which is that the reaction runs backwards unless the halogen is removed. The liberated halogen is an oxidiser sitting next to freshly made silver. In your tube there is no halogen acceptor to speak of, which is why the darkening slows and stops instead of running to black — Ware puts the ceiling for a pure crystal at an optical density of about 0.02 and the particle size at about 10 nm. A photographic emulsion works because gelatin, and in printing-out papers free silver ion and water, take the halogen away.

The dissolution in thiosulfate:

AgX + 2 S₂O₃²⁻ → [Ag(S₂O₃)₂]³⁻ + X⁻

Fixing, in a test tube

Silver ion is removed into a complex as fast as it dissolves, so the ion product never reaches the solubility product and the solid keeps going. The overall constant is the solubility product multiplied by the formation constant, which is why the halide with the smallest Ksp is the slowest to clear.

The dissolution in ammonia:

AgCl + 2 NH₃ → [Ag(NH₃)₂]⁺ + Cl⁻

The diammine complex: the same idea, a weaker ligand

And the quench, which is the same reaction driven backwards by mass action:

[Ag(NH₃)₂]⁺ + Cl⁻ → AgCl + 2 NH₃

Saturated chloride takes the silver back out of solution

The gelatin comparison. Nothing chemical differs: the same equation, the same product. What differs is that the gelatin adsorbs onto the crystals as they form, keeping them small and keeping them apart, so they cannot aggregate into curds heavy enough to fall. That is peptisation, and gelatin doing it is why Kodak’s own emulsion patent calls gelatin the peptiser and the gelatin added afterwards the protective colloid. The same patent’s example precipitates at 70 °C into rapidly stirred gelatin and gets crystals of about 0.2 micrometres. You have just done, badly and by hand, the first step of making an emulsion.

Data to record

For every tube: the halide, whether it was exposed or a control, the exact volumes used and the concentrations you actually made (from the masses you actually weighed), the time of mixing, and the temperature of the room.

Column	What goes in it
Tube	e.g. Br-E, Br-C
Colour and texture at t = 0	in words, against white card
Time of first visible change	in seconds or minutes, or “none”
Colour at each reading	1, 2, 5, 10, 20, 40, 60 min
Final colour at 60 min	exposed and control, side by side
Time to clear in thiosulfate	or “not cleared at 15 min”
Behaviour in ammonia	dissolved / partly / not, and how fast
Gelatin tube at 0, 5, 15, 30 min	settled or suspended, and how far

Also record, because they are the things you will wish you had: the weather and the light source, the distance from the window, whether the sun was direct, and any tube you disturbed.

Analysis

1. Write the balanced equation for every change you saw — nine at least: three precipitations, three photolyses, three dissolutions, plus the ammonia reactions and the quench if you ran them. This is the deliverable that matters most.
2. Rank the halides by rate of darkening and say whether the ranking matches the absorption-edge prediction, the halogen-acceptor prediction, both or neither.
3. Rank them by time to clear in thiosulfate and plot that ranking against the logarithm of their solubility products. Three points is not a curve, but the direction is a result.
4. State what the controls license. Write the sentence: “the exposed tube changed and the control did not, therefore …” — and then write what you can not conclude from it.
5. Estimate your concentration honestly. From the mass you weighed and the volume you made up, and the uncertainty of your balance and glassware, give your silver nitrate concentration to the number of significant figures your instruments support.

Troubleshooting

The control tube went grey too. Three causes, and the foil distinguishes them. If the foil was torn or the top open, it is light and the control was never a control — check it. If the foil was intact, suspect contamination: a pipette used for two solutions, or a trace of a reducing agent on the glass. And if both tubes of a pair went grey equally, suspect the stock: silver nitrate solution standing in a clear bottle reduces on its own, which is why the bottle is foil-wrapped.

No visible precipitate. Check the order of addition and that both solutions really are what the label says. A tube that stays clear when silver nitrate meets a halide is almost always two aliquots of the same solution.

Everything went brown at once. Silver nitrate meeting organic matter. A dirty tube, a wooden spatula, a fingerprint inside the glass.

The gelatin tube set solid. It was too cool, or the gelatin too strong. Warm it and repeat; the comparison needs both tubes fluid.

The iodide will not clear in hypo. That is very probably the result rather than a fault. Record the time as greater than 15 minutes and say why you expect it.

Clean-up

Everything wet goes into the silver waste container: tubes, supernatants, quench beaker, rinses, and the paper towel you wipe with. Rinse the glassware before anything dries, twice into the waste container and then under the tap, because a dried silver residue is far harder to remove and is the one way this procedure can produce a dust. Wipe the tray last. Gloves off, then goggles, then wash your hands.

Storage

The four stock solutions keep in labelled bottles, the silver nitrate in amber glass or foil in a dark cupboard, away from ammonia and from every reducing agent. The thiosulfate keeps for weeks. The gelatin does not keep: it is a protein solution at room temperature and it will spoil; make it fresh. Nothing ammoniacal is stored, ever.

Disposal considerations

Every liquid from this page is silver-bearing and none of it is poured away. The chemistry is straightforward: silver is recoverable, and Kodak’s own recovery literature describes both the electrolytic route and metallic replacement, in which iron reduces silver from solution and takes its place. Domestically, the practical route is the one ILFORD gives United Kingdom users — bottle wastes separately, label them, and take them to a household waste and recycling centre.

The reason this matters is on the label of the bottle you started from: silver nitrate’s harmonised classification carries H400 and H410, very toxic to aquatic life with long-lasting effects. Check your local regulations; they govern, they differ between jurisdictions, and they change.

Questions

1. You mixed 2.0 mL of 0.100 mol/L silver nitrate with 2.0 mL of 0.100 mol/L potassium bromide. How many moles of silver bromide could form, what mass is that, and how much of it would dissolve in the 4 mL of liquid left in the tube?
2. Why does the halide go into the tube first and the silver nitrate second? What would be different about the crystals if you reversed it?
3. Your chloride tube cleared in thiosulfate in 40 seconds and your iodide had not cleared at 15 minutes. Express that as a ratio, and compare it with the ratio of the two solubility products. Why are the two numbers so different?
4. The control tubes are wrapped in foil rather than put in a drawer. Give two reasons why that is a better control.
5. What single change to this procedure would make the precipitates finer, and what evidence from the gelatin tube supports your answer?

Further experiments

• Vary the order and rate of addition. Add the silver nitrate all at once to one tube and dropwise over two minutes to another, and compare the texture and the settling rate.
• Vary the excess. Make one bromide tube with the silver in excess and one with the bromide in excess, and see whether they darken at the same rate. Ware’s account says the two crystals carry different adsorbed ions and are chemically different objects; this is the cheapest way to test that.
• Look for Abney’s colours. He reported that white light transmitted through a gelatine film of silver bromide can look ruby, orange, green, purple or grey depending on the state of aggregation. Hold your gelatin tube up to a window against a bright sky and record what colour the light coming through it is, before and after exposure.
• Repeat one halide at a tenth of the concentration. Does it darken faster or slower, and is the question about the amount of silver or the size of the crystals?