Acids, Bases and Salts

A grandmother reaching for lemon juice, vinegar, or a spoon of baking soda after a heavy meal is doing real chemistry without knowing the words for it. She is choosing between two families of substances — acids and bases — that cancel each other out, and picking the base (baking soda solution) to calm the acid churning in an overfull stomach. That single kitchen instinct is the doorway into this entire chapter: acids and bases are not just "sour" and "bitter" tastes, they are substances defined by the ions they release in water, and the reaction between them, neutralisation, is one of the most useful ideas in all of chemistry. This chapter builds the whole picture from the ground up. It starts with how we detect acids and bases without tasting them — litmus, phenolphthalein, methyl orange, and even the smell of onion or vanilla changing in their presence. It then works through how acids and bases behave with metals, with metal carbonates, with each other, and with metal and non-metal oxides, always circling back to one repeated pattern: acids donate hydrogen ions, bases donate hydroxide ions, and when the two meet, they form a salt and water. From there the chapter asks a sharper question — how strong is a given acid or base — and answers it with the pH scale, a single number from 0 to 14 that quietly governs digestion, tooth decay, farming, and the survival of fish in a river. Finally it turns to salts themselves: how ordinary table salt becomes the raw material for sodium hydroxide, bleaching powder, baking soda, and washing soda, and why some crystals that look bone dry are secretly holding water inside their structure. By the end, you should be able to look at any everyday acid-base situation — an ant bite, a rusting brass vessel, a doctor's plaster cast, a fizzy antacid tablet — and explain it in the language of ions, pH, and neutralisation.

Key takeaways

Acids taste sour and turn blue litmus red; bases taste bitter, feel soapy, and turn red litmus blue; phenolphthalein and methyl orange give sharper colour changes than taste.
Acids react with active metals to release hydrogen gas and form a salt: Metal + Acid gives Salt + Hydrogen gas, confirmed by the pop sound test.
Metal carbonates and hydrogencarbonates react with acids to give a salt, water, and carbon dioxide gas, which turns lime water milky.
Acids and bases neutralise each other: Acid + Base gives Salt + Water, because H-plus(aq) and OH-minus(aq) ions combine into neutral water.
Metallic oxides are basic oxides (react with acids like a base); non-metallic oxides are acidic oxides (react with bases like an acid).
Acidic behaviour appears only when H-plus ions are freed in the presence of water; dry HCl gas alone does not turn dry litmus paper red.
The pH scale (0 to 14) measures hydrogen ion concentration: below 7 is acidic, exactly 7 is neutral, above 7 is basic; lower pH means more H-plus ions.
Common salt is the raw material, via the chlor-alkali process, for sodium hydroxide, bleaching powder, baking soda, and washing soda; some salts carry a fixed water of crystallisation, as in copper sulphate and gypsum.
Key idea: acids and bases are defined by the ions they release in water (H-plus and OH-minus respectively), so their reactions with metals, carbonates, and each other always follow the same few patterns, and the pH scale simply measures how much of that defining ion is present.

The core insight

Acids release H-plus(aq) ions and bases release OH-minus(aq) ions in water, so every reaction in this chapter — with metals, carbonates, each other in neutralisation, and the pH scale measuring their strength — is the story of these two ions meeting, being produced, or being counted.

What this chapter covers

The chapter moves in four connected stages, and this article follows the same order. Section 2.1 shows how to recognise acids and bases using indicators, then walks through their five signature reactions: with metals, with metal carbonates and hydrogencarbonates, with each other (neutralisation), and with metallic and non-metallic oxides. Section 2.2 asks what makes something acidic or basic at the ion level, using a conductivity experiment to show H-plus and OH-minus ions are the common thread, and that water is essential for this ionisation. Section 2.3 introduces the pH scale to measure the strength of that ion concentration, and surveys where pH matters daily — digestion, tooth decay, soil, acid rain. Section 2.4 zooms in on salts: their families, their own pH, the industrial chemicals derived from common salt, and why some "dry" crystals secretly contain fixed water.

The core ideas

Detecting acids and bases: indicators

An indicator is a substance that changes colour (or, in a few cases, smell) depending on whether it is placed in an acidic or a basic solution. Litmus, extracted from lichen, is the most familiar one: it turns red in acid and blue in base, staying purple in a neutral solution. Synthetic indicators give sharper changes — phenolphthalein is colourless in acid and pink in base; methyl orange is red-orange in acid and yellow in base. A few substances change smell instead of colour and are called olfactory indicators; onion juice and clove oil lose their characteristic smell in base but keep it in acid, and vanilla essence loses its smell in base too.

How acids and bases react with metals

When an acid reacts with a reactive metal, hydrogen gas is released and a salt is formed. Zinc granules dropped into dilute sulphuric acid fizz visibly, and the gas collected turns a lit matchstick out with a small pop, confirming hydrogen.

\text{Metal} + \text{Acid} \rightarrow \text{Salt} + \text{Hydrogen gas}

where: the metal must be reactive enough to displace hydrogen (copper does not react this way with dilute acids). A few bases also release hydrogen with certain metals — zinc reacts with hot sodium hydroxide to give sodium zincate and hydrogen, $2\text{NaOH(aq)} + \text{Zn(s)} \rightarrow \text{Na}_2\text{ZnO}_2\text{(s)} + \text{H}_2\text{(g)}$ — but this is far less general than the acid-metal reaction.

How acids react with metal carbonates and hydrogencarbonates

Sodium carbonate and sodium hydrogencarbonate both fizz vigorously when dilute hydrochloric acid is added, releasing a gas that turns lime water milky — the standard test for carbon dioxide.

\text{Metal carbonate/Metal hydrogencarbonate} + \text{Acid} \rightarrow \text{Salt} + \text{Carbon dioxide} + \text{Water}

where: this pattern holds for every metal carbonate or hydrogencarbonate. Passing excess carbon dioxide through the milky lime water clears it again, since the insoluble calcium carbonate precipitate converts to soluble calcium hydrogencarbonate, $\text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)} \rightarrow \text{Ca(HCO}_3)_2\text{(aq)}$ .

Neutralisation: acids and bases reacting with each other

Adding dilute hydrochloric acid, drop by drop, to sodium hydroxide solution containing phenolphthalein makes the pink colour fade to colourless, showing the base has been used up.

\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}

where: this is a neutralisation reaction, and it is why an antacid (a mild base) relieves the pain of excess stomach acid. At the ion level, since acids give H-plus(aq) and bases give OH-minus(aq), the reaction is simply $\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$ , with the leftover ions combining to form the salt.

Metallic and non-metallic oxides

Copper oxide dissolves in dilute hydrochloric acid to give a blue-green solution of copper(II) chloride, behaving exactly like a base reacting with an acid — so metallic oxides are classified as basic oxides. Carbon dioxide, a non-metallic oxide, reacts with calcium hydroxide (lime water) to give a salt and water, the same pattern as an acid reacting with a base — so non-metallic oxides are classified as acidic oxides.

\text{Metal oxide} + \text{Acid} \rightarrow \text{Salt} + \text{Water}

What acids and bases have in common: ions in water

A simple conductivity test explains why acids and bases behave the way they do. Dip two electrodes connected to a battery and bulb into dilute hydrochloric acid, and the bulb glows brightly — the acid solution conducts electricity. Glucose and alcohol solutions, which also contain hydrogen atoms, do not make the bulb glow at all. The difference is that acids release H-plus ions (written H-plus(aq), or the hydronium ion H₃O⁺, since a bare proton cannot exist alone in water) that carry current through the solution, while glucose and alcohol do not ionise. Bases behave the same way with hydroxide ions: $\text{NaOH(s)} \xrightarrow{\text{H}_2\text{O}} \text{Na}^+(aq) + \text{OH}^-(aq)$ . This ionisation needs water — dry HCl gas does not turn dry litmus paper red, because H-plus can only separate from chlorine in the presence of water, $\text{HCl} + \text{H}_2\text{O} \rightarrow \text{H}_3\text{O}^+ + \text{Cl}^-$ . Water-soluble bases are called alkali; every alkali is a base, but not every base (insoluble copper hydroxide, for example) is an alkali. Mixing a concentrated acid or base into water is strongly exothermic, which is why acid must always be added to water — never water to concentrated acid — since local heating near a small volume of acid can make it splash and cause burns.

Figure

Electrical conductivity apparatus showing a beaker of dilute hydrochloric acid with two electrodes connected through a battery a switch and a bulb which glows to show that ions in the acid solution carry the electric current

The pH scale: measuring the strength of acids and bases

Even after knowing something is acidic or basic, we often need to know how acidic or basic it is. The pH scale does this using a universal indicator, a mixture of dyes that turns a different colour at every hydrogen ion concentration, matched against a reference chart.

\text{pH} < 7 \Rightarrow \text{acidic}, \qquad \text{pH} = 7 \Rightarrow \text{neutral}, \qquad \text{pH} > 7 \Rightarrow \text{basic}

where: the scale runs from 0 (strongly acidic) to 14 (strongly basic), and a higher hydrogen ion concentration always corresponds to a lower pH value. Acids that ionise almost completely, giving many H-plus ions, are called strong acids (hydrochloric acid); acids that ionise only partly, giving fewer H-plus ions at the same concentration, are called weak acids (acetic acid, or vinegar). The same distinction of strong and weak applies to bases based on how completely they release OH-minus ions. pH governs an enormous amount of everyday life. The human body works only within pH 7.0 to 7.8, so even a small pH shift in blood is dangerous. Rain water below pH 5.6 is called acid rain, and it damages aquatic life by lowering river and lake pH. The stomach produces hydrochloric acid to digest food, but excess acid causes indigestion, relieved by a mild base called an antacid — commonly magnesium hydroxide — which neutralises the surplus. Tooth enamel starts to corrode once mouth pH falls below 5.5, because bacteria convert leftover sugar into acid; basic toothpaste neutralises this and helps prevent decay. Farmers test soil pH before choosing crops and treat overly acidic soil with quicklime, slaked lime, or chalk.

Figure

The pH scale from zero to fourteen coloured from red for strongly acidic through green for neutral to deep blue for strongly basic with common substances marked at their approximate positions including lemon juice tomato juice pure water baking soda solution and sodium hydroxide solution

Salts: families, pH, and chemicals from common salt

A salt is the product formed when an acid reacts with a base. Salts sharing a common positive or negative part belong to the same family — sodium chloride and sodium sulphate are both sodium salts, while sodium chloride and potassium chloride are both chloride salts. A salt's own pH depends on which acid and base formed it: salts of a strong acid and a strong base (sodium chloride) are neutral, pH 7; salts of a strong acid and a weak base are acidic, pH less than 7; salts of a strong base and a weak acid, such as sodium hydrogencarbonate, are basic, pH greater than 7, and can neutralise excess acid. Ordinary sodium chloride, mined as rock salt or extracted from seawater, is the raw material for several everyday chemicals through the chlor-alkali process: electrolysing brine gives sodium hydroxide, chlorine gas, and hydrogen gas together. From this come bleaching powder (chlorine reacted with slaked lime), baking soda (NaHCO₃, used in cooking and antacids), and washing soda (sodium carbonate, used for cleaning and softening hard water). Some salt crystals that look dry actually hold a fixed number of water molecules locked into their structure, called water of crystallisation — blue copper sulphate, $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ , turns white on heating as this water is driven off, and blue again when moistened. Heating gypsum, $\text{CaSO}_4 \cdot 2\text{H}_2\text{O}$ , gives Plaster of Paris, $\text{CaSO}_4 \cdot \tfrac{1}{2}\text{H}_2\text{O}$ , which sets back into hard gypsum on mixing with water — why it supports fractured bones and must be stored away from moisture.

Derivation AWhy acids and bases neutralise each other: the ionic pathway

The goal of this derivation is to show, using only the ionisation facts already established for acids and bases, exactly why mixing an acid with a base always ends in a salt and water, rather than treating "Acid + Base gives Salt + Water" as a rule to memorise. The chain runs from how an acid and a base separately behave in water, through what happens when their ions meet, to the general pattern covering every acid-base pair in this chapter.

Step 1. Write what the acid alone does in water. Hydrochloric acid ionises completely: $\text{HCl(aq)} \rightarrow \text{H}^+(aq) + \text{Cl}^-(aq)$ . The H-plus(aq) ion is what makes the solution acidic and turns blue litmus red.

Step 2. Write what the base alone does in water. Sodium hydroxide dissociates completely: $\text{NaOH(aq)} \rightarrow \text{Na}^+(aq) + \text{OH}^-(aq)$ . The OH-minus(aq) ion is what makes the solution basic and turns red litmus blue.

Step 3. Mix the two solutions and track every ion present: H-plus, Cl-minus, Na-plus, and OH-minus are all now in the same beaker. The H-plus and OH-minus ions combine immediately into the stable neutral molecule water: $\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$ . This is why the pink phenolphthalein colour disappears in Activity 2.6 — the OH-minus ions responsible for the base's colour and character are being used up.

Step 4. Look at what remains once H-plus and OH-minus have combined: only Na-plus and Cl-minus ions are left, and evaporating the water leaves solid sodium chloride. Generalising for any acid HX and any base MOH gives the pattern seen throughout the chapter, $\text{HX} + \text{MOH} \rightarrow \text{MX} + \text{HOH}$ , or in words, Acid + Base gives Salt + Water — the salt is whatever is left once the defining ions of the acid and base have combined into water.

Go deeper · Derivation B · 1 minScaling a neutralisation: the volume-concentration reasoning behind acid-base titration

Derivation BScaling a neutralisation: the volume-concentration reasoning behind acid-base titration

The goal here is to turn "acid and base neutralise each other" into arithmetic that predicts exactly how much acid is needed to neutralise a given amount of base, using nothing beyond the idea that neutralisation is complete when the H-plus ions supplied exactly match the OH-minus ions supplied. This is the reasoning behind a classic exercise: if 10 mL of an NaOH solution is exactly neutralised by 8 mL of a given HCl solution, how much of that same HCl is needed to neutralise 20 mL of the same NaOH solution?

Step 1. State the neutralisation condition in terms of amount, not formula. For a one-to-one acid and base like HCl and NaOH, the mixture is exactly neutral when the moles of H-plus supplied equal the moles of OH-minus present.

Step 2. Since moles equal concentration multiplied by volume, write this balance as $c_a V_a = c_b V_b$ , where $c_a, V_a$ and $c_b, V_b$ are the concentration and volume of the acid and base used. Because both solutions are unchanged between trials, $c_a$ and $c_b$ stay fixed, so the ratio $V_a/V_b$ is fixed too.

Step 3. Read off that fixed ratio from the given data: 10 mL of base needs 8 mL of acid, so $V_a/V_b = 8/10 = 0.8$ . This does not need the actual concentrations — only that the same two solutions are used every time.

Step 4. Apply the same ratio to the new volume of base: for 20 mL of the same NaOH solution, $V_a = 0.8 \times 20\ \text{mL} = 16\ \text{mL}$ . Doubling the base doubles the acid needed — this is why titration volumes always scale in direct proportion when the two solutions themselves are unchanged.

Go deeper · Derivation C · 1 minFrom common salt to three useful products: the chlor-alkali reasoning chain

Derivation CFrom common salt to three useful products: the chlor-alkali reasoning chain

The goal of this derivation is to trace, step by step, how passing an electric current through ordinary salt solution produces three separate, independently useful chemicals — sodium hydroxide, chlorine gas, and hydrogen gas — from one starting material, and how two of these combine further into bleaching powder and baking soda.

Step 1. Start with brine, a concentrated solution of sodium chloride, containing Na-plus(aq), Cl-minus(aq), and the ions of water itself. Passing electricity through it — electrolysis — forces these ions apart at the two electrodes instead of staying mixed.

Step 2. At the cathode, hydrogen gas is released and sodium hydroxide accumulates nearby; at the anode, chlorine gas bubbles off. The overall reaction is $2\text{NaCl(aq)} + 2\text{H}_2\text{O(l)} \rightarrow 2\text{NaOH(aq)} + \text{Cl}_2\text{(g)} + \text{H}_2\text{(g)}$ — the chlor-alkali process, named for its two main products, chlorine and alkali.

Step 3. Pass the chlorine produced over dry slaked lime: $2\text{Ca(OH)}_2 + 2\text{Cl}_2 \rightarrow \text{Ca(ClO)}_2 + \text{CaCl}_2 + 2\text{H}_2\text{O}$ , giving bleaching powder — used for bleaching cloth and paper pulp, as an industrial oxidising agent, and for disinfecting drinking water.

Step 4. Separately, treat sodium chloride with water, carbon dioxide, and ammonia to get baking soda: $\text{NaCl} + \text{H}_2\text{O} + \text{CO}_2 + \text{NH}_3 \rightarrow \text{NH}_4\text{Cl} + \text{NaHCO}_3$ . Heating baking soda gives sodium carbonate, $2\text{NaHCO}_3 \xrightarrow{\text{heat}} \text{Na}_2\text{CO}_3 + \text{H}_2\text{O} + \text{CO}_2$ , and recrystallising it gives washing soda, $\text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}$ . So common salt branches, through electrolysis and a separate carbonation route, into four household and industrial chemicals.

Worked examples

Board level - Zinc with dilute sulphuric acid: balanced equation and confirming the gas

A student adds a few pieces of zinc granules to about 5 mL of dilute sulphuric acid in a test tube. Write the balanced chemical equation for the reaction, name the gas evolved, and describe how to confirm its identity.

Step 1. Identify the pattern: Metal + Acid gives Salt + Hydrogen gas. The metal is zinc and the acid is sulphuric acid, so the products are zinc sulphate and hydrogen gas.

Step 2. Balance: $\text{Zn(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{ZnSO}_4\text{(aq)} + \text{H}_2\text{(g)}$ — one Zn, two H, one S, four O on each side.

Step 3. Confirm the gas by passing it through soap solution and bringing a burning candle near a bubble. Answer: Zn(s) + H₂SO₄(aq) gives ZnSO₄(aq) + H₂(g); the gas is hydrogen, confirmed by the pop sound when a gas-filled soap bubble meets a flame.

Board level - Scaling a neutralisation: volume of HCl needed for double the NaOH

10 mL of a solution of NaOH is found to be completely neutralised by 8 mL of a given solution of HCl. If 20 mL of the same NaOH solution is taken instead, what volume of the same HCl solution is required to neutralise it?

Step 1. Use the fixed-ratio reasoning above: since the same two solutions are used both times, $V_a/V_b$ stays constant, $8/10 = 0.8$ .

Step 2. Apply it to the new base volume: $V_a = 0.8 \times 20\ \text{mL} = 16\ \text{mL}$ . Answer: 16 mL of the HCl solution is required — doubling the base doubles the acid needed, since both solutions are unchanged.

Board HOTS / foundation JEE-NEET - pH shift when milk turns to curd, and why baking soda delays souring

Fresh milk has a pH of about 6. Bacteria present in milk slowly convert milk sugar into lactic acid, and this souring is what turns milk into curd. (a) How does the pH of milk change as it sets into curd? (b) A milkman adds a very small amount of baking soda to fresh milk before it is sold. Why does this shift the pH from 6 to slightly alkaline, and why does this milk take longer than usual to set as curd?

Step 1. Connect souring to acid production: bacteria converting milk sugar into lactic acid raises the H-plus ion concentration as milk turns to curd, so the pH must fall from about 6 to a more acidic value.

Step 2. Reason about baking soda chemically: sodium hydrogencarbonate is a mild base, so a small amount added to milk at pH 6 neutralises some natural acidity, pushing the pH up past 7 to slightly alkaline.

Step 3. Combine both steps: curdling appears only once bacterial acid brings the pH back down past the souring threshold. Starting from a higher, slightly alkaline pH means more acid — and more time — is needed to fall that far. Answer: pH falls as milk sours into curd; baking soda raises the starting pH above 6, so bacterial souring takes longer to bring it back down to the curdling point. Timing note: about 90 seconds. Recognise immediately that souring equals acid production (so pH must fall), then treat baking soda purely as a mild base raising the starting pH — the marks come from acid-base reasoning, not microbiology.

Common traps

Mixing up strength (how completely an acid or base ionises) with concentration (how much is dissolved) is the single most common slip — a dilute strong acid can still release more H-plus ions than a concentrated weak acid. Students often write a free H-plus ion floating in solution; hydrogen ions cannot exist alone and must always be shown as H-plus(aq) or H₃O⁺. Another frequent error is adding water to a concentrated acid rather than acid to water — the safe order is always acid added slowly to water. Not every base is an alkali: an alkali is specifically a base that dissolves in water. On the pH scale, a lower number means more acidic, which some students get backwards. Baking soda (NaHCO₃) and washing soda (Na₂CO₃·10H₂O) are frequently confused, even though baking soda converts into sodium carbonate only on heating. Finally, water of crystallisation is a fixed, exact number of water molecules built into a salt's crystal structure, not general dampness the salt has absorbed.

Common trapStrength is not the same as concentration

The biggest misconception in this chapter is treating "strong acid" and "concentrated acid" as the same thing, or "weak acid" and "dilute acid" as the same thing. Strength describes how completely an acid or base ionises in water — a fixed chemical property (hydrochloric acid is always strong; acetic acid is always weak). Concentration describes how much acid or base is dissolved per unit volume, and can be changed simply by adding or removing water. A dilute solution of a strong acid can still contain more H-plus ions, and so have a lower pH, than a concentrated solution of a weak acid — the two properties are independent, and conflating them leads to wrong predictions about pH, metal reactivity, and how vigorously a reaction fizzes.

You should be able to…

Identify whether a solution is acidic, basic, or neutral using litmus, phenolphthalein, methyl orange, or an olfactory indicator.
Balance the equation for a metal reacting with a dilute acid, and describe the pop-sound test for hydrogen gas.
Balance the equation for a metal carbonate or hydrogencarbonate reacting with an acid, and describe the lime-water test for carbon dioxide.
Explain, at the level of H-plus and OH-minus ions, why an acid and a base neutralise each other into a salt and water.
Explain why metallic oxides are basic oxides and non-metallic oxides are acidic oxides.
Explain, using the conductivity experiment, why acids and bases need water present to show their ionic behaviour.
Read the pH scale, including why lower pH means more acidic, and why the body, soil, and rivers depend on a narrow pH range.
Trace the chlor-alkali process from common salt to sodium hydroxide, bleaching powder, baking soda, and washing soda.
Explain water of crystallisation and predict what happens when a hydrated salt is heated.

Ask the tutor: Want to work through more neutralisation and pH numericals, or see how this chapter's salts connect to the metal reactivity you will meet in the next chapter?

Why must we always add acid to water and never water to a concentrated acid?

Dissolving a concentrated acid or base in water releases a large amount of heat, since the process is strongly exothermic. Adding acid slowly to a large volume of water lets that heat spread out and dissipate safely. But adding water to a small volume of concentrated acid instead means the sudden heat generated at the point of contact can boil the water almost instantly, splashing the mixture out violently and causing burns, or even cracking the glass. So the rule is always "acid into water, slowly, with stirring" — never the reverse.

Is baking soda the same thing as washing soda?

No, though they are closely related. Baking soda is sodium hydrogencarbonate, NaHCO₃, a mild base used in cooking and antacids. Washing soda is sodium carbonate with its water of crystallisation, Na₂CO₃·10H₂O, used for cleaning and softening hard water. You actually get from one to the other: heating baking soda gives sodium carbonate (plus water and carbon dioxide), and recrystallising that with water gives washing soda — so washing soda is a further product made from baking soda, not the same compound.

Why is pH 7 chosen as the neutral point rather than some other number?

The pH scale runs from 0 to 14 because pure water at room temperature contains a very small but exactly equal number of H-plus and OH-minus ions, and this balance sits at the halfway point of the scale, pH 7. Any solution with more H-plus than OH-minus (an acid) has a pH below 7, and any solution with more OH-minus than H-plus (a base) has a pH above 7. So pH 7 is not an arbitrary label — it is the pH of pure water itself, the reference point where acidic and basic character are exactly balanced, which is why it is called neutral.

Why should curd and other sour foods not be stored in brass or copper vessels?

Curd and other sour foods contain acids (lactic acid in curd), and acids react with reactive metals to release hydrogen gas and form a salt of that metal. Brass and copper are reactive enough to be slowly attacked by these food acids, forming metal salts that dissolve into the food. These compounds can be toxic in enough quantity and also spoil taste and colour — curd stored this way can develop a greenish tinge from copper salts. This is why traditional kitchens store acidic and fermented foods in glass, ceramic, or stainless steel vessels, which do not react with the acid.