DoAssignment study guide

C3.3 · Link s-, p-, and d-block properties to electron configurations

Learn to link s-, p-, and d-block properties to electron configurations through clear examples and targeted practice.

Ontario Grade 12 Chemistry

Structure and Properties of Matter

How subshell patterns help explain element properties

Some elements commonly lose electrons to form positive ions. Others may gain or share electrons. These observable patterns are related to how electrons are arranged in atoms. An electron configuration is a notation that shows this arrangement. The periodic table’s s-, p-, and d-blocks are named for the subshell that receives the last electron in an element’s ground-state configuration. In this lesson, you will use that connection to explain broad patterns in element properties. Block membership is useful, but it does not determine every property by itself.

What you will learn

Identify the s-, p-, and d-blocks using electron configurations.
Explain how block membership is linked to the subshell receiving the last electron.
Connect outer-electron patterns with broad chemical properties of elements.
Distinguish useful patterns from claims that would apply to every element.

1. Review: electrons, energy levels, and subshells

An atom has a small, positively charged nucleus and negatively charged electrons around it. The atomic number tells you the number of protons in the nucleus. A neutral atom has the same number of electrons as protons, so its electron count equals its atomic number.

Electrons occupy energy levels. Each energy level contains subshells, which are labelled s, p, or d in the configurations used here. For example,

3p^4

means that four electrons occupy the p subshell in the third energy level. The raised number is called a superscript; it tells you the number of electrons in that subshell.

For many main-group elements, the outer electrons are especially useful for explaining chemical patterns. These are called valence electrons. For d-block elements, electrons in the outer s subshell and a nearby d subshell can both be involved in chemical behaviour. This is why group patterns are not as simple for the d-block as they are for Groups 1 and 2.

electrons in a neutral atom = atomic number

A neutral atom has an electron count equal to its atomic number.
The superscript in an electron configuration gives the number of electrons in a subshell.
Valence electrons are outer electrons that help explain many chemical patterns.

2. The s- and p-blocks

The periodic table is divided into blocks based on the type of subshell that receives the last electron in the ground-state configuration. Ground state means the atom’s usual, lowest-energy arrangement. The s-block is on the left side of the table. Its configurations end in an s subshell. The p-block is on the right, excluding helium, and its configurations end in a p subshell.

Helium is a useful exception to remember. It is placed in Group 18 because its properties fit that group, but its configuration is

1s^2

. It is therefore an s-block element by configuration. It has two outer electrons. The other Group 18 elements have eight outer electrons.

The s-block includes the Group 1 alkali metals and Group 2 alkaline earth metals, as well as hydrogen and helium. Group 1 metals have one outer s electron; Group 2 metals have two. These metals commonly lose their outer electrons to form positive ions. Losing one electron gives a

1+

charge; losing two gives a

2+

charge. The tendency to lose these electrons helps explain the usual ion charges for these groups.

The p-block contains metals, metalloids, and non-metals. A metalloid has some properties of metals and some properties of non-metals. Because the p-block includes several kinds of elements, it does not have one uniform set of properties. Its elements have different numbers of outer s and p electrons. Depending on the element, atoms may commonly lose electrons, gain electrons, or share electrons when forming compounds.

Elements in the same main-group column often have similar chemical behaviour because they have the same number of outer electrons. Their electrons are in different energy levels as you move down a group. Helium is the exception to the usual outer-electron count in Group 18: it has two, while the other Group 18 elements have eight.

\text{s-block: } ns^{1\text{–}2};\quad \text{p-block: } ns^2np^{1\text{–}6}

The s- and p-block labels identify the subshell receiving the last electron.
Group 1 and Group 2 metals commonly lose one and two outer electrons, respectively.
The p-block includes metals, metalloids, and non-metals.
Helium is in Group 18 but has two outer electrons and an s-block configuration.

3. The d-block and its properties

The d-block lies between the s- and p-blocks. Its elements are metals. Their configurations involve electrons entering a d subshell in the energy level just inside the outermost level. A common simplified pattern is a partly filled d subshell together with electrons in the outer s subshell. Some atoms do not follow the simplest pattern exactly, so use a reliable periodic table or reference when you need an exact configuration.

Many d-block elements form ions with different positive charges. Their outer s electrons and nearby d electrons can both be involved when ions form. As a result, different numbers of electrons may be removed in different compounds. For example, iron can form both

\mathrm{Fe^{2+}}

and

\mathrm{Fe^{3+}}

. This illustrates a common d-block pattern, not a rule that every d-block element forms the same ions.

Many d-block elements also form coloured compounds. The presence of d electrons is linked to this pattern, but not every d-block ion or compound is coloured. The specific substance matters. The important link for this lesson is that d-electron arrangements help explain patterns that distinguish d-block metals from many s-block metals.

Properties across the d-block often change more gradually than properties between some neighbouring main-group families. Still, d-block elements are not all chemically identical. Use block membership to identify a broad electron-configuration pattern, then use the element and its compounds to make more specific claims.

\text{d-block: } (n-1)d^{1\text{–}10}ns^{0\text{–}2}

The d-block consists of metals whose configurations involve a d subshell.
Many d-block elements form ions with more than one positive charge.
Coloured compounds are common for some d-block elements, but colour is not universal.

4. A careful method for connecting configuration and properties

Start by locating the subshell at the end of the electron configuration. This identifies the block. Next, find the outer-electron pattern and, where useful, the element’s group. Then make a claim that matches the pattern. For example, a Group 1 s-block metal commonly loses one outer electron.

Use words such as commonly and often when describing patterns. These words show that a trend is not a guarantee for every atom or compound. Do not infer all properties from block membership alone. The p-block has a wide range of element types, and d-block elements do not all form the same ions or coloured compounds.

In the example, the atom is neutral, so its configuration contains the same number of electrons as its atomic number. Reading the configuration identifies both the element and a typical ion-forming pattern. The equation shows the electron loss and conserves the atom and total charge.

Use the final subshell to identify the block.
Use the group and outer-electron pattern to make a more specific prediction.
Describe broad trends without treating them as universal rules.

Worked example

From configuration to block and common ion

A neutral atom has the electron configuration

1s^2 2s^2 2p^6 3s^1

. Identify its block and group. Predict the charge of its common ion and explain your reasoning.

Identify the block
The last occupied subshell shown is the 3s subshell. The configuration therefore places the element in the s-block.
$3s^1$
Identify the group pattern
There is one electron in the outermost energy level. This is the outer-electron pattern of a Group 1 element.
$3s^1$
Predict the common ion
The configuration contains 11 electrons, so the neutral atom has atomic number 11. It is sodium. A Group 1 metal commonly loses its one outer electron. The resulting ion has one more proton than electrons, so its charge is $1+$ . The equation represents this electron loss.
$\mathrm{Na} \rightarrow \mathrm{Na^+} + e^-$

Answer: The atom is sodium, an s-block element in Group 1. Its common ion is

\mathrm{Na^+}

because it commonly loses its one outer electron.

Check: The equation has one sodium atom on each side. Its total charge is zero on each side: the product charges sum to

+1-1=0

Common mistakes and how to avoid them

Treating block membership as a complete prediction of an element’s properties.

Correction: A block identifies a subshell pattern. Use group and outer-electron information for more specific chemical patterns, and remember that individual elements and compounds can differ.

Assuming that every p-block element is a non-metal.

Correction: The p-block includes metals, metalloids, and non-metals.

Assuming that every d-block element forms the same ions or coloured compounds.

Correction: Many d-block elements form ions with different positive charges, and many form coloured compounds. These are common patterns, not universal rules.

Calling helium a p-block element because it appears on the right side of the periodic table.

Correction: Block identity follows the configuration. Helium ends in

1s^2

, so it is an s-block element. It is in Group 18 and has two outer electrons.

Lesson summary

The s-, p-, and d-block labels are linked to the subshell receiving the last electron in an electron configuration.
Outer-electron patterns help explain common chemical behaviour, including typical ion formation in Groups 1 and 2.
Helium is in Group 18 but has two outer electrons; the other Group 18 elements have eight.
The p-block includes metals, metalloids, and non-metals. The d-block consists of metals with configurations involving d electrons.
Use block and group patterns as explanations for trends, not as claims that every element behaves identically.

Check your understanding

Question 1

An element’s configuration ends in

4p^3

. Which block does it belong to?

s-block
p-block
d-block
The block cannot be identified from the ending subshell

Show answer and explanation

p-block

The p-block is identified by a configuration ending in a p subshell.

Question 2

Why does a Group 2 s-block metal commonly form a

2+

ion?

It commonly loses its two outer electrons.
It commonly gains two protons.
It commonly gains two outer electrons.
It loses two neutrons.

Show answer and explanation

It commonly loses its two outer electrons.

Group 2 atoms have two outer electrons and commonly lose them. Losing two negatively charged electrons leaves an ion with a

2+

charge.

Question 3

Which statement about the d-block is most accurate?

Every d-block ion has a $2+$ charge.
Every d-block compound is coloured.
Many d-block elements can form ions with different positive charges.
The d-block contains only non-metals.

Show answer and explanation

Many d-block elements can form ions with different positive charges.

Many d-block elements form ions with different positive charges. The other choices wrongly turn common patterns into universal rules or misidentify the elements.

Key terms

Electron configuration: A notation showing how an atom’s electrons are arranged among energy levels and subshells.
Subshell: A subdivision of an energy level, labelled s, p, or d in the configurations used here.
Valence electrons: Outer electrons that are especially useful for explaining many patterns in chemical behaviour.
Block: A region of the periodic table identified by the subshell receiving the last electron in an element’s ground-state configuration.
Ground state: An atom’s usual, lowest-energy electron arrangement.

Continue through SCH4U

View the complete SCH4U Ontario Grade 12 Chemistry curriculum and lessons

About this lesson and its review

Published by DoAssignment. This reviewed lesson follows Ontario Grade 12 Chemistry (SCH4U), expectation C3.3. It is a study resource, not an official curriculum publication.

Before publication, content is checked for structure, mathematical or chemical notation, calculations, course boundaries, and readability. Errors can still occur, so corrections are welcomed.

Official curriculum reference

Report a correction or ask a question