Mix a 6% Salt Solution: How to Calculate and Prepare a Given Mass Percent

This is a required experiment in ninth-grade chemistry and one of the most frequently examined lab tasks anywhere. It looks like three moves — weigh some salt, measure some water, stir — but every step hides a trap that will pull the result off target. This guide walks the real procedure from start to finish, laying out the five steps and the five sources of error, then uses an interactive simulation to pull apart what adding water, evaporating water, and adding solute each do to the mass percent.

Safety first (read this section before you start)

• This experiment uses table salt and plain water, which are harmless in themselves — but never taste anything in a laboratory, even when you are certain it is salt. That rule matters far more than this one experiment.
• Glassware (beaker, graduated cylinder, stirring rod) breaks easily. Handle it gently and check for chips or cracks after use.
• When stirring with a glass rod, don't knock it against the beaker's walls or bottom, and never tap it hard — you can crack the glass.
• On a balance, pick up the weights with forceps, and never put chemicals directly on the pan — put a clean sheet of weighing paper under them (corrosive chemicals go in a glass container instead).
• If you upgrade this experiment to preparing dilute hydrochloric acid, sodium hydroxide solution or similar, a teacher must be present and goggles are required. When diluting a concentrated acid, always pour the acid slowly down the wall of the beaker into the water — never the other way around.
• Wash the equipment when you finish, and pour waste liquids into the designated container, not casually down the sink.

Materials and equipment

• A balance (with its weights and forceps) and two sheets of weighing paper;
• A graduated cylinder (50 mL or 100 mL) and a dropper;
• A beaker (100 mL) and a glass stirring rod;
• Table salt (sodium chloride) and distilled or clean water;
• A narrow-mouth reagent bottle, a label sticker and a pen;
• A lab spoon, a dry cloth, and paper and pen for recording.

Step 1: calculate first, then touch the equipment

The target is 50 g of solution at a mass percent of 6%. The defining formula for mass percent is

ω = mass of solute ÷ mass of solution × 100%

Rearrange it, and the solute you need is

m_solute = 50 g × 6% = 3 g

Everything else is water: m_water = 50 g − 3 g = 47 g. Water's density is 1 g/cm³, so 47 g of water is exactly 47 mL and can be measured directly with the graduated cylinder. Write both numbers in your lab notebook before doing anything else — calculate first, prepare second is this experiment's first rule of discipline.

Step 2: weigh, measure, dissolve

Weigh 3 g of salt. After leveling the balance, put a sheet of identical weighing paper on each pan (so the papers' masses cancel out), set 3 g on the right pan, and add salt to the left pan with the lab spoon until the balance levels. A 3 g mass usually needs the sliding rider: set the rider to 3 g first, then add salt until balanced.

Measure 47 mL of water. Stand the cylinder on a flat surface, pour water to just below 47 mL, then switch to the dropper and add the last of it drop by drop up to the mark. Read with your eye level with the lowest point of the meniscus: looking down makes the reading too high (you actually get less water), looking up makes it too low (you get more) — either way the concentration drifts.

Dissolve. Pour the measured water into the beaker, add the weighed salt, and stir with the glass rod slowly in one direction until the solid has completely disappeared and the liquid is clear and transparent. Stirring only speeds up dissolving — it does not let more dissolve.

Bottle and label. Pour the solution into the narrow-mouth reagent bottle, stopper it, and stick on a label giving the name of the chemical and the mass percent of solute — for example "Sodium chloride solution, 6%." An unlabeled bottle is a chemical with no identity; this step is not optional.

What to record

Check the observation notes against these points:

• After 3 g of salt goes into 47 mL of water, ten-odd seconds of stirring makes it vanish completely, leaving a colorless, clear liquid — a solution doesn't have to have a color.
• The liquid level barely rises after dissolving, yet the solution's mass really is 50 g, not 47 g. Mass is conserved; volumes don't simply add.
• Let the finished solution stand for a whole day: no sediment appears at the bottom, and the concentration is identical top to bottom. That is what it means for a solution to be uniform and stable.
• If you keep adding salt to this 50 g of solution, at roughly 17 g more it stops dissolving — salt starts piling up at the bottom, and the solution is saturated.
• Leave the saturated salt water on a windowsill to evaporate slowly, and small white cube-shaped crystals grow on the bottom of the glass.
• Repeat the whole thing with copper sulfate and the solution is blue, deeper as it gets stronger; with potassium permanganate, a tiny pinch turns it purple-black. Color depth is no measure of concentration across substances.

How it works

A solution is a uniform, stable mixture made of a solute and a solvent, obeying one simple mass relationship: mass of solution = mass of solvent + mass of dissolved solute. Mass percent is the solute mass divided by that total. The word "dissolved" in the denominator is the most important qualifier in the whole chapter: solid sitting undissolved at the bottom of the beaker counts neither as solute nor as part of the solution.

Why does the volume barely change when 3 g of salt joins 47 g of water? Because there are gaps between water molecules to begin with, and the sodium and chloride ions — each surrounded by water molecules — slip into those gaps. Masses add; volumes don't simply add. That is also why solutions are prepared by mass rather than by volume.

The maximum number of grams that 100 g of water can dissolve at a given temperature is the substance's solubility, and a solution at that limit is a saturated solution. At 20 °C the solubility of sodium chloride is 36 g, which works out to a mass percent of 36 ÷ 136 × 100% ≈ 26.5%. Solubility and mass percent have different denominators — one divides by the solvent, the other by the solution — so the two numbers were never supposed to match.

Put those two ideas together and you can predict the direction of every operation. Add water: solute unchanged, solution heavier, so ω must fall — that is dilution; and because the solute mass stays fixed, you get the dilution formula m₁ω₁ = m₂ω₂. Evaporate water: solute unchanged, solution lighter, ω rises; once saturated, crystals come out in step with the water lost and ω holds at the saturation value. Add solute: below saturation ω rises; past saturation everything extra piles up on the bottom and ω doesn't move. Heat it: most solids dissolve better when warm, so bottom crystals dissolve and ω rises. The simulation above exists precisely so students can press out these four rules one by one.

How sharply solubility responds to temperature varies enormously between substances, and that decides how each one is purified. Sodium chloride's solubility creeps only from 36 g at 20 °C to about 38 g at 80 °C — almost temperature-blind — so getting salt from seawater has to rely on evaporation crystallization. Potassium nitrate and copper sulfate, whose solubility climbs steeply with temperature, can instead be dissolved into a concentrated hot solution and then pulled out by cooling crystallization.

Tips for teachers and parents

• Timing and grouping: 30 minutes, in pairs; the scale and graduated cylinders are shared between groups.
• Before hands go on: no chemical is ever tasted.

The step most often skipped. Have different groups prepare the same concentration but with different amounts of water, then compare the ω values at the end — one comparison settles "mass percent doesn't depend on how much you make." Grade the calculation and the hands-on technique separately.

How to know it has really landed. A student meets the bar when they can use ω = mass of solute ÷ mass of solution × 100% correctly and predict whether a given slip in procedure pushes the concentration high or low.

Going further, and common misconceptions

• Misconception 1: "50 g of a 6% solution means 3 g of salt plus 50 g of water" — wrong. The water should be 47 g; otherwise the solution weighs 53 g and the concentration is only 5.7%.
• Misconception 2: "The undissolved salt at the bottom still counts as solute" — it doesn't. It enters neither the numerator nor the denominator; this is where the most marks are lost.
• Misconception 3: "A saturated solution is always more concentrated than an unsaturated one" — not necessarily. Saturated potassium permanganate solution at 20 °C is only about 6% — far weaker than an unsaturated 20% salt solution.
• Misconception 4: "Stirring makes more dissolve" — it doesn't. Stirring only speeds up the rate of dissolving; the ceiling is set by solubility.
• Misconception 5: "It doesn't matter whether you look down or up at the cylinder" — it does. Looking down reads high and delivers less water, so the solution comes out too concentrated; looking up does the opposite.
• Misconception 6: "The solvent is always water" — not always. Iodine barely dissolves in water but dissolves readily in alcohol, and then the alcohol is the solvent.
• Work it out: to dilute the finished 50 g of 6% salt water down to 3%, how much water must be added? By m₁ω₁ = m₂ω₂ the solution must become 100 g — so add another 50 g of water.
• To see what the crystals look like when a saturated solution dries out, read on in Saturated Solutions and Evaporation Crystallization.

When a student can look at any single step and say whether it pushes the concentration high or low, this experiment is passed.