Manufacturing Ferric Oxalate
This article explains how to make high quality ferric oxalate for platinum/palladium printing.
It is my attempt to document a consistent process for manufacturing ferric oxalate. The process is based upon the one described by Giuseppe (Josef) Pizzighelli and Baron Arthur von Hübl in Die Platinotypie/The Platinotype (1883). It was modernised by Pierre Brocher (I am not sure when), and was the basis of Jean-Claude Mougin’s book: Short Treatise on Iron Salts. I am indebted to Jean-Pascal Laux for his support and encouragement to learn this method.
The instructions assume that you are familiar with basic chemistry. Some chemistry topics are covered in my book about platinum/palladium printing: The Platinum Printing Workshop (Amazon.com or ebook) which also covers the full platinum/palladium printing process.
If you want to read Pizzighelli’s original instructions, you can find them here: The Platinotype (English Translation).
I would love to hear from you, so if you have questions or feedback then please do not hesitate to reach out.
Chapter 1: Equipment
The following items of equipment are required for the manufacture of ferric oxalate solution. All are easy to find on Amazon and other retailers.
- Scales suitable for weighing 50-100g of chemicals ±1g
- Scales suitable for weighing 1-50g of chemicals ±0.1g
- Safety glasses
- Nitrile or other suitable gloves
- Glass stirring rod
- 2L (or larger) plastic heat proof mixing jug
- A nylon fine mesh food strainer bag
- A 2L plastic funnel – sufficiently large to fit the nylon strainer bag when it is full of water and ferric hydrate ‘mud’
- A lab stand strong enough to support the nylon strainer bag when it is full of water and ferric hydrate sediment
- A Büchner funnel, flask, and vacuum pump
- Filter paper (Whatman Number One or equivalent)
- pH indicator paper
- Large plastic wastewater bucket
- Chapter 2: Chemicals and Chemical Safety
Ferric oxalate is properly called iron(III) oxalate, Fe2(C2O4)3. It is formed from two iron ions, Fe3+, and three oxalate ions, C2O42-. Its close cousin ferrous oxalate is properly called iron(II) oxalate, FeC2O4, and is formed of one iron ion, Fe2+, and one oxalate ion, C2O42-. Ferric oxalate converts (reduces) to ferrous oxalate when heated, exposed to ultraviolet light, or when mixed with certain chemicals. It is important to avoid these reactions during manufacture.
This manufacturing process uses the following chemicals. All are easily available.
- Ferric Chloride – iron(III) chloride (FeCl3) which is available in anhydrous form (black powder) or hexahydrate form (yellow chunks). These instructions assume you are using the hexahydrate (FeCl3·6H2O), but the anhydrous form is also acceptable
- Caustic Soda – which more correctly is called sodium hydroxide (NaOH). It is a white powder. These instructions assume you are using the anhydrous form, but the monohydrate (NaOH·H2O) is also acceptable
- Oxalic Acid – ideally anhydrous oxalic acid (C2H2O4) but a hydrate will do if that is all you can find, for example oxalic acid dihydrate (C2H2O4·2H2O). Oxalic acid is a white powder
- Water (H2O) – at least 10 litres. My water supply is clean enough to use for this process, but you may want to consider using distilled water
Ferric Chloride Hexahydrate
All darkroom chemicals require careful handling. In general:
- Read and understand the safety data sheets (SDS)
- Wear suitable protective equipment and clothing when handling and working with chemicals. Normally safety glasses, gloves and a long-sleeved lab coat will be sufficient
- Work in a well-ventilated space to avoid build-up of harmful fumes
- Ensure you store all chemicals in suitable containers, locked away from curious children
Although used in small quantities, the ferric oxalate ingredients are toxic and require special caution:
Ferric chloride (SDS) is harmful, highly corrosive and acidic. Wear gloves and suitable clothing to protect your skin. If you get it on your skin then remove contaminated clothing and wash immediately with plenty of water. Wear safety glasses. If you get it in your eyes then rinse immediately with plenty of water for several minutes (removing contact lenses if easily possible) and seek medical assistance. Use in a well-ventilated space. If you inhale fumes then go to fresh air and rest. Seek medical assistance if coughing or a sore throat persist.
Sodium hydroxide (SDS) is harmful and highly corrosive. It can cause severe burns and permanent blindness. It has been used by serial killers to dispose of dead bodies. Wear gloves and suitable clothing to protect your skin. If you get it on your skin then remove contaminated clothing and wash immediately with large quantities of water for ten to fifteen minutes and seek medical assistance. Wear safety glasses. If you get it in your eyes then rinse immediately with plenty of water for several minutes (removing contact lenses if easily possible) and seek medical assistance. Use in a well-ventilated space. Inhaling fumes leads to coughing, sore throat, burning sensation and shortness of breath: go to fresh air and seek medical assistance.
Oxalic acid (SDS) is highly toxic. It can be harmful if ingested or inhaled. It causes kidney failure, and ingesting as little as 15-30g may be lethal. It is a skin irritant. If you get it on your skin then wash immediately with plenty of water, and if irritation persists then seek medical assistance. If you get it in your eyes then flush eyes with running water for 15 minutes and seek medical assistance. If you inhale oxalic acid then move to a well-ventilated area and seek medical assistance. If ingested seek immediate medical assistance (do not induce vomiting).
Chapter 3: The Manufacturing Process
This method for manufacturing ferric oxalate has three parts:
- Make some ferric hydrate
- Convert the ferric hydrate into a ferric oxalate solution
- Standardise the ferric oxalate solution for printing
The method described in this chapter typically makes about 80ml of 20% ferric oxalate solution.
MANUFACTURING PART ONE: FERRIC HYDRATE
Ferric hydrate is made by mixing ferric chloride (FeCl3) and sodium hydroxide (NaOH). When dissolved in water they react to produce ferric hydrate (Fe(OH)3) plus some sodium chloride (NaCl). Ferric hydrate is more correctly called iron(III) hydroxide but I continue to call it ‘ferric hydrate’ because that is the name used by Pizzighelli.
The chemical equation for the reaction is:
FeCl3 + 3NaOH → Fe(OH)3↓ + 3NaCl
In theory, this will produce about 30g of ferric hydrate for every 50g of ferric chloride hexahydrate, but in practice the yield will be less because of inefficiencies and loss in the manufacturing process.
You can make a batch of ferric hydrate in an evening, and it can be made in normal artificial light.
Although ferric hydrate is not particularly toxic, it is a skin irritant so handle it carefully. If you get it in your eyes then rinse with plenty of water and seek medical assistance. It will stain anything it touches so wear gloves, eye protection and old clothes for this next step.
Ferric Hydrate (before drying)
- Heat about 2000ml of distilled water to 80ºC in a plastic mixing jug. Do not use your family’s heirloom porcelain jug because you will probably never get it clean again once it has been used to make ferric hydrate
- Slowly add 50g of ferric chloride hexahydrate (or 30g of anhydrous ferric chloride) to the hot water with stirring until fully dissolved. This produces a very hot strong acid, so wear safety glasses and gloves
- Carefully add about 25g of anhydrous sodium hydroxide (or 36g of sodium hydroxide monohydrate) with stirring. The precise amount varies depending on the specific ingredients you are using, so add a little at a time then check the solution’s pH with indicator paper. Stop adding sodium hydroxide when you get a distinct alkaline reaction (pH 8 or higher). Don’t worry if you add slightly too much, because any excess will be washed out in later steps. The reaction produces a thick brown suspension, which is your ferric hydrate
- Place the nylon mesh food strainer inside the 2L funnel, and hang it from the lab stand over your waste bucket. Then slowly pour the jug of ferric hydrate suspension into the bag, and allow to drain into the waste bucket
- Wash the ferric hydrate repeatedly by pouring hot water over it until the wastewater has neutral pH. 8-10L of hot water should be sufficient. Do not attempt to rush this by squeezing the bag because that will probably cause you to lose a lot of ferric hydrate and significantly reduce yield
- Leave the ferric hydrate to stand so that most of the remaining water drains away
Ferric Chloride Solution
Ferric Hydrate Suspension
After standing for 30-60 minutes, your strainer bag should contain somewhere between 160-200g of brown, incredibly sticky, ferric hydrate. Most of this mass is water – in the next step you’ll remove as much of the remaining water as possible.
In the old days, the ferric hydrate was dried by putting it in a fabric pouch and squeezing out the excess water. You still see instructions on the internet for this method. However, this is inefficient, wastes a lot of ferric hydrate, and is filthy-messy too. Using a Büchner funnel and vacuum pump is much faster, cleaner and gives a higher yield.
First cover the funnel’s ceramic floor with a disc of filter paper. With a spoon, transfer the drained ferric hydrate into the funnel. Switch on the vacuum pump and watch it suck water out of the ferric hydrate, as if by magic. Regularly stir the ferric hydrate to help extract the water and close any air holes that appear.
Eventually the ferric hydrate will solidify and take on the consistency of a dry, cracked block. The time this takes will vary considerably depending on your vacuum pump and the volume of water to extract. My vacuum pump achieves this in 20-30 minutes.
The block of ferric hydrate is simple to remove when dry, although it crumbles easily. The protective filter paper should peel off cleanly. Don’t worry about little bits of paper that may remain attached to it: these debris will be removed later.
The dried ferric hydrate should weigh about 80-90g.
I have been told that ferric hydrate can be kept for prolonged periods if sealed from air and refrigerated, but have not tested this.
MANUFACTURING PART TWO: FERRIC OXALATE
Now you have some dry-ish ferric hydrate you just need to add some oxalic acid to convert it into ferric oxalate. The chemical equation is:
2Fe(OH)3 + 3C2H2O4 → Fe2(C2O4)3 + 6H2O
This equation shows that even if the ferric hydrate is perfectly dry, adding oxalic acid will create water (H2O). Given that the ferric hydrate will still be somewhat wet and that the oxalic acid may also be a hydrate (i.e. it has extra water bound into its structure), you will quickly find plenty of water sloshing around. That is why it is so important to first dry the ferric hydrate as much as you can.
The ferric oxalate produced in this recipe is sensitive to ultraviolet light so the process must be completed in the lowest level of artificial light that you can comfortably work in. Avoid daylight and fluorescent tubes (which tend to produce UV).
- Place your dry-ish ferric hydrate into a 400ml glass beaker
- Add 20g of anhydrous oxalic acid (or 28g of oxalic acid dihydrate) for every 50g ferric chloride used to produce the ferric hydrate. Add it bit by bit, stirring continuously. As you mix it the oxalic acid and ferric hydrate will react, creating water, and gradually liquefying the mixture. The reaction generates heat, and if you add the oxalic acid too quickly you can actually boil off the water and ruin the ferric hydrate/oxalate
- Leave the beaker in the dark at normal room temperature overnight. During this time, the ferric oxalate will mature and separate out. You can speed this up by gently warming it, but only at the risk of converting some of your ferric oxalate into ferrous oxalate which will later have to be removed by adding hydrogen peroxide
- Filter the ferric oxalate solution into a 100ml glass cylinder and leave it to settle for a few more days (ideally a week). Some sediment is likely to fall to the bottom of the cylinder during maturation – this is mostly ferric hydrate and oxalic acid crystals. Initially the solution will have an olive-green colour. The colour of mature ferric oxalate depends upon the amount of oxalic acid it contains. Most likely the ferric oxalate made following these instructions will be a pure apple green colour, but with very low levels of oxalic acid it will be a gold/yellow colour
- Once the ferric oxalate has matured, filter it through Whatman #1 filter paper into a brown glass bottle, removing all remaining solids
Congratulations: you have now made some ferric oxalate solution!
Matured Ferric Oxalate Solution (With 5% Free Oxalic Acid)
MANUFACTURING PART THREE: STANDARDISING FOR PRINTING
In The Platinotype, Pizzighelli recommended a standard solution containing 20g of ferric oxalate for every 100ml of solution – a 20% solution. But nowadays most published recipes refer to 27% solution, which is 27g of ferric oxalate for every 100ml of solution.
Confusing as it may seem, Pizzighelli’s 20% solution and the modern 27% solution are actually about the same. Pizzighelli was measuring his ferric oxalate concentration using analytical chemistry which determine how much pure ferric oxalate is in the solution. Most dry ferric oxalate powders contain considerable quantities of other substances as well as their ferric oxalate: mostly water and oxalic acid, but sometimes other things too. 26-27g of the dry powders actually contains about 20g of pure ferric oxalate.
When I refer to a percent ferric oxalate solution in these instructions, I mean pure ferric oxalate (i.e. Pizzighelli’s definition). My target is therefore 20% not the 27% you may see in contemporary recipes.
Assuming that you removed sufficient water from your ferric hydrate then your manufactured solution should already be close to 20%; most likely a little bit lower. It can be easily adjusted upwards (by evaporation) or downwards (by adding water). But first you need to find out the concentration of the solution you have just made.
Assuming you do not have access to a chemistry lab, the easiest way to determine the percent concentration is by measuring the ferric oxalate solution’s specific gravity. This is not highly precise, but is generally sufficient for photographic purposes.
Measuring Specific Gravity
First find a small syringe. The syringes that come with children’s medicine bottles are good for this. The syringe I use has a fancy mechanism that stops the plunger coming out so I can fill it in the dark and know I have exactly the same quantity in it (which for this syringe is 5.2ml).
Weigh the syringe so you know its mass. Draw 5ml of distilled water into the syringe and weigh the syringe again. Empty the syringe and draw exactly the same volume of your ferric oxalate solution into the syringe and weigh it a third time. It is a good idea to repeat each measurement a few times, averaging the results.
The specific gravity of your ferric oxalate solution can be calculated as follows:
Mass of the distilled water = Mass of the syringe containing distilled water − Mass of the empty syringe
Mass of the ferric oxalate solution = Mass of the syringe containing ferric oxalate solution − Mass of the empty syringe
Specific gravity of the ferric oxalate solution = Mass of the ferric oxalate solution ÷ Mass of the distilled water
For example, when I measured the last batch of ferric oxalate I manufactured, I found:
Mass of the empty syringe = 6.55g
Mass of the syringe containing distilled water = 11.70g
Mass of the distilled water = 11.70 − 6.55 = 5.15g
Mass of the ferric oxalate solution = 12.57 − 6.55 = 6.02g
Specific gravity of the ferric oxalate solution = 6.02 ÷ 5.15 = 1.17
Converting Specific Gravity to a Percent Concentration
The following chart can be used to convert the specific gravity to a percent concentration.
Conversion of Specific Gravity to Percent Concentration
Please note that a similar chart was published in Dick Stevens’ Making Kallitypes but these data are different. This chart is based upon my chemical analyses of ferric oxalate produced using this method. You can, of course, use Stevens’ chart if you prefer. It produces results that are about 40-50% higher than mine, so presumably was calibrated for ferric oxalate hexahydrate solutions (the normal yellow-green powder you can buy for platinum/palladium printing).
To use the chart:
- Look along the horizontal axis until you find the solution’s specific gravity
- From that point, draw a line vertically upwards until it crosses the diagonal line
- Then draw a horizontal line back until it crosses the vertical axis
- Your solution’s percent concentration is the value where the horizontal line crosses the vertical axis
Using the Conversion Chart
In my example, the ferric oxalate solution’s specific gravity of 1.17 converts to a percent concentration of 19%. This is a little lower than my 20% target.
Increasing the Concentration of Ferric Oxalate Solution
The easiest way to increase the ferric oxalate solution’s concentration is to evaporate some of the water. This is usually quite a slow process, but can be speeded up if you pour the solution into a large flat-bottomed dish and blow air over it using a fan. This should be done at room temperature to avoid cooking the ferric oxalate.
I have a large bank of fans that I use to dry paper. When I blow these over a dish of ferric oxalate they can evaporate about 3ml of water per hour.
To work out how much water to evaporate, use the following maths:
Volume of solution at 20% = Starting volume of ferric oxalate solution × Starting percent concentration ÷ Target percent concentration (which is 20%)
Volume of water to evaporate = Starting volume of ferric oxalate solution − Volume of solution at 20%
For example, to increase my ferric oxalate concentration from 19% to 20% I needed to evaporate 4ml of water:
Volume of solution at 20% = 83ml × 19% ÷ 20% = 79ml
Volume of water to evaporate = 83ml
Starting volume of ferric oxalate solution = 83ml − 79ml = 4ml
Because 1ml of water weighs 1g, you can check progress by weighing the dish of ferric oxalate solution.
With accelerated evaporation you will probably overshoot your target concentration somewhat. In my example, I blew my fans over the ferric oxalate solution for a few hours. After evaporation I had 66ml of ferric oxalate (I spilled a bit); measuring its specific gravity showed it had a concentration of 23%.
Decreasing the Concentration of Ferric Oxalate Solution
It is easy to decrease the ferric oxalate solution’s concentration by adding distilled water.
Volume of solution at 20% = Starting volume of ferric oxalate solution × Starting percent concentration ÷ Target percent concentration (20%)
Volume of water to add = Volume of solution at 20% − Starting volume of ferric oxalate solution
For example, to reduce my 23% solution to 20% I needed to add 10ml of water:
Volume of solution at 20% = 66ml × 23% ÷ 20% = 76ml
Volume of water to add = 76ml − 66ml = 10ml
After adding the distilled water, I rechecked the specific gravity and verified that I had a 20% solution of ferric oxalate.
Achieving More Rigorous Standards
Measuring specific gravity will only allow you to estimate the solution’s concentration. It tells you nothing about what else besides the ferric oxalate is in the solution.
To understand more about your solution and achieve more rigorous standards you need to use analytical chemistry techniques such as titration. These are fairly straightforward once you understand the tools and techniques, but they are beyond the scope of this article. If you want to explore them, then I recommend that you consult a friendly chemist.
Chapter 4: The Proof is in the Printing
The ultimate test of your ferric oxalate is to make a print. The ferric oxalate you manufacture using these instructions should give you rich dark blacks and good clean highlights.
You may find it helpful to make step wedge prints to help track any inconsistencies from batch to batch. However, do not try to use this as a short cut to avoid doing at least some degree of analysis or it is highly unlikely that you will be able to achieve a standard ferric oxalate solution from batch to batch. At a minimum, you should measure your solution’s specific gravity and use this to estimate its concentration.
Enjoy your printing!