# Acids, Bases, and Solutions, Starting From a Pot of Purple Cabbage URL: https://sciencestruck.com/chemistry/acids-bases-solutions Category: Chemistry Published: 2026-08-22T11:03:59 Updated: 2026-08-22T11:03:59 Image: https://sciencestruck.com/_astro/acids-bases-solutions.0BgJcR7L_neNR1.webp Boil a few leaves of red cabbage, strain off the purple liquid, and pour it into three glasses. Splash vinegar into the first and it goes pink. Stir a spoonful of [baking soda](https://sciencestruck.com/chemistry/chemical-formula-for-baking-soda) into the second and it drifts toward green, then blue. The third stays purple. Nothing was created or destroyed in any of those glasses. One very small particle, a single proton, moved from one molecule to another, and the cabbage pigment reported where it went. That transfer is the entire subject: an acid hands protons out, a base takes them in, and a solution is the water that lets the two find each other. Everything else on this page is a consequence of that one idea. The pH scale is a way of counting the loose protons. Strength describes how willingly a substance lets go of them. Neutralization is what happens when a donor meets an acceptor. Buffers are chemical shock absorbers that keep the count steady. Start with the cabbage and the rest follows. ## What actually makes something an acid or a base? An acid donates protons (hydrogen ions, written H+); a base accepts them. That definition, worked out independently by Johannes Nicolaus Brønsted and Thomas Lowry in 1923, is the one that survives contact with real chemistry, and it is the one worth memorizing. The older description came from Svante Arrhenius in the 1880s: an acid is something that releases hydrogen ions into water, and a base is something that releases hydroxide ions (OH\-). It works beautifully for the two classroom standards. Hydrochloric acid, HCl, dropped into water, splits into H+ and Cl\-. [Sodium hydroxide](https://sciencestruck.com/chemistry/difference-between-potassium-hydroxide-sodium-hydroxide), NaOH, splits into Na+ and OH\-. Two clean equations, and every introductory chapter still opens with them for good reason. Then a student asks about ammonia. NH3 has no hydroxide anywhere in its formula, and yet dissolve it in water and the solution is unmistakably basic. Under the Arrhenius rule that should be impossible. Under the proton rule it is obvious: ammonia pulls a proton off a water molecule to become NH4+, and the leftover OH\- is what makes the solution basic. The bottle’s label lists nothing but nitrogen and hydrogen, and the solution still turns cabbage juice blue, because the hydroxide was assembled in the glass rather than poured in. That example also exposes the most useful habit in this whole topic: acids and bases come in pairs. When HCl gives up its proton, what remains (Cl\-) is technically a base, because it could in principle take that proton back. Chemists call it the conjugate base. Ammonia’s partner, NH4+, is its conjugate acid. Water plays both roles depending on the company it keeps, which is why it turns up on both sides of so many equations. You may also meet the Lewis definition, which widens the net to electron pairs rather than protons and catches reactions that never involve hydrogen at all. Keep it in your back pocket. For nearly everything in a kitchen, a garden, or a first-year lab, proton transfer is the story. Sour, slippery, and corrosive all sit downstream of proton transfer, which is why none of them appears in the definitions above. Treat them as symptoms rather than as tests, because plenty of things are sour without being useful acids, and because a lab is the one place where tasting a sample to identify it has never been an acceptable method. Your tongue carries receptors that fire when loose protons arrive, so an acid announces itself on the palate before you can name it, and the slippery feel of a base is the same kind of clue: it is dissolving the oils on your skin into soap. ## What does the pH scale really measure? pH counts hydrogen ions, on a logarithmic scale, so every whole-number step means a factor of ten. Written out, pH is the negative base-ten logarithm of the hydrogen ion concentration. Søren Sørensen introduced it in 1909 while working on beer at the Carlsberg Laboratory, which tells you something about where practical chemistry problems come from. The logarithm is there because the numbers are otherwise unmanageable. Pure water at 25 °C contains hydrogen ions at a concentration of 0.0000001 moles per liter, which is 10\-7. Nobody wants to write that on a lab notebook forty times. Take the negative logarithm and you get 7. Stomach acid, roughly ten million times more concentrated in hydrogen ions, becomes a tidy number between 1.5 and 3.5 instead of a decimal with a runway of zeros. The practical consequence is the part that trips people up. A solution at pH 4 holds ten times the hydrogen ions of one at pH 5, and ten times again at pH 3. Between pH 2 and pH 6 there is a factor of ten thousand. Drop a stream from pH 6.5 to 5.5 and the acidity has multiplied by ten; the fish notice. A rough ladder, from most acidic to most basic: - Battery acid and other concentrated strong acids: 0 to 1 - Stomach contents: about 1.5 to 3.5 - Lemon juice, roughly 2; household vinegar, a little under 3 - Black coffee and tomatoes: around 5 - Pure water at room temperature: 7, the neutral point - Blood plasma: tightly held between 7.35 and 7.45 - Seawater: about 8.1 - Baking soda dissolved in water: around 8 to 9 - Household ammonia: roughly 11 - Lye-based drain and oven cleaners: 13 and up Two honest footnotes that most summaries skip. First, the scale is not fenced in at 0 and 14; concentrated solutions can sit outside that range, and negative pH values are real. The 0-to-14 window is a convenience, not a law. Second, neutral does not always mean 7. The pH of neutrality depends on temperature, because water’s own tendency to split into ions increases as it warms. Water at 60 °C is neutral at a pH near 6.5, because neutrality means the hydrogen and hydroxide counts still match each other, and at that temperature both counts have risen together. The decibel comparison comes up every time a class meets this scale, since both compress enormous ranges into small numbers, and it gets you about eighty percent of the way there before it quits. The two scales part company over what they are anchored to: decibels are calibrated against how loudness reaches an ear, while pH counts particles in a liquid and would read exactly the same in a sealed jar nobody ever opened. Borrow the compression from the analogy and leave the rest of it on the shelf. ## Is a “strong” acid the same as a concentrated one? No, and this is the single most common mix-up in the whole subject. Strength describes how completely a substance breaks apart in water. Concentration describes how much of it you dissolved. They are independent, and a solution can be any combination of the two. Drop hydrochloric acid into water and essentially every molecule surrenders its proton. That is what strong means: complete ionization. Drop acetic acid, the sour part of vinegar, into water and the great majority of molecules hold on to their hydrogens. In a 0.1-molar solution, well under two percent of the acetic acid molecules have let go at any given instant. The rest are sitting there intact, in constant exchange with the few that have dissociated. Line those two up at the same concentration and the difference is loud. A 0.1-molar hydrochloric acid solution lands near pH 1. A 0.1-molar acetic acid solution sits near pH 2.9, roughly eighty times fewer loose protons, from the same amount of dissolved acid. Same concentration, different strength. Now run it the other way. Dilute that hydrochloric acid a thousandfold and you have a strong acid at pH 4, milder than vinegar. Glacial acetic acid, meanwhile, is a weak acid at very high concentration and will burn skin on contact. The four combinations all exist: - Strong and concentrated: lab-grade hydrochloric acid, sulfuric acid in a car battery - Strong and dilute: a few drops of HCl in a liter of water, gentler than orange juice - Weak and concentrated: glacial acetic acid, genuinely hazardous - Weak and dilute: the vinegar in your cupboard Chemists put a number on strength with the acid dissociation constant, Ka, or its logarithmic cousin pKa, which compresses those values the way pH compresses ion counts. Smaller pKa, stronger acid. Converting between the two is a logarithm in one direction and an exponent back in the other. A borderline case gets settled by that number rather than by argument: [formic acid](https://sciencestruck.com/chemistry/formic-acid-strong-or-weak), with a pKa near 3.75, is a good deal stronger than the acetic acid in vinegar at about 4.76, and still nowhere near hydrochloric acid, which sits below zero. The same split applies on the basic side. Sodium hydroxide and potassium hydroxide dissociate completely and are the strong bases you meet first. On paper they behave identically, and on a bench they do not, which is why soapmakers reach for potassium hydroxide when they want a soft or liquid soap and sodium hydroxide when they want a bar that holds its shape. Ammonia and the amines are [weak bases](https://sciencestruck.com/chemistry/weak-bases), accepting protons only partially, which is why an ammonia solution and a lye solution behave nothing alike despite both turning cabbage juice blue. One warning that has to go here rather than at the bottom of the page, because students consistently draw the wrong conclusion: weak does not mean safe. Hydrofluoric acid is a weak acid by the strict chemical definition and one of the most dangerous liquids in any building that stocks it. Strength describes how completely a substance ionizes, and the safety data sheet is what tells you how to handle the bottle. Read both, and give the second one authority over your gloves. ## What happens when an acid and a base meet? They cancel each other out and leave behind a salt and water. That is the acid base neutralization reaction, and in its cleanest form it fits on one line: > HCl + NaOH → NaCl + H2O Hydrochloric acid plus sodium hydroxide gives sodium chloride and water. Both starting materials were aggressive; the products are table salt and water. The proton the acid was carrying met the hydroxide the base was carrying, and the two combined into H2O. The sodium and chloride ions were spectators, and they end up paired as the salt. “Salt” in chemistry names a whole category: any compound built from the leftover ions of an acid and a base. Combine sulfuric acid with sodium hydroxide and the salt is sodium sulfate, a reaction worth watching closely because sulfuric acid can give up two protons rather than one. Combine nitric acid with potassium hydroxide and you get potassium nitrate. Same pattern, different cast. The reaction also releases heat, and more of it than beginners expect. Mixing concentrated acid and concentrated base releases enough energy to bring the liquid to a boil and throw it back out of the flask, which is a great deal of violence for a reaction whose products are salt and water. This is why every neutralization in a school lab happens with dilute solutions, slowly, with the flask on the bench rather than in your hand. The kitchen version is friendlier and slightly more interesting than it looks. Pour vinegar onto baking soda and the acetic acid hands a proton to the bicarbonate ion. Baking soda’s chemical formula, NaHCO3, tells you the hand-off is coming: there is a hydrogen sitting in the middle of it, and a carbonate waiting to take one more. That produces sodium acetate, water, and carbonic acid, and carbonic acid is unstable enough that it immediately falls apart into water and carbon dioxide gas. The foam is the CO2 escaping. When the fizzing stops and you evaporate the liquid, the white residue left behind is sodium acetate: a salt, made exactly the way the equation says. Swap in [washing soda](https://sciencestruck.com/chemistry/washing-soda-vs-baking-soda), sodium carbonate, and the reaction still finishes in carbon dioxide, but the solution you started from was far more alkaline, because a carbonate carries no hydrogen of its own to hand over and pulls protons off water instead. Two refinements worth carrying forward. First, neutralization is counted in moles, not in volumes or spoonfuls. Matching a cup of acid with a cup of base means nothing until you know the concentrations of both. That mole-counting is the whole basis of titration. When the endpoint refuses to show itself, say the sample is a solid that dissolves too slowly to watch, you add a known excess of one reagent and titrate what survives. That is [back titration](https://sciencestruck.com/chemistry/back-titration), and it is the standard move for the awkward cases. Second, an exactly neutralized solution is not always at pH 7. Neutralize acetic acid with sodium hydroxide and the sodium acetate left behind is itself mildly basic, so the finished solution lands above 7. Only a strong acid paired with a strong base gives you a neutral salt. Whenever a student’s titration result reads 8.7 and they assume they have made an error, that is often the chemistry behaving correctly. ## How do chemists put a number on a solution? With molarity: moles of dissolved substance per liter of solution. Written as M, so a 2 M solution has two moles in every liter. A mole is a counting unit, 6.022 × 1023 particles, in the same way a dozen is a counting unit. Chemists count in moles because reactions happen particle by particle. When the equation says one molecule of HCl neutralizes one unit of NaOH, that ratio holds at the scale of a beaker only if you are counting particles rather than weighing scoops. Sodium hydroxide weighs about 40 grams per mole and sodium chloride about 58.4, so equal masses of the two contain very different numbers of particles. The definition contains a trap that has cost more lab points than any other detail on this page. Molarity is moles per liter of finished solution, not per liter of water you started with. To make one liter of 1 M sodium chloride you dissolve 58.44 grams in some water, then top up to the one-liter mark, because the salt itself takes up room. Both of the classic errors, weighing into a full liter and topping up past the mark, come from reading “per liter” as a description of the water rather than of the finished solution. Molarity is what you will meet in nine problems out of ten. The other measures exist because a liter is not always trustworthy: warm a solution and it expands, so its molarity drops without a single particle leaving the flask. Molality counts moles per kilogram of solvent instead, and a kilogram weighs the same at 4 °C as it does at 80 °C. That is the whole reason freezing-point and boiling-point work is done in molality. Those calculations run across a temperature change, so a volume-based number would be quietly wrong at both ends of the problem. Percent by mass is the one printed on things sold to the public. Household vinegar is typically around five percent acetic acid, stated that way because a shopper can weigh a bottle and cannot count moles. Biology needs yet another count, and the reason is one line of arithmetic. One mole of glucose dissolves into one mole of particles; one mole of sodium chloride dissolves into two, a sodium ion and a chloride ion, and a cell membrane responds to the particle count rather than the compound count. Osmolarity is that particle count, per liter rather than per compound. Dissolve 0.1 moles of glucose in a liter and you have a 0.1 osmolar solution; dissolve 0.1 moles of table salt and you have a 0.2 osmolar one, twice the pull on water across a membrane from the same single scoop of moles. Osmolality makes the molality swap all over again, per kilogram of solvent, for the same temperature reason: a sample gets warmed, cooled, and read hours apart, and a kilogram holds still while a liter does not. Tonicity asks the further question of what those particles do to one particular cell, which is physiology rather than arithmetic. It is the only term on this list you cannot get from a balance and a flask. Diluting a solution is the other half of this arithmetic, and it is the operation you will actually perform most often, since stockrooms sell concentrates. Two lines cover both jobs, and the second one falls out of the first. The definition: > M = n ÷ V M is the molarity in moles per liter, n is the number of moles of dissolved substance, and V is the volume of finished solution in liters. Weigh out 58.44 grams of sodium chloride, which is one mole, make it up to 0.500 liters, and M = 1 ÷ 0.5 = 2 M. Push real numbers through that once by hand and dividing moles by liters stops feeling like a rule to memorize and starts feeling like counting things. The dilution relationship follows, because adding water changes V and leaves n exactly as it was: > M1V1 = M2V2 The 1s describe the concentrated stock you started with and the 2s describe what you ended up with. Each side of the equals sign is the same fixed mole count, written two different ways, so any three of the four numbers give you the fourth by one rearrangement. Take 25 mL of that 2 M stock up to 250 mL and M2 = (2 × 25) ÷ 250 = 0.2 M: ten times the volume, one tenth the concentration, the same moles the whole way down. Milliliters and liters mixed in a single line is the most common way a correct method still produces a wrong answer, and it costs a point every year. What the number is telling you never changes: how many particles are waiting in every liter you pour. ## What keeps a solution’s pH from swinging wildly? A buffer: a weak acid and its conjugate base sitting in the same solution in comparable amounts, each waiting to intercept whatever gets added. The mechanism is genuinely simple once you have the conjugate pair idea. Add acid to a buffered solution and the incoming protons are absorbed by the conjugate base before they can float free. Add base and the weak acid donates protons to neutralize the incoming hydroxide. Either way the population of loose hydrogen ions barely moves, and the pH holds. Take an unbuffered glass of pure water and a single drop of strong acid will send it from 7 to somewhere near 3. Do the same to a well-made buffer and the reading may not move by a tenth of a unit. A mixture of acetic acid and sodium acetate is the classic teaching example, and if you have ever made the vinegar-and-baking-soda reaction and stopped partway through, you have made one by accident. Both partners are present at once, which is exactly the condition a buffer requires. Blood plasma is the example everyone remembers, and rightly so. Human blood is held between 7.35 and 7.45, a window tight enough that each direction of departure has its own name, acidosis when the blood pH falls below that band and alkalosis when it climbs above. The main chemical mechanism is the carbonic acid and bicarbonate pair, dissolved carbon dioxide on one side and bicarbonate ion on the other, trading protons back and forth. What makes it more capable than a beaker of buffer is that the body can adjust both partners: breathing rate changes the carbon dioxide side within seconds, and the kidneys manage the bicarbonate side over hours. A rise in dissolved carbon dioxide pushes the blood pH down, a fall pushes it up, and it is the same conjugate pair arithmetic as in the beaker, running on a timescale of seconds. Buffers do have limits, and skipping over this is how people end up trusting one too far. A buffer works best when the two partners are present in roughly equal amounts and the target pH sits close to the weak acid’s pKa. Push enough acid or base into it and one partner is consumed entirely, after which the buffer has nothing left to trade and the pH falls off a cliff. Chemists call that quantity the buffer capacity, and it is finite in every case. Once you know the pattern you start seeing it on labels. Shampoo formulated to a specific pH, swimming pool chemistry that fights back when you shock it, the sodium bicarbonate in an aquarium, the buffered growth media in every biology lab. All the same trick: keep both halves of a conjugate pair on hand and let them absorb whatever arrives. ## How can you tell an acid from a base without a meter? With an indicator, a dye whose color depends on how many protons are floating around it. The indicator molecule is itself a weak acid, and its protonated and deprotonated forms absorb different wavelengths of light, so gaining or losing a proton literally changes what color the molecule is. What you are watching in the glass is a molecule changing identity, one proton at a time. Litmus is the oldest of the household names, extracted from lichens and soaked into paper. Red in acid, blue in base, and nothing more. It answers one question, which side of neutral you are on, and refuses to answer anything else. Perfectly adequate when that is the question. Red cabbage does considerably better, because the anthocyanin pigments in it shift through several distinct forms across the range rather than flipping between two. That is why your three glasses gave three answers instead of a yes and a no. Approximate pH Red cabbage juice Litmus Universal indicator 2 and below bright red to pink red red 4 pink to light purple red orange 7 purple neutral, no strong change green 9 blue to blue-green blue blue 12 and above green to yellow blue violet Universal indicator, the one that comes as a bottle of dropper solution or as strips, is a deliberate blend of several dyes chosen so that their individual transitions overlap into a continuous rainbow. That gives you a number rather than a direction, accurate to roughly a pH unit if you read it against the chart in decent light. Know what indicators cannot do. A colored or cloudy sample drowns the signal, which is why they are useless on coffee and hopeless on soil until you have settled and filtered it. Homemade cabbage juice fades within a few days in the fridge and loses its range. And an indicator gives you a band, never a decimal place. When you need a real number, a pH meter reads the electrical potential across a thin glass membrane that responds to hydrogen ion activity, and it has to be calibrated against known standards before it can be trusted. Your glass of cabbage juice is running a crude version of what a laboratory spectrophotometer does on purpose: light goes through the sample, something measures how much of it survives the trip, and the color that comes out the far side gets read back as a concentration. The instrument uses a photodetector and a set of known solutions to calibrate against. You are using an eyeball and a printed chart, which is the same procedure with worse optics and somebody else’s calibration. That comparison is also the honest limit of the method, and the reason it is still worth doing. A color you match by eye is a measurement, good to roughly one pH unit, and knowing which digit you are entitled to claim is most of what separates a result from a guess. For a kettle, a fish tank, or a flower bed, one pH unit is the answer you needed anyway. ## Where do acids and bases turn up outside the classroom? Everywhere, so the list below groups them by what the chemistry is being asked to do in each case: dissolve something, preserve something, hold a plant’s nutrients within reach. ### Cleaning, where the target dictates the chemistry Alkaline cleaners attack fats and proteins. Sodium hydroxide in a drain or oven product converts grease into soap in place, which is the same reaction soapmakers have used for centuries, running in your pipes. Acidic cleaners attack mineral deposits instead. Limescale, rust stains, and the chalky ring in a kettle are largely carbonate compounds, and an acid dissolves them by protonating the carbonate. No single cleaner does both jobs well, which is why reaching for the wrong one leaves you scrubbing. I used to send a class home with one question: find the strongest-sounding cleaner under your sink and work out from the label which of those two targets it is built for. They came back convinced the answer was on the front of the bottle. It never is. The front says powerful; the ingredient list says sodium hydroxide or citric acid, and that one word tells you whether the bottle is any use on the kettle at all. A kettle full of limescale will sit there and mock a drain cleaner all afternoon. ### Food, where acid and base are ingredients Lowering pH preserves food, because most spoilage microbes cannot grow in strongly acidic conditions. That is the principle behind pickling, and behind the acidity requirements written into canning guidelines. Acid also does flavor work, cutting richness on the tongue. Bases do texture work: baking soda in a cookie dough raises the pH, encourages browning, and softens the crumb, and the alkaline solutions used for pretzels and for certain noodles change the protein structure at the surface. If you have ever wondered why a home-baked pretzel comes out pale and a bakery one comes out mahogany, that is the whole answer, and it is a pH difference of a couple of units on the surface of the dough. Chemical leavening is neutralization harnessed on purpose, an acid and a bicarbonate meeting in a wet batter and releasing carbon dioxide exactly where you want the bubbles. ### Soil and water, where pH controls availability Soil pH works on a plant at one remove, deciding which nutrients stay dissolved and reachable by roots, which is why blueberries insist on acidic ground. Then there is the hydrangea, a living indicator growing in somebody’s front yard. Pink flowers mean the soil is closer to neutral. Blue flowers mean it is acidic, because aluminum only stays dissolved and available to the roots at low pH, and the plant needs that aluminum to make the blue pigment. The flower color is a reading of the ground the bush stands in, taken over a season rather than a second, and it depends on the aluminum being there in the first place: acidify a bed of soil that holds none and you will get pink flowers in thoroughly acidic dirt. Walk a street in June and you can read the front gardens. Growers work that dial on purpose: lime, a carbonate base, raises soil pH, and sulfur or an acidic amendment lowers it. Change the soil and you change the flower, though the bush takes its own time about agreeing. Municipal water treatment adjusts pH for a plainer reason: water that is too acidic corrodes the pipes it travels through. ### Industry and the body A lead-acid car battery is the neatest chemistry lesson in any driveway. Discharging it consumes the sulfuric acid: the acid leaves the solution and ends up on the plates as lead sulfate, so the liquid that remains is genuinely lighter than it was when the battery was full. Float a hydrometer in a cell and the depth it settles to reports that density, around 1.27 in a fully charged cell and closer to 1.12 in a flat one, which is how a mechanic reads state of charge without touching a meter. It is the same instrument a brewer uses to watch sugar disappear from a fermentation, and it costs about as much as a sandwich, which made it my favorite thing to hand a class that had just learned what concentration means. [Phosphoric acid](https://sciencestruck.com/chemistry/phosphoric-acid-structure-chemical-formula-properties) holds down two jobs that sound unrelated until you look at the mechanism. It gives colas their sharp edge, and it turns rust into a stable phosphate layer that paint will grip. Two facts about the molecule cover both jobs. It carries three acidic hydrogens and surrenders them one at a time, which is why the tartness in a cola is measured and mild rather than a single sharp hit. And the phosphate left behind grips iron hard enough to lock it in place, turning a flaking red oxide into a dull grey layer that stays where it is and gives paint something to hold. H3PO4: three acidic hydrogens and one phosphate, and both jobs are sitting right there in the formula. Your stomach keeps itself somewhere between pH 1.5 and 3.5 to break protein apart, and the antacids sold against it are bases, usually carbonates, that neutralize a portion of that acid and give off carbon dioxide while doing it. That is the vinegar and baking soda reaction from earlier on this page, running in a warmer room, which is also why the relief arrives with a burp. Carbonates differ in how fast they go to work, and speed costs duration: a compound that neutralizes stomach acid quickly is also spent quickly, which is why one box on the shelf promises fast relief and the box beside it promises long relief. Which of them suits a particular person is a question for a pharmacist rather than a chemistry page. ## What should you know before handling any of this yourself? Read this section before you mix anything, including the cabbage demonstration at the top of the page. None of what follows is medical advice; it is the conservative handling baseline. - **Eye protection, always.** Splash goggles, not reading glasses, for anything stronger than kitchen vinegar. Eyes do not recover from chemical burns the way skin does. - **Acid into water, never the reverse.** Diluting concentrated acid releases a lot of heat, and if you pour water onto acid that heat concentrates at the surface and can spit the mixture back out. Add the acid slowly to a larger volume of water, stirring. The old classroom line is “do as you oughta, add acid to water,” and it has stuck because it works. - **Never mix bleach with ammonia, and never mix bleach with an acid.** The first combination produces chloramine vapors, the second releases chlorine gas, and both are genuinely dangerous in an ordinary kitchen or bathroom. As a rule, do not combine household cleaning products at all. - **Ventilate.** Open a window or work outdoors for anything that smells sharp. If a demonstration calls for a fume hood, it belongs in a school lab with a teacher present, and no home substitute is adequate. - **Adults with children, every time.** Cabbage juice, vinegar, baking soda, and lemon juice are fine for a supervised kitchen project. Drain cleaner, pool acid, and anything from a hardware store are not household experiment materials at any age. - **If something contacts skin or eyes,** flush immediately with plenty of cool running water and keep flushing while someone seeks medical help. Do not try to neutralize a chemical burn on skin with another chemical; that adds heat to an injury. - **Label everything and never use food containers.** A solution in an unmarked drinking glass is an accident with a delay built in. The cabbage juice, meanwhile, is waiting in the fridge and good for a few days. Line up a row of clear glasses and work along a shelf: lemon juice, vinegar, tap water, a dissolved antacid tablet, baking soda solution, a drop of dish soap. Match each color against the table above and you will have built the pH scale by hand, in the right order, from things you already own. Then leave a clean jar out in the next rain and test what lands in it against the tap water beside it. Rain that has fallen through nothing but clean air still arrives mildly acidic, near pH 5.6, because carbon dioxide dissolves into the drop on the way down and makes carbonic acid inside it. That is the same conjugate pair doing the same trade as in the buffer section, assembled midair, in a jar on your step.