2.1.1 Development of the atomic model

Dalton's atom: a solid sphere with nothing inside it

1. An atomic model is a scientific description of what an atom contains and how its parts are arranged.
2. In the early 1800s John Dalton described each atom as a tiny solid sphere that could not be split.
3. Dalton's model treated atoms as indivisible, so it allowed nothing smaller to exist inside them.
4. It fitted the evidence of the time, because elements combined in fixed ratios as though built from identical unsplittable units.
5. Every later change to the model came from finding a subatomic particle that Dalton's solid sphere had no room for.

Thomson: the electron turns the sphere into a plum pudding

1. In 1897 J. J. Thomson used cathode-ray experiments to show that atoms contain electrons, which carry a negative charge.
2. An electron is far smaller than an atom, so atoms could no longer be described as indivisible.
3. Thomson proposed the plum pudding model, with electrons embedded in a ball of positive charge.
4. The positive charge balanced the negative electrons, so the atom as a whole stayed electrically neutral.

Rutherford: alpha scattering puts the positive charge in a nucleus

1. In the alpha-scattering experiment, positively charged alpha particles were fired at a very thin sheet of gold foil.
2. Most alpha particles passed straight through, which showed that most of the atom is empty space.
3. A small number were deflected through large angles, which showed a concentrated region of positive charge.
4. A very small number bounced almost straight back, which showed that this region is both tiny and dense.
5. Rutherford concluded that nearly all the positive charge and nearly all the mass sit in a tiny central nucleus.
6. That replaced the spread-out positive charge of the plum pudding model and left the electrons somewhere outside the nucleus.

Bohr and Chadwick: shells for the electrons, neutrons in the nucleus

1. Niels Bohr showed that electrons orbit at fixed distances from the nucleus rather than anywhere outside it.
2. Calculations based on fixed energy levels matched the experimental observations, which Rutherford's model could not explain.
3. Later work identified the positively charged particles inside the nucleus as protons.
4. James Chadwick provided evidence for the neutron in 1932, explaining why nuclei are heavier than their protons alone can account for.
5. The current model therefore has a tiny nucleus of protons and neutrons, with electrons in shells around it.

Why the model changed: each discovery exposed a limit in the one before

1. Dalton's solid sphere became Thomson's plum pudding once the electron was found inside the atom.
2. The plum pudding became Rutherford's nuclear model once alpha scattering showed the positive charge was concentrated.
3. The nuclear model became Bohr's shell model once fixed energy levels were needed to match what was observed.
4. Bohr's model gained neutrons once Chadwick accounted for the missing nuclear mass.
5. A model survives only while it explains every observation, so new evidence is what forces each change.

2.1.2 Structure of the atom and subatomic particles

Inside the atom: a nucleus of protons and neutrons, electrons in shells

1. The nucleus sits at the centre of the atom and holds the protons and the neutrons.
2. The electrons move around the nucleus, occupying shells at fixed distances from it.
3. Protons, neutrons and electrons are the three subatomic particles that make up an atom.

1. An atom carries no overall charge, so it holds as many electrons as protons.

Relative charge and relative mass of each particle

1. Relative charge compares a particle's charge with the charge on a proton, which is set at $+1$.
2. Relative mass compares a particle's mass with the mass of a proton, which is set at $1$.
3. Proton: relative charge $+1$, relative mass $1$, found in the nucleus.
4. Neutron: relative charge $0$, relative mass $1$, found in the nucleus.
5. Electron: relative charge $-1$, relative mass $\frac{1}{1840}$, found in a shell around the nucleus.
6. A proton and a neutron therefore contribute about the same mass, while an electron contributes almost none.

Why an atom is neutral: protons and electrons cancel exactly

1. A proton carries $+1$ and an electron carries $-1$, so one of each cancels to zero.
2. Equal numbers of protons and electrons give equal amounts of positive and negative charge.
3. Those charges cancel, so the atom has an overall charge of zero.
4. Neutrons carry no charge, so they take no part in this balance.
5. A particle whose protons and electrons are unequal is an ion, not an atom.

Size and mass: a tiny nucleus that carries nearly all the mass

1. The nucleus has a radius of roughly $\frac{1}{10000}$ that of the whole atom.
2. Almost all of the atom is therefore empty space through which the electrons move.
3. Each proton and each neutron has relative mass $1$, while each electron has relative mass $\frac{1}{1840}$.
4. Adding in the electrons changes the total mass by an amount too small to matter.
5. Nearly all of the atom's mass therefore sits in the nucleus, even though the nucleus occupies almost none of its volume.
6. The electrons fix the size of the atom while the nucleus fixes its mass.

2.1.3 Mass number, isotopes and relative atomic mass

Atomic number and mass number: counting the particles in an atom

1. Every atom of a given element has the same number of protons, and no other element shares that number.
2. Carbon always has 6 protons, so carbon's atomic number is 6.
3. A neutral atom holds as many electrons as protons, so the atomic number also gives the electron count.
4. The mass number counts the protons and neutrons together, because those two particles carry essentially all the mass.
5. On a periodic table entry the larger number is the relative atomic mass, not the mass number, and the smaller is the atomic number.

Isotopes: same element, different numbers of neutrons

1. Isotopes of one element share an atomic number, because they contain the same number of protons.
2. They have different mass numbers, because they contain different numbers of neutrons.
3. Chlorine-35 and chlorine-37 are both chlorine, so each atom has 17 protons.
4. A chlorine-35 atom has $35-17=18$ neutrons.
5. A chlorine-37 atom has $37-17=20$ neutrons.
6. Extra neutrons change the mass but not the chemical reactions, because the electron arrangement is unchanged.

Counting the particles from the two numbers

1. The number of protons is equal to the atomic number.
2. The number of electrons in a neutral atom is also equal to the atomic number.
3. The number of neutrons is the mass number minus the atomic number: $$neutrons=mass\; number-atomic\; number$$
4. Sodium with atomic number 11 and mass number 23 has 11 protons, 11 electrons and $23-11=12$ neutrons.
5. Chlorine-37 with atomic number 17 has 17 protons, 17 electrons and $37-17=20$ neutrons.
6. Check that your proton count matches the element named in the question before you accept the answer.

Why relative atomic mass is often not a whole number

1. Every individual atom has a whole-number mass number, because an atom cannot hold part of a neutron.
2. A real sample of an element contains a mixture of isotopes in fixed proportions.
3. Relative atomic mass averages those isotopes, weighting each one by how common it is.
4. Chlorine is about three-quarters chlorine-35 and one-quarter chlorine-37, which averages to $35.5$.
5. The value sits closer to $35$ than to $37$ because chlorine-35 is the more abundant isotope.
6. A value such as $35.5$ describes the sample, never any single atom in it.

2.1.4 Calculating relative atomic mass from isotopic abundance

The two numbers a question gives you

1. A question supplies the mass of each isotope and how common that isotope is.
2. An isotope's relative isotopic mass is the same number as its mass number.
3. Abundance may be given as a percentage, as a decimal fraction, or as a count of atoms detected.
4. The more abundant isotope pulls the average towards its own mass.
5. Relative atomic mass has no unit, because it compares one mass with another.

The method: multiply, add, then divide by the total abundance

1. Write each isotopic mass and its abundance side by side before calculating anything.
2. Multiply each isotopic mass by its own abundance.
3. Add all of those products together.
4. Divide by the total abundance: $$A_r=\frac{(m_1\times a_1)+(m_2\times a_2)}{a_1+a_2}$$
5. The denominator is $100$ when the abundances are percentages that add to $100$.
6. Extend the same pattern for three or more isotopes by adding one product per isotope.

Worked example: weighting three isotopes

1. An element has isotopes of mass $24$, $25$ and $26$, with abundances of $79\%$, $10\%$ and $11\%$.
2. The abundances add to $100$, so the denominator is $100$.
3. Work out the three products separately before adding, so that a slip in one does not hide inside the total.

Checking your answer

1. Your answer must lie between the lowest and the highest isotopic mass in the data.
2. It must sit nearer the mass of the most abundant isotope.
3. Do not round the intermediate products, because rounding early shifts the final value.
4. Do not take a plain average of the masses, because the isotopes are not equally abundant.
5. Give $A_r$ with no unit, and to the precision the question asks for.