Hello! Welcome to your first lesson on the Group 1 elements. Given your exam is just around the corner, we'll work through the material in a focused and efficient way. This session will establish the fundamental concepts that everything else in this topic builds upon.
Our goal for this lesson is to identify the Group 1 elements, learn how to write their electron configurations, and understand why they consistently form ions with a +1 charge. Mastering these three points will provide a solid foundation for understanding their chemical properties and reactions, which we'll cover later.
Let's begin by identifying exactly which elements we're talking about.
Identifying the Group 1 Elements
The Group 1 elements, also known as the alkali metals, occupy the first column on the left-hand side of the periodic table.
As you can see, this group includes:
- Lithium (Li)
- Sodium (Na)
- Potassium (K)
- Rubidium (Rb)
- Caesium (Cs)
- Francium (Fr)
You might notice that Hydrogen (H) is also positioned at the top of Group 1. While it shares a key electronic feature with the alkali metals (which we'll discuss next), it is a non-metal gas at room temperature and is generally not considered an alkali metal due to its vastly different properties.
For a quick clarification on this point, please watch this short segment from the video "Alkali Metals".
This video explains which elements are considered alkali metals and clarifies the special case of hydrogen.
Watch from the beginning until the end of the explanation about hydrogen. Pay attention to why hydrogen, despite its position, is not classified as an alkali metal.
Electron Configuration: The Key to Group 1 Behaviour
The single most important fact about Group 1 elements is their electron configuration. This arrangement of electrons dictates all of their chemical properties, including their reactivity and the type of ions they form.
All elements in Group 1 have one electron in their outermost shell. This is why they are in "Group 1". This single outer electron is often called a valence electron.
To see this pattern, let's examine the electron configurations in more detail. A comprehensive table in the "Doc Brown" resource lays this out perfectly.
This table provides a concise summary of the key data for the alkali metals, including their electron configurations in two formats.
In the table under "7.2. Information and Data Table GROUP 1 ALKALI METALS", focus on the rows for electron configuration. First, look at the "simple electron configuration" row. For Sodium (Na), it's 2,8,1. This means it has 2 electrons in the first shell, 8 in the second, and 1 in the outer shell. Notice that Lithium (2,1) and Potassium (2,8,8,1) also end with a 1. Next, examine the "electron configuration" row. This is the standard A-Level notation. For example, Sodium (Na) is [Ne]3s¹. [Ne] represents the stable electron core of the preceding noble gas, Neon. 3s¹ indicates the single valence electron is in the 's' orbital of the 3rd shell. Notice how all the elements have a configuration ending in ns¹, where n is the period number (the row in the periodic table). This ns¹ configuration is the defining electronic feature of Group 1.
Forming +1 Ions
Atoms are most stable when they have a full outer shell of electrons, like the noble gases (e.g., Neon, Argon). For a Group 1 atom with its single, lonely valence electron, the most straightforward way to achieve this stability is to lose that one outer electron.
When a neutral atom loses a negatively charged electron (e⁻), it is left with an excess positive charge. Since Group 1 elements lose one electron, they form a positive ion, or cation, with a charge of +1.
- Sodium atom (Na):
[Ne] 3s¹(neutral, 11 protons, 11 electrons) - Sodium ion (Na⁺):
[Ne](positive, 11 protons, 10 electrons)
The process can be written as: , where M represents any alkali metal.
Why don't they form +2 ions? Once the first electron is lost, the resulting ion (M⁺) has a stable, full outer shell. Removing a second electron would mean breaking into this stable arrangement, which requires a very large amount of energy. It is so energetically unfavorable that it effectively never happens in chemical reactions. Therefore, the only stable ion for Group 1 elements is the +1 ion, and their oxidation state in compounds is always +1.
The following resources reinforce this fundamental concept.
This text explicitly explains why Group 1 elements form +1 ions and not +2 ions.
Read the short paragraph that begins with this sentence. The key takeaway is the prohibitive energy cost of removing a second electron.
For a final, quick summary, this video segment ties together the single outer electron and the formation of +1 ions very clearly.
GCSE Chemistry - Group 1 Elements - Alkali Metals | Properties | Reactivity (2026/27 exams)
This clip from a GCSE-level video provides a simple, direct explanation of why alkali metals form +1 ions.
Watch the segment from "full outermost shell". This is a great recap of the core idea.
Conclusion
Let's summarize the essential foundations we've covered in this lesson.
- Identification: The Group 1 alkali metals are Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Caesium (Cs), and Francium (Fr). Hydrogen is in the same column but is not an alkali metal.
- Electron Configuration: The defining characteristic of Group 1 elements is their electron configuration. They all have one electron in their outermost shell, which can be written generally as
ns¹. - Ion Formation: To achieve a stable, full outer shell, alkali metals lose their single valence electron. This results in the formation of a positive ion (cation) with a charge of +1 (e.g., Li⁺, Na⁺, K⁺). This is their only stable ion.
These three facts are the bedrock of Group 1 chemistry. In our next lesson, we will build directly on this by exploring how the ease of losing this outer electron changes as you move down the group. This will allow us to explain the trend in ionisation energy and, ultimately, the trend in their reactivity.