Lemon juice and vinegar taste sour because each contains an 'acid'.

Acids share three observable properties.

First property: foods that contain an acid taste sour.

Second property: acids change the color of 'litmus paper' — a test strip coated with a plant dye.

Dip blue litmus paper into an acid and the paper turns red.

Third property: an acid reacts with an active metal, such as zinc or magnesium, and the reaction releases hydrogen gas.

You will see the hydrogen as streams of gas bubbles rising off the metal.

The taste facts here describe everyday foods only — laboratory acids are never tasted or touched, because they can burn skin and eyes.

Soapy water feels slippery because it contains a 'base'.

Bases share three observable properties.

First property: foods that contain a base taste bitter — unsweetened cocoa owes its bitterness to bases.

Second property: a solution of a base feels slippery between the fingers, the way soapy water does.

Third property: a base turns red litmus paper blue — the reverse of the acid response.

As with acids, the taste and feel facts describe everyday foods and household materials — laboratory bases are never tasted or touched, because they can burn skin and eyes.

Classify an unknown by matching its observed properties against the two profiles.

A sour taste in a food points to an acid.

A bitter taste in a food points to a base.

Turning blue litmus paper red points to an acid.

Turning red litmus paper blue points to a base.

A slippery feel points to a base.

Releasing hydrogen gas from an active metal points to an acid.

One clear profile observation is enough — a solution that turns blue litmus paper red is an acid.

'Hydrochloric acid' is HCl — the acid your stomach makes to digest food.

'Sulfuric acid' is H₂SO₄ — the acid inside car batteries.

'Nitric acid' is HNO₃ — used in huge amounts to make fertilizers.

'Acetic acid' is CH₃COOH — the acid that gives vinegar its sour taste.

Three of the four formulas begin with H; acetic acid's formula is the odd one out, written CH₃COOH.

Bottled laboratory acids like these are never tasted or touched — the vinegar fact is about food, not lab bottles.

'Sodium hydroxide' is NaOH — the base in drain cleaners.

'Potassium hydroxide' is KOH — used to make liquid soaps.

'Calcium hydroxide' is Ca(OH)₂ — sold as slaked lime and spread on fields to treat overly acidic soil.

'Ammonia' is NH₃ — the base in many glass cleaners.

Three of the four are metal hydroxides — a metal paired with OH; ammonia is the odd one out, with no metal and no OH in its formula.

Drain cleaner and the other products owe their power to these bases — never touch them, because they burn skin and eyes.

Look at where hydrogen sits in the formula.

An acid's formula is written with H first — HCl, H₂SO₄, and HNO₃ all lead with hydrogen.

A common base's formula ends in OH — NaOH, KOH, and Ca(OH)₂ all finish with the hydroxide unit.

A formula like Ca(OH)₂ still counts as ending in OH — the subscript outside the parentheses just counts the OH units.

Two formulas break the patterns, and you memorize them.

H₂O is written with H first but is not an acid — it is just water.

NH₃ does not end in OH, but it is a base — ammonia.

Dissolve HCl in water and it splits into two ions: H⁺ and Cl⁻.

The equation is HCl → H⁺ + Cl⁻.

Dissolve any acid in water and the same kind of thing happens: it releases hydrogen ions, H⁺.

A hydrogen ion is a hydrogen atom that has lost its one electron, leaving a particle with a 1+ charge.

The Swedish chemist Svante Arrhenius made this the definition: an 'Arrhenius acid' is a substance that releases hydrogen ions (H⁺) when it dissolves in water.

The released H⁺ is the same particle no matter which acid it came from.

That is why every acid shows the same property profile — the H⁺ itself is the source of the shared properties.

Dissolve NaOH in water and it splits into two ions: Na⁺ and OH⁻.

The equation is NaOH → Na⁺ + OH⁻.

OH⁻ is the hydroxide ion you met among the polyatomic ions — it is also called the 'hydroxyl' ion.

Dissolve any common base in water and it releases hydroxide ions, OH⁻.

An 'Arrhenius base' is a substance that releases hydroxide ions (OH⁻) when it dissolves in water.

The released OH⁻ is the same particle no matter which base it came from.

That is why every base shows the same property profile — the OH⁻ itself is the source of the shared properties.

Read the dissolving equation and find the ions released.

A substance that releases H⁺ in water is an Arrhenius acid.

A substance that releases OH⁻ in water is an Arrhenius base.

A substance that releases neither ion is neither an Arrhenius acid nor an Arrhenius base.

HI → H⁺ + I⁻, so HI is an Arrhenius acid.

KOH → K⁺ + OH⁻, so KOH is an Arrhenius base.

CaCl₂ → Ca²⁺ + 2Cl⁻ — no H⁺ and no OH⁻ — so CaCl₂ is neither.

Watch the hydrogens: KOH contains hydrogen, but that hydrogen leaves inside the OH⁻ unit, not as H⁺.

A hydrogen ion in water does not float free.

The positive H⁺ attaches to the slightly negative oxygen end of a water molecule.

The joined particle is the 'hydronium ion', H₃O⁺ — three hydrogens, one oxygen, charge 1+.

The equation is H⁺ + H₂O → H₃O⁺.

H⁺ and H₃O⁺ are two ways of writing the same dissolved particle.

Chemists write H⁺ as the shorthand and H₃O⁺ as the honest picture — this course uses both.

The H⁺ released by HCl in water is therefore carried as H₃O⁺.

Look again at HCl dissolving, written the honest way: HCl + H₂O → H₃O⁺ + Cl⁻.

The hydrogen ion does not simply leave the HCl — it is passed to a water molecule.

The passed particle has a one-word name: a hydrogen atom is one proton with one electron, so the H⁺ left after losing the electron is just a 'proton'.

In the Brønsted-Lowry definition, a 'Brønsted-Lowry acid' is a particle that donates a proton — an H⁺ — to another particle.

HCl acts as the proton donor when it passes H⁺ to water, so HCl is a Brønsted-Lowry acid.

Any particle caught donating an H⁺ is, in that moment, acting as a Brønsted-Lowry acid.

In the Brønsted-Lowry definition, a 'Brønsted-Lowry base' is a particle that accepts a hydrogen ion from another particle.

Ammonia shows the role clearly: NH₃ gains an H⁺ and becomes NH₄⁺.

The NH₃ acts as the proton acceptor, so NH₃ is a Brønsted-Lowry base.

Notice what accepting does to a formula: it grows by one H, and its charge rises by one.

Water can play the role too: when an acid donates an H⁺, the water molecule that accepts it becomes H₃O⁺.

Whenever a water molecule accepts an H⁺, that water molecule is acting as a Brønsted-Lowry base.

Compare each particle on the left of the equation with the particle it becomes on the right.

The particle that loses an H⁺ between the two sides is the proton donor — the Brønsted-Lowry acid.

The particle that gains an H⁺ between the two sides is the proton acceptor — the Brønsted-Lowry base.

In HCl + H₂O → H₃O⁺ + Cl⁻, HCl becomes Cl⁻ — it lost an H⁺, so HCl is the donor, the acid.

H₂O becomes H₃O⁺ — it gained an H⁺, so H₂O is the acceptor, the base.

Water does not always play the acceptor: whichever particle gains the H⁺ in that equation is the base.

The Arrhenius definition classifies by what a substance releases in water: an acid releases H⁺, and a base releases OH⁻.

The Brønsted-Lowry definition classifies by proton transfer: an acid donates an H⁺, and a base accepts one.

Every Arrhenius acid is also a Brønsted-Lowry acid, because releasing H⁺ into water is donating H⁺ to water molecules.

The base definitions do not match up the same way, because a particle can accept H⁺ without ever containing OH⁻.

NH₃ contains no OH⁻, so the Arrhenius definition cannot call it a base.

But NH₃ accepts an H⁺ to form NH₄⁺, so the Brønsted-Lowry definition classifies it as a base.

So the Brønsted-Lowry definition extends the Arrhenius definition: it keeps everything Arrhenius covers and adds the bases that contain no hydroxide.

When an acid donates its H⁺, the particle left behind differs from the acid by exactly one H⁺.

Two particles whose formulas differ by exactly one H⁺ form a 'conjugate acid-base pair'.

The member with the extra H⁺ is the conjugate acid; the member without it is the conjugate base.

In NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, NH₃ and NH₄⁺ differ by one H⁺ — NH₄⁺ is the conjugate acid formed from the base NH₃.

H₂O and OH⁻ also differ by one H⁺ — the equation's second conjugate pair.

Every Brønsted-Lowry equation contains exactly two conjugate pairs, and each pair has one member on each side of the arrow.

Look at what an acid's formula contains besides hydrogen.

A 'binary acid' contains hydrogen and exactly one other element.

An 'oxyacid' contains hydrogen and a polyatomic ion that contains oxygen.

HBr contains hydrogen and only bromine — a binary acid.

HNO₃ contains hydrogen and the nitrate ion, NO₃⁻ — nitrate contains oxygen, so HNO₃ is an oxyacid.

The quick check: no oxygen in the formula means binary; hydrogen plus an oxygen-containing ion means oxyacid.

The name of a binary acid has three parts: hydro-, the other element's root, and -ic acid.

Start with the prefix 'hydro-'.

Add the root of the other element's name.

Finish with -ic and the word acid.

For HBr the other element is bromine, root 'brom': hydro + brom + ic acid gives hydrobromic acid.

The subscripts play no part in the name — however many hydrogens the formula carries, the pattern is the same.

A name of the form hydro-…-ic acid announces a binary acid: hydrogen plus one other element.

The root between hydro- and -ic names that element.

Write hydrogen with that element's negative ion, balancing the charges the same way as for any ionic formula.

Hydrogen counts as 1+ here, so the ion's charge number tells you how many H to write.

For hydrosulfuric acid the root is sulfur, and the sulfide ion is S²⁻.

Balancing 1+ hydrogens against one 2− sulfide takes two hydrogens: H₂S.

Find the polyatomic ion inside the oxyacid's formula.

If the ion's name ends in -ate, the acid's name ends in -ic acid.

If the ion's name ends in -ite, the acid's name ends in -ous acid.

There is no hydro- prefix on an oxyacid's name.

H₂SO₄ contains the sulfate ion — sulfate becomes sulfuric acid.

H₂SO₃ contains the sulfite ion — sulfite becomes sulfurous acid.

The element root sometimes stretches so the name says smoothly: sulfate gives sulfURic, and phosphate gives phosphORic.

Reverse the suffix first: -ic acid points back to an -ate ion, and -ous acid points back to an -ite ion.

Write down that ion's formula and charge.

Add enough H, counting 1+ each, to cancel the ion's charge.

For nitrous acid, -ous points to the nitrite ion, NO₂⁻.

One 1− charge takes one H: HNO₂.

Check the finish: the H count equals the ion's charge number, so the formula is neutral.

A metal hydroxide is an ionic compound, so its name is its ordinary ionic name.

Name the metal, then the hydroxide ion: metal name + hydroxide.

Ca(OH)₂ is calcium hydroxide — the subscript 2 appears nowhere in the name.

Base names never take hydro-, -ic, or -ous — those patterns belong to acids.

One common base breaks the pattern: NH₃ carries the special name 'ammonia'.

NH₃ contains no metal and no hydroxide, so the ionic routine cannot name it — the special name is all there is.

When an acid or base dissolves in water, its particles can split into ions — chemists call this splitting 'ionizing'.

A 'strong' acid or base ionizes completely: every dissolved particle splits into ions.

A 'weak' acid or base ionizes only partially: most of its dissolved particles stay whole, and only a few split.

The two beakers show the contrast at the particle level: every HCl has split into H⁺ and Cl⁻, while most CH₃COOH particles are still intact.

HCl is a strong acid, and complete ionization is written with a single arrow: HCl → H⁺ + Cl⁻.

Acetic acid is weak, and partial ionization is written with a double arrow: CH₃COOH ⇌ CH₃COO⁻ + H⁺.

The double arrow shows the change also runs backward, so both forms are present.

More free ions carry more current — the bright bulb was the strong acid's beaker, and the dim bulb the weak acid's.

Strong and weak mean exactly the same for bases: a strong base ionizes completely, and a weak base only partially.

Six common acids are strong — each ionizes completely in water.

Three are binary acids of halogens: HCl, HBr, and HI.

The fourth halogen acid, HF, is NOT on the list.

Three are oxyacids: nitric acid HNO₃, sulfuric acid H₂SO₄, and HClO₄, called perchloric acid.

An acid that is not on this list can be assumed weak in this course.

So HF — and any other unlisted acid you meet here — counts as weak.

The hydroxides of the Group 1 metals lithium, sodium, and potassium are strong bases: LiOH, NaOH, and KOH.

So KOH is a strong base — every dissolved KOH splits into K⁺ and OH⁻.

The hydroxides of the heavier Group 2 metals are also strong: Ca(OH)₂, Sr(OH)₂, and Ba(OH)₂.

The common weak base is ammonia, NH₃ — most dissolved NH₃ stays whole.

An acid's 'strength' reports the fraction of its dissolved particles that ionize in water.

An acid's 'concentration' reports how much acid is dissolved per liter of solution — the molarity you have seen.

Strength is fixed by which acid it is; concentration is chosen by whoever mixes the solution.

Hydrochloric acid ionizes completely at 6.0 M and it ionizes completely at 0.10 M — diluting it changes the concentration, never the strength.

'Concentrated' and 'dilute' describe concentration; 'strong' and 'weak' describe strength.

So a dilute solution of hydrochloric acid is still a strong-acid solution — a small amount of acid, every particle of it ionized.

And a concentrated solution of acetic acid is still a weak-acid solution — a large amount of acid, most of it staying whole.

Whether an acid or base is strong or weak can be read from data about its solution.

Evidence one — the fraction ionized: a strong acid or base ionizes close to 100%, while a weak one ionizes only a few percent.

A 0.10 M HCl solution is essentially 100% ionized; a 0.10 M HF solution is about 8% ionized.

Evidence two — a particle diagram: a strong acid's diagram shows every dissolved unit split into ions, while a weak acid's diagram shows mostly whole molecules with only a few split.

Evidence three — conductivity: compared with a known strong acid at the same concentration, a weak acid lights a conductivity bulb dimly, because it supplies fewer moving ions.

Whatever the data type, the question is always the same: what fraction of the dissolved particles has split into ions?

Watch the boundary you have seen: a small total number of dissolved particles means dilute, not weak — only the fraction split reports strength.

Chemists write the hydrogen-ion concentration of a solution as '[H⁺]' — read the square brackets as 'the concentration of', measured in molarity.

'pH' is a single number, usually between 0 and 14, that reports a solution's [H⁺].

The scale runs in reverse: the lower the pH, the higher the [H⁺].

A pH 3 solution has a higher [H⁺] than a pH 10 solution.

Read the other way, the higher the pH, the lower the [H⁺].

The pH scale sorts every solution into three bands.

A solution with pH below 7 is 'acidic'.

A solution with pH exactly 7 is 'neutral' — pure water sits here.

A solution with pH above 7 is 'basic'.

The boundary is sharp: pH 6 is acidic and pH 8 is basic, even though both sit close to 7.

Lemon juice at pH 2 is acidic, pure water at pH 7 is neutral, and household ammonia at pH 11 is basic.

Each change of one pH unit means a tenfold change in [H⁺].

Step down one pH unit, and the [H⁺] is ten times higher.

Step up one pH unit, and the [H⁺] is ten times lower — one tenth as much.

A pH 6 solution has ten times the [H⁺] of a pH 7 solution.

Take one more step down, and a pH 5 solution has ten times the [H⁺] of the pH 6 solution.

One small step on the scale stands for a whole factor of ten in [H⁺] — that is how pH swallows all those zeros.

To compare two solutions' [H⁺], count the pH units between them.

Each pH unit of difference multiplies the [H⁺] by ten.

So the factor between the two solutions is a ten for every unit: two units apart means 10 × 10 = 100 times.

The solution with the lower pH is the one with the greater [H⁺].

A pH 4 solution and a pH 6 solution sit two units apart, so the pH 4 solution has 100 times the [H⁺] of the pH 6 solution.

Every water solution contains both H⁺ ions and OH⁻ ions — the acidic ones, the neutral ones, and the basic ones alike.

The two concentrations trade off: when [H⁺] rises, [OH⁻] falls, and when [H⁺] falls, [OH⁻] rises.

In an acidic solution, [H⁺] is greater than [OH⁻].

In a neutral solution, [H⁺] and [OH⁻] are equal.

In a basic solution, [OH⁻] is greater than [H⁺].

So a basic solution still contains some H⁺ — just less of it than OH⁻.

pH is defined by the equation pH = −log[H⁺].

The 'log' of a power of ten is simply its exponent: the log of 1 × 10⁻⁴ is −4.

The minus sign in the equation then flips that exponent's sign: pH = −(−4) = 4.

So for [H⁺] = 1 × 10⁻ⁿ M, the pH is simply n.

In this course the equation is used only when [H⁺] is a whole-number power of ten — concentrations written 1 × 10⁻ⁿ M, never values in between.

The tenfold-per-unit behavior falls straight out: each step of n is one more power of ten in [H⁺].

To calculate a pH, write down the value of [H⁺], write down the equation, substitute, and calculate.

In the calculate step, the log of 1 × 10⁻ⁿ is −n, and the equation's minus sign turns −n into n.

If the concentration arrives as a decimal, first rewrite it as a power of ten: 0.00001 M is 1 × 10⁻⁵ M.

pH carries no unit — the answer is a bare number.

So a solution with [H⁺] = 0.00001 M has [H⁺] = 1 × 10⁻⁵ M, and its pH is 5.

The pH equation also runs backwards: from a known whole-number pH, you can recover [H⁺].

Make [H⁺] the subject of pH = −log[H⁺]: the concentration is [H⁺] = 1 × 10⁻ᵖᴴ M — ten to the power of MINUS the pH.

The pH becomes the exponent's number, and the exponent's sign is minus.

A solution with pH 4 has [H⁺] = 1 × 10⁻⁴ M.

The answer is a concentration, so it carries the unit M.

An 'acid-base indicator' is a dye whose color depends on the pH of the solution it is in.

'Phenolphthalein' is an indicator with two states: colorless in acidic and in neutral solutions, and pink in basic solutions.

'Universal indicator' is a mixture of dyes that runs through a whole range of colors as pH changes.

Its color is matched against a chart to read an approximate pH.

On the chart: red near pH 0–2, orange at pH 3–4, yellow at pH 5–6, green at pH 7, blue at pH 8–11, and purple at pH 12–14.

Read an indicator result in two steps: name the indicator, then match the observed color against that indicator's facts.

Blue litmus turning red is the acid signal; red litmus turning blue is the base signal.

When neither litmus paper changes color, the solution is near-neutral.

Phenolphthalein turning pink identifies a basic solution.

Phenolphthalein staying colorless only rules out basic — acidic and neutral solutions both leave it colorless, so colorless cannot separate those two.

Universal indicator is matched against its color chart: red sits at the strongly acidic end, orange and yellow are weakly acidic, green is near-neutral, blue is weakly basic, and purple sits at the strongly basic end.

A solution that turns universal indicator red is therefore identified as acidic.

When an acid and a base react, each cancels the properties of the other — chemists call the reaction 'neutralization'.

Neutralization produces the same two kinds of product every time: water and a salt.

A 'salt' is an ionic compound made from the base's positive ion and the acid's negative ion.

Table salt is only one member of the family — every neutralization makes some salt.

Hydrochloric acid and sodium hydroxide react to form water and the salt sodium chloride: HCl + NaOH → NaCl + H₂O.

The salt NaCl took its Na⁺ from the base and its Cl⁻ from the acid.

In water, hydrochloric acid exists as separate H⁺ and Cl⁻ ions, and sodium hydroxide as separate Na⁺ and OH⁻ ions.

Neutralization produces water and a salt because the H⁺ ions from the acid combine with the OH⁻ ions from the base to form water, leaving the remaining ions as the salt.

Each H⁺ joins one OH⁻ to form one H₂O molecule.

The Na⁺ and Cl⁻ take no part in the joining — they stay dissolved, and they are the salt's ions.

When exactly enough base has been added, every free H⁺ has been locked into water, so the acid's properties disappear.

What remains is simply the salt's ions in water — here, a solution of NaCl.

One product never needs predicting: water, H₂O, forms in every acid–metal hydroxide neutralization.

The salt takes the base's metal ion and the acid's negative ion.

Find the acid's negative ion by removing the H⁺ from its formula: HCl gives Cl⁻, HNO₃ gives NO₃⁻, and H₂SO₄ gives SO₄²⁻.

Read the base's metal ion straight from its formula: KOH gives K⁺, and Ca(OH)₂ gives Ca²⁺.

Write the salt's formula by criss-crossing the two charges, exactly as for any ionic compound.

For nitric acid reacting with potassium hydroxide: K⁺ and NO₃⁻ pair one-to-one, so the products are H₂O and the salt KNO₃.

Write the acid and the base on the left, and the predicted salt and water on the right.

Balance by adjusting coefficients one element at a time — never change a subscript.

The quick route: count the H⁺ each acid unit gives and the OH⁻ each base unit gives, and pick coefficients that make the two counts equal.

Each H⁺–OH⁻ pair becomes one H₂O, so the water's coefficient equals that shared count.

Sulfuric acid gives two H⁺ but each NaOH gives one OH⁻, so two NaOH are needed: H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O.

Check by counting atoms: 4 H, 1 S, 6 O, and 2 Na on each side.

Concentration leaves no visible trace — a 0.1 M and a 1.0 M hydrochloric acid solution look identical.

A 'titration' finds an unknown concentration by reacting the solution with one whose concentration is known.

A measured volume of the unknown solution is placed in a flask.

The known solution is added slowly from a thin volume-marked tube, a little at a time.

The additions stop at the moment the two have exactly reacted — no acid and no base left over.

The volume of known solution delivered is then read off the tube.

From the known concentration, the two measured volumes, and the balanced equation for the reaction — which supplies how many of each react together — the unknown concentration can be calculated.

A few drops of an indicator are added to the flask before the titration begins.

The 'endpoint' of a titration is the moment the indicator changes color.

The color change signals that the acid and the base have exactly reacted — or as close to it as one drop can get: a careful titration lands the endpoint within a single drop of that moment.

At the endpoint, the additions stop and the delivered volume is recorded.

With phenolphthalein in an acid-filled flask, the endpoint is the first faint pink that stays after swirling.

This pattern covers titrations in which the acid and the base react one-to-one.

At the endpoint, the two solutions have delivered exactly matching amounts of acid and base.

A solution that delivered its matching amount in less volume packs more into every milliliter.

So the solution that needed the smaller volume is the more concentrated one.

Equal volumes at the endpoint mean equal concentrations.

When 10.0 mL of a base neutralizes 20.0 mL of an acid, the base delivered the same amount in half the volume — the base is the more concentrated.

This equation is for titrations in which the acid and the base react one-to-one.

At the endpoint of such a titration, the moles of acid that have reacted equal the moles of base.

Each solution's delivered moles are counted by the relationship you have seen: n = M × V.

Setting the two mole counts equal gives the titration equation: M(acid) × V(acid) = M(base) × V(base).

The two volumes must be in the same unit — both in mL or both in L — and then the volume unit cancels.