# Periodic table and element classification

> IB Chemistry HL · IB Chemistry HL
> Source: https://www.owlsprep.com/study/ib-chemistry-hl-u3-periodic-table-and-element-classification/

This subtopic covers the development, structure, and organizing principles of the modern periodic table. You will learn to classify elements by group, period, block, and broad elemental class, building a foundation for all subsequent periodicity topics.

**Prerequisites:** [Atomic structure and electron configuration](https://www.owlsprep.com/study/ib-chemistry-hl-u2-atomic-structure-electron-config/)

## Learning objectives

- Recall the structure and organizing principle of the modern periodic table
- Classify elements by group, period, and s/p/d/f block based on electron configuration
- Distinguish between metals, non-metals, and metalloids
- Name and recall key properties of common element groups

## Development of the Modern Periodic Table

Early attempts to classify elements included Dobereiner's triads and Newlands' Law of Octaves. Dmitri Mendeleev created the first widely accepted periodic table, ordering elements by atomic mass and leaving gaps for undiscovered elements, successfully predicting their properties.

**Modern Periodic Law** — The physical and chemical properties of elements are a periodic function of their atomic number (proton number).

*Example:* All group 1 elements have 1 valence electron and share similar reactivity, regardless of their mass.

**Worked example:** Explain why tellurium (atomic mass 127.6, $Z=52$) is placed before iodine (atomic mass 126.9, $Z=53$) in the modern periodic table.

1. Recall the modern periodic table is ordered by increasing atomic number, not atomic mass.
2. Compare the atomic numbers: tellurium has $Z=52$, iodine has $Z=53$. Since $52 < 53$, tellurium comes first.
3. This order also matches chemical properties: tellurium fits group 16, iodine fits group 17, confirming the arrangement.

> **Exam tip:** Multiple choice questions often trick students into choosing atomic mass as the modern ordering principle — always select atomic number.

## Groups and Periods

The periodic table is arranged into vertical groups and horizontal periods, each with a clear meaning linked to electron structure:

- **Groups (vertical):** Same number of valence electrons, so similar chemical properties. Common named groups: Group 1 = Alkali metals, Group 2 = Alkaline earth metals, Group 17 = Halogens, Group 18 = Noble gases, Groups 3-12 = Transition metals.
- **Periods (horizontal):** Same number of occupied electron shells. Properties change predictably across a period from metallic to non-metallic.

**Worked example:** State the number of occupied electron shells and valence electrons for an element in period 4, group 16.

1. Period number equals the number of occupied electron shells, so period 4 = 4 occupied shells.
2. For main group elements, group number equals the number of valence electrons, so group 16 = 6 valence electrons.
3. Confirm with selenium ($Z=34$): electron configuration $[Ar] 4s^2 3d^{10} 4p^4$, which has 4 shells and 6 valence electrons.

## s, p, d and f Block Classification

Elements are divided into blocks based on which sub-level holds their highest energy valence electron. This classification directly connects electron configuration to the arrangement of the periodic table.

**Block Classification** — Grouping of elements by the sub-orbital containing the highest energy valence electron.

*Example:* The highest energy electron in calcium is in 4s, so calcium is an s-block element.

> **tip**
>
> Remember that $(n-1)d$ is higher energy than $ns$, even though $ns$ fills first. This means d-block classification is based on the d sub-level, not the highest numbered shell.

**Worked example:** Classify each element by block: (a) Chlorine ($Z=17$), (b) Vanadium ($Z=23$), (c) Uranium ($Z=92$).

1. Write the electron configuration for each element, identify the sub-level with the highest energy electron.
2. (a) Chlorine: $1s^2 2s^2 2p^6 3s^2 3p^5$. Highest energy electron is in 3p → p-block.
3. (b) Vanadium: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^3$. Highest energy electron is in 3d → d-block.
4. (c) Uranium: $[Rn] 7s^2 5f^3 6d^1$. Highest energy electron is in 5f → f-block.

## Metals, Non-Metals and Metalloids

Elements are broadly classified into three categories based on physical and chemical properties, which helps predict bonding and reactivity.

| Class | Key Properties | Location |
| --- | --- | --- |
| Metals | Good conductors, malleable, ductile, form cations | Left and lower region of the periodic table |
| Non-metals | Poor conductors, brittle, often gaseous, form anions | Right and upper region of the periodic table |
| Metalloids | Intermediate properties, semiconductors | Along the diagonal between groups 13 and 16 |

**Worked example:** Classify arsenic (As) and state one practical use based on its classification.

1. Locate arsenic on the periodic table: it falls on the diagonal line between metals and non-metals.
2. Elements on this diagonal are classified as metalloids.
3. Arsenic is a semiconductor, so it is used in semiconductor electronic devices.

## Common pitfalls

- **Wrong:** Ordering the modern periodic table by increasing atomic mass.
  - Why it fails: This matches Mendeleev's original table but is incorrect for the modern definition.
  - Correct: Always use atomic number (proton number) as the ordering principle for the modern periodic table.
- **Wrong:** Confusing the meaning of group and period numbers: stating period number equals valence electron count.
  - Why it fails: This mixes up the definition of groups (vertical) and periods (horizontal).
  - Correct: Period number = number of occupied electron shells; main group number = number of valence electrons.
- **Wrong:** Classifying potassium ($[Ar] 4s^1$) as a d-block element because it has a 4th shell.
  - Why it fails: Blocks are classified by the highest energy valence electron, not the highest shell number.
  - Correct: Potassium's highest energy electron is in 4s, so it is an s-block element.
- **Wrong:** Claiming all d-block elements are transition metals per IB definition.
  - Why it fails: IB defines transition metals as elements with a partially filled d sub-level, so scandium and zinc are excluded.
  - Correct: Only d-block elements with incomplete d sub-levels are classified as transition metals.

## Cheatsheet

| Feature | Key Fact |
| --- | --- |
| Ordering principle | Increasing atomic number |
| Groups (vertical) | Same number of valence electrons, similar properties |
| Periods (horizontal) | Same number of occupied electron shells |
| s-block | Groups 1-2, highest energy electron in s sub-level |
| p-block | Groups 13-18, highest energy electron in p sub-level |
| d-block | Groups 3-12, highest energy electron in d sub-level |
| f-block | Lanthanides/actinides, highest energy electron in f sub-level |
| Metalloids | B, Si, Ge, As, Sb, Te, At |
| Transition metal (IB def) | Partially filled d sub-level |

## What's next

This subtopic is the foundation for all periodicity and inorganic chemistry in IB Chemistry HL. The classification system you learned here allows you to predict periodic trends in atomic radius, ionization energy, and electronegativity, which are frequently tested in both Paper 1 and Paper 2. It also underpins the study of transition metal chemistry and bonding, so mastering these basics will make all subsequent topics much easier to understand. Review the key definitions here before moving on to more advanced periodicity concepts.

- [Periodic Trends](https://www.owlsprep.com/study/ib-chemistry-hl-u3-periodic-trends/)
- [d-block elements and transition metals](https://www.owlsprep.com/study/ib-chemistry-hl-u3-d-block-elements-and-transition/)
- [Acid-base classification: Bronsted-Lowry and Lewis](https://www.owlsprep.com/study/ib-chemistry-hl-u3-acid-base-classification-bronsted-lowry/)

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