The last chapter said that everything is made of atoms. This one opens them up. An atom is a tiny, heavy, positive nucleus with electrons around it, and three numbers say almost everything chemistry needs: how many protons, how many neutrons, how many electrons. The protons decide which element it is, the neutrons which isotope, and the electrons whether it is a neutral atom or an ion. By the end you will also know why the periodic table says chlorine weighs 35.45 when no chlorine atom does.
By the end you can
- Give the charge, mass and place of protons, neutrons and electrons.
- Read and write a nuclide symbol such as ³⁵Cl and count its particles.
- Explain what an ion and an isotope are, and how they differ.
- Compute an element's average atomic mass from its isotopes, and read a mass spectrum.
In 1808 John Dalton turned the old Greek guess about atoms into a scientific theory, built to explain the laws of the last chapter. Its four claims, and how they held up:
| Dalton's claim (1808) | Today |
|---|---|
| every element is made of tiny atoms | true |
| all atoms of one element are identical | same chemistry, but they can differ in mass: isotopes |
| compounds join atoms in small whole-number ratios | true: H₂O, CO₂, NH₃ |
| reactions rearrange atoms; they never create or destroy them | true for chemistry (nuclear reactions do change atoms) |
Dalton's atoms had no parts. Then J. J. Thomson found in 1897 that every atom contains electrons, light particles with a negative charge, and in 1911 Ernest Rutherford's team showed that the positive charge and almost all the mass sit in a tiny nucleus at the centre. The nucleus is about 100 000 times smaller than the atom: if an atom were a football stadium, its nucleus would be a pea at the centre spot, and the rest would be the electrons' territory. How the electrons are arranged around it is quantum mechanics, worked out in Physics, "The Hydrogen Atom"; chemistry mostly needs to count them.
| Particle | Charge | Mass (u) | Where |
|---|---|---|---|
| proton (p⁺) | +1 | 1.007 276 | in the nucleus |
| neutron (n) | 0 | 1.008 665 | in the nucleus |
| electron (e⁻) | −1 | 0.000 549 | around the nucleus |
Charges are given in units of e = 1.602 × 10⁻¹⁹ coulomb, the size of the electron's charge, so a proton is +1 and an electron −1 exactly. Masses are given in u, the atomic mass unit (defined below; 1 u ≈ 1.66 × 10⁻²⁷ kg). A proton and a neutron weigh almost the same, about 1 u each, while an electron is 1836 times lighter. So nearly all of an atom's mass is in its nucleus, and nearly all of its volume, and all of its chemistry, belongs to the electrons.
The number of protons in the nucleus is the atomic number, Z. It is the element's identity card: every atom with 6 protons is carbon, every atom with 79 is gold, and changing Z means changing the element. The total number of protons and neutrons, the particles of the nucleus (together called nucleons), is the mass number, A. It is a count, not a mass, and always a whole number. A particular kind of nucleus, with given Z and A, is called a nuclide, written with A on top and Z below, to the left of the symbol.
A neutral atom has as many electrons as protons, so the charges cancel. In chemical reactions the nucleus is never touched (that would take the energies of nuclear physics), but electrons can be moved. An atom that loses electrons is left with more protons than electrons: a positive ion, or cation. One that gains electrons becomes a negative ion, or anion. The charge is written as a superscript: Na⁺, Mg²⁺, Cl⁻, O²⁻.
| Ion | Z | Electrons | How it formed |
|---|---|---|---|
| Na⁺ | 11 | 10 | sodium lost 1 electron |
| Mg²⁺ | 12 | 10 | magnesium lost 2 |
| Cl⁻ | 17 | 18 | chlorine gained 1 |
| O²⁻ | 8 | 10 | oxygen gained 2 |
Notice that Na⁺, Mg²⁺ and O²⁻ all end up with 10 electrons, the same number as the noble gas neon. That is no accident, and it is the key to which ions form; the bonding chapters build on it. For now, build some atoms and ions yourself.
Add and remove particles with the buttons. Change the protons and the element itself changes: 6 is always carbon, 7 always nitrogen. Change the neutrons and you get another isotope of the same element: same chemistry, different mass, and only a few combinations last (carbon-12 and carbon-13 are stable, carbon-14 decays slowly, carbon-8 cannot hold together at all). Change the electrons and you get an ion: the nucleus stays, only the charge moves. Notice the mass line: the parts add up to slightly MORE than the measured mass, because binding the nucleus releases energy (Physics, "The Nucleus").
Atoms of the same element always have the same number of protons, but they can have different numbers of neutrons. These variants are the isotopes of the element (Greek for "same place": they sit in the same box of the periodic table). Isotopes have the same electrons, so the same chemistry: carbon-12 and carbon-13 burn, bond and taste the same. They differ in mass, and sometimes in the stability of the nucleus.
| Isotope | p / n | Share in nature | Note |
|---|---|---|---|
| hydrogen-1 | 1 / 0 | 99.988 % | the only nuclide with no neutron |
| hydrogen-2, deuterium (D) | 1 / 1 | 0.012 % | "heavy water" D₂O is 11 % denser than water |
| hydrogen-3, tritium (T) | 1 / 2 | traces | radioactive, half-life 12.3 years |
| carbon-12 | 6 / 6 | 98.93 % | defines the atomic mass unit |
| carbon-13 | 6 / 7 | 1.07 % | stable |
| carbon-14 | 6 / 8 | traces | radioactive, half-life 5730 years: radiocarbon dating |
Atoms are far too light for kilograms to be convenient. Chemists use the atomic mass unit u (also called the dalton, Da), defined as exactly one twelfth of the mass of a carbon-12 atom. On that scale a hydrogen-1 atom weighs 1.007 825 u and an oxygen-16 atom 15.994 915 u. The masses are close to the mass numbers but never exactly equal (except carbon-12, by definition), for two reasons: a proton and a neutron each weigh a bit more than 1 u, and binding a nucleus releases energy, which takes away a little mass (E = mc², Physics, "Relativistic Energy").
Natural chlorine is a mixture: about 75.76 % of its atoms are chlorine-35 (34.969 u) and 24.24 % are chlorine-37 (36.966 u). The mix is the same in practically every sample on Earth, so a handful of chlorine atoms always weighs, on average, the same. That average, weighted by how common each isotope is, is the atomic mass (or standard atomic weight) printed in the periodic table.
The isotope masses and shares above were measured with a mass spectrometer. It works in four steps. Ionise: knock an electron off each atom, so it becomes a positive ion that electric and magnetic fields can steer. Accelerate: send all the ions through the same voltage, so they all get the same energy. Deflect: pass them through a magnetic field, which bends each path into a circle. Detect: count the ions where they land. Ions with the same charge and energy follow a circle whose radius depends only on their mass, so each isotope lands in its own place.
Every ion leaves the source with the same energy and enters the field. The magnetic force bends its path into a circle, and a heavier ion, harder to turn, follows a wider circle and lands farther along the detector. Watch the bars: they grow in proportion to how common each isotope is, and the weighted average of the counts settles on the atomic mass printed in the periodic table. That is what the table's mass is: not the mass of any single atom, but the average over the natural mix. Chlorine has no atom of mass 35.45; it is about three atoms of 35 for every one of 37. (Real paths differ by a few percent; the drawing spreads them out.)
1. How many protons, neutrons and electrons are in ⁵⁶Fe³⁺? Iron has Z = 26, so 26 protons; N = 56 − 26 = 30 neutrons; the charge is +3, so 26 − 3 = 23 electrons.
2. An ion has 16 protons, 18 neutrons and 18 electrons. Z = 16 is sulfur, A = 16 + 18 = 34, charge = 16 − 18 = −2: it is ³⁴S²⁻.
3. Magnesium is 78.99 % ²⁴Mg (23.985 u), 10.00 % ²⁵Mg (24.986 u) and 11.01 % ²⁶Mg (25.983 u). Average: 0.7899 × 23.985 + 0.1000 × 24.986 + 0.1101 × 25.983 = 18.946 + 2.499 + 2.861 = 24.305 u, the table's value.
Working backwards: copper's isotopes
4. Copper's atomic mass is 63.546 u, and it has two isotopes, ⁶³Cu (62.930 u) and ⁶⁵Cu (64.928 u). What share is copper-63? Call it x; then copper-65 is 1 − x. So 63.546 = 62.930 x + 64.928 (1 − x) = 64.928 − 1.998 x, giving x = (64.928 − 63.546) / 1.998 = 0.692: about 69 % copper-63 and 31 % copper-65.
| Wrong | Right | Why |
|---|---|---|
| a different number of neutrons makes a different element | it makes a different isotope | only the number of protons fixes the element |
| Na⁺ gained a positive charge | Na⁺ lost an electron | ions form by moving electrons; protons never move in chemistry |
| mass number = atomic mass | A is a whole-number count; the mass in u is not | ³⁵Cl has A = 35 but a mass of 34.969 u |
| average the isotope masses directly | weight each by its abundance | (35 + 37)/2 = 36 ignores that ³⁵Cl is three times as common |
| Cl⁻ has 17 electrons | it has 18 | a negative charge means extra electrons: 17 + 1 |
Key ideas
- An atom is a tiny nucleus of protons (+1) and neutrons (0), about 1 u each, surrounded by electrons (−1), 1836 times lighter.
- The atomic number Z (protons) fixes the element; the mass number A = Z + N counts the nucleons.
- Ions form when atoms lose electrons (cations) or gain them (anions): charge = Z − electrons.
- Isotopes have the same Z and different N: same chemistry, different mass.
- 1 u is one twelfth of the mass of a carbon-12 atom.
- The atomic mass in the table is the abundance-weighted average of the isotope masses, measured with a mass spectrometer.