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1.2 Atomic structure

1.2.1 Development of the atomic model

Models of the atom change with new evidence

Definition

Atomic model

A scientific description of what an atom contains and how its parts are arranged.

  1. Atoms are far too small to see, so scientists build models to explain how atoms behave in experiments.
  2. A model is kept for as long as it explains all the experimental evidence.
  3. When a new experiment gives a result the model cannot explain, the model is changed or replaced.
  4. The main steps ran from Dalton in the early 1800s, through Thomson, Rutherford and Bohr, to Chadwick in 1932.
  5. Each new model kept the ideas that still fitted the evidence and changed only the parts that did not.
Key Idea

Every change to the atomic model followed a new experimental result, so the reason for each change is the evidence behind it.

Thomson's electron discovery replaced the solid sphere model

Definition

Electron

A tiny, negatively charged particle found in shells around the nucleus of an atom.

Definition

Plum pudding model

Thomson's model of the atom as a ball of spread-out positive charge with negative electrons embedded through it.

  1. In the early 1800s, John Dalton proposed that elements are made of tiny atoms that cannot be divided.
  2. Dalton pictured each atom as a solid sphere, with every atom of one element identical and different from the atoms of other elements.
  3. In 1897, J. J. Thomson found that cathode rays are streams of tiny negative particles, far lighter than any atom.
  4. These particles came from inside atoms, so atoms had to contain smaller parts.
  5. Atoms have no overall charge, so Thomson reasoned that positive charge must balance the negative electrons, and in 1904 he proposed the plum pudding model.
Example
  • In a currant bun, the dough stands for the spread-out positive charge and the currants stand for the electrons.
  • Nothing in the bun is concentrated in one place, which is the idea the next experiment tested.

Rutherford's gold foil experiment revealed a tiny nucleus

Definition

Alpha particle

A small, fast-moving, positively charged particle given out by some radioactive substances.

Definition

Nucleus

The tiny, dense centre of an atom, which carries its positive charge and almost all of its mass.

  1. In 1909, Geiger and Marsden, working with Ernest Rutherford, fired alpha particles at a very thin sheet of gold foil.
  2. The plum pudding model predicted that every alpha particle would pass through with at most a tiny deflection, because its positive charge was spread thinly.
  3. Most alpha particles did pass straight through, which showed that most of the atom is empty space.
  4. A few were deflected through large angles, and about 1 in 8000 bounced back, which only a small, concentrated positive charge could cause.
  5. In 1911, Rutherford replaced the plum pudding model with the nuclear model, in which electrons surround a central nucleus.
Common Mistake
  • Do not say that most alpha particles were repelled or bounced back, because most passed straight through the foil.
  • Do not credit the large deflections to electrons, because electrons are far too light to turn back a fast alpha particle.

Bohr placed electrons in shells at fixed distances

Definition

Electron shell

A fixed distance from the nucleus at which electrons orbit, with each shell at its own energy level.

  1. Rutherford's nuclear model placed the electrons around the nucleus but did not describe the electron arrangement.
  2. In 1913, Niels Bohr suggested that electrons are not free to sit anywhere around the nucleus.
  3. Bohr's theoretical calculations agreed with measurements of the light given out by heated atoms, which strongly supported the shell idea.
  4. His model kept Rutherford's nucleus and added a definite structure for the electrons around it.
  5. Shell diagrams, with electrons drawn on rings around a central nucleus, are still used to show how electrons are arranged in atoms.
Key Idea

Bohr's shells give each element a definite electron arrangement, which later links an element's electrons to its place in the Periodic Table.

Chadwick's discovery of the neutron completed the model

Definition

Neutron

A particle in the nucleus with no charge and a relative mass of 1.

  1. By about 1920, Rutherford had shown that the positive charge of the nucleus is carried by particles called protons.
  2. Nuclei turned out to be heavier than their protons alone could explain, which suggested that the nucleus held another particle.
  3. In 1932, James Chadwick provided experimental evidence for the neutron.
  4. The resulting model has a tiny nucleus of protons and neutrons, with electrons in shells around it.
Exam technique
  • When a question asks why the model changed, name the experiment or discovery first.
  • Then say what the new evidence showed that the earlier model could not explain, such as large deflections that a spread-out positive charge could not produce.
Self review
  • What did Dalton think atoms were like?
  • Which discovery showed that atoms contain smaller particles?
  • What did most alpha particles passing straight through the gold foil show?
  • Why could the plum pudding model not explain the alpha particles that bounced back?
  • What did Bohr and then Chadwick add to the nuclear model?

1.2.2 The nuclear model and the size of atoms

An atom has a positive nucleus and negative electrons

Definition

Atom

The smallest part of an element that can exist.

Definition

Nucleus

The tiny, dense centre of an atom, which carries its positive charge and almost all of its mass.

  1. The nucleus is made of protons and neutrons packed tightly together.
  2. The nucleus has a positive charge because protons are positive and neutrons have no charge.
  3. Electrons move around the nucleus in shells, and each one carries a negative charge.
  4. The positive nucleus attracts the negative electrons, and this attraction holds the atom together.
  5. An atom has no overall charge, because the positive charge of its nucleus is balanced by the negative charge of its electrons.
Common Mistake
  • Do not place electrons inside the nucleus, because only protons and neutrons are found there.
  • Do not describe the nucleus as neutral, because its protons give it an overall positive charge.

The nucleus is tiny compared with the atom

Definition

Atomic radius

The distance from the centre of an atom's nucleus to its outermost electrons.

  1. The nuclear radius is less than a ten-thousandth of the atomic radius.
  2. The electrons move in the large region around the nucleus, so almost all of an atom's volume is empty space.
  3. This is why most alpha particles passed straight through the gold foil in Rutherford's experiment.
  4. An atom drawn to scale would show the nucleus as a dot far too small to see.
  5. Textbook diagrams draw the nucleus much larger than its true size so that it shows up at all.
Example
  • If an atom were enlarged to the size of a football stadium, its nucleus would be about the size of a pea on the centre spot.
  • The electrons would be moving around out in the stands, with empty space in between.

An atom's mass is concentrated in its nucleus

  1. Protons and neutrons are much heavier than electrons.
  2. The electrons add so little that more than 99.9% of an atom's mass is in the nucleus.
  3. Because so much mass sits in so little space, the nucleus is extremely dense.
  4. The electrons take up nearly all of the atom's volume but contribute almost none of its mass.
  5. The mass of an atom is therefore set almost entirely by the number of protons and neutrons in its nucleus.
Key Idea

The nucleus holds almost all of the mass but very little of the volume, while the electrons fill almost all of the volume but add almost no mass.

The typical size of atoms and small molecules

Definition

Order of magnitude

The power of ten closest to a value, used to give its approximate size.

  1. A typical atom has a radius of about 1×10−10 m1 \times 10^{-10}\,\text{m}1×10−10m, which is 0.1 nm0.1\,\text{nm}0.1nm.
  2. The diameter of that atom, straight across its centre, is twice the radius, about 2×10−10 m2 \times 10^{-10}\,\text{m}2×10−10m.
  3. Small molecules, such as water, carbon dioxide or methane, contain only a few atoms, so their size is also of the order of 10−10 m10^{-10}\,\text{m}10−10m.
  4. Different atoms have different radii, but nearly all of them fall within the same order of magnitude.
  5. Using 1×10−10 m1 \times 10^{-10}\,\text{m}1×10−10m as the radius lets you estimate how many atoms fit across an everyday object.
Example
  • Estimate how many magnesium atoms fit across a strip of magnesium ribbon 3.0 mm3.0\,\text{mm}3.0mm wide.
  • Convert the width to metres: 3.0 mm=3.0×10−3 m3.0\,\text{mm} = 3.0 \times 10^{-3}\,\text{m}3.0mm=3.0×10−3m.
  • Double the typical atomic radius to get the width of one atom: 2×(1×10−10 m)=2×10−10 m2 \times (1 \times 10^{-10}\,\text{m}) = 2 \times 10^{-10}\,\text{m}2×(1×10−10m)=2×10−10m.
  • Divide the ribbon width by the width of one atom: 3.0×10−3 m2×10−10 m=1.5×107\dfrac{3.0 \times 10^{-3}\,\text{m}}{2 \times 10^{-10}\,\text{m}} = 1.5 \times 10^{7}2×10−10m3.0×10−3m​=1.5×107 atoms.
Self review
  • What charge does the nucleus carry, and why?
  • How does the radius of a nucleus compare with the radius of its atom?
  • Where is almost all of an atom's mass found?
  • What is the typical radius of an atom, in metres?
  • Roughly how many atoms would fit across a line 1 mm1\,\text{mm}1mm long?

1.2.3 Subatomic particles in atoms and ions

Relative charge and mass of subatomic particles

Definition

Relative charge

The charge on a particle compared with the charge on a proton, which is given a relative charge of +1.

Definition

Relative mass

The mass of a particle compared with the mass of a proton, which is given a relative mass of 1.

  1. A proton has a relative charge of +1+1+1 and a relative mass of 111.
  2. A neutron has a relative charge of 000 and a relative mass of 111.
  3. An electron has a relative charge of −1-1−1 and a relative mass of about 12000\frac{1}{2000}20001​, or 0.00050.00050.0005.
  4. Protons and neutrons sit in the nucleus, while electrons are found in shells around it.
  5. Almost all of an atom's mass comes from its protons and neutrons, because each electron is so light.
Common Mistake
  • Do not give the electron a relative mass of 000, because it is very small but not zero.
  • Write the electron's relative charge as −1-1−1, since a minus sign on its own is not a charge value.
  • Relative charges and relative masses have no units, because each one is a comparison with a proton.

Atomic and mass numbers give protons and neutrons

Definition

Atomic number

The number of protons in the nucleus of an atom.

Definition

Mass number

The total number of protons and neutrons in the nucleus of an atom.

  1. In the symbol 1123Na{}^{23}_{11}\text{Na}1123​Na, the mass number 232323 is written at the top left and the atomic number 111111 at the bottom left.
  2. Every atom of an element has the same atomic number, so the atomic number alone tells you which element it is.
  3. The Periodic Table lists the elements in order of atomic number, starting with hydrogen at 111.
  4. The number of neutrons is the difference between the two numbers: neutrons=mass number−atomic number\text{neutrons} = \text{mass number} - \text{atomic number}neutrons=mass number−atomic number
  5. In a neutral atom, the number of electrons equals the number of protons, because the charges must balance.
Example
  • Find the numbers of protons, neutrons and electrons in an atom of 919F{}^{19}_{9}\text{F}919​F.
  • The atomic number is 999, so the atom has 999 protons.
  • Subtract the atomic number from the mass number: 19−9=1019 - 9 = 1019−9=10 neutrons.
  • The atom is neutral, so it has 999 electrons.

Isotopes differ in their number of neutrons

Definition

Isotope

An atom of an element with the same number of protons as other atoms of that element but a different number of neutrons.

  1. Carbon-12 and carbon-14, written 612C{}^{12}_{6}\text{C}612​C and 614C{}^{14}_{6}\text{C}614​C, both have 666 protons.
  2. Carbon-12 has 12−6=612 - 6 = 612−6=6 neutrons, while carbon-14 has 14−6=814 - 6 = 814−6=8 neutrons.
  3. The number after the name, as in carbon-14, is the mass number of that isotope.
  4. Chlorine-35 and chlorine-37 both have 171717 protons, but they have 181818 and 202020 neutrons respectively.
  5. Isotopes of an element have the same number of electrons, so they react in the same way.
Common Mistake

Do not say that isotopes have different numbers of electrons or protons, because only the number of neutrons differs.

Ions form when atoms lose or gain electrons

Definition

Ion

An atom or group of atoms that carries an overall electric charge.

  1. A positive ion forms when an atom loses electrons, leaving more protons than electrons.
  2. A negative ion forms when an atom gains electrons, giving more electrons than protons.
  3. The number of protons never changes when an ion forms, because a different proton number would mean a different element.
  4. The number of neutrons also stays the same, so an ion has the same mass number as its atom.
  5. To find the electrons in an ion, start from the number of protons, then subtract the size of the charge for a positive ion or add it for a negative ion.
Example
  • Find the numbers of protons, neutrons and electrons in the ion 1327Al3+{}^{27}_{13}\text{Al}^{3+}1327​Al3+.
  • The atomic number is 131313, so the ion has 131313 protons.
  • Subtract the atomic number from the mass number: 27−13=1427 - 13 = 1427−13=14 neutrons.
  • The 3+3+3+ charge means three electrons were lost: 13−3=1013 - 3 = 1013−3=10 electrons.
Self review
  • What are the relative charge and relative mass of a neutron?
  • Why is the relative mass of an electron not written as zero?
  • How do the atoms of chlorine-35 and chlorine-37 differ?
  • How many protons, neutrons and electrons are in 1632S2−{}^{32}_{16}\text{S}^{2-}1632​S2−?
  • Why does forming an ion never change the number of protons?

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An atomic model is a scientific description of what an atom contains and how its parts are arranged. Because atoms are too small to see directly, models are tested against experimental evidence.

Dalton described atoms as indivisible solid spheres. In 1897, Thomson discovered negatively charged electrons inside atoms, proving that atoms contain smaller particles.

Because atoms are neutral overall, Thomson proposed the plum pudding model: electrons embedded in a ball of spread-out positive charge. New evidence later showed that this positive charge was not spread throughout the atom.

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1.2 Atomic structure Revision Guide

  1. GCSE
  2. /Chemistry
  3. /1.2 Atomic structure

Revision notes for OCR GCSE Chemistry 1.2 Atomic structure: explanations and worked examples on 1.2.1 Development of the atomic model, 1.2.2 The nuclear model and the size of atoms, and 1.2.3 Subatomic particles in atoms and ions.

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