- The structure of the atom has fascinated scientists for centuries.
- Over time, as new experiments and technologies were developed, our ideas about atoms have changed and improved.
- The current quantum model of the atom is based on over a hundred years of experimental evidence.
- It gives us a clearer picture of the tiny building blocks of matter and how they behave.
In the early 19th century, John Dalton proposed one of the first scientific models of the atom.
Dalton suggested that:
- All matter is made up of tiny particles called atoms.
- Atoms are indivisible – they cannot be split into smaller parts.
- Atoms of the same element are identical and have the same mass and properties.
- Atoms of different elements are different and have different masses.
- A chemical reaction is simply a rearrangement of atoms; atoms are not created or destroyed in the process.
Although we now know that atoms can be split into smaller subatomic particles, Dalton’s ideas were an essential starting point for modern atomic theory.
- Later experiments in the late 19th and early 20th centuries showed that Dalton’s model was incomplete:
- Electrons were discovered (showing atoms are not indivisible).
- Protons and neutrons were discovered in the nucleus.
- We now know that a typical atom contains three main subatomic particles:
- Protons
- Neutrons
- Electrons
- These particles are the key building blocks of atomic structure.
- The proton has a positive charge and is found in the nucleus.
- The neutron has no charge (it is neutral) and is also found in the nucleus.
- The electron has a negative charge and is found outside the nucleus, in regions we call energy levels or electron shells.
- In simple diagrams, electrons are often shown as “orbiting” the nucleus like planets around the Sun.
- In reality, in the quantum model, electrons occupy regions of space called orbitals, but the shell model is still useful at MYP/IB level for basic structure.
The masses of subatomic particles are extremely small, so we use a special unit called the atomic mass unit, symbol u.
| Subatomic particle | Symbol | Relative charge | Relative mass | Approximate mass (g) | Location |
|---|
| Proton | p⁺ | +1 | 1 | 1.6726 × 10$^{-24}$ | Nucleus |
| Neutron | n | 0 | 1 | 1.6749 × 10$^{-24}$ | Nucleus |
| Electron | e⁻ | -1 | 1/1836 | 9.1094 × 10$^{-28}$ | Electron shells |
Two key numbers describe an atom:
- Atomic number (Z)
- Mass number (A)
- If two atoms have different numbers of protons, they are different elements.
- In a neutral atom, the number of electrons is equal to the number of protons, so: $$\text{Neutral atom} \implies \text{electrons} = \text{protons} = Z$$
$$A=\text { number of protons }+ \text { number of neutrons }$$
We often call the number of neutrons $N$, so:
$$A=Z+N \quad \text { or } \quad N=A-Z$$
If you know A and Z, you can always find the number of neutrons using
$$N=A-Z$$
Knowing $Z$ and $A$ helps chemists to:
- Describe the composition of an atom (how many protons, neutrons and electrons it has).
- Predict how atoms might behave in chemical reactions.
Let’s take the isotope sodium-23.
- The atomic number of sodium is 11, so every sodium atom has:
- 11 protons in its nucleus.
- Sodium-23 has a mass number of 23.
- Mass number: $A=23$
- Atomic number: $Z=11$
- To find the number of neutrons: $$N=A−Z=23−11=12$$ So sodium-23 has 12 neutrons.
- In a neutral sodium atom, the number of electrons equals the number of protons:
- So a neutral sodium-23 atom has:
- 11 protons
- 12 neutrons
- 11 electrons
An oxygen atom has a mass number of 16 and an atomic number of 8. How many neutrons does this oxygen atom have?
At the centre of every atom is a tiny region called the nucleus.
- The nucleus contains protons and neutrons only.
- These particles together are called nucleons.
- The nucleus is very small compared with the whole atom, but it contains almost all of the atom’s mass.
- Because it contains positively charged protons, the nucleus has an overall positive charge.
- Electrons are found outside the nucleus in energy levels (shells).
- The number of protons in the nucleus (the atomic number, Z) decides which element the atom is.
- The total number of protons and neutrons in the nucleus (the mass number, A) decides which nuclide/isotope it is.
- A nuclide is a specific kind of atom defined by:
- its number of protons (atomic number, Z)
- its number of neutrons (which together with protons gives the mass number, A)
- So a nuclide is a specific isotope of an element with a particular combination of protons and neutrons.
- Atoms of the same element have the same number of protons, but can have different numbers of neutrons.
- These different versions are called isotopes, and each isotope is one or more specific nuclides.
We often write nuclides using the notation:
$$^A_Z E$$
where
- $E$ is the chemical symbol of the element
- $Z$ is the atomic number (number of protons)
- $A$ is the mass number (total number of nucleons in the nucleus, namely, protons + neutrons)
Chlorine has two common stable isotopes:
$$^{35}_{17}\text{Cl} \quad \text{and} \quad ^{37}_{17}\text{Cl}$$
Both are chlorine because they both have 17 protons.
- Nuclide 1: $^{35}_{17}\text{Cl}$
- Protons: 17
- Neutrons: $35−17=18$
- Electrons (neutral atom): 17
- Nuclide 2: $^{37}_{17}\text{Cl}$
- Protons: 17
- Neutrons: $37−17=20$
- Electrons (neutral atom): 17
These two nuclides are isotopes of chlorine. They have:
- the same number of protons and electrons (in the neutral atoms)
- different numbers of neutrons (18 vs 20)
Now consider the chloride ion:
$$^{37}_{17}\text{Cl}^-$$
- Protons: 17 (this never changes for chlorine)
- Neutrons: 20 (still the isotope with mass number 37)
- Electrons: 18 (one extra electron gives the −1 charge)
The chemical properties of isotopes of the same element are almost identical because:
- Chemical properties are mainly determined by the number and arrangement of electrons.
- For a given element, all isotopes have the same number of protons and (in neutral atoms) the same number of electrons.
- Both have 6 protons and (in neutral atoms) 6 electrons.
- They form the same types of chemical bonds and take part in the same reactions.
- However, their nuclei are different (different numbers of neutrons), so their nuclear properties (like stability and radioactivity) differ.
Because elements usually exist as a mixture of isotopes, the relative atomic mass takes into account:
- the mass of each isotope
- the abundance (how common) each isotope is
We often use percentage abundance data to calculate $A_r$.
Step 1 – Convert percentage abundance to decimal (fractional) abundance
Divide each percentage by 100.
For example, if an element has isotopes with abundances:
Then the fractional abundances are:
Step 2 – Multiply each isotope’s mass by its fractional abundance
For each isotope:
$$\text { (fractional abundance) × (isotopic mass) }$$
This gives how much that isotope contributes to the overall average mass.
Step 3 – Add all the contributions
Add up all the results from Step 2:
$$A_r=\left(\text { fractional abundance }_1 \times \text { mass }_1\right)+$$
$$+\left(\text { fractional abundance }_2 \times \text { mass }_2\right)+\ldots$$
This total is the relative atomic mass of the element.
Magnesium has three naturally occurring isotopes:
| Isotope | % abundance | Fractional abundance | Mass number |
|---|
| $^{24}_{12}\text{Mg}$ | 80.5% | 0.805 | 24 |
| $^{25}_{12}\text{Mg}$ | 10.0% | 0.100 | 25 |
| $^{26}_{12}\text{Mg}$ | 9.5% | 0.095 | 26 |
Using the mass numbers as approximate isotopic masses:
$$A_r(\mathrm{Mg}) \approx(0.805 \times 24)+(0.100 \times 25)+(0.095 \times 26)$$
$$A_r(\mathrm{Mg}) \approx 19.32+2.50+2.47=24.29 \approx 24.3$$
- Spectroscopy is the study of how light (electromagnetic radiation) interacts with matter.
- Electrons in atoms can absorb or emit energy in the form of light when they change energy levels.
- When an electron absorbs energy, it moves from a lower energy level to a higher one (an excited state). This can produce an absorption spectrum.
- When an electron falls back from a higher energy level to a lower one, it releases energy as light. This produces an emission spectrum.
- Each element produces a unique set of spectral lines because it has a unique set of allowed energy levels.
- These spectral lines are like atomic fingerprints.
- They allow scientists to identify elements in substances, including in distant stars.
The emission spectrum of hydrogen was crucial evidence for quantized (discrete) energy levels in atoms.
- When hydrogen gas is excited (for example, by an electric discharge), it emits light.
- When this light is passed through a prism or diffraction grating, it splits into a series of sharp colored lines at specific wavelengths.
- These lines show that the electrons in hydrogen can only exist at certain energy levels, and the energy differences between these levels correspond to the energies (and wavelengths) of the emitted photons.
Niels Bohr used the hydrogen emission spectrum to propose his model of the atom, in which:
- Electrons move in fixed energy levels (shells) around the nucleus.
- Electrons can move between these levels by absorbing or emitting fixed amounts (quanta) of energy.