Atomic Foundations of Matter

Why This Matters

Mix baking soda into a glass of vinegar. It fizzes wildly. Bubbles rush up and burst. Now weigh the glass before and after. The balance shows less mass after the fizzing stops. Did some matter just vanish into thin air?

This little kitchen puzzle bothered scientists for a very long time. When wood burns, the ash left behind weighs far less than the log. When iron rusts, it seems to gain weight from nowhere. For centuries, people had no clear rule for what happens to matter when it changes.

Then two simple, powerful rules were discovered. Matter is never made or destroyed in a chemical change — it only rearranges. And when elements join to make a compound, they always join in the same fixed ratio by mass. These two rules pointed straight to a stunning idea: all matter is built from tiny, unbreakable building blocks called atoms.

In this chapter you will learn how atoms stick together to make everything around you — the water you drink, the salt in your food, the air you breathe. You will learn to write the secret “code” (the chemical formula) for any compound. By the end, that fizzing glass of vinegar will hold no mystery at all.

The Big Idea

The Big Idea: In a chemical reaction, atoms are never created or destroyed — they only swap partners and rearrange. So the total mass stays the same (Law of Conservation of Mass), and every pure compound always has its elements in the same fixed ratio by mass (Law of Constant Proportions). Atoms join in two main ways: by sharing electrons (a covalent bond) or by giving and taking electrons to form charged ions (an ionic bond). Once you know which atoms join and how many, you can write the compound’s chemical formula — and from the formula, you can work out its mass.

Before we start, let us quickly refresh the idea this whole chapter is built on: the atom, from the last chapter.

Mass in a Change — Does Anything Get Lost?

Let us settle the fizzing-glass puzzle first. We will start gently with a physical change, then move to a chemical change.

Physical changes keep their mass

Dissolve a spoon of salt in water and the salt seems to “disappear”. But put the beaker on a digital balance, set it to zero, add water, add salt, and stir. The reading does not change. The salt is still there — its tiny particles have just spread out between the water particles. Tear a sheet of paper into bits, weigh before and after: same mass. In a physical change, nothing is added or removed, so mass stays the same.

Chemical changes — the trick of the escaping gas

Now the vinegar and baking soda reaction. In words:

Vinegar + Baking soda → Carbon dioxide + Other substances

Here is the catch that fooled people for years. One of the products is carbon dioxide gas. In an open flask, that gas floats away into the room before you read the balance. So the balance shows less mass — not because matter was destroyed, but because some of it (the gas) silently left and was never weighed.

Seal the same reaction with a balloon, and the gas has nowhere to go. It inflates the balloon but stays on the balance. Now the reading does not change at all. Figure 9.1 shows both set-ups side by side.

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The Law of Conservation of Mass

This gives us our first big law, proposed by Antoine Lavoisier in 1789 (he is called the Father of Modern Chemistry):

Law of Conservation of Mass: Mass can neither be created nor destroyed in a chemical reaction.

In plain words: total mass of reactants = total mass of products. Always. The “Other substances” in our reaction include water and sodium acetate; add up everything (gas included) and the mass matches perfectly.

But why must this be true? Here is the deep reason, which the next sections will build up to. A chemical reaction is just atoms letting go of old partners and grabbing new ones. No atom is ever made from nothing, and no atom is ever destroyed. The exact same atoms are present before and after — only their arrangement changed. Same atoms means same mass. That is the whole secret.

Let us put the law to work on a real calculation.

Same Compound, Same Recipe — Constant Proportions

Now a different question. When hydrogen and oxygen join to make water, can they join in any amounts, or must it always be the same recipe?

Soon after Lavoisier, Joseph Proust answered this. Take water from a river, a borewell, or the ocean. Purify it and split it apart. You always get hydrogen and oxygen in a mass ratio of exactly 1 : 8. So 9 g of pure water always gives 1 g of hydrogen and 8 g of oxygen — no matter the source.

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Law of Constant Proportions (also called the Law of Definite Proportions, or Proust’s Law): A pure compound always contains the same elements combined in a fixed ratio by mass, no matter how it is made or where it comes from.

Why does this happen? Again it points to atoms. A water “unit” is always 2 hydrogen atoms locked to 1 oxygen atom. You can never have “two-and-a-half” hydrogen atoms — atoms come in whole numbers. So the recipe can never drift. The mass ratio is fixed because the atom count is fixed.

This law is true for compounds, where elements are chemically joined. It is not true for mixtures, where you can mix things in any amount you like (you can make tea weak or strong).

Let us use this law to find an unknown mass.

Dalton’s Atomic Theory — Why Both Laws Are True

Both laws kept pointing at the same hidden idea. In 1808, John Dalton put it into words as his atomic theory. He stated it as a set of postulates — a postulate is a starting idea we accept as true, and then build a theory on it.

Dalton’s main postulates were:

• All matter is made of very tiny particles called atoms.
• Atoms are indivisible — they cannot be created or destroyed in a chemical reaction.
• Atoms of the same element are identical in mass and chemical properties.
• Atoms of different elements have different masses and properties.
• Atoms combine in simple whole-number ratios to form compounds.
• The number and kind of atoms in a given compound is always the same.

Figure 9.3 shows these ideas as pictures, and explains how they neatly account for both laws.

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This is the satisfying part. The two laws were observations — things people measured. Dalton’s atoms explain them:

• Atoms are never made or destroyed → so total mass can never change → Conservation of Mass.
• A compound always uses the same whole-number count of identical atoms → so its mass ratio is always the same → Constant Proportions.

How Atoms Combine

An atom of an element can join with other atoms to become stable. When it does, it forms a molecule.

A molecule is an electrically neutral group of two or more atoms joined together, which can exist on its own and shows all the properties of that substance.

Atoms of the same element can join — like two hydrogen atoms forming a hydrogen molecule (H₂). Atoms of different elements can join — like one hydrogen and one chlorine forming a hydrogen chloride molecule (HCl). A few elements, like helium, are already stable as single atoms, so they do not bother joining.

Atoms join to reach a full, stable outer shell. They do this in two main ways:

• Sharing electrons — two atoms share one or more pairs of electrons. This makes a covalent bond.
• Transferring electrons — one atom gives electrons away and another takes them. This makes charged ions held by an ionic bond.

In both cases, the joined-up arrangement has lower energy than the separate atoms, which is exactly why it is more stable. The force that holds atoms together is called a chemical bond. Let us look at each kind.

Covalent Bonds — Atoms That Share

When two atoms share a pair of electrons, both atoms get to count that pair as their own. The shared pair sits between the two nuclei and is pulled by both — that pull is what holds them together. This is a covalent bond.

Molecules of elements (same atom joins same atom)

Start with hydrogen. A hydrogen atom has 1 electron in its K-shell, but the K-shell can hold 2. So it needs just 1 more. Two hydrogen atoms each share their single electron, making one shared pair. Now both atoms “see” 2 electrons — full and stable. We write it H—H (one line = one shared pair = a single bond).

Chlorine has 7 valence electrons and needs 1 more for an octet. Two chlorine atoms share one pair, so each reaches 8. We write it Cl—Cl — also a single bond.

Oxygen has 6 valence electrons and needs 2 more. So two oxygen atoms share two pairs of electrons. Two shared pairs is a double bond, written O=O (two lines). Figure 9.4 shows all three.

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This also answers an old question: what is the difference between O and O₂? A single O is one lone oxygen atom, which is unstable on its own. O₂ is the stable oxygen molecule — two atoms double-bonded. The air we breathe is O₂.

Molecules of compounds (different atoms join)

Now different elements. Hydrogen chloride (HCl): hydrogen needs 1 electron, chlorine needs 1 electron. Perfect — they share one pair, written H—Cl.

Water (H₂O): oxygen needs 2 electrons, but each hydrogen can only offer 1. So oxygen pairs up with two hydrogen atoms — each hydrogen shares one pair with the oxygen. That uses up oxygen’s 2 needs and completes both hydrogens. The result has 2 hydrogen atoms and 1 oxygen atom, so we write H₂O. Figure 9.5 shows both.

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Naming covalent compounds

Covalent compounds are named by counting the atoms of each element, using Greek prefixes: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6). The first element keeps its name; the second element ends in -ide. A few rules:

• mono- is dropped for the first element but kept for the second (CO is carbon monooxide → “carbon monoxide”, not “monocarbon dioxide”).
• If a prefix ending in ‘o’ or ‘a’ meets a vowel, drop that vowel: mono + oxide → “monoxide”; penta + oxide → “pentoxide”.
• Keep an ‘i’ for sound: di + oxide → “dioxide”.

Some examples: CO = carbon monoxide, CO₂ = carbon dioxide, CS₂ = carbon disulfide, PCl₃ = phosphorus trichloride, SF₆ = sulfur hexafluoride, N₂O₅ = dinitrogen pentoxide. When hydrogen is the first element, we skip its prefix: H₂S is just “hydrogen sulfide”. And a few have famous common names — H₂O is “water”, NH₃ is “ammonia”.

Ionic Bonds — Atoms That Give and Take

Some atoms do not share — they hand electrons over completely. This usually happens when one atom has few valence electrons (easy to give away) and the other has many (eager to grab a few more).

Take common salt, sodium chloride (NaCl). Sodium has just 1 valence electron — far easier to lose 1 than to gain 7. Chlorine has 7 — far easier to gain 1 than to lose 7. So sodium gives its 1 electron to chlorine.

But now think carefully. A sodium atom had 11 protons (+) and 11 electrons (−), perfectly balanced. After it loses 1 electron, it has 11 protons but only 10 electrons. The positives now outnumber the negatives by one. So it carries a +1 charge. A charged atom is called an ion. A positive ion is a cation, written Na⁺.

Chlorine had 17 protons and 17 electrons. After gaining 1 electron, it has 17 protons and 18 electrons — one extra negative. So it carries a −1 charge. A negative ion is an anion, written Cl⁻.

Now the two ions have opposite charges, and opposite charges pull together. That electrostatic pull is the ionic bond. Figure 9.6 traces the whole story.

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Some atoms give or take two electrons. Sulfur has 6 valence electrons and gains 2 to become S²⁻. Magnesium loses 2 to become Mg²⁺.

Ionic compounds form crystals, not molecules

Here is an important surprise. Ionic compounds do not float around as single “NaCl molecules”. Instead, billions of Na⁺ and Cl⁻ ions pack into a giant, repeating 3-D grid called a crystal lattice. In it, every Na⁺ is surrounded by 6 Cl⁻, and every Cl⁻ by 6 Na⁺.

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So when we write NaCl, we are not naming one molecule. We are giving the simplest whole-number ratio of ions (1 Na⁺ to 1 Cl⁻). This ratio-unit is called a formula unit.

Naming ionic compounds, and polyatomic ions

For ionic compounds, write the cation first, then the anion, and the simple anion ends in -ide (sodium chloride, calcium oxide, magnesium sulfide). Metals usually form cations; non-metals usually form anions.

Some ions are made of several atoms joined together but carrying one overall charge. These are polyatomic ions — like sulfate (SO₄²⁻), nitrate (NO₃⁻), carbonate (CO₃²⁻), hydroxide (OH⁻) and ammonium (NH₄⁺). Their names usually do not end in -ide.

You will need the common ions and their valencies (the charge number) to write formulae. Here are the ones to know.

Writing Chemical Formulae the Quick Way

You can always work out a formula by counting shared or transferred electrons. But there is a faster trick: the criss-cross method.

The idea is simple. Each part of the compound brings a number — its valency (for covalent) or its charge number (for ions). You write the symbols, write each number below, then swap the numbers so each becomes the subscript of the other symbol. Finally, reduce the subscripts to the simplest ratio.

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A few rules to remember:

• A subscript that comes out as 1 is not written (so we write CaCl₂, not Ca₁Cl₂).
• Always reduce the subscripts by any common factor (Mg₂O₂ becomes MgO).
• For a polyatomic ion taken more than once, wrap it in brackets: magnesium hydroxide is Mg(OH)₂ (two OH⁻ ions), not MgOH₂.
• The charges/valencies do not appear in the final formula.

Let us write a couple together.

Telling Ionic and Covalent Compounds Apart

Because their bonds are so different, ionic and covalent compounds behave very differently. You can sort them by simple tests: do they dissolve in water or in petrol/kerosene, and do they conduct electricity?

The most useful test is electricity. This finally answers a puzzle from earlier classes: why does salt water conduct electricity, but sugar water does not?

Electric current is just moving charges. Salt is ionic: in water its Na⁺ and Cl⁻ ions break apart and float free, ready to carry charge — so the bulb glows. Sugar is covalent: it dissolves but stays as neutral molecules, making no free ions — so no current flows. And solid salt does not conduct either, because its ions are locked tight in the crystal and cannot move. Free, moving charges are the whole point.

Working Out the Mass of a Compound

Once you have the formula, finding the mass of a compound is just careful adding. Each atom has an atomic mass measured in u (the unified atomic mass unit — the standard unit for weighing atoms). Add up the masses of all the atoms in the formula.

For a covalent compound we call this the molecular mass. For an ionic compound (which has no molecules, only a formula unit) we call the same total the formula unit mass. The method is identical. Figure 9.9 shows both.

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Let us do one with a polyatomic ion, where brackets matter.

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Summary

You can now explain:

• The Law of Conservation of Mass — that total mass of reactants equals total mass of products, and why (atoms only rearrange, they are never made or destroyed). You can spot the “escaping gas” trick that makes mass seem to change.
• The Law of Constant Proportions — that a pure compound always has its elements in the same fixed ratio by mass, and use that ratio to find unknown masses.
• Dalton’s atomic theory — its postulates, and how the idea of indivisible atoms combining in whole-number ratios explains both laws.
• How atoms join to become stable: by sharing electrons (covalent bonds, single or double) or by transferring electrons to form ions (cations and anions) held by ionic bonds.
• Why ionic compounds form crystal lattices (not molecules), and why they conduct electricity in water but not as solids.
• How to write a chemical formula with the criss-cross method, and how to name covalent and ionic compounds.
• How to calculate molecular mass and formula unit mass by adding up atomic masses.

What’s Next

You have seen how tiny atoms join, vibrate and arrange themselves into everything around you. When those particles vibrate in a special back-and-forth way, they can pass energy through the air as something you can hear. That is the topic of Chapter 10 — Sound Waves: Characteristics and Applications, where we explore how sound is made, how it travels, and why it has pitch and loudness.