Chemistry · Atomic structure and the periodic law · Grade 7

Atomic structure and isotopes

Why is the relative atomic mass of chlorine not a whole number, but 35.45?

Open the experiment

The situation

Every grain of table salt contains sodium and chlorine atoms. Yet in the periodic table the relative atomic mass of chlorine is 35.45 — not a whole number. What is an atom made of? What changes if one proton is added to it — and what if one neutron is added? And why does the uranium mined in Navoi region slowly turn into lead over billions of years?

Topic

Atomic structure and isotopes. Atom builder: add a proton and the element changes, add a neutron and you get an isotope — carbon-12 and carbon-14, chlorine-35 and chlorine-37: why Ar(Cl) = 35.45

Key ideas

• Proton, neutron, electron
• Isotopes
• Relative atomic mass

Formulas

Mass number: the sum of the numbers of protons ($Z$) and neutrons ($N$); so $N=A-Z$.

Number of electrons: $Z$ in a neutral atom; $Z-q$ in an ion of charge $q$ (10 in Na⁺, 18 in Cl⁻).

Relative atomic mass — the average of the isotope masses weighted by their natural abundances ($x_i$).

α-decay: the nucleus emits a ⁴He nucleus — $A$ decreases by 4 and $Z$ by 2.

β⁻-decay: a neutron turns into a proton and an electron is emitted — $Z$ increases by 1, $A$ does not change.

Theory

The theory of the atom

In the 5th century BC Democritus suggested that matter is made of indivisible particles — atoms. Dalton (1803–1808) linked this to chemical experiments: each element consists of its own atoms, and they differ in mass. Thomson discovered the electron in 1897 and proposed the “plum pudding” model (1904). From the scattering of α-particles by gold foil (the Geiger–Marsden experiment, 1909), Rutherford found that at the centre of the atom there is a small, heavy nucleus (1911). Bohr (1913) showed that electrons move in certain levels, and the present quantum model describes the electron as a “cloud” (orbital). In “Models” mode compare the six models; the scattering experiment in detail is “Rutherford's experiment” in the physics section.

The structure of the atom: proton, neutron, electron

At the centre of the atom is the nucleus: positive protons ($p^{+}$, charge $+1.602\cdot {10}^{-19}$ C, mass $1.6726\cdot {10}^{-27}$ kg ≈ 1.0073 u) and uncharged neutrons ($n^0$, ≈ 1.0087 u). Around it are negative electrons ($e^{-}$, $9.109\cdot {10}^{-31}$ kg — 1836 times lighter than the proton). The size of an atom is ≈ ${10}^{-10}$ m, that of the nucleus ${10}^{-15}$–${10}^{-14}$ m: the nucleus is about 100 000 times smaller than the atom, yet it holds more than 99.9 % of the mass. The atomic number $Z$ = the number of protons; in a neutral atom there are also $Z$ electrons. The mass number is $A=Z+N$.

Element, isotopes, isobars, ions

A chemical element is a kind of atom with a given number of protons: add a proton and the element changes. Isotopes are atoms with the same $Z$ and different $N$: the hydrogen isotopes protium ¹H, deuterium ²H (D) and tritium ³H (T); ³⁵Cl and ³⁷Cl. Because they have the same electrons, their chemical properties are almost the same. Isobars have the same mass number and different $Z$: ⁴⁰Ar, ⁴⁰K, ⁴⁰Ca. Isotones have the same number of neutrons: ¹³C and ¹⁴N (7 each). If an atom gives up an electron it forms a positive cation (Na⁺); if it gains an electron, a negative anion (Cl⁻); the nucleus does not change. On the $N$–$Z$ map in the builder, isotopes form a horizontal row, isotones a column and isobars a diagonal.

Relative atomic mass and the mass spectrometer

The atomic mass unit is $1\mathrm{u}=m({}^{12}\mathrm{C})/12=1.6605\cdot {10}^{-27}$ kg. The relative atomic mass $A_r$ is the ratio of the mass of an atom to 1 u; the $A_r$ of an element in the table is the average over the mixture of natural isotopes: $A_r(\mathrm{Cl})=0.7576\cdot 34.969+0.2424\cdot 36.966=35.45$. The fraction of each isotope is measured by the mass spectrometer: ions are accelerated through a voltage $U$ and in a magnetic field $B$ trace an arc of radius $r=\frac{1}{B}\sqrt{2mU/q}$ — heavy ions travel along a larger arc and land at a different place on the detector.

Radioactivity and nuclear reactions

Some nuclei are unstable: they spontaneously emit a particle and turn into another nucleus — radioactivity (A. Becquerel, 1896). In α-decay a ⁴He nucleus is emitted: $A$ decreases by 4 and $Z$ by 2. In β⁻-decay a neutron in the nucleus turns into a proton and an electron is emitted: $Z$ increases by 1 and $A$ does not change (the displacement rule). The time in which half of the nuclei decay is the half-life $T_{1/2}$: ≈ 5730 years for ¹⁴C, 4.47 billion years for ²³⁸U. ²³⁸U passes through 8 α-decays and 6 β⁻-decays to become stable ²⁰⁶Pb. There are also artificial nuclear reactions: Rutherford (1919) ${\phantom{X}}_{\phantom{7}}^{\phantom{14}}{\phantom{X}}_{\phantom{2}\phantom{7}7}^{\phantom{2}\phantom{14}14}\mathrm{N}+{\phantom{X}}_{\phantom{2}}^{\phantom{4}}{\phantom{X}}_{\phantom{2}\phantom{2}2}^{\phantom{2}\phantom{4}4}\mathrm{He}\stackrel{}{\to }{\phantom{X}}_{\phantom{8}}^{\phantom{17}}{\phantom{X}}_{\phantom{2}\phantom{8}8}^{\phantom{2}\phantom{17}17}\mathrm{O}+{\phantom{X}}_{\phantom{1}}^{\phantom{1}}{\phantom{X}}_{\phantom{2}\phantom{1}1}^{\phantom{2}\phantom{1}1}\mathrm{H}$, Chadwick (1932) ${\phantom{X}}_{\phantom{4}}^{\phantom{9}}{\phantom{X}}_{\phantom{2}\phantom{4}4}^{\phantom{2}\phantom{9}9}\mathrm{Be}+{\phantom{X}}_{\phantom{2}}^{\phantom{4}}{\phantom{X}}_{\phantom{2}\phantom{2}2}^{\phantom{2}\phantom{4}4}\mathrm{He}\stackrel{}{\to }{\phantom{X}}_{\phantom{6}}^{\phantom{12}}{\phantom{X}}_{\phantom{2}\phantom{6}6}^{\phantom{2}\phantom{12}12}\mathrm{C}+{\phantom{X}}_{\phantom{0}}^{\phantom{1}}{\phantom{X}}_{\phantom{2}\phantom{0}0}^{\phantom{2}\phantom{1}1}\mathrm{n}$ — the discovery of the neutron. In detail: “Radioactivity, half-life” and “The atomic nucleus and binding” in the physics section.

About the sources and the model

Particle masses — CODATA 2022; isotope masses — AME2020; natural abundances — IUPAC (2013, representative composition); half-lives — NUBASE2020. Model: the sizes are schematic (the nucleus is really 100 000 times smaller than the atom), and the Bohr shells are a simplified picture. In the builder $Z\le 20$; a combination of protons and neutrons that is not in the table is shown as “unstable”. In the gold foil scene the bounce-backs are shown more often than in the real experiment (in 1909, with platinum — about one in 8000), in the mass spectrometer the motion of the ions is greatly slowed down, and the decay chain has no time scale — one decay per press.

Did you know?

• The neutron was discovered in 1932 by the English physicist James Chadwick by bombarding beryllium with α-particles; for this discovery he received the Nobel Prize in Physics in 1935.
• The radiocarbon method (Willard Libby, Nobel Prize in Chemistry, 1960) rests on the fact that after an organism dies the share of ¹⁴C in it halves about every 5730 years: it can date remains up to about 50 thousand years old.

Inside the experiment

The experiment starts from a real-life situation: you predict the result, change the controls, run the experiment and draw a conclusion. Inside are missions, theory, a quiz and a guided tour in which the camera moves and explains the key points.

More experiments in this section

• Elements and the periodic table — Grade 7
• Electron shells and orbitals — Grade 8