# S- and p-block element chemistry

> IB Chemistry SL · Reactivity 1: Stoichiometry and Periodicity
> Source: https://www.owlsprep.com/study/ib-chemistry-sl-u4-s-and-p-block-element/

This subtopic covers the classification of elements into s- and p-blocks, key periodic and group trends in physical and chemical properties, and characteristic reactions of common main group elements for IB SL Chemistry.

**Prerequisites:** [Electron configuration](https://www.owlsprep.com/study/ib-chemistry-sl-electron-configuration/); [Periodic table structure](https://www.owlsprep.com/study/ib-chemistry-sl-periodic-table-introduction/)

## Learning objectives

- Classify elements as s-block or p-block from their electron configuration
- Explain trends in physical and chemical properties down groups and across periods
- Predict reactions of common s- and p-block elements
- Describe the acid-base trend in period 3 oxides

## Classification of s- and p-block Elements

**s-block elements** — Elements with their outermost (valence) electron occupying an s-orbital. All s-block elements are highly reactive metals that readily lose valence electrons to form cations.

*Notation:* Groups 1 and 2

*Example:* Sodium (Na) has electron configuration $[Ne]3s^1$, a reactive group 1 s-block metal.

**p-block elements** — Elements with their outermost valence electrons occupying p-orbitals, after the valence s-orbital is filled. This block contains metals, non-metals and metalloids with diverse properties.

*Notation:* Groups 13 to 18

*Example:* Chlorine (Cl) has electron configuration $[Ne]3s^2 3p^5$, a reactive non-metal p-block element.

An element's block is determined exclusively by the subshell of its highest energy valence electron, not its atomic number or position. This classification allows us to predict properties based on periodic position.

**Worked example:** Classify magnesium (Z=12) and arsenic (Z=33) as s-block or p-block.

1. Step 1: Write the electron configuration for each element
2. $$Mg\ (Z=12): 1s^2 2s^2 2p^6 3s^2$$
3. $$As\ (Z=33): 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^3$$
4. Step 2: Identify the highest energy valence subshell
5. For Mg: valence electrons are in 3s → s-block. For As: highest energy valence electrons are in 4p → p-block.

## Group Trends in s-block Elements

Group 1 (alkali metals) and group 2 (alkaline earth metals) show consistent trends moving down the group. Atomic radius increases due to additional electron shells, ionization energy decreases due to increased shielding and larger distance from the nucleus, and reactivity increases as a result.

> **tip**
>
> Reactivity of s-block metals depends on ability to lose valence electrons, so lower ionization energy = higher reactivity down the group.

**Worked example:** Explain why potassium is more reactive than lithium in water.

1. Step 1: Compare atomic structure: K has 4 electron shells, Li has 2. K has larger atomic radius and more inner electron shielding.
2. Step 2: Nuclear attraction for the valence electron is weaker in K than Li, so K has a lower first ionization energy.
3. Step 3: Since s-block reactivity relies on losing the valence electron, K loses its electron more easily, so it is more reactive than Li.

**Check your understanding**

Test your understanding of group 1 trends

1. Which property decreases down group 1?

   - Atomic radius
   - First ionization energy
   - Reactivity
   - Relative atomic mass

   *Why:* Correct! Ionization energy decreases down the group. All other options listed increase moving down group 1.

## Group Trends in Selected p-block Elements

For IB SL, the most commonly tested p-block groups are 17 (halogens) and 18 (noble gases). Group 17 are reactive non-metals that gain one electron to form -1 anions. Unlike s-block metals, reactivity decreases down group 17.

**Halogen displacement reaction** — A reaction where a more reactive halogen (higher up group 17) displaces the halide ion of a less reactive halogen from aqueous solution.

*Example:* Chlorine disolves bromide ions from solution to form bromine.

**Worked example:** Predict if a reaction occurs when chlorine water is added to aqueous potassium iodide, write the balanced equation.

1. Step 1: Order halogen reactivity: reactivity decreases down group 17, so Cl > I, chlorine is more reactive than iodine.
2. Step 2: A more reactive halogen displaces a less reactive halide ion, so reaction occurs.
3. $$Cl_2(aq) + 2KI(aq) \rightarrow 2KCl(aq) + I_2(aq)$$
4. The solution turns brown due to formation of elemental iodine.

## Acid-Base Trends Across Period 3 Oxides

Across period 3, elements change from metallic to non-metallic, so their oxides change from ionic (basic) to covalent (acidic), with one amphoteric oxide in the middle. This is a very common exam question.

**Worked example:** Classify $Na_2O$, $Al_2O_3$ and $SO_2$ as basic, amphoteric or acidic.

1. Step 1: Recall the trend: left (metals) = basic, middle = amphoteric, right (non-metals) = acidic.
2. (a) Sodium is a group 1 metal on the far left of period 3: $Na_2O$ is basic. It reacts with water to form sodium hydroxide: $Na_2O + H_2O \rightarrow 2NaOH$
3. (b) Aluminium is in the middle of period 3: $Al_2O_3$ is amphoteric, it reacts with both acids and strong bases.
4. (c) Sulfur is a non-metal on the right of period 3: $SO_2$ is acidic. It reacts with water to form sulfurous acid: $SO_2 + H_2O \rightarrow H_2SO_3$

## Common pitfalls

- **Wrong:** Assuming all reactivity trends down groups are the same
  - Why it fails: Reactivity increases down s-block metal groups but decreases down p-block non-metal groups like halogens
  - Correct: Remember: metals lose electrons to react (higher reactivity down), non-metals gain electrons (lower reactivity down for halogens)
- **Wrong:** Confusing amphoteric and amphiprotic
  - Why it fails: Amphoteric means reacts with both acids and bases, while amphiprotic means can donate and accept protons. These terms are not interchangeable
  - Correct: For period 3 oxides, use amphoteric (not amphiprotic) to describe aluminium oxide
- **Wrong:** Classifying d-block elements as p-block
  - Why it fails: Block is determined by the highest energy valence subshell, not the highest numbered filled subshell
  - Correct: For d-block elements, the highest energy valence subshell is d, so they are not classified as p-block
- **Wrong:** Claiming all noble gases are completely unreactive
  - Why it fails: Heavier noble gases like xenon form stable covalent compounds, though IB SL focuses on their general inertness
  - Correct: For exams, state noble gases are generally unreactive due to their full valence shell, which meets marking criteria

## Cheatsheet

| Property | s-block (Groups 1-2) | p-block (Groups 13-18) |
| --- | --- | --- |
| Location | Left of periodic table | Right of periodic table |
| Valence electrons | Occupy s-orbitals | Occupy p-orbitals |
| General classification | Reactive metals | Metals, metalloids, non-metals |
| Group 1 reactivity | Increases down group | - |
| Group 17 reactivity | - | Decreases down group |
| Period 3 oxide property | Basic (left) | Amphoteric/acidic (right) |

## What's next

Understanding s- and p-block element chemistry is the foundation for all inorganic chemistry topics in IB SL. This subtopic connects fundamental periodic trends to real observed chemical behavior, and the patterns you learn here help you predict reaction outcomes and properties of unfamiliar elements in exam questions. Most exam questions on periodicity require you to explain trends, so mastering the reasoning behind each trend (rather than just memorizing the direction of change) is critical to earn full marks. This topic also underpins all subsequent work on bonding, acid-base chemistry and inorganic reactions in the IB syllabus.

- [Reactivity 2: Energetics and Kinetics](https://www.owlsprep.com/study/ib-chemistry-sl-u5-overview/)
- [Temperature change and heat capacity](https://www.owlsprep.com/study/ib-chemistry-sl-u5-temperature-change-and-heat-capacity/)
- [Enthalpy of reaction](https://www.owlsprep.com/study/ib-chemistry-sl-u5-enthalpy-of-reaction/)

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