Topic 1 Atomic structure and the periodic table
How atomic structure dictates the periodic table’s organisation and trends. Ignite your curiosity and test your knowledge!
Quiz
An atom is the smallest part of an element that can exist. Think of it as the ultimate LEGO brick. Everything in the universe the air you breathe, the water you drink, and even you is made of atoms.
Did you know? Atoms are incredibly empty! If an atom were the size of a football stadium, its nucleus (the central part) would only be the size of a marble.
An element is a pure substance that is made up of only one type of atom. There are over 100 known elements, which are all organised in the Periodic Table.
Common Examples:
Hydrogen (H)
Carbon (C)
Oxygen (O)
Sodium (Na)
Exam Tip: You are given a periodic table in your exam, but learning the names and symbols of the first 20 elements (from Hydrogen to Calcium) will save you valuable time.
A compound is a substance formed when two or more different elements are chemically bonded together. These elements combine in a fixed ratio. For example, a molecule of water always consists of two hydrogen atoms bonded to one oxygen atom.
Common Examples:
Water (H₂O) – made from hydrogen and oxygen.
Table Salt (NaCl) – made from sodium and chlorine.
Carbon Dioxide (CO₂) – made from carbon and oxygen.
Calcium Carbonate (CaCO₃) – made from calcium, carbon, and oxygen.
In chemistry, we use shorthand to represent elements and compounds.
A symbol is a one or two-letter code for an element. The first letter is always uppercase, and if there is a second letter, it’s always lowercase. For example, Ca is the symbol for calcium.
A formula shows which atoms are in a compound and how many of each there are. The small number written after a symbol (a subscript) tells you how many of that atom are in the molecule. For example, in CO₂, the ‘2’ indicates there are two oxygen atoms for every one carbon atom.
A binary compound is made of only two different elements. There is a simple rule for naming them:
Name the first element (usually the metal).
Name the second element, but change the ending to “-ide”.
Examples:
NaCl: Sodium chloride
CaS: Calcium sulfide
Al₂O₃: Aluminium oxide
Heads up! The “-ide” ending is for simple compounds with only two elements. Later, you’ll meet compounds with endings like “-ate” (e.g., copper sulfate), which tells you that oxygen is also present.
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Chemical reactions can be represented in two main ways:
Word Equation: Uses the names of the substances involved. The chemicals that react are called reactants, and the new substances formed are called products.
Magnesium + Oxygen ⟶ Magnesium oxide (Reactants) ⟶ (Product)
Symbol Equation: Uses the chemical formulae instead of names. A key rule for symbol equations is that they must be balanced. This means there must be the same number of atoms of each element on both sides of the arrow.
2Mg + O₂ ⟶ 2MgO
In this balanced equation, there are 2 magnesium atoms and 2 oxygen atoms on the left (reactants) and the right (products).
A mixture consists of two or more different elements or compounds that are not chemically bonded together.
Because no chemical reaction occurs, each substance in the mixture keeps its own original properties. This also means we can separate the components of a mixture using physical methods. A great example of a mixture is salt water; it’s just salt and water mixed together, and you can separate them again.
Exam Tip: Remember to state that in a mixture, no new chemical substance is formed. The properties of the mixture are simply the combined properties of the substances it contains.
Since mixtures aren’t chemically combined, we can separate them using different physical techniques. The method you choose depends on the properties of the substances in the mixture.
This technique is used to separate an insoluble solid (a solid that doesn’t dissolve) from a liquid.
How it works: The mixture is poured through filter paper. The tiny pores in the paper allow the liquid particles to pass through, but they are too small for the larger solid particles, which get trapped.
The liquid that passes through is called the filtrate.
The solid left behind in the paper is called the residue.
Common Example: Separating sand from water.
This method is used to separate a soluble solid (a solid that has dissolved) from a solution, such as separating salt from salt water.
How it works:
The solution is gently heated in an evaporating basin to boil off some of the solvent (the liquid), making the solution more concentrated.
Once small crystals start to form, the heat is removed.
The solution is left to cool slowly. As it cools, the solid becomes less soluble, and pure crystals form.
The crystals are then filtered out and left to dry.
Use this technique to separate a liquid solvent from a dissolved solid when you want to keep the liquid.
How it works:
Evaporation: The solution is heated, and the liquid with the lower boiling point turns into a vapour.
Condensation: The vapour travels into a condenser (a cold tube), where it cools down and turns back into a pure liquid.
Collection: The pure liquid drips into a separate beaker for collection.
Common Example: Producing pure water from salt water.
This method is used to separate a mixture of two or more miscible liquids (liquids that mix together) with different boiling points.
How it works: The setup is similar to simple distillation but includes a fractionating column filled with glass beads. This column provides a temperature gradient (hottest at the bottom, coolest at the top). The liquid with the lowest boiling point evaporates first, rises highest up the column, and passes into the condenser to be collected. Liquids with higher boiling points will condense at lower levels and fall back down.
Key Example: Separating the different components of crude oil.
Chromatography is used to separate different soluble, coloured substances, like the different dyes in an ink pen.
How it works:
A spot of the mixture (e.g., ink) is placed on a pencil line near the bottom of a piece of chromatography paper. This paper is the stationary phase.
The bottom edge of the paper is placed into a solvent (e.g., water). This solvent is the mobile phase.
As the solvent soaks up the paper, it dissolves the ink spot and carries the different coloured dyes with it.
Dyes that are more soluble and less attracted to the paper will travel further up, separating them out.
You can calculate an Rf value for each spot to identify the substance:
Our modern picture of the atom wasn’t thought of overnight. It was built up over centuries, with each new model revising the last based on new experimental evidence. This is a story of scientific discovery, from solid spheres to electron shells.
The journey began with an idea. The Greek philosopher Democritus proposed that if you kept cutting matter into smaller and smaller pieces, you would eventually reach a tiny, indivisible particle. He called this an “atomos”, which means “uncuttable”.
In the early 19th century, John Dalton brought the idea of the atom into modern science. Based on his experiments, he proposed that:
All matter is made of atoms, which he imagined as tiny, solid, indivisible spheres.
Atoms of a specific element are all identical to each other.
Atoms of different elements have different masses.
J.J. Thomson’s experiments with cathode ray tubes led to a major discovery: the electron. He found that atoms contained tiny, negatively charged particles. This meant atoms were not indivisible after all.
He adapted the atomic model to include this new discovery, suggesting the atom was a sphere of positive charge with negative electrons dotted inside it, like plums in a pudding. This became known as the plum pudding model.
Ernest Rutherford conducted the famous gold foil experiment, which completely changed our understanding of the atom’s structure. He fired tiny, positive alpha particles at a very thin sheet of gold foil.
The Evidence and Conclusions:
Most alpha particles passed straight through the foil. This proved that the atom is mostly empty space.
Some alpha particles were deflected at small angles. This showed that the centre of the atom had a positive charge that repelled the positive alpha particles.
A very small number bounced straight back. This meant there must be a tiny, dense, and positively charged nucleus at the centre of the atom where most of its mass is concentrated.
Exam Tip: When answering questions about the gold foil experiment, you must mention both key findings:
Most alpha particles passed through, showing the atom is mostly empty space.
A few were deflected or rebounded, showing a small, dense, positive nucleus.
Niels Bohr refined Rutherford’s model. He realised that if electrons were just orbiting the nucleus, they would eventually spiral inwards.
He proposed that electrons must exist in fixed energy levels, or shells, at set distances from the nucleus. An electron can “jump” to a higher shell if it absorbs energy, and it will emit that energy (often as light) when it falls back to a lower shell.
The Evidence: The unique line spectra produced by elements (like hydrogen emitting light only at specific wavelengths) supported the idea that electrons could only exist at specific energy levels. Bohr’s model was a crucial step towards our modern understanding of atomic structure.
We now know that atoms are not the smallest things that exist. They are made of even smaller, fundamental sub-atomic particles. There are three you need to know: protons, neutrons, and electrons.
The structure of an atom is actually very simple. It’s mostly empty space, with two key regions: a tiny, dense nucleus at the centre, and the surrounding electron shells.
The nucleus is found at the very centre of the atom and contains almost all of its mass. It is made up of two types of particles:
Protons: These have a positive (+) electric charge.
Neutrons: These have no electric charge; they are neutral.
Electrons are tiny particles with a negative (-) electric charge. They are not in the nucleus; instead, they orbit the nucleus in specific energy levels called shells. Electrons have a mass so small that it is considered negligible when calculating the mass of an atom.
It is essential to learn the relative mass and charge of each particle.
After identifying the sub-atomic particles, the next step is to understand their arrangement and the scale of the atom. The key takeaway is that an atom is almost entirely empty space, with its mass packed into an incredibly small centre.
Atoms are unimaginably small, and the nucleus is even smaller in comparison.
The atomic radius (the distance from the nucleus to the outermost electron shell) is approximately 1 × 10⁻¹⁰ metres.
The radius of the nucleus is about 1 × 10⁻¹⁴ metres.
This means the nucleus is roughly 10,000 times smaller than the entire atom. To put that in perspective, if an atom were the size of a football stadium, the nucleus would be the size of a pea placed in the centre.
Despite its tiny size, the nucleus contains nearly all of the atom’s mass.
This is because protons and neutrons (found in the nucleus) each have a relative mass of 1.
Electrons, which orbit the nucleus, have a negligible mass (approximately 1/2000th of a proton or neutron).
Therefore, when calculating the mass of an atom, the mass of the electrons is considered to be zero.
Exam Tip: Be prepared to use standard form to compare the sizes of the atom and its nucleus. For example, an atom (radius 1 × 10⁻¹⁰ m) is 10,000 (or 10⁴) times larger than its nucleus (radius 1 × 10⁻¹⁴ m).
So far, we’ve assumed that all atoms of an element are identical. However, that’s not always the case. This is where the concept of isotopes comes in.
Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons.
Because isotopes of an element have the same number of protons, they also have the same number of electrons. Since electrons determine how an atom reacts, isotopes have identical chemical properties.
The key difference between isotopes is their mass number.
Atomic (Proton) Number: The number of protons. This defines the element.
Mass Number: The total number of protons and neutrons in the nucleus.
Let’s look at carbon as an example. All carbon atoms have 6 protons.
Carbon-12 has 6 protons and 6 neutrons (Mass number = 12).
Carbon-13 has 6 protons and 7 neutrons (Mass number = 13).
Carbon-14 has 6 protons and 8 neutrons (Mass number = 14).
They are all carbon atoms, but they have different masses due to the different number of neutrons.
If elements can have atoms with different masses, how do we list a single mass for them on the periodic table? We use the relative atomic mass ().
The relative atomic mass is the weighted mean mass of an atom of an element, taking into account its naturally occurring isotopes.
This is why some elements on the periodic table have mass numbers that are not whole numbers. Chlorine, for example, has a relative atomic mass of 35.5. This is because it is a weighted average of its two main isotopes: Chlorine-35 and Chlorine-37.
You can calculate the using this formula:
Ar = Isotope abundance X Isotope mass number / 100
Worked Example: Chlorine
In nature, chlorine exists as two isotopes:
75% is Chlorine-35
25% is Chlorine-37
Step 1: Multiply the abundance of each isotope by its mass number.
(75 × 35) = 2625
(25 × 37) = 925
Step 2: Add these values together.
2625 + 925 = 3550
Step 3: Divide by 100.
3550 / 100 = 35.5
So, the relative atomic mass () of chlorine is 35.5
The way electrons are arranged in an atom is fundamental to chemistry. It dictates how an element behaves, why it reacts, and where it sits in the periodic table. The rules for this arrangement are simple but powerful.
Electrons orbit the nucleus in specific energy levels called shells. These shells are filled from the inside out, starting with the one closest to the nucleus.
Electrons always occupy the lowest available energy level first (the innermost shell).
Each shell has a maximum number of electrons it can hold. For GCSE, you need to know the capacity for the first four shells:
Shell 1: holds a maximum of 2 electrons.
Shell 2: holds a maximum of 8 electrons.
Shell 3: holds a maximum of 8 electrons.
The electron configuration is a way of writing down how many electrons are in each shell. You can work it out from the atomic number (the small number on the periodic table), as this tells you the number of electrons in a neutral atom.
Example 1: Sodium (Na)
Sodium has an atomic number of 11, so it has 11 electrons.
The first shell fills with 2 electrons.
The second shell fills with 8 electrons. (2 + 8 = 10 electrons used so far).
The remaining 1 electron goes into the third shell.
The electron configuration for Sodium is written as 2.8.1.
Example 2: Chlorine (Cl)
Chlorine has an atomic number of 17, so it has 17 electrons.
The first shell fills with 2 electrons.
The second shell fills with 8 electrons.
The third shell gets the remaining 7 electrons.
The electron configuration for Chlorine is written as 2.8.7
| Z | Element | Symbol | Electronic structure (by shells) |
|---|---|---|---|
| 1 | Hydrogen | H | 1 |
| 2 | Helium | He | 2 |
| 3 | Lithium | Li | 2,1 |
| 4 | Beryllium | Be | 2,2 |
| 5 | Boron | B | 2,3 |
| 6 | Carbon | C | 2,4 |
| 7 | Nitrogen | N | 2,5 |
| 8 | Oxygen | O | 2,6 |
| 9 | Fluorine | F | 2,7 |
| 10 | Neon | Ne | 2,8 |
| 11 | Sodium | Na | 2,8,1 |
| 12 | Magnesium | Mg | 2,8,2 |
| 13 | Aluminium | Al | 2,8,3 |
| 14 | Silicon | Si | 2,8,4 |
| 15 | Phosphorus | P | 2,8,5 |
| 16 | Sulfur | S | 2,8,6 |
| 17 | Chlorine | Cl | 2,8,7 |
| 18 | Argon | Ar | 2,8,8 |
| 19 | Potassium | K | 2,8,8,1 |
| 20 | Calcium | Ca | 2,8,8,2 |
The electrons in the outermost shell are the most important. They are called valence electrons.
The number of valence electrons determines the chemical properties of an element.
Elements in the same group in the periodic table have the same number of valence electrons, which is why they have similar chemical properties.
An atom is most stable when it has a full outer shell. The noble gases (Group 0) have full outer shells, which is why they are so unreactive.
Other atoms will react by gaining, losing, or sharing electrons to achieve a stable, full outer shell. This is the basis for all chemical bonding.
The periodic table is a masterfully organised chart of all the known chemical elements. It’s not just a random list; it’s arranged to reveal patterns in the properties of elements, a concept known as periodicity.
The modern periodic table arranges elements in order of increasing atomic number (Z).
Atomic Number (Z): This is the number of protons in an atom’s nucleus. It is the fundamental characteristic that defines an element—think of it as the element’s unique ID number.
Relative Atomic Mass (): This is the larger number shown for an element. It represents the weighted average mass of its naturally occurring isotopes. While crucial for calculations, it is not used to determine the order of the elements in the table.
Exam Tip: A common mistake is to say the table is ordered by relative atomic mass. Always state that the modern periodic table is ordered by atomic number.
The table’s layout in rows and columns gives us vital information about each element’s atomic structure.
The rows across the periodic table are called periods. The period number tells you how many occupied electron shells an element has.
Period 1 elements (H, He) have 1 electron shell.
Period 2 elements (Li to Ne) have 2 electron shells.
Period 3 elements (Na to Ar) have 3 electron shells, and so on.
The columns down the periodic table are called groups. For the main groups, the group number tells you the number of electrons in the element’s outermost shell (valence electrons).
Key Fact: Elements in the same group have similar chemical properties because they have the same number of outer shell electrons.
Examples:
Group 1 (Alkali Metals): Have 1 outer electron, making them very reactive.
Group 7 (Halogens): Have 7 outer electrons. They are also very reactive as they only need one more electron to get a full outer shell.
Group 0 (Noble Gases): Have a full outer shell, making them stable and very unreactive.
Exam Tip: You can use an element’s group number to quickly predict how it will react and what kind of bonds it will form.
Periodicity is the repeating pattern of chemical and physical properties that you see as you move across each period. Trends in properties like reactivity, melting point, and boiling point repeat in a predictable way from one period to the next. This happens because the electron configurations of the elements change in a regular, repeating pattern.
The periodic table we use today is built on the foundations laid by earlier scientists. The most important of these was the Russian chemist Dmitri Mendeleev, who published his version in 1869.
Mendeleev’s great insight was to arrange the known elements in order of increasing relative atomic mass. While doing this, he noticed that a periodic pattern emerged in their properties.
His genius was demonstrated in two key ways:
He left gaps: Mendeleev was so confident in his patterns that he left gaps for elements that he believed had not yet been discovered but should exist.
He made predictions: He predicted the properties and masses of these missing elements. For example, he predicted the existence of “eka-aluminium” and “eka-silicon”.
When elements like Gallium (discovered in 1875) and Germanium (discovered in 1886) were found, their properties almost perfectly matched Mendeleev’s predictions. This provided powerful evidence that his arrangement was correct.
Despite its success, Mendeleev’s table had some inconsistencies. When ordered strictly by atomic mass, some pairs of elements, like Argon (Ar) and Potassium (K), appeared to be in the wrong order to fit the pattern of chemical properties.
The puzzle was solved by the discovery of isotopes—atoms of the same element with different numbers of neutrons. This showed that using atomic mass for ordering was flawed because an element’s chemical identity comes from its proton number, not its mass.
This led to the modern arrangement.
The modern periodic table orders the elements by increasing atomic number (Z), which is the number of protons in the nucleus.
This principle is fundamental and has no exceptions.
Ordering by proton number resolves all the inconsistencies from Mendeleev’s table.
It correctly places elements with similar chemical properties in the same groups.
Exam Tip: When asked why the modern table is ordered by atomic number instead of atomic mass, you should mention that the existence of isotopes meant that ordering by mass could incorrectly place some elements. Ordering by atomic number, which is unique to each element, fixes this problem.
The periodic table has a fundamental dividing line that separates elements into two broad categories: metals and non-metals. This distinction is crucial as it reflects major differences in their physical and chemical properties.
Generally, metals are found on the left-hand side and in the centre of the periodic table, while non-metals are located on the right-hand side.
The best way to understand the differences is to compare their typical properties side-by-side.
| Property | Metals | Non-Metals |
|---|---|---|
| Appearance | Lustrous (shiny) | Dull (if solid) |
| Mechanical | Malleable, ductile | Brittle (if solid) |
| Conductivity | Good conductors of heat & electricity | Poor conductors |
| Density & Hardness | Generally high density & hardness | Generally lower density & variable |
Exam Tip: When asked to classify an element, first state its position on the periodic table (e.g., “It’s on the left, so it’s a metal”). Then, support this by citing two typical physical or chemical properties from the list above.
The type of oxide an element forms is a key chemical indicator of whether it is a metal or a non-metal.
Metal oxides react with acids in neutralisation reactions to produce a salt and water. For example, magnesium oxide (a white powder) reacts with hydrochloric acid to form magnesium chloride and water.
Non-metal oxides typically react with water to form acidic solutions. For instance, sulfur dioxide (a gas produced when sulfur burns) dissolves in water to form sulfurous acid.
The elements in the far right-hand column of the periodic table form Group 0. They are known as the noble gases. This group includes elements like Helium (He), Neon (Ne), and Argon (Ar).
Their defining characteristic is that they are chemically inert, meaning they are extremely unreactive.
The key to the noble gases’ stability lies in their electron arrangement. They all have a full outer shell of electrons.
Helium has an electron configuration of 2. Its first and only shell is full.
Neon has an electron configuration of 2.8. Its second shell is full.
Argon has an electron configuration of 2.8.8. Its third shell is full.
A full outer shell is the most stable arrangement for an atom. Because of this, noble gases have no tendency to lose, gain, or share electrons to form bonds, which is what happens in a chemical reaction.
This stability also means they don’t join up with other atoms. They exist as individual, single atoms, a property known as being monatomic.
Exam Tip: Whenever you are asked to explain why noble gases are unreactive, you must mention that they have a stable electron arrangement because they have a full outer shell.
All noble gases are colourless gases at room temperature. Although they don’t react chemically, they do show a clear trend in their physical properties.
Trend in Boiling Point:
The boiling points of the noble gases increase as you go down the group.
Explanation:
As you move down the group, the atoms become larger and have more electrons.
This leads to stronger, weak intermolecular forces of attraction between the atoms.
More energy is needed to overcome these stronger forces to turn the liquid into a gas.
Therefore, the boiling point increases.
The elements in Group 1 of the periodic table, located on the far left-hand side, are known as the alkali metals. This group includes elements like Lithium (Li), Sodium (Na), and Potassium (K).
Their chemical behaviour is defined by their electron structure. All Group 1 elements have just one electron in their outermost shell.
Lithium (Li): 2.1
Sodium (Na): 2.8.1
Potassium (K): 2.8.8.1
Alkali metals have physical properties that are very different from typical metals like iron or copper.
They are very soft and can be easily cut with a knife.
They have low densities (lithium, sodium, and potassium are all less dense than water, so they float).
They have low melting and boiling points, and these decrease as you go down the group.
The chemistry of alkali metals is all about that single outer electron. In reactions, they lose this electron to form a stable, full outer shell and a positive ion (with a +1 charge).
Trend in Reactivity:
Reactivity increases as you go down the group.
Explanation:
As you go down the group, the outer electron is in a shell that is further away from the nucleus.
There are more inner shells, which shield the outer electron from the positive pull of the nucleus.
This weaker attraction means the outer electron is lost more easily, making the atom more reactive.
Exam Tip: When explaining the trend in reactivity for Group 1, always refer to the single outer electron and how easily it is lost due to increasing atomic size and shielding.
The alkali metals all react vigorously with water to produce a metal hydroxide and hydrogen gas. You can clearly see the increase in reactivity as you go down the group.
Lithium floats and moves slowly across the surface, fizzing steadily as it produces hydrogen gas. It eventually disappears, leaving a solution of lithium hydroxide.
2Li + 2H₂O ⟶ 2LiOH + H₂
Sodium’s reaction is more vigorous. It melts into a silvery ball that darts quickly across the water’s surface, fizzing intensely.
2Na + 2H₂O ⟶ 2NaOH + H₂
The reaction with potassium is very rapid and vigorous. The heat produced is enough to ignite the hydrogen gas, which burns with a characteristic lilac flame.
2K + 2H₂O ⟶ 2KOH + H₂
The elements in Group 7, the second column from the right on the periodic table, are known as the halogens. This group includes elements like Fluorine (F), Chlorine (Cl), and Bromine (Br).
Their chemical behaviour is determined by their electron structure. All halogen atoms have seven electrons in their outermost shell.
Fluorine (F): 2.7
Chlorine (Cl): 2.8.7
Bromine (Br): 2.8.18.7
To achieve a stable, full outer shell, they share one electron with another halogen atom. This is why they exist as diatomic molecules (X₂), with the two atoms joined by a single covalent bond.
The halogens show clear trends in their physical properties as you move down the group.
| Element | Formula | State at RTP | Colour | Relative molecular mass | Melting & boiling points |
|---|---|---|---|---|---|
| Fluorine | F₂ | Gas | Pale yellow | Lightest | Lowest |
| Chlorine | Cl₂ | Gas | Green | ↑ | ↑ |
| Bromine | Br₂ | Liquid | Red-brown | ↑ | ↑ |
| Iodine | I₂ | Solid | Grey-black | Heaviest | Highest |
Trend in Melting and Boiling Points:
The melting points and boiling points of the halogens increase as you go down the group.
Explanation:
As you move down the group, the diatomic molecules become larger and have more electrons.
This leads to stronger intermolecular forces of attraction between the molecules.
More energy is needed to overcome these stronger forces to change state.
Therefore, the melting and boiling points increase.
In reactions, halogens need to gain one electron to achieve a stable, full outer shell. They do this by gaining an electron to form a negative ion (with a -1 charge) or by sharing electrons in covalent bonds.
Trend in Reactivity:
Reactivity decreases as you go down the group.
Explanation:
As you go down the group, the outer electron shell is further away from the nucleus.
There is more shielding by the inner electron shells.
This makes it harder for the nucleus to attract an extra electron to complete the outer shell.
Therefore, reactivity decreases.
The trend in reactivity can be shown by displacement reactions. A more reactive halogen will displace a less reactive halide from an aqueous solution of its salt.
For example, chlorine is more reactive than bromine. If you bubble chlorine gas through a solution of potassium bromide, the chlorine will displace the bromide ions.
While both alkali metals (Group 1) and transition metals are classified as metals, they have very different physical and chemical properties.
The table below summarises the key differences between these two groups of metals.
Exam Tip: When comparing the properties of two metals, try to include specific examples with data if you can. For example, “Sodium is a soft, low-density metal with a low melting point of 98 °C, whereas iron is a hard, dense metal with a high melting point of 1538 °C.”
Transition metals exhibit a set of key properties that distinguish them from the metals in Group 1 and Group 2. These properties arise from their unique electron structures.
Transition metals have strong metallic bonding. The delocalised electrons in their structure create powerful forces of attraction that hold the atoms tightly together. As a result, they typically have:
High melting points and boiling points.
High densities.
A distinctive feature of transition metals is that their compounds are often coloured. This is in sharp contrast to Group 1 compounds, which are almost always white. The colour of the solution can often help identify which transition metal ion is present.
Common Examples:
Solutions containing Copper(II) ions (Cu²⁺) are a characteristic blue colour.
Solutions containing Iron(III) ions (Fe³⁺) are a pale yellow-brown colour.
Unlike Group 1 metals which only ever form +1 ions, most transition metals can form positive ions with different charges. This is known as having variable oxidation states.
Common Examples:
Iron can form the Fe²⁺ ion (Iron(II)) and the Fe³⁺ ion (Iron(III)).
Copper can form the Cu⁺ ion (Copper(I)) and the Cu²⁺ ion (Copper(II)).
Many transition metals and their compounds are excellent catalysts, meaning they speed up chemical reactions without being used up themselves. They are vital in many industrial processes.
Common Examples:
Iron is used as the catalyst in the Haber process for making ammonia. N₂ + 3H₂ ⇌ 2NH₃
Copper(II) ions can catalyse the decomposition of hydrogen peroxide.