Common Acids and Bases · Intro Chemistry

Sour and slippery

Lemons and vinegar taste sour, and your stomach contains hydrochloric acid. Soapy water feels slippery, and the lime mortar used by builders hardens as it soaks up carbon dioxide from the air. These substances fall into two big families: acids and bases.

This lesson has two parts:

Indicators and common acids: acid–base indicators, hydrochloric and sulfuric acid, diluting concentrated sulfuric acid, and the properties all acids share.
Common bases: sodium hydroxide and calcium hydroxide, and the properties all bases share.

A short practice set follows Part 1; do it before you start Part 2.

Part 1: Indicators and common acids

Acid-base indicators

An acid-base indicator酸碱指示剂 is a substance that shows different colors in acidic and basic solutions:

	Acidic solution	Neutral solution	Basic solution
Litmus (purple)	Turns red	Stays purple	Turns blue
Phenolphthalein (colorless)	Stays colorless	Stays colorless	Turns pink

Choose an indicator and add it to the solutions below:

Try it: choose an indicator, then click a test tube to add it

Predict the color first, then click a tube.

Look at sodium carbonate: it’s a salt, but its solution is basic. So a basic solution isn’t necessarily a solution of a base.

Common acids

Hydrochloric acid (a solution of $\ce{HCl}$ in water): concentrated hydrochloric acid gives off hydrogen chloride gas, so a mist forms at the mouth of an open bottle (tiny droplets of acid formed with water vapor in the air). Stomach acid contains hydrochloric acid, which helps digestion.

Sulfuric acid ( $\ce{H2SO4}$ ): concentrated sulfuric acid absorbs water strongly, so it’s used as a drying agent. It’s extremely corrosive and pulls water out of paper and wood, charring them black.

Diluting concentrated sulfuric acid: always pour the acid slowly down the side of the container into the water, stirring constantly, and never pour water into the acid (“add acid to water”). If acid gets on your skin, rinse immediately with plenty of water, then apply a dilute (3%–5%) sodium bicarbonate solution.

Reactions of acids:

They turn litmus red and leave phenolphthalein colorless.
With reactive metals, they form a salt and hydrogen gas (lesson 1.8.3).
With metal oxides, they form a salt and water. When hydrochloric acid removes rust, the rust dissolves and the solution turns yellow:

\ce{Fe2O3(s) + 6HCl(aq) -> 2FeCl3(aq) + 3H2O(l)}

\ce{CuO(s) + H2SO4(aq) -> CuSO4(aq) + H2O(l)}

With bases, they form a salt and water (lesson 1.8.5).
With some salts, they form a new salt and a new acid (lesson 1.8.6):

\ce{CaCO3(s) + 2HCl(aq) -> CaCl2(aq) + H2O(l) + CO2(g)}

Different acids behave alike because they all release $\ce{H+}$ ions in water: in an acid solution, all the positive ions released are $\ce{H+}$ .

Part 1 practice

Practice: what color change does the indicator show?

Litmus added to hydrochloric acid
Litmus added to sodium hydroxide solution
Phenolphthalein added to sulfuric acid
Phenolphthalein added to limewater
Litmus added to salt water
Phenolphthalein added to sodium carbonate solution
Litmus added to white vinegar
Phenolphthalein added to distilled water

Part 2: Common bases

Sodium hydroxide ( $\ce{NaOH}$ , also called lye or caustic soda): a white solid that dissolves easily in water and releases a lot of heat as it does. It absorbs moisture from the air until it dissolves in it (it’s deliquescent), so it can dry some gases. It’s highly corrosive, and it must be stored tightly sealed.

Calcium hydroxide ( $\ce{Ca(OH)2}$ , also called slaked lime): a white powder, only slightly soluble in water; its solution is called limewater. It’s made by reacting quicklime with water, which releases a lot of heat:

\ce{CaO(s) + H2O(l) -> Ca(OH)2(s)}

Calcium hydroxide is used in construction, to make acidic soil less acidic, and in the Bordeaux mixture that protects grapevines from fungus.

Reactions of bases:

They turn litmus blue and phenolphthalein pink.
With some nonmetal oxides, they form a salt and water:

\ce{2NaOH(aq) + CO2(g) -> Na2CO3(aq) + H2O(l)}

\ce{Ca(OH)2(aq) + CO2(g) -> CaCO3(s) + H2O(l)}

With acids, they form a salt and water (lesson 1.8.5).
With some salts, they form a new salt and a new base (lesson 1.8.6):

\ce{2NaOH(aq) + CuSO4(aq) -> Cu(OH)2(s) + Na2SO4(aq)}

In a base solution, all the negative ions released are $\ce{OH-}$ , which is why different bases behave alike.

Example 1: Why never pour water into concentrated sulfuric acid?

Step-by-step solution

Heat. Concentrated sulfuric acid releases a lot of heat when it mixes with water.
Density. Water is less dense than the acid, so water poured in floats on top.
The result. That thin layer of water heats up so fast that it boils, spraying droplets of acid everywhere.
The right way. Pour the acid slowly into the water while stirring, so the heat spreads out through plenty of water.

Example 2: Removing rust with hydrochloric acid

A rusty iron nail is placed in excess dilute hydrochloric acid. What do you see first? What happens if the nail stays in the acid after the rust is gone?

Step-by-step solution

First. The rust (mostly $\ce{Fe2O3}$ ) dissolves, and the solution turns from colorless to yellow: $\ce{Fe2O3 + 6HCl -> 2FeCl3 + 3H2O}$
After the rust is gone. The bare iron reacts with the acid, and bubbles form on it: $\ce{Fe + 2HCl -> FeCl2 + H2}$ .
Conclusion. Don’t leave iron objects in acid too long, or the acid will eat away the iron itself.

Common mistakes

Mistake 1: Pouring water into concentrated sulfuric acid. Always add acid to water, slowly, while stirring.

Mistake 2: “Anything that turns phenolphthalein pink is a base.” Sodium carbonate is a salt, yet its solution turns phenolphthalein pink. All you can say is that the solution is basic.

Mistake 3: “Open concentrated hydrochloric acid gives off white smoke.” It’s a mist, not smoke. A mist is made of tiny liquid droplets; smoke is made of tiny solid particles.

Quick check

Which substance, left open to the air, gains mass AND changes into a different substance?

What is the correct way to dilute concentrated sulfuric acid?

Summary

Litmus turns red in acids and blue in bases; phenolphthalein turns pink in bases and stays colorless in acids. A basic solution isn’t necessarily a solution of a base.
Concentrated hydrochloric acid gives off hydrogen chloride; concentrated sulfuric acid absorbs water and is highly corrosive. Always add acid to water.
Acids react with indicators, reactive metals, metal oxides, bases, and some salts, because they all contain $\ce{H+}$ .
Sodium hydroxide absorbs moisture and is corrosive, so store it sealed. Bases react with indicators, some nonmetal oxides, acids, and some salts, because they all contain $\ce{OH-}$ .

On to AP (for later)

AP topic 8.1 defines acids and bases the Brønsted–Lowry way: an acid donates a proton ( $\ce{H+}$ ) and a base accepts one. In water, $\ce{H+}$ actually exists as the hydronium ion, $\ce{H3O+}$ . AP also distinguishes strong acids, such as hydrochloric and sulfuric acid, which ionize almost completely, from weak acids, such as acetic acid, which ionize only slightly. Indicators return in topic 8.5, where they mark the end point of a titration.

Common Acids and Bases常见的酸和碱

Learning goals

Before you start

Sour and slippery

Part 1: Indicators and common acids

Acid-base indicators

Common acids

Part 1 practice

Part 2: Common bases

Example 1: Why never pour water into concentrated sulfuric acid?

Example 2: Removing rust with hydrochloric acid

Common mistakes

Quick check

Summary

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