Lab: Commissioning Your Laboratory

The first thing you will make in this course is not a photograph and not a developer. It is a set of numbers about your own instruments, written down and dated, and a bottle of salt water made twice — once with care and once by eye — so that the difference between the two is a figure rather than an opinion.

Purpose

To commission the laboratory: to lay it out, to find out how wrong each of your three measuring instruments is and to record it, to make and label a stock solution and a working dilution from it correctly, to measure the cost of not bothering, and to open the notebook that runs for the rest of the course.

Nothing hazardous is handled, and that is deliberate. This session is about discipline and measurement, and putting a corrosive or an oxidiser on the bench would divide your attention at the one moment when the whole point is technique. The substance is ordinary table salt.

Learning objectives

By the end of this session you should be able to:

• lay out a bench with a dry area and a wet area and run a session in one direction through it;
• check a balance against a known mass and against itself, and say which of those tests answers which question;
• check a thermometer at the ice point, state the deviation, and decide what to do about it;
• find the volume your own graduate really delivers, by weighing water;
• make a solution up to a stated final volume, calculate a dilution from it with C₁V₁ = C₂V₂, and check the arithmetic by an independent route;
• state the concentration of what you made to the number of significant figures your instruments support, and justify that number;
• complete every field of the course’s labelling scheme and the first three entries of the notebook.

Prerequisites

All five of the pages this part has already given you, and this lab assumes them rather than repeating them: laying out the laboratory, protecting yourself, storage and incompatibilities, measurement and uncertainty and concentration and dilution. Read the measurement page again the evening before; every check here is one it explains.

Safety classification

Level A. The criteria of the classification rubric that decided it, taken in the order the rubric states them:

• Substances and quantities. One substance is handled: sodium chloride, CAS 7647-14-5, as a dry solid and as a solution of about 10 per cent w/v, in quantities under 100 g. The ECHA Classification and Labelling Inventory aggregation on PubChem records that 89.2 per cent of 2,063 reports classify it as not meeting GHS hazard criteria, with 10.8 per cent notifying H319, causes serious eye irritation. That is comfortably inside the Level A criterion, which admits substances classified at most as irritant, harmful if swallowed, or corrosive at the concentrations actually handled.
• Energy. Water is brought to the boil for one optional check, which is the only heat on the page. Nothing else is heated, there is no flame beyond a domestic hob or kettle, and no solvent.
• Electrical. Nothing is built, wired or modified. The only mains equipment is a kettle or hob you already own and a purchased balance.
• Waste. Dilute salt solution and rinse water. Nothing requires specialist neutralisation, and nothing carries silver.

What is not a hazard here, and why

This section exists because a Level A page that says nothing about absent hazards reads as evasive, and because naming an absence teaches you how the assessment was made.

There is no corrosive, no oxidiser and no sensitiser on this bench. Sodium chloride carries no harmonised classification for skin corrosion, no oxidising classification, and no skin-sensitisation statement in the aggregated notifications. That is a statement about this substance at these concentrations, not about photographic chemistry: the sodium sulfite you will weigh in Part VIII carries GHS05 and GHS07 with H302, H314, H315 and H319 in the same aggregation, and the silver nitrate of Part IV carries H272, H314, H400 and H410 in its harmonised entry. The instruments are the same; the substances are not.

There is no vapour, dust or mist worth controlling. Sodium chloride is not volatile, and the quantities weighed are small and coarse-grained. This is why the Ventilation section below says what it says.

There is no silver and no thiosulfate, so the waste on this page is the only waste in the course that raises no question about aquatic toxicity or recovery. The waste page still asks you to route it deliberately, because the habit is the point.

What is left is what the Hazards section names: hot water, breakable glass, and a wet floor. Those are real, they are the only real ones, and they are the reason this page is Level A rather than unclassified.

Hazards

Scalding, from the boiling-point check. This is the one genuine injury risk on the page. Water at 100 °C is the hottest thing you will handle in Part II. Use a small pan or kettle, a stable surface, and put the thermometer in from the side rather than leaning over the vessel, so that your face is not over the steam. If you do scald yourself, NHS first-aid guidance for burns and scalds is to cool the burn under cool or lukewarm running water for 20 minutes as soon as possible, to remove clothing and jewellery near the area but not anything stuck to it, and not to burst blisters. If you would rather not do this check at all, do not do it: the ice point is the check that matters, and the Procedure says which parts are optional.

Glass breakage. A graduate and a thermometer are the two most breakable things you own, and a glass thermometer in a pan of boiling water is both hazards at once. Work over a tray, do not put a graduate down near an edge, and if a thermometer breaks in hot water, turn the heat off, leave the room until the steam clears, and pick up the glass with a dustpan and a wet paper towel rather than with your fingers.

Slips. Spilled water on a hard floor, in a session that involves carrying full vessels. Keep the tray under everything, keep a cloth to hand, and wipe as you go rather than at the end.

Ingestion of the wrong thing. The salt is edible in a kitchen and it is not food on this bench, because it is being weighed with equipment that will later touch photographic chemistry. The moment you open this session, everything on the bench is laboratory material. Nothing on the bench goes anywhere near your mouth.

Eye irritation, at the margin. About one supplier notification in nine classifies sodium chloride as H319, causes serious eye irritation, and a splash of a 10 per cent solution in the eye would sting whatever the classification says. Eye protection is worn.

Required PPE

Two items, and neither is here because the salt demands it.

• Eye protection: safety glasses with side protection, to a stated standard. The Level A controls require eye protection whenever solutions are handled, and the boiling-water check adds a splash of something at 100 °C to the list of things that could reach your face. Ordinary spectacles do not count; the protective equipment page explains why.
• Nitrile gloves. The Level A controls require them whenever solutions are handled. On this page their real work is habit: putting them on before the first cap comes off, and taking them off by the sequence on the protective equipment page, is a rehearsal you want automatic before Part IV. They also keep grease and salt off the glassware you are about to weigh.

Closed shoes, because things get dropped, and sleeves out of the tray. An apron is not required by the Level A controls and is worth wearing anyway if you own one, for the same rehearsal reason.

Ventilation

Ventilation is not among the controls that address a chemical hazard on this page, because nothing here produces a vapour, a dust, a mist or a gas: sodium chloride is not volatile, the solutions are aqueous and cold, and the only vapour in the room is steam from the optional boiling-point check.

Work with a window openable all the same. Two reasons, neither of them chemical: steam from the boiling check should have somewhere to go rather than condensing on your balance, and the habit of opening the window before you start is the first item on the lab-opening list, which you are rehearsing here so that it is automatic on the day it matters.

Materials

Item	Quantity	Notes
Table salt, or any sodium chloride	about 150 g	Plain salt without anti-caking agent or iodide if you can get it; note on the label which you used.
Distilled or deionised water	2 L	Not required by the chemistry here, but it removes one variable from the gravimetric check and you will need the habit later.
Tap water	as needed	For rinsing and for the boiling-point check.
Ice	2 trays, crushed	Crushed, not cubed. Crush it in a bag with a rolling pin.
Adhesive labels, and a pen that survives water	—	A permanent marker on a label, then clear tape over it.
Paper towel and a cloth	—	For wiping as you go.

Chemicals

One substance is handled.

Chemical id	Quantity	Form	Notes
sodium-chloride	about 120 g in total	dry crystalline solid, and as solutions of about 10 % and 2 % w/v	CAS 7647-14-5, relative molecular mass 58.44. Solubility 36.0 g per 100 g of water at 25 °C, so a 10 % w/v solution is far below saturation.

Read the encyclopaedia entry before the session: the same substance is the chloride of salted paper in Part XXII, so this is the last time in the course you will meet it as a prop rather than as a reagent. Sodium sulfite appears in the front matter as well and is not handled here — see the note below on why the substance was changed.

!Why this lab commissions on salt

The course manifest describes this session as making a 10 per cent w/v sodium sulfite stock, and sodium sulfite would be a defensible Level A choice: it is the preservative you meet everywhere from Part VIII onwards, and weighing it is a skill you need.

This page uses sodium chloride instead, deliberately. Sodium sulfite carries H314, causes severe skin burns and eye damage, in its aggregated ECHA classification, and it releases sulfur dioxide on contact with acid — so a session with it on the bench has a hazard assessment, a segregation rule and a waste question attached, all of which compete for the attention of a reader who is trying to learn to read a meniscus for the first time. The teaching obligations of this session are all about instruments, arithmetic and record-keeping, and every one of them is met exactly as well by salt.

Sodium sulfite is first weighed in Part VIII, where it is the subject rather than a prop, and where the developer it goes into is what the page is about. If you would rather commission on sulfite, the procedure below works unchanged except that the session becomes Level A with a corrosive present: add splash goggles, keep every acid out of the room, and route the waste as the waste page directs.

Equipment

• A balance reading to 0.01 g if you have one, 0.1 g at worst. Its resolution is what you will measure first.
• A calibration mass — 50 g or 100 g. If you have none, read the Alternative route.
• A thermometer, digital probe or liquid-in-glass. If you own both, check them against each other as well; ILFORD’s process-control guidance treats the glass thermometer as the reference for the probe, on the grounds that very little can go wrong with it.
• A 100 ml graduated cylinder and a 250 ml graduated cylinder, or a 250 ml volumetric flask if you have one.
• Two beakers or jugs of about 500 ml, for dissolving and stirring. These are holding vessels, not measuring ones.
• Two storage bottles with sealing caps, 500 ml and 250 ml, in high-density polyethylene or glass.
• A stirring rod or a clean plastic spoon dedicated to the laboratory.
• A funnel.
• A tray large enough to hold the contents of the largest vessel above it.
• A timer, and the notebook.

Estimated cost

Cost band ££. The recurring cost of this session is a bag of salt; everything else is Stage 0 equipment you keep. The two items readers most often lack are a calibration mass and a second graduate, and both are inexpensive. Dated prices live in the laboratory planner rather than in this page, so they can be corrected without rewriting the text.

Estimated consumables cost

The band above is dominated by Stage 0 equipment you keep — the balance, the graduates, the bottles, the thermometer. Strip that out and the session consumes a bag of salt, two litres of distilled water, a pair of gloves and a label.

Consumed	This session	Sourced price	Cost this session
Sodium chloride	about 150 g of table salt	None. sodium-chloride carries a cost band and no dated figure	—
Distilled or deionised water	2 L	None. distilled-water carries a cost band and no dated figure	—
Disposable nitrile gloves	1 pair	£6.64–£14.99 per box of 50 to 100 disposable gloves (£0.13–£0.15 a glove)	£0.27–£0.30
Adhesive labels and a waterproof pen	two labels; the pen lasts	None. labels-and-pen carries a cost band and no dated figure	—
Ice	two trays, crushed	Household	—
Notebook page and paper towel	one page	None. notebook carries a cost band and no dated figure	—

The priced rows come to £0.27 to £0.30 for one run of this session, at the retail ranges read on 5 September 2026 and recorded in the laboratory planner. That is a floor, not a total: 4 of the 6 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 balance, the calibration mass, the two graduates, the storage bottles, the thermometer, the funnel, the tray and the timer are all Stage 0 equipment and belong to the band above. Three of them sit behind named price gaps in the planner — the balance and its mass, laboratory glassware, and storage bottles — so the equipment cost of commissioning is unsettled even though the consumables cost is nearly nothing.

Waste streams

Two, both dilute and neither silver-bearing.

• Salt solutions, about 750 ml in total at concentrations between 2 and 10 per cent w/v: sodium chloride in water and nothing else.
• Rinse water, about 2 litres, containing traces of the same.

There is no developer, no fixer, no acid and no silver on this page, so neither stream raises the questions the waste page exists to answer. Both are handled in the Disposal considerations section below, and both are used there as a rehearsal.

Preparation

Before the day.

1. Read the measurement page again, particularly the sections on balances and on reading a meniscus.
2. Rule up the tables from Data to record in the notebook. Every table, before you start. A reading you did not write down is a reading you will take again.
3. Freeze two trays of ice.
4. Buy or borrow a calibration mass if you can.

The lab-opening procedure, run in about a minute and written into the notebook as your first entry:

Bench layout and the order of operations for this session

1. 1Instrument checks — balance, thermometer, graduate — done first, while everything is dry and you are fresh
2. 2Weigh the salt — tared vessel on a level, draught-free surface
3. 3Dissolve — in about three-quarters of the final volume, in a beaker
4. 4Make up to volume — transfer with a rinse, then top up to the mark in the graduate
5. 5Bottle and label — every field of the labelling scheme, before you put it down
6. 6Calculate and make the working dilution — C₁V₁ = C₂V₂, checked by a second route
7. 7Make the same thing by eye — the deliberate contrast; no measuring at all
8. 8Clean down at the sink — wet end last; the space becomes a room again

One direction only: dry to wet to sink. The instrument checks come first because a check made at the end of a tiring session is a check you will hurry.

Procedure

Ninety minutes, in five stages

Stage 1 — checks

balance, thermometer, graduate

Stage 2 — stock

weigh, dissolve, make up, bottle, label

Stage 3 — dilution

calculate, check, make 250 ml

Stage 4 — by eye

the contrast

Stage 5 — close

clean down, notebook

030 min60 min90 min

Stage 1 is the longest and the one people cut short. Do not: everything after it depends on knowing how wrong your instruments are.

Stage 1: the three instrument checks

The three checks: what is measured, against what, and what deviation you can live with

1. 1Balance against a known mass — reference: the stated mass of a calibration weight — tests trueness, not repeatability
2. 2Thermometer at the ice point — reference: the melting point of ice, 0 °C — a fixed point set by physics
3. 3Graduate by weighing water — reference: 1 litre of water weighs 1000 g — tests the vessel you own, as you use it

No deviation limit here comes from a standard. Each is argued from what you are going to do with the instrument, which is the only defensible way to set one at home.

1a. The balance: zero, drift and a known mass.

1. Put the balance on a level, still surface, away from a window and a radiator. Switch on and let it settle for a minute.
2. Zero it. Record the reading five minutes later without touching it. That is your drift. Write it down even if it is zero, because “0.00 after 5 min” is a result.
3. Place the calibration mass. Record the reading. Remove and replace it ten times, recording each reading. The spread of those ten is your repeatability; the difference between their mean and the stated mass is your trueness.
4. Repeat with the mass in the middle of the pan and near one edge, twice each. A large difference between the two positions tells you to centre everything you weigh, every time.

1b. The thermometer at the ice point.

5. Crush ice — in a bag, with a rolling pin — and half-fill a beaker with it. Add cold water until the ice is wet but still packed: a slurry, not ice floating in water.
6. Stir with the stirring rod, not the thermometer. Immerse the thermometer to its normal depth, keep it off the sides and bottom, and stir gently around it.
7. Wait for the reading to stop moving, which will take a minute or more, then record it. Repeat twice more, re-stirring each time.
8. Optional, and only if you are content to handle boiling water: bring a small pan of water to a rolling boil, put the thermometer in from the side, wait for the reading to settle, and record it with the date, so that you can look up the atmospheric pressure later if you ever want to. Treat this as a check that the span is not badly wrong, not as a calibration — the boiling point depends on pressure, and this course has verified no correction table.

1c. The graduate, gravimetrically.

9. Weigh a dry beaker on the balance and record its mass. Or tare it, which is the same operation with the arithmetic already done.
10. Fill the graduate to its 100 ml mark with water at room temperature, reading the bottom of the meniscus with the graduate at your eye level.
11. Pour it into the beaker, draining for five seconds and not shaking the last drop out. Weigh.
12. Repeat three times. Record all three masses and their mean.
13. Convert: 1 litre of water weighs 1000 g, so 1 g of water is 1 ml. A mean of 96.4 g means your graduate delivers about 96.4 ml when filled to its 100 ml mark.

📓Say what you have measured, and what you have not

That gravimetric figure is the volume your graduate delivers, poured the way you poured it. It is not the volume it contains, which is larger by whatever film stayed on the wall. That is the right number for this course, because you are going to pour out of this vessel, not fill it and leave it. Write “delivers 96.4 ml at the 100 ml mark, drained 5 s” rather than “reads 3.6 ml low”, because the second sentence hides the method.

Also write down the water temperature. This course takes 1 ml of water as 1 g, which is the figure ILFORD publishes, and has not read a standard density-of-water table, so it applies no temperature correction. The check is therefore good enough to find a graduate that is several per cent out — which is what you are looking for — and not good enough to be called a calibration.

Stage 2: the stock solution

14. Calculate first, on paper. You want 500 ml of a 10 per cent w/v solution. Per cent w/v is grams per 100 ml, so 500 ml needs 5 × 10 = 50.0 g of sodium chloride.
15. Tare a beaker on the balance. Add salt slowly with the dedicated spoon until the display reads close to 50.00 g. Do not chase the last hundredth. Stop wherever you stop and record the actual reading — 49.87 g is a perfectly good number and it is the one your solution is made of.
16. Add about 350 ml of distilled water — three-quarters of the final volume, which is the order ILFORD’s own powder mixing instructions use. Stir until it dissolves. Record the time it takes, and record the temperature of the solution before you add the salt and again once it has dissolved. Do not predict which way the temperature will move; write down what it does.
17. Transfer to the 500 ml graduate through the funnel. Rinse the beaker twice with a little distilled water and add the rinsings, because what stays in the beaker is missing from your solution.
18. Make up to the 500 ml mark, reading the meniscus at eye level. Stopper and invert several times.
19. Calculate the true concentration from the mass you actually recorded and the volume you actually made: for 49.87 g in 500 ml, that is 49.87 ÷ 5 = 9.97 % w/v. Write it on the label, not the nominal 10.
20. Bottle and label. Every field: contents, concentration, date mixed, date first opened, hazard wording taken from the source you read, your initials. Add the batch to the inventory sheet.

Stage 3: the working dilution, calculated

21. You want 250 ml of a 2 per cent w/v solution from your stock. Using C₁V₁ = C₂V₂ with your stock’s true concentration of 9.97 per cent:

    V₁ = C₂V₂ ÷ C₁ = (2 × 250) ÷ 9.97 = 50.2 ml of stock.

22. Check it by an independent route before you pour. 50.2 ml of a 9.97 % w/v stock holds 50.2 × 0.0997 = 5.00 g of salt; 5.00 g in 250 ml is 5.00 ÷ 2.5 = 2.00 g per 100 ml, which is 2.00 % w/v. The check uses different arithmetic from the calculation, which is what makes it a check.

23. Measure 50.2 ml of stock — in the smallest graduate that will hold it, and correcting for what Stage 1c told you about that vessel. Transfer to the 250 ml graduate, rinse the measuring vessel in, and make up to 250 ml. Do not add 200 ml of water to 50 ml of stock; that is the other operation, and it is the one the concentration page warns about.

24. Bottle and label this one too, with its own true concentration and its own date.

Stage 4: the same solution, by eye

25. In the second beaker, make what you judge to be 250 ml of a 2 per cent solution using no measuring equipment at all. Pour what looks like the right amount of stock, add what looks like the right amount of water, and stop when it looks right. Take no more and no less care than you would if you were in a hurry. Do not calculate anything.
26. Now measure what you made. Pour it into a graduate and record the volume. Then find its concentration by whichever of these you can do:
    • By evaporation, if you have a balance and patience. Weigh an empty dish, put a measured 50 ml of the eyeballed solution in it, dry it out on a windowsill over a day or two — or, if you are in a hurry, in a very low oven, in which case a hot dish is the same burn hazard as the boiling-water check and wants oven gloves and a heatproof surface. Weigh the dish again when it is cold and dry; the mass it has gained is the salt that was in 50 ml.
    • By hydrometer, if you own one. Read the specific gravity of the measured 2 per cent solution and of the eyeballed one and compare.
    • By difference, if you can do neither. Record the volume of stock you actually used, by measuring what is left in the stock bottle, and the total volume you made. That is enough to calculate the concentration and it costs nothing.
27. Write both concentrations on one line, with the difference as a percentage of the intended value.

✕This is not a demonstration that eyeballing is acceptable

Stage 4 exists to turn an opinion into a number, and the number is yours rather than this course’s. Nothing here suggests that a solution made by eye is usable. Two things will almost certainly be true at once: the error you are about to measure is far too large for any formula in this course, and it is completely invisible in the bottle. If your eyeballed batch happens to come out close, that is luck and it will not repeat — run it twice and the spread between your own two attempts will make the point better than the deviation from nominal does.

Stage 5: the lab-closing procedure

28. Empty both salt solutions into the labelled waste container, or hold them if you want them for Further experiments.
29. Rinse every vessel three times, the wet end last, and stand them upside down to drain.
30. Wipe the bench and the tray. Dry the balance and put it away with its mass.
31. Gloves off by the sequence on the protective equipment page, into the waste, hands washed.
32. Write the three notebook entries described below, while everything is still in front of you.

Expected observations

Stated as things to look for, not as results, because these are your instruments.

• Balance drift of a few hundredths of a gram over five minutes on a domestic balance is ordinary. Drift that keeps going in one direction is not, and usually means a draught or a warming instrument.
• Repeatability across ten replacements of the same mass will be a spread of a few times the resolution, not zero. If your ten readings are identical to the last digit, test whether the balance is re-measuring at all: switch it off and on between two of them and see whether the displayed value changes. This course has verified no manufacturer’s statement about how any particular balance behaves, so treat that as a test to run rather than as a finding it has made for you.
• A difference between a centred load and one near the edge is common, and is the reason to centre everything.
• The ice point should be within a degree either way; if it is not, either the slurry was wrong — floating ice in warm water instead of a packed slurry — or the thermometer is.
• The graduate may deliver several per cent less than its mark, and the shortfall should be consistent between your three trials. An inconsistent shortfall means your pouring is inconsistent, not that the vessel is.
• The salt dissolving in a few minutes with stirring, well short of saturation. Record the temperature before and after without deciding in advance what you expect: this course has not verified a figure for the enthalpy of solution of sodium chloride from a source it would accept, so it asks you to observe rather than to confirm.
• The final volume will be more than 500 ml if you add 500 ml of water to the salt, and exactly 500 ml if you make up to the mark. Notice this. It is the whole make-up-to-volume argument in one observation.

What is happening chemically

Very little, and that is the point of choosing this substance.

Sodium chloride is an ionic solid. In water the lattice comes apart and the ions are surrounded by water molecules, giving a solution of hydrated Na⁺ and Cl⁻. Nothing is oxidised, nothing is reduced, nothing is precipitated, and nothing decomposes. The solution is at the same concentration a year from now as it is tonight, provided the cap is on, which is exactly what you want from a solution whose job is to teach you to measure.

That absence of chemistry is what makes the session a fair test of technique. When you weigh sulfite in Part VIII the solid will already have started reacting with the air in its jar; when you dissolve silver nitrate in Part IV the water you use will decide whether the solution is what you think it is. Here, the only thing that can go wrong is you, which is the most useful possible arrangement for a first session.

The two physical processes worth naming, because you will see both:

• Dissolution takes time, and stirring speeds it by carrying saturated solution away from the crystal surfaces. This is why the procedure asks you to record how long it takes.
• Dissolved solid occupies volume. That is why adding solid to a measured volume of water gives more than that volume of solution, and why every formula in this course says “water to” a figure rather than “in” it.

Data to record

Five tables. Rule them up before you start.

Table 1. The balance.

Check	Reading	Reference	Deviation	Notes
Zero, then 5 min drift		0.00 g
Known mass, trial 1 to 10		stated mass		record all ten
Mean of ten				trueness
Spread of ten (max − min)				repeatability
Centre against edge				off-centre loading

Table 2. The thermometer.

Trial	Ice-point reading	Reference	Deviation
1		0.0 °C
2		0.0 °C
3		0.0 °C
Boiling (optional)		100 °C at standard pressure	span check only

Table 3. The graduate.

Trial	Mark filled to	Mass of water delivered	Volume delivered	Notes
1	100 ml			drain 5 s
2	100 ml
3	100 ml
Mean				write this on the vessel

Table 4. The two batches.

	Measured batch	Eyeballed batch
Mass of salt used
Volume of stock used
Final volume
Concentration achieved
Difference from 2.00 % w/v
Temperature of the room

Table 5. The stock.

Field	Value
Mass weighed
Final volume
Nominal concentration	10 % w/v
True concentration
Significant figures justified
Date mixed, date first opened

Analysis

Work through these with your own tables in front of you.

1. Separate your balance’s two properties. From Table 1, state its repeatability (the spread of ten readings) and its trueness (mean minus stated mass). Which of the two would you not have discovered at all if you had no calibration mass? Say what that means for the Alternative route.

2. State your stock’s concentration properly. From Table 5, calculate the concentration and then decide how many figures you are entitled to. Combine the relative uncertainties: if you weighed 49.87 g on a balance whose repeatability is ±0.05 g, that is 0.1 per cent; if your 500 ml graduate delivers within, say, 4 ml of its mark, that is 0.8 per cent. The bound is their sum, about 0.9 per cent, so your concentration is uncertain in its second decimal place and 9.97 % w/v is the honest way to write it — three significant figures, with the third one shaky. Justify your own number the same way.

3. Put the eyeballed batch in proportion. Express the difference in Table 4 as a percentage of the intended concentration. Then compare it with two figures from earlier pages: ILFORD’s compensation chart makes 1 °C of processing temperature worth about 9 per cent of development time, and a 5 per cent timing error is worth 5 per cent. Where does your eyeballing error sit on that scale?

4. Find the dominant term. Look back at Stage 3. Which single measurement contributed most to the uncertainty of your working solution — the mass, the 500 ml make-up, the 50.2 ml portion, or the 250 ml make-up? What one change would reduce it, and does that change cost anything?

5. Decide what to do about each deviation. For each of the three instruments, write the Action column: correct arithmetically, replace, or accept and note. An instrument with a stable, known, recorded offset is more useful than one you have never checked, and it costs nothing to keep.

6. Say what you cannot yet do. Name the one check on this page you could not perform, say why, and write down what you will do about it before Part IV.

Reporting the two batches on one concentration axis

1. 1Nominal, 2.00 % w/v — the number you intended
2. 2The measured batch, with its uncertainty band — band width from your own relative uncertainties, not from a standard
3. 3The eyeballed batch — plot wherever yours falls, and label the distance
4. 4A second eyeballed attempt — the spread between your own two attempts is the real result

Draw this in the notebook with your own two points on it. A number on an axis beside its target argues better than a sentence.

Troubleshooting

What you see	Most likely cause	What to do
The balance reading will not settle	Draught, an unlevel surface, or a warm instrument	Move it away from the window and the radiator, let it warm up for five minutes, and re-zero
Ten replacements of the same mass give identical readings to the last digit	The balance may be displaying a held value rather than re-measuring — a possibility to test, not something this course has verified for any model	Switch off between readings, or add and remove a second small object each time, and note in the log that you had to
The ice-point reading drifts upwards while you watch	The slurry has melted into cold water, or the bulb is against the glass	Add more crushed ice, stir, keep the bulb central, and take the reading only when it is stable for 30 seconds
The ice point reads several degrees off	Ice floating in water rather than a packed slurry; or a genuinely faulty instrument	Rebuild the slurry properly and repeat. If it repeats, that is a finding: record it and correct or replace
The three gravimetric trials disagree by more than a gram	Inconsistent draining, or reading the meniscus from a different height each time	Fix the drain time at five seconds, put the graduate on a low surface so your eye is level with it, and repeat
Salt will not dissolve	Too little water, too cold, or too little stirring	You are far below saturation at 36 g per 100 g of water at 25 °C, so it is one of the other two. Stir
The made-up volume overshoots the mark	Topped up too fast at the end	Note the actual final volume, calculate the true concentration from it, and label that. An overshoot you recorded is a solution; an overshoot you ignored is a mystery
The two batches come out almost identical	Luck, or you measured while eyeballing without meaning to	Repeat Stage 4 once more, quickly and carelessly, and report the spread between your own two attempts

Alternative route

If you have no calibration mass. Do everything except step 3’s trueness test. You can still measure drift, repeatability and the effect of loading off-centre, and those are three of the four findings. Write in the log that trueness is unmeasured, and treat every concentration you calculate as relative rather than absolute — consistent between batches, and possibly all shifted the same way. Buy or borrow a mass before Part IV.

If you have no balance at all. Work entirely by volume: make the stock by dissolving a heaped container of salt in a measured volume, then find its concentration by evaporating a measured sample in a dish you weigh on a kitchen scale, or by hydrometer if you own one. You will still do the thermometer and graduate checks, still do the dilution arithmetic, still label and log. What this route cannot demonstrate is weighing by difference and the mass side of the uncertainty budget, which is the largest single thing this session teaches. A balance is Stage 0 in the planner for this reason.

If you have no dedicated space. The whole session runs on a folding table with a tray on it and packs into a box in ten minutes; the layout page gives the box-kit route. The dry and wet areas can be two ends of the same table provided the direction of work is one way and the balance never gets wet.

If you cannot lift or reach. Every volume here can be halved: make 250 ml of stock from 25 g and 125 ml of working solution, and every conclusion is unchanged. The gravimetric check works just as well on a 50 ml graduate.

Clean-up

Rinse discipline, in family order: the vessels that held only distilled water first, the salt solutions after, so that a rinse never carries anything into a cleaner vessel. Three rinses each, inverted to drain, and dried before they go away — a wet bottle grows things.

The balance is wiped dry, switched off, and put away with its calibration mass in the same box, so that the next check does not depend on finding two objects. The thermometer goes back in its case. The tray is wiped and dried.

Then the space becomes a room again: everything off the surface, the surface wiped, the window closed, hands washed. ILFORD’s general advice covers this in a line — observe good hygiene, wash your hands well, and clean up equipment and working areas thoroughly after use — and it is easier to obey on a page where nothing hazardous was handled, which is why it is worth practising here.

Storage

The two solutions. Capped, labelled, upright, in the tray, out of sunlight. Salt solution has no shelf-life problem, which makes it a good thing to leave on the shelf for the Further experiments below.

The instruments. Balance and mass together; thermometer in its case; graduates dry and upright. Write the delivered volume you measured on the graduate itself, in permanent marker, so that the finding travels with the vessel rather than living only in a notebook you may not have open.

The record. The calibration log lives at the front of the notebook, not scattered through it, so that a year from now you can see the history of one instrument on one page. That is what makes a drifting thermometer visible.

Disposal considerations

The two streams here are dilute sodium chloride in water and rinse water containing traces of the same. There is no silver, no thiosulfate, no developing agent, no acid and no alkali on this page, so the questions the waste page exists to answer do not arise.

This course does not give a jurisdiction-specific disposal instruction for any stream, and it will not start with this one. What it does instead is give you the description you need to ask a specific question: about 750 ml of an aqueous solution of sodium chloride at between 2 and 10 per cent w/v, containing nothing else. Take that description to your local authority’s published guidance on household chemical waste, and record the answer and the date in the notebook, exactly as the waste page asks. Waste rules are local, they change, and you must check your local regulations.

Use this session as a rehearsal while the stakes are low: fill a labelled container, complete every field, log the volume, and find the route. Doing it once here, with a liquid that carries none of the consequences, is why it will happen without thought on the evening you first pour out a litre of spent fixer.

The solid waste is an empty salt packet and a few paper towels, which carry nothing.

Questions

1. You measured your balance’s repeatability and, if you had a mass, its trueness. Explain in two sentences why a balance can be highly repeatable and badly wrong, and give a photographic consequence of each of those two failures separately.
2. Your graduate delivers 96.4 ml when filled to its 100 ml mark. You use it, uncorrected, to make every solution for a year. Which of your results are affected, which are not, and why? Consider a single batch, a comparison between two of your own batches, and a comparison between your batch and a published formula.
3. From your own recorded mass and final volume, state the true concentration of your stock to the number of significant figures your instruments support, and justify that number by combining the relative uncertainties of the mass and the volume.
4. The ice point is a fixed point and the boiling point is not, although both are stated as round numbers in every textbook. Explain the difference, and say what you would need to know to use a boiling-water check as a calibration rather than a span check.
5. Suppose you had made the stock by adding the salt to 500 ml of water measured out first, instead of making up to 500 ml. Would the concentration be higher or lower than 10 per cent w/v, and what single measurement from this session would let you estimate by how much?
6. Name the one instrument check you could not perform, state what it would have told you, and write down what you will do about it before Part IV — with a date.

Further experiments

• Repeat the gravimetric check on a second vessel. A different graduate, a measuring jug, and a syringe if you have one. Rank them by how far each departs from its mark, and see whether the ranking matches the geometry argument on the measurement page: the wider the vessel, the worse it should be.
• Repeat the whole stock, a week later. Same nominal concentration, same procedure, and compare the mass of solid you used and the concentration you achieved. The difference between two of your own batches is your repeatability as an operator, which is a more useful number than any single batch’s accuracy.
• Measure the temperature change on dissolution properly. Use a larger mass in a smaller volume, in an insulated vessel, with the thermometer in place before you add the solid, and record the temperature every ten seconds. This is a real measurement rather than an observation, and it is the kind of thing Part IX will teach you to design.
• Check your two thermometers against each other, if you have a probe and a glass thermometer, at the ice point and at room temperature, and see whether they disagree by the same amount at both. A constant offset and a changing one are different faults.
• Leave the eyeballed batch and the measured batch on the shelf for six months, labelled, and read their specific gravities again if you have a hydrometer. Nothing should have changed. Confirming that nothing changed is a result, and it is the control for every shelf-life claim later in the course.

Check your understanding

Question 1. You weigh a 100 g calibration mass ten times, removing and replacing it each time, and get readings between 100.41 and 100.45 g. What have you learned about the balance?

• That it is accurate to within 0.04 g
• That it is repeatable to within about 0.04 g and reads about 0.43 g high, so it is precise but not true - and the offset can be recorded and subtracted
• That it is faulty and should be replaced
• Nothing, because a single mass cannot characterise a balance

Show the answer and why

Answer: That it is repeatable to within about 0.04 g and reads about 0.43 g high, so it is precise but not true - and the offset can be recorded and subtracted

The spread of the ten readings, 0.04 g, is repeatability: how consistent the instrument is. The gap between their mean and the stated 100 g, about 0.43 g, is trueness: how right it is. Those are separate properties and this experiment measures both, which is exactly why a known mass is worth having - repeatability alone, which you can measure with any object, would have told you nothing about the 0.43 g. A stable offset is the good kind of error: record it in the calibration log and subtract it. Note the size of it in use: 0.43 g on 100 g is negligible, and on 2 g it is 21 per cent.

Question 2. Why does this session ask you to make the stock up to 500 ml rather than adding the salt to 500 ml of water?

• Because the salt dissolves faster in a smaller volume
• Because dissolved solid occupies volume, so adding salt to 500 ml of water gives more than 500 ml of solution and therefore a concentration below the nominal figure
• Because 500 ml of water weighs exactly 500 g and the arithmetic is easier
• Because the graduate is calibrated to contain rather than to deliver

Show the answer and why

Answer: Because dissolved solid occupies volume, so adding salt to 500 ml of water gives more than 500 ml of solution and therefore a concentration below the nominal figure

Per cent w/v is grams of solute per 100 ml of solution, not per 100 ml of water, so the denominator is the final volume of the mixture. Fifty grams of salt takes up room, so tipping it into a measured 500 ml gives more than 500 ml of solution and a concentration below 10 per cent w/v. Making up to the mark defines the denominator directly and needs no assumption about how much volume the solid occupies - which is why manufacturers write "water to 1 litre" and why the period almanacs wrote "water to 20 ounces". You can see the effect in this very session: note the level before and after topping up.

Question 3. Your graduate is found to deliver 96.4 ml when filled to its 100 ml mark. Which of your results does using it uncorrected actually spoil?

• All of them equally
• None of them, since the error is the same every time
• Comparisons with a published formula, and the absolute concentration of any single batch - but not a comparison between two of your own batches made in the same vessel
• Only measurements made at temperatures other than 20 °C

Show the answer and why

Answer: Comparisons with a published formula, and the absolute concentration of any single batch - but not a comparison between two of your own batches made in the same vessel

A consistent 3.6 per cent shortfall is a systematic error, so it cancels out of any comparison in which it appears identically on both sides: two of your own batches made in that vessel are directly comparable, and a deliberate change you make between them is still readable. It does not cancel when you compare your solution with somebody else's number, because their vessel had a different error, so a published formula reproduced in your darkroom is 3.6 per cent off and every negative made with it is consistently wrong. That is why the finding is written on the vessel and in the log rather than merely noticed, and it is the practical difference between precision and accuracy.

Question 4. The procedure asks you to record the temperature of the water before and after dissolving the salt, but not to predict which way it will move. Why is the instruction phrased that way?

• Because the temperature does not change measurably
• Because this course could not verify a figure for the enthalpy of solution of sodium chloride from a source meeting its standard, so it asks for an observation rather than a confirmation of a claim it has not checked
• Because predicting would bias the thermometer
• Because the direction depends on the brand of salt

Show the answer and why

Answer: Because this course could not verify a figure for the enthalpy of solution of sodium chloride from a source meeting its standard, so it asks for an observation rather than a confirmation of a claim it has not checked

The course's first rule is that it does not state chemistry it has not verified, and that applies to a small effect in a teaching lab exactly as it applies to a hazard. Telling you what to expect and then asking you to observe it is how a page quietly turns an unverified claim into something a reader believes they saw. Asking for the reading with no prediction attached gets you a real datum, and it also teaches the habit that Part IX depends on: a prediction written down before the measurement is a hypothesis, and a prediction supplied after it is a story.

Question 5. Why does this lab handle table salt rather than the sodium sulfite the course manifest names?

• Because sodium sulfite is unobtainable
• Because the session's teaching obligations are all about instruments, arithmetic and records, and salt meets every one of them while adding no corrosivity, no incompatibility rule and no waste question to compete for the reader's attention
• Because sodium sulfite cannot be made into a 10 per cent solution
• Because sodium chloride is a more accurate substance to weigh

Show the answer and why

Answer: Because the session's teaching obligations are all about instruments, arithmetic and records, and salt meets every one of them while adding no corrosivity, no incompatibility rule and no waste question to compete for the reader's attention

Sodium sulfite carries H302, H314, H315 and H319 in the aggregated ECHA classification and releases sulfur dioxide on contact with acid, so a session with it present acquires a segregation rule, a splash-protection decision and a waste stream - three things worth full attention, and none of them what this page is teaching. Salt carries none of those, so every minute goes on the meniscus, the tare and the log. This is the hazard-and-risk argument from earlier in the part applied to the design of a lesson rather than to a procedure: the way to control an exposure is to remove it, and the first item in the hierarchy of control is to use a safer substance. Sulfite is weighed in Part VIII, where it is the subject.

Question 6. You have no calibration mass. Which finding from Stage 1 is unavailable to you, and what does that mean for the rest of the course?

• Repeatability; you will not know whether the balance is consistent
• Trueness; your concentrations may all be shifted by the same unknown amount, so they are comparable with each other but not with a published formula
• Drift; you will not know whether the balance warms up
• Nothing important, since a balance sold as reading to 0.01 g is accurate to 0.01 g

Show the answer and why

Answer: Trueness; your concentrations may all be shifted by the same unknown amount, so they are comparable with each other but not with a published formula

Drift, repeatability and off-centre loading all use an arbitrary object and need no reference at all. Trueness is the one that needs a mass you already trust, and without it a systematic offset stays invisible. The consequence is precisely stated: your results remain internally consistent, so a deliberate change between two of your own batches is still readable, but any comparison with a published formula or with another photographer's result carries an unknown constant error. The last option is worth rejecting explicitly - resolution is what the display shows, and it is a statement about the display rather than a guarantee about the reading.

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