Welcome to the Periodic Table!
Imagine walking into a massive supermarket where nothing is organised—milk next to shampoo, apples next to batteries. Finding anything would be a nightmare! Scientists once faced the exact same problem with the chemical elements. That is, until the Periodic Table was created.
The Periodic Table is the ultimate chemist’s map. In this chapter, you will learn how it was developed, how it is arranged today, and the fascinating patterns (or trends) shown by different families of elements. Don’t worry if this seems like a lot to take in at first—once you learn the simple rules, the Periodic Table actually gives you the answers!
1. The Development of the Periodic Table
The Periodic Table did not just appear overnight. It took decades of brilliant detective work by scientists, the most famous of whom was Russian chemist Dmitri Mendeleev.
Mendeleev's Periodic Table (1869)
Mendeleev arranged the known elements in order of increasing atomic mass (the weight of the atoms known at the time). However, he did two genius things that made his table legendary:
• He left gaps: Mendeleev realised that some elements had not yet been discovered. Instead of forcing elements into the wrong place, he left blank spaces and boldly predicted their properties (such as eka-silicon, which we now know as Germanium). When these elements were discovered later, his predictions were almost spot on!
• He swapped pairs of elements: Sometimes, arranging purely by atomic mass placed an element into a group with completely different chemical behaviour. Mendeleev swapped them (such as Tellurium, \( \text{Te} \), and Iodine, \( \text{I} \)) so that elements with similar chemical properties stayed in the same vertical column.
The Modern Periodic Table
Today, our table is slightly different because we now know about subatomic particles (protons, neutrons, and electrons):
• Elements are arranged in order of increasing atomic number (the number of protons), NOT atomic mass.
• There are no gaps—all elements up to atomic number \( 118 \) have been discovered or synthesised.
• Noble gases (Group 0) are included (they were completely unknown in Mendeleev's time because they are so unreactive).
• The Transition Metals are placed in a central block between Groups 2 and 3.
Quick Review: Mendeleev vs Modern Table
• Mendeleev: Arranged by atomic mass, left gaps, no noble gases.
• Modern: Arranged by atomic number (protons), no gaps, includes Group 0 and the transition block.
Key Takeaway: Mendeleev prioritised grouping elements with similar chemical properties together, paving the way for our modern table arranged by atomic number.
2. Layout of the Modern Periodic Table
Understanding the layout makes navigating the table easy. Let's look at the basic terms:
Groups (Vertical Columns)
• The columns running down the table numbered from \( 1 \) to \( 7 \), plus Group \( 0 \), are called Groups.
• Elements in the same group have the same number of electrons in their outer shell (e.g., Group 1 elements have \( 1 \) outer electron; Group 7 elements have \( 7 \) outer electrons).
• Because they have the same number of outer electrons, elements in the same group have similar chemical properties (they react in very similar ways).
Periods (Horizontal Rows)
• The rows running across the table are called Periods.
• The period number tells you the total number of electron shells an atom has (e.g., elements in Period 3 have \( 3 \) occupied electron shells).
Metals vs Non-Metals
• A stepped line (the "staircase") on the right side of the table separates metals from non-metals.
• Metals lie to the left and centre of the staircase (making up about \( 80\% \) of all elements). Metals generally conduct electricity, are shiny, malleable, and lose electrons to form positive ions.
• Non-metals lie to the far right (plus Hydrogen on the top left). Non-metals are poor conductors, brittle when solid, and gain or share electrons in reactions.
Key Takeaway: Group number = outer electrons (determines chemical behaviour); Period number = number of electron shells.
3. Group 1: The Alkali Metals
Group 1 elements include Lithium (\( \text{Li} \)), Sodium (\( \text{Na} \)), and Potassium (\( \text{K} \)). They are known as the Alkali Metals.
Physical Properties
Unlike typical everyday metals (like iron), alkali metals are unusual:
• They are very soft and can easily be cut with a butter knife.
• They have low densities (Lithium, Sodium, and Potassium actually float on water!).
• They have relatively low melting and boiling points for metals.
• They are shiny when freshly cut, but tarnish rapidly in air as they react with oxygen to form a dull metal oxide layer.
Storage and Safety
Because they react violently with oxygen and moisture in the air, alkali metals are stored under oil. When handling them, always wear safety goggles and use tongs—never bare hands!
Reactions with Water
When an alkali metal is dropped into water, it reacts vigorously to produce a metal hydroxide and hydrogen gas:
\( 2\text{M} + 2\text{H}_2\text{O} \rightarrow 2\text{MOH} + \text{H}_2 \) (where \( \text{M} \) is \( \text{Li} \), \( \text{Na} \), or \( \text{K} \))
The metal hydroxide dissolves to create an alkaline solution (turns universal indicator purple, \( \text{pH } 12\text{–}14 \)).
Observations for each metal with water:
• Lithium (\( \text{Li} \)): Floats, fizzes (effervescence of \( \text{H}_2 \) gas), moves steadily across the surface, and gradually disappears.
• Sodium (\( \text{Na} \)): Floats, fizzes rapidly, melts into a silvery ball (due to the heat released), moves quickly across the surface, and disappears.
• Potassium (\( \text{K} \)): Floats, moves very violently, burns with a characteristic lilac flame, fizzes loudly, and may end with a small pop/spark.
Trend in Reactivity (Group 1)
Reactivity increases as you go down Group 1 (\( \text{Li} < \text{Na} < \text{K} < \text{Rb} < \text{Cs} \)).
Why does reactivity increase down Group 1? (Exam Gold Standard Explanation!)
1. Group 1 atoms react by losing their \( 1 \) outer electron to get a stable full outer shell.
2. As you go down the group, atoms have more electron shells, so the atom gets larger (atomic radius increases).
3. The single outer electron is further away from the positively charged nucleus and is shielded by more inner electron shells.
4. This means there is a weaker electrostatic attraction between the nucleus and the outer electron.
5. Therefore, the outer electron is lost more easily, making the element more reactive.
Key Takeaway: Group 1 metals are soft, low-density metals that react vigorously with water to form alkaline hydroxides and hydrogen gas. Reactivity increases down the group because the outer electron is further away and easier to lose.
4. Group 2: The Alkaline Earth Metals
Group 2 includes Beryllium (\( \text{Be} \)), Magnesium (\( \text{Mg} \)), and Calcium (\( \text{Ca} \)).
• They have \( 2 \) electrons in their outer shell.
• They are harder, denser, and have higher melting points than Group 1 metals.
• They are less reactive than Group 1 elements in the same period because losing \( 2 \) electrons requires more energy than losing \( 1 \).
• Reactivity still increases down the group (Calcium reacts much faster with water than Magnesium does).
5. Group 7: The Halogens
Group 7 elements are known as the Halogens. They are reactive non-metals existing as diatomic molecules (pairs of atoms joined by a covalent bond, such as \( \text{F}_2 \), \( \text{Cl}_2 \), \( \text{Br}_2 \), \( \text{I}_2 \)).
States and Colours at Room Temperature
You must memorise the colours and physical states at room temperature (\( 20^\circ\text{C} \)):
• Fluorine (\( \text{F}_2 \)): Toxic yellow gas.
• Chlorine (\( \text{Cl}_2 \)): Toxic pale green / yellow-green gas.
• Bromine (\( \text{Br}_2 \)): Dense red-brown liquid that readily gives off an orange-brown vapour.
• Iodine (\( \text{I}_2 \)): Shiny grey-black solid that sublimes (turns straight from solid to gas) into a stunning purple vapour when heated.
Trend in Reactivity (Group 7)
Reactivity decreases as you go down Group 7 (\( \text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2 \)).
Why does reactivity decrease down Group 7?
1. Halogens react by gaining \( 1 \) electron into their outer shell to form negative halide ions (\( \text{F}^- \), \( \text{Cl}^- \), \( \text{Br}^- \), \( \text{I}^- \)).
2. As you go down the group, atoms have more electron shells, so the atom is larger.
3. The outer shell is further from the nucleus and experiences more shielding.
4. There is a weaker electrostatic attraction between the nucleus and the incoming electron.
5. Therefore, it is harder to attract and gain an electron, making the element less reactive.
Halogen Displacement Reactions
A more reactive halogen will displace (push out) a less reactive halogen from a solution of its salt (halide).
Example 1: Chlorine + Potassium Bromide
Chlorine is more reactive than bromine, so it takes the electron from the bromide ion:
\( \text{Cl}_2\text{ (aq)} + 2\text{KBr}\text{ (aq)} \rightarrow 2\text{KCl}\text{ (aq)} + \text{Br}_2\text{ (aq)} \)
• Observation: The colourless solution turns orange/orange-brown (because free Bromine, \( \text{Br}_2 \), is produced).
Example 2: Chlorine + Potassium Iodide
\( \text{Cl}_2\text{ (aq)} + 2\text{KI}\text{ (aq)} \rightarrow 2\text{KCl}\text{ (aq)} + \text{I}_2\text{ (aq)} \)
• Observation: The colourless solution turns brown / dark brown (because free Iodine, \( \text{I}_2 \), is produced).
Example 3: Iodine + Potassium Bromide
\( \text{I}_2\text{ (aq)} + \text{KBr}\text{ (aq)} \rightarrow \text{No reaction} \)
• Reason: Iodine is less reactive than bromine, so it cannot displace it.
Key Takeaway: Halogens are diatomic non-metals. Their reactivity decreases down the group because larger atoms cannot attract incoming electrons as strongly. More reactive halogens displace less reactive halide ions.
6. Group 0: The Noble Gases
Group 0 includes Helium (\( \text{He} \)), Neon (\( \text{Ne} \)), Argon (\( \text{Ar} \)), Krypton (\( \text{Kr} \)), and Xenon (\( \text{Xe} \)).
Properties and Stability
• They are all colourless, odourless gases.
• They exist as single individual atoms—they are monatomic.
• They are completely unreactive (inert). Why? Because they already have a full outer shell of electrons (Helium has \( 2 \); all others have \( 8 \)). They do not need to lose, gain, or share electrons.
Physical Trends
• As you go down Group 0, the boiling points and densities increase because the atoms get bigger, leading to stronger van der Waals forces between atoms that require more energy to break.
Key Takeaway: Group 0 gases are unreactive monatomic elements because they have fully occupied, stable outer electron shells.
7. The Transition Metals
The Transition Metals sit in the large central block of the Periodic Table (including Iron, \( \text{Fe} \); Copper, \( \text{Cu} \); and Nickel, \( \text{Ni} \)).
Comparison: Transition Metals vs Group 1 Alkali Metals
Compared to the soft, highly reactive Group 1 metals, Transition Metals:
• Have much higher melting and boiling points (except mercury, which is liquid).
• Have much higher densities.
• Are harder and stronger.
• Are much less reactive (e.g., iron reacts very slowly with water/air to form rust, whereas sodium reacts in seconds).
Special Characteristics of Transition Metals
• They form coloured compounds: For example, Copper(II) compounds are typically blue/green, and Iron(III) compounds are reddish-brown.
• They form ions with different charges: For example, iron can form \( \text{Fe}^{2+} \) or \( \text{Fe}^{3+} \) ions.
• They act as excellent catalysts: They speed up industrial reactions without being used up (e.g., Iron is used in the Haber process to make ammonia).
Key Takeaway: Transition metals are strong, dense metals with high melting points. They form coloured compounds, have variable ion charges, and are widely used as catalysts.
Common Pitfalls to Avoid in Exams!
• Mass number vs Atomic number: Remember, Mendeleev ordered by atomic mass; the modern table is ordered by atomic number (proton number).
• Opposite reactivity trends: Group 1 reactivity goes UP as you go down (easier to lose an electron). Group 7 reactivity goes DOWN as you go down (harder to gain an electron). Do not mix them up!
• Bromine vs Bromide: Bromine (\( \text{Br}_2 \)) is the orange halogen molecule; Bromide (\( \text{Br}^- \)) is the colourless ion in a compound. Use the correct name in explanations!