Acids and Bases: pH, Indicators, and Neutralization

What Makes Something an Acid or a Base?

In everyday life, you meet acids and bases constantly: vinegar tastes sour (acidic), baking soda feels slippery in water (basic), and stomach acid helps digestion (acidic). In chemistry, acids and bases are defined by how they behave in water and how they affect the concentration of hydrogen ions.

The key idea: hydrogen ions and hydroxide ions

When many substances dissolve in water, they can produce ions. Acids increase the amount of hydrogen ions in water. In water, a “free” hydrogen ion (H+) does not float around by itself for long; it attaches to a water molecule (H2O) to form hydronium (H3O+). For beginner-level work, you will often see H+ used as a shorthand for “hydrogen ion in water,” even though H3O+ is the more accurate picture.

Bases increase the amount of hydroxide ions (OH−) in water or remove H+ from water (which indirectly increases OH−). So, in a practical sense:

• Acidic solution: higher concentration of H+ (or H3O+)
• Basic solution: higher concentration of OH− (and lower H+)
• Neutral solution: balanced H+ and OH−

Two beginner-friendly definitions (Arrhenius and Brønsted–Lowry)

You may see two common definitions. They overlap a lot in water-based chemistry.

• Arrhenius acid: produces H+ in water (example: hydrochloric acid, HCl, produces H+ and Cl−).
• Arrhenius base: produces OH− in water (example: sodium hydroxide, NaOH, produces Na+ and OH−).

Another definition is broader:

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• Brønsted–Lowry acid: donates a proton (H+).
• Brønsted–Lowry base: accepts a proton (H+).

This helps explain why some bases do not contain OH− in their formula but still make solutions basic. For example, ammonia (NH3) accepts a proton from water, forming NH4+ and leaving behind OH− in solution.

pH: A Number That Describes Acidity

pH is a scale used to describe how acidic or basic a solution is. It is based on the concentration of hydrogen ions in water. Because H+ concentrations can vary over many powers of ten, pH uses a logarithmic scale.

The pH scale (0 to 14 in most beginner contexts)

• pH 7: neutral (pure water at room temperature is close to pH 7)
• pH less than 7: acidic
• pH greater than 7: basic

Common reference points (approximate):

• Stomach acid: pH 1–2
• Lemon juice: pH ~2
• Vinegar: pH ~3
• Milk: pH ~6–7
• Seawater: pH ~8
• Household ammonia cleaner: pH ~11
• Drain cleaner (strong base): pH ~13–14

Why pH changes “fast”: the logarithmic idea

Each whole-number step on the pH scale represents a tenfold change in H+ concentration. That means a solution at pH 3 has 10 times more H+ than pH 4, and 100 times more H+ than pH 5. This is why small pH changes can matter a lot in biology, cleaning, and corrosion.

Connecting pH to concentration (simple math)

The definition is:

pH = -log10[H+]

Here, [H+] means “the concentration of hydrogen ions” in moles per liter (mol/L). You do not need advanced math to use this in simple cases if you remember common powers of ten:

• If [H+] = 1 × 10^-7, then pH = 7
• If [H+] = 1 × 10^-3, then pH = 3
• If [H+] = 1 × 10^-11, then pH = 11

Example: A solution has [H+] = 1 × 10^-5 mol/L. The pH is 5.

pOH and the link between H+ and OH−

Because water contains both H+ and OH−, acidity and basicity are connected. Another scale is pOH:

pOH = -log10[OH-]

In many beginner problems at room temperature, you can use:

pH + pOH = 14

Example: If pH = 9, then pOH = 5. That means the solution has relatively low [H+] and higher [OH−].

Strong vs. Weak Acids and Bases (Not the Same as Concentrated vs. Dilute)

People often confuse “strong” with “concentrated,” but they describe different ideas.

Strength: how completely it forms ions in water

Strong acids ionize almost completely in water, producing lots of H+. Weak acids only partially ionize; many acid molecules remain un-ionized in solution.

Similarly, strong bases produce OH− very effectively (often by dissociating completely if they are ionic hydroxides), while weak bases only partially react with water to form OH−.

Concentration: how much acid/base is present per volume

A solution can be:

• Concentrated: a lot of solute per liter
• Dilute: a small amount of solute per liter

So you can have a dilute strong acid (few acid molecules, but each ionizes well) or a concentrated weak acid (many molecules, but each only partly ionizes).

Practical examples

• Hydrochloric acid (HCl) is strong. Even a small amount in water can make pH drop sharply.
• Acetic acid (in vinegar) is weak. Vinegar can still be quite acidic because there is a lot of acetic acid present, but each molecule does not fully ionize.
• Sodium hydroxide (NaOH) is a strong base. It produces OH− readily and can be dangerous.
• Ammonia (NH3) is a weak base. It makes OH− by reacting with water, but not completely.

Indicators: Seeing Acids and Bases

An indicator is a substance that changes color depending on pH. Indicators work because their molecules have different structures at different H+ concentrations, and those structures absorb and reflect light differently.

Common indicators and what they tell you

• Litmus: simple acid/base test. Blue litmus turns red in acid; red litmus turns blue in base.
• Universal indicator: a mixture that shows a range of colors across many pH values (often red/orange for acids, green near neutral, blue/purple for bases).
• Phenolphthalein: colorless in acidic/neutral solutions and pink in basic solutions (commonly used in titrations).

How to use indicator paper (step-by-step)

This is a practical way to estimate pH without a meter.

• Step 1: Place a small sample of the solution in a clean cup or on a watch glass. Do not dip the whole indicator container into the solution.
• Step 2: Dip a strip of pH paper briefly into the sample (or use a clean stirring rod to place a drop on the paper).
• Step 3: Wait the amount of time recommended (often a few seconds).
• Step 4: Compare the color to the chart provided with the paper to estimate pH.
• Step 5: Record the pH and rinse/dispose of materials safely.

Using a pH meter (step-by-step)

A pH meter gives a more precise reading, but it must be used correctly.

• Step 1: Rinse the electrode with distilled water (not tap water if you can avoid it), then gently blot dry (do not rub).
• Step 2: Calibrate the meter using standard buffer solutions (commonly pH 4, pH 7, and pH 10). Follow the device instructions.
• Step 3: Place the electrode into the sample so the sensing tip is fully immersed.
• Step 4: Stir gently or wait for the reading to stabilize.
• Step 5: Record the pH, then rinse the electrode again and store it in the proper storage solution.

Practical tip: pH electrodes are sensitive. Letting them dry out or storing them in pure water can damage them.

Neutralization: When Acids and Bases React

Neutralization is a reaction in which an acid and a base react to form water and (usually) a salt. The simplest picture is that H+ from the acid combines with OH− from the base to form water:

H+ + OH- → H2O

What remains (the “spectator ions”) often forms an ionic compound called a salt. “Salt” here does not only mean table salt; it means an ionic product formed from an acid-base reaction.

Example: hydrochloric acid and sodium hydroxide

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

In this reaction, the solution can move toward neutral pH if the acid and base are mixed in the right proportions.

What “neutral” really means in practice

Neutralization does not always produce a solution with pH exactly 7. The final pH depends on the strength of the acid and base and whether one is left over.

• Strong acid + strong base: often ends near pH 7 if mixed in exact stoichiometric amounts.
• Strong acid + weak base: can end acidic even at the equivalence point.
• Weak acid + strong base: can end basic even at the equivalence point.

At a beginner level, it is enough to understand that “neutralization” means acid and base are reacting, and the pH moves toward the middle compared with where it started.

Neutralization in everyday contexts

• Antacids: basic compounds neutralize excess stomach acid, reducing acidity.
• Soil treatment: acidic soil can be treated with lime (basic materials) to raise pH.
• Spills and safety: neutralization can reduce hazards, but it can also release heat and cause splattering if done carelessly.

Titration: Measuring an Unknown Concentration Using Neutralization

A titration is a controlled neutralization used to determine the concentration of an unknown acid or base. You slowly add a solution of known concentration (the titrant) to a measured volume of unknown solution until you reach the endpoint, which is detected by an indicator color change or a pH meter.

Key terms used in titration

• Titrant: the solution with known concentration (in the burette).
• Analyte: the solution with unknown concentration (in the flask).
• Endpoint: the indicator color change you observe.
• Equivalence point: the point where the acid and base have reacted in exact stoichiometric proportions (moles based on the reaction). Endpoint and equivalence point are close but not always identical.

Simple titration setup (what you typically use)

• Burette (to deliver titrant accurately)
• Erlenmeyer flask (holds analyte)
• Pipette or graduated cylinder (to measure analyte volume)
• Indicator (like phenolphthalein) or pH meter

Step-by-step: titrating an acid with a base using phenolphthalein

This example assumes the unknown is an acid and the titrant is a base (often NaOH). Phenolphthalein is commonly used because it changes to pink in basic conditions.

• Step 1: Rinse and fill the burette. Rinse the burette with a small amount of the titrant solution, then fill it. Remove air bubbles from the tip and record the initial burette reading.
• Step 2: Measure the analyte. Use a pipette (best) or a graduated cylinder to place a known volume of the acid into an Erlenmeyer flask.
• Step 3: Add indicator. Add 2–3 drops of phenolphthalein to the acid. The solution should remain colorless if it is acidic.
• Step 4: Titrate slowly. Add base from the burette while swirling the flask. At first you can add faster, but as you approach the endpoint, add drop-by-drop.
• Step 5: Identify the endpoint. The endpoint is the first faint pink color that persists for about 20–30 seconds after swirling.
• Step 6: Record final burette reading. Subtract initial from final to get the volume of titrant used.
• Step 7: Repeat for accuracy. Do at least two or three trials and look for consistent volumes (within a small difference).

How the calculation idea works (without overcomplicating)

The core logic is: at equivalence, the reacting amounts match the balanced reaction ratio. For a simple 1:1 reaction (one H+ reacts with one OH−), the moles of acid equal the moles of base at equivalence.

moles = M × V

Here M is molarity (mol/L) and V is volume in liters.

Example (1:1 reaction): You titrate 25.0 mL of an unknown HCl solution with 0.100 M NaOH. It takes 18.6 mL of NaOH to reach the endpoint.

• Convert volume of base to liters: 18.6 mL = 0.0186 L
• Moles of NaOH used: 0.100 mol/L × 0.0186 L = 0.00186 mol
• For HCl + NaOH (1:1), moles of HCl = 0.00186 mol
• Volume of acid in liters: 25.0 mL = 0.0250 L
• Molarity of acid: 0.00186 mol ÷ 0.0250 L = 0.0744 M

This is the practical power of neutralization: you can measure “how much acid” by seeing how much base it takes to neutralize it.

Polyprotic Acids and “More Than One H+”

Some acids can donate more than one proton per molecule. These are called polyprotic acids. This matters in neutralization because one mole of acid may require more than one mole of base to fully neutralize.

Example idea: sulfuric acid

Sulfuric acid (H2SO4) can provide two H+ ions (in steps). In a simplified neutralization with sodium hydroxide, the overall reaction can be:

H2SO4 + 2 NaOH → Na2SO4 + 2 H2O

Notice the 1:2 ratio. This means if you have 0.0100 mol of H2SO4, you need 0.0200 mol of NaOH for full neutralization.

Acid-Base Reactions Without Obvious OH−: Carbonates and Bicarbonates

Not all neutralizations involve a base that contains OH− in its formula. Carbonates (CO3^2−) and bicarbonates (HCO3−) act as bases because they can react with H+.

Example: acid + bicarbonate produces carbon dioxide

If you mix vinegar (acetic acid) with baking soda (sodium bicarbonate, NaHCO3), you see bubbling. That gas is carbon dioxide (CO2). The bubbling is a clue that an acid-base reaction is happening.

A simplified ionic idea is that bicarbonate reacts with H+ to form carbonic acid (H2CO3), which decomposes into water and carbon dioxide:

H+ + HCO3- → H2CO3 → H2O + CO2(g)

This is why baking soda can neutralize acids and also why it can act as a leavening agent in baking (gas formation makes batter rise).

Safety and Practical Handling (Acids/Bases Are Often Corrosive)

Many acids and bases can damage skin, eyes, and materials. Even household products can be hazardous at high concentration.

• Wear eye protection when working with unknown solutions or strong cleaners.
• Add acid to water, not water to acid when diluting (this reduces splashing risk from heat release).
• Do not mix cleaners unless you know the chemistry; dangerous gases can form.
• Neutralization releases heat in many cases, especially with strong acids and bases, so add slowly and stir.

Practice: Predicting pH Changes During Mixing

You can often predict the direction of pH change without calculations by thinking about whether you are adding H+ (acid) or removing H+ / adding OH− (base).

Scenario practice

• You add lemon juice to water: pH decreases (more acidic).
• You add a small amount of baking soda to vinegar: pH increases compared with vinegar alone (less acidic), and CO2 bubbles form.
• You add a strong base to a weak acid: pH rises; if enough base is added to react with most of the acid, the solution can become basic.
• You dilute an acid with water: pH increases toward 7 (less acidic), but it may still be acidic.

Quick pH estimation with powers of ten

If a problem gives [H+] in scientific notation, you can often read pH directly from the exponent when the leading number is 1.

• [H+] = 1 × 10^-2 → pH 2
• [H+] = 1 × 10^-8 → pH 8

If the leading number is not 1 (for example 3.2 × 10^-4), the pH is not a whole number. At this stage, it is enough to know it will be close to 4 but slightly lower (because 3.2 is greater than 1, meaning more H+ than 1 × 10^-4).