The Periodic Table
Layout, Periods, Groups, Metals, Nonmetals & Metalloids
History of the Periodic Table
By the mid-1800s, chemists had identified dozens of elements and were searching for patterns in their properties. Several scientists noticed that elements seemed to repeat certain characteristics at regular intervals — a concept called periodicity.
In 1869, Russian chemist Dmitri Mendeleev published the first widely accepted periodic table. He arranged elements by increasing atomic mass and grouped those with similar properties into vertical columns. Crucially, he left gaps for elements he predicted had not yet been discovered — and his predictions proved correct when gallium, scandium, and germanium were found shortly after, matching his descriptions almost exactly.
The modern periodic table, refined in the early 20th century, arranges elements by increasing atomic number (not mass). It contains 118 confirmed elements. The periodic law states: the properties of elements are a periodic function of their atomic numbers.
Layout of the Periodic Table
The periodic table is arranged in rows called periods and columns called groups (or families). Elements in the same group share similar chemical properties because they have the same number of electrons in their outermost shell.
A horizontal row. There are 7 periods. The period number tells you how many electron shells the atom has. Period 1 has 1 shell, Period 3 has 3 shells, etc.
A vertical column. There are 18 groups. Elements in the same group have the same number of valence electrons (outer electrons) and thus similar reactivity and bonding behaviour.
Electrons in the outermost shell. These are the electrons that participate in chemical bonding. The group number (for main-group elements) tells you how many valence electrons the element has.
The repeating pattern of properties across periods. Properties like reactivity, atomic radius, and electronegativity follow predictable trends that repeat with each new period.
Reading an Element Cell
Each element in the periodic table is represented by a cell containing key information. Being able to read this cell fluently is an essential skill.
26 — Atomic number (Z). Iron always has 26 protons. This number defines the element.
Fe — Chemical symbol. From the Latin Ferrum. Symbols are 1–2 letters, capitalised first letter only.
Iron — Full element name.
55.85 — Relative atomic mass (weighted average of all isotopes). Not a whole number because it reflects the natural mixture of isotopes.
Key Groups
Certain groups in the periodic table have special names and distinct properties that you will encounter throughout chemistry. These are the most important to know at this stage.
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Grp 1Alkali MetalsLi, Na, K, Rb, Cs, Fr. Highly reactive metals. 1 valence electron — readily donated to form +1 ions. React vigorously with water to produce hydrogen gas. Reactivity increases down the group. Stored in oil to prevent reaction with air.
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Grp 2Alkaline EarthBe, Mg, Ca, Sr, Ba, Ra. Reactive metals, less so than Group 1. 2 valence electrons — form +2 ions. Calcium is essential for bones and teeth. Magnesium is the central atom in chlorophyll.
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Grp 17HalogensF, Cl, Br, I, At. Highly reactive nonmetals. 7 valence electrons — one short of a full shell, so they readily gain 1 electron to form −1 ions. React with metals to form salts. Reactivity decreases down the group.
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Grp 18Noble GasesHe, Ne, Ar, Kr, Xe, Rn. Essentially inert — full outer shell (8 electrons, except He with 2). Extremely stable, colourless, odourless gases. Used in lighting (neon signs), welding (argon), and balloons (helium).
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Grps 3–12Transition MetalsThe d-block elements (Fe, Cu, Zn, etc.). Good conductors of heat and electricity. Often form coloured compounds. Can have multiple oxidation states (e.g. Fe²⁺ and Fe³⁺). Many are used as catalysts.
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Grp 14Carbon GroupC, Si, Ge, Sn, Pb. Notably diverse — carbon is the basis of all life; silicon is the basis of semiconductors and computing. 4 valence electrons enable rich bonding chemistry.
Element Categories
Beyond groups, elements are classified into broader categories based on their properties. These categories reflect the block structure of the periodic table.
Soft, shiny metals. 1 valence electron. Highly reactive, especially with water. Reactivity increases down the group. Never found free in nature.
2 valence electrons. Harder and less reactive than alkali metals. Form +2 ions. Calcium and magnesium are biologically essential.
D-block elements. High melting points, good conductors, multiple oxidation states, form coloured compounds, many are catalysts.
Properties intermediate between metals and nonmetals. Semiconductors — silicon and germanium power all modern electronics and computing.
Poor conductors. Brittle as solids. Tend to gain electrons (form anions) or share electrons in covalent bonds. Include the building blocks of life.
7 valence electrons. Most reactive nonmetals. Form salts with metals. Exist as diatomic molecules (F₂, Cl₂, etc.). Used in disinfectants, pharmaceuticals, and PVC.
Full outer shells — chemically inert under normal conditions. Monatomic gases. Extremely stable. Their stability is the model for why other elements form bonds.
P-block metals. Softer and have lower melting points than transition metals. Aluminium is the most abundant metal in Earth’s crust.
Metals, Nonmetals & Metalloids
The most fundamental division in the periodic table is between metals, nonmetals, and metalloids. The dividing line (the “staircase”) runs diagonally from boron to astatine — metals to the left, nonmetals to the right, metalloids along the boundary.
- Good conductors of heat and electricity
- Lustrous (shiny) surface
- Malleable — can be hammered into sheets
- Ductile — can be drawn into wire
- Solid at room temperature (except Hg)
- Tend to lose electrons (form cations)
- About 80% of all elements are metals
- E.g. iron, copper, gold, sodium, calcium
- Poor conductors (insulators)
- Dull appearance (no metallic lustre)
- Brittle as solids — shatter when struck
- Not malleable or ductile
- Many are gases at room temperature
- Tend to gain electrons (form anions)
- Include the elements of life: C, H, O, N
- E.g. carbon, oxygen, sulfur, chlorine
- Intermediate electrical conductivity
- Semiconductors — conduct under some conditions
- Some metallic appearance
- Properties vary between metal and nonmetal
- Silicon is the basis of all computer chips
- Germanium used in fibre optics
- Found along the periodic table “staircase”
- E.g. silicon, germanium, arsenic, antimony
Hydrogen sits alone at the top of Group 1, but it is not an alkali metal. It is a nonmetal gas that can either donate or gain an electron depending on conditions. It forms H⁺ (like a metal) or H⁻ (like a nonmetal). Its placement in Group 1 reflects its 1 valence electron, but its properties are unique.
Worked Examples
Practice Questions
Q1. Elements in the same group of the periodic table have similar chemical properties because they:
Q2. Which group contains the noble gases?
Q3. An element is located in Period 3, Group 2. How many electron shells does it have, and how many valence electrons?
Q4. Which of the following is a property of metals?
Q5. Mendeleev’s key contribution to the periodic table was:
Key Takeaways
- Mendeleev (1869) arranged elements by atomic mass and predicted gaps. The modern table uses atomic number.
- Horizontal rows are periods (tell you the number of electron shells). Vertical columns are groups (tell you valence electrons).
- Elements in the same group have the same number of valence electrons and similar chemical behaviour.
- Key named groups: Alkali metals (1), Alkaline earth metals (2), Halogens (17), Noble gases (18).
- Metals (left/centre): conductors, lustrous, malleable, ductile, lose electrons. ~80% of all elements.
- Nonmetals (right): insulators, brittle, dull, gain electrons.
- Metalloids (along the “staircase”): semiconductors with intermediate properties. Silicon powers modern electronics.
- Hydrogen is a unique nonmetal placed in Group 1 due to its 1 valence electron, but it shares few properties with alkali metals.