To read the periodic table, start with the element box. The top number is the atomic number, the central letters are the chemical symbol, and the bottom number is the relative atomic mass. The row the element sits in, called its period, tells you how many electron shells it has. The column labelled its group tells you how many valence electrons it has.
The table is provided in every Singapore GCE O-Level Chemistry exam paper. The skill being tested is the ability to read and apply it quickly and accurately.
For students working through the 6092 syllabus, the periodic table appears in questions on bonding, reactivity, ionic formulae, and group properties. Building fluency with it early makes the rest of the course considerably more manageable.
What is the Periodic Table of Elements?
The periodic table of elements is a systematic chart of all 118 known chemical elements, arranged in order of increasing atomic number. Each element occupies a fixed position, and elements with similar properties are placed together.
The word ‘periodic’ refers to a repeating pattern. As the atomic number increases, elements with similar properties appear at regular intervals. This regularity is known as periodicity, and it explains why the table carries its name.
The first widely accepted version was published by Russian chemist Dmitri Mendeleev in 1869. Working with 63 elements known at the time, he arranged them by atomic mass and identified the pattern of recurrence. He left deliberate gaps for elements not yet discovered, and many of those gaps were filled accurately in the decades that followed.
For students, the practical value is direct. The periodic table allows you to identify element properties, predict chemical behaviour, and work out ionic charges and formulae from an element’s position, without memorising data for each element individually.
How to Read an Element Box
Each element on the periodic table has its own cell, called an element box. It contains four key pieces of information.
The atomic number sits at the top of the element box. It is the number of protons in the nucleus of the atom. In a neutral atom, the number of electrons always equals the number of protons, so the atomic number gives you the electron count as well.
No two elements share an atomic number. It is the single value that uniquely identifies an element on the table.
The chemical symbol is the one or two letter abbreviation shown at the centre of the box. Many symbols follow directly from the English name: C for Carbon, O for Oxygen, N for Nitrogen. Others come from Latin names: Na for Sodium (from Natrium), K for Potassium (from Kalium), Fe for Iron (from Ferrum).
The element name appears directly below the chemical symbol. It is the full English name of the element.
The relative atomic mass sits at the bottom of the element box. It represents the average mass of the element’s atoms, weighted across all naturally occurring isotopes. Because most elements exist as a mix of isotopes with slightly different masses, this figure is usually a decimal rather than a whole number.
Worked example: Carbon (C)
The element box for Carbon shows:
- Atomic number: 6 (Carbon has 6 protons and, in a neutral atom, 6 electrons)
- Symbol: C
- Name: Carbon
- Relative atomic mass: 12.01 (a weighted average across Carbon’s naturally occurring isotopes, primarily carbon-12 and carbon-13)
The rows on the periodic table are known as periods, and there are 7 periods in total, each containing a different number of elements.
Period 1 contains just two elements: hydrogen (H) and helium (He). Periods 2 and 3 each contain eight elements. Periods 4 and 5 each contain 18 elements.
Periods 6 and 7 are the longest. Each includes one of the inner transition metal series, displayed separately beneath the main table to keep the layout compact.
Lanthanide and actinide series
The lanthanide series (atomic numbers 57 to 71) and actinide series (atomic numbers 89 to 103) are presented as two rows beneath the main periodic table. Both represent the inner transition metals.
The lanthanide series includes elements such as lanthanum (La), europium (Eu) and ytterbium (Yb), while the actinide series includes elements such as uranium (U), plutonium (Pu) and even einsteinium (Es)!
The key O-Level rule: period number equals the number of electron shells an element has. Sodium (Na) is in Period 3 and has three electron shells. Carbon (C) is in Period 2 and has two.
Moving left to right across a period, the number of protons increases by one at each step. The character of elements also shifts: metals appear on the left and non-metals on the right. The number of valence electrons increases by one with each step across.
The columns of the periodic table are called groups. There are 18 groups, and elements in the same group share similar chemical properties because they have the same number of electrons in their outer shell.
For Groups 1 to 7, the group number tells you directly how many valence electrons an element has. Group 1 elements have one valence electron; Group 7 elements have seven. This relationship between group number and valence electron count is one of the most useful shortcuts available in an O-Level Chemistry exam.
Elements in the same group form ions with the same charge. Sodium (Na) and potassium (K) are both in Group 1, each losing one electron to form a +1 ion: Na⁺ and K⁺.
Group 18 is the exception: noble gases have a completely full outer electron shell, which explains their characteristic low reactivity.
Metals, Non-Metals and Metalloids
Elements in the periodic table fall into three broad categories. Metals occupy the left and centre. Non-metals sit on the right, and a diagonal zigzag line runs between them; the elements along this boundary are called metalloids.
Metalloids, including boron (B), silicon (Si) and antimony (Sb), display properties of both metals and non-metals depending on conditions. The practical exam implication: metals tend to lose electrons to form positive ions, and non-metals tend to gain electrons to form negative ions.
Elements in the same group react in similar ways. Group 1 alkali metals all react with water to produce a metal hydroxide and hydrogen gas. The reaction becomes more energetic moving down the group: lithium reacts steadily, sodium reacts more vigorously, and potassium ignites the hydrogen gas produced.
This is because the electrons from the alkali metal transfer into the surrounding water in an extremely fast, exothermic reaction. Other groups follow the same principle: shared electron configuration produces shared chemical behaviour.
Looking for Patterns in the Periodic Table
The periodic table is organised systematically to enable you to glean information at a glance. Here are some important patterns that can help you refer to the table more quickly and easily.
Metals, metalloids and non-metals
With the exception of hydrogen, elements on the left side of the table are metals and elements on the right are non-metals. Between them runs the diagonal zigzag boundary, and elements along this line are metalloids: boron (B), silicon (Si) and antimony (Sb) are the most commonly encountered at O-Level.
Metals lose electrons to form positive ions. Non-metals gain electrons to form negative ions. Metalloids can behave as either, depending on the reaction.
The six noble gases, helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn), sit in Group 18 at the far right of the table. They are colourless and odourless at room temperature with extremely low chemical reactivity.
Noble gases have a completely full outer electron shell, giving them a stable electronic configuration. With no tendency to gain or lose electrons, they rarely form chemical bonds. Helium fills balloons and airships, valued for its low density and non-flammable properties. Neon glows in advertising signs when electrically excited. Argon fills incandescent light bulbs to prevent the tungsten filament from oxidising.
The five halogens, fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At), are found in Group 17. The name comes from the Greek for ‘salt-producing’; all five react with metals to form a range of salts.
Halogens have seven valence electrons and tend to gain one more to form -1 ions. They exist as diatomic molecules: F₂, Cl₂, Br₂ and I₂. Fluorine is the most reactive halogen; reactivity falls with each step down the group. This trend, and the halogen displacement reactions that demonstrate it, features regularly in O-Level exam questions.
Elements with atomic numbers 1 to 94 occur naturally. Elements 95 to 118 were produced in laboratories, nuclear reactors or nuclear explosions and do not exist in nature. In the context of the 6092 O-Level Chemistry syllabus, only naturally occurring elements are assessed.
Key Groups You Need to Know for O-Level Chemistry
The Singapore O-Level Chemistry syllabus (6092) tests periodic trends and group properties in detail. The four groups below account for the majority of group-related exam questions.
Group 1 contains lithium (Li), sodium (Na) and potassium (K), among others.
Physical properties:
- Soft solids that can be cut with a knife
- Low density (lithium, sodium and potassium all float on water)
- Shiny when freshly cut, but tarnishes quickly on exposure to air
- Low melting and boiling points relative to most metals
- Good conductors of heat and electricity
Chemical properties:
- One valence electron, easily lost to form a +1 ion
- Highly reactive, particularly with water and oxygen
- Stored under oil to prevent reaction with air and moisture
Key reactions:
- With water: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
- With oxygen: metal + oxygen → metal oxide
- With chlorine: metal + chlorine → metal chloride
Trend to know: reactivity increases from lithium to caesium; melting and boiling points decrease.
Watching these reactions in sequence makes the reactivity trend immediately obvious. For students curious about chemistry beyond the classroom, science experiments at home can bring group trends like these to life far more effectively than reading about them.
Group 17 contains fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At).
Physical properties:
- Coloured: fluorine is pale yellow, chlorine is yellow-green, bromine is red-brown, iodine is grey-black
- Poor conductors of electricity
- State at room temperature changes down the group: fluorine and chlorine are gases, bromine is a liquid, iodine is a solid
Chemical properties:
- Seven valence electrons; gain one electron to form a -1 ion
- Exist as diatomic molecules (Cl₂, Br₂, I₂)
- React with metals to form salts
- Reactivity decreases down the group: F > Cl > Br > I
Halogen displacement reactions: a more reactive halogen displaces a less reactive one from its salt solution. Chlorine, for example, displaces bromine from potassium bromide solution, turning the solution orange-brown. This reaction type is directly tested in O-Level Chemistry papers.
Group 18 contains helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn).
Physical properties:
- Colourless gases at room temperature
- Very low melting and boiling points
- Exist as monoatomic molecules (single atoms, not bonded pairs)
Chemical properties:
- Fully filled outer electron shell
- Extremely low reactivity; do not typically form ions or chemical bonds
- Essentially inert for O-Level Chemistry purposes
Practical uses: helium in balloons (low density, non-flammable), neon in advertising signs (glows when electrically excited), argon in light bulbs (prevents filament oxidation).
Transition elements occupy the central block of the periodic table, between Groups 2 and 13. Common O-Level examples include iron (Fe), copper (Cu) and zinc (Zn).
Physical properties:
- High melting and boiling points
- High density
- Hard, malleable and ductile
- Good conductors of heat and electricity
Chemical properties:
- Variable oxidation states: iron can form Fe²⁺ or Fe³⁺ ions depending on conditions
- Form coloured compounds and ions in solution (copper sulphate solution is blue; iron(III) chloride solution is yellow-brown)
- Act as catalysts: iron in the Haber process, manganese(IV) oxide in the decomposition of hydrogen peroxide
How to Use the Periodic Table in Your O-Level Chemistry Exam
The periodic table is provided in every O-Level Chemistry exam paper. No element data needs to be memorised separately. Knowing how to navigate it quickly and accurately is one of the more practical ways to manage exam stress on the day, as the table answers a significant portion of questions if you know where to look.
Five practical ways to use it:
- Find the electron shell count. Period number equals the number of electron shells. An element in Period 3 has three electron shells; an element in Period 4 has four.
- Determine valence electrons. For Groups 1 to 7, the group number equals the number of outer electrons. This tells you directly how an element will behave when bonding or forming ions.
- Predict ion charge. Metals in Groups 1, 2 and 3 lose their valence electrons to form positive ions (+1, +2 and +3 respectively). Non-metals in Groups 5, 6 and 7 gain electrons to form negative ions (-3, -2 and -1, respectively).
- Write ionic formulae. Once you know two ions’ charges, balance them to write the formula. Sodium (Na⁺) and chlorine (Cl⁻) combine 1:1 to give NaCl. Magnesium (Mg²⁺) and chlorine (Cl⁻) combine 1:2 to give MgCl₂.
- Compare reactivity. Group 1 metals become more reactive moving down the group. Group 17 halogens become less reactive. Use an element’s position to rank or compare reactivity without additional data.
O-level chemistry tuition at The Science Academy includes guided practice applying the periodic table across past-paper questions from the 6092 syllabus.
Frequently Asked Questions
How do you read the periodic table?
Start with the element box. The top number is the atomic number (number of protons), the central letters are the chemical symbol, and the bottom number is the relative atomic mass. The row an element sits in, its period, tells you how many electron shells it has. The column, its group, tells you how many valence electrons it has. Together, these components answer most periodic table questions in O-Level Chemistry.
How do you use the periodic table in O-Level Chemistry?
The periodic table is provided in the exam paper. Use it to find atomic numbers, determine electron shell count from the period number, and identify valence electrons from the group number. From there, you can predict ion charge, write ionic formulae, and compare reactivity trends across groups and periods.
What do the rows and columns of the periodic table mean?
Rows are called periods. All elements in the same period have the same number of electron shells. Columns are called groups. Elements in the same group have the same number of valence electrons and share similar chemical properties. The periodic table has seven periods and 18 groups, with 118 elements distributed across them.
What are the most important groups for O-Level Chemistry?
The four groups tested most consistently in the 6092 syllabus are Group 1 (Alkali Metals), Group 17 (Halogens), Group 18 (Noble Gases), and the Transition Elements. Each has distinct physical and chemical properties, and each appears in questions covering reactivity trends, ion charge, group behaviour, and practical applications.
What is the difference between atomic number and atomic mass on the periodic table?
The atomic number is the number of protons in an atom’s nucleus. In a neutral atom, it also equals the number of electrons. The relative atomic mass is the average mass of the element’s atoms across all naturally occurring isotopes. Atomic number is always a whole number; relative atomic mass is usually a decimal.
Master the Periodic Table With The Science Academy
The Science Academy’s O-Level Chemistry tutors have over 20 years of combined teaching experience working with students across Singapore.
At The Science Academy, science tuition goes beyond notes and theory. Students work through past-paper questions from the 6092 syllabus, building the speed and accuracy the exam requires.
To find out about current intake and schedule a trial lesson, get in touch with The Science Academy.
