Chemistry (0620) | Group I properties

👋 Welcome to Group I: The Alkali Metals!

Hello future chemist! This chapter is all about one of the most exciting and reactive families on the Periodic Table: Group I, also known as the Alkali Metals.
We will explore their strange physical properties (they are soft enough to cut with a knife!) and why they are so eager to react with everything around them. Understanding Group I helps you master the structure of the Periodic Table itself. Let's dive in!

1. Defining the Alkali Metals

1.1 What are Group I Elements?

Group I is the first vertical column (group) of the Periodic Table (excluding hydrogen, H). The core elements we focus on are:

• Lithium (Li)
• Sodium (Na)
• Potassium (K)

1.2 Electronic Structure: The Reason for Reactivity

All elements in Group I share a fundamental similarity: they each have exactly one electron in their outermost electron shell.

Example:
Lithium (Proton Number 3): 2, 1
Sodium (Proton Number 11): 2, 8, 1
Potassium (Proton Number 19): 2, 8, 8, 1

Chemically, atoms want a full outer shell (like the Noble Gases). It is much easier for Group I elements to lose this single outer electron than to gain seven.

When they lose the electron, they form a stable positive ion (a cation) with a charge of +1 (e.g., \(Na^+\)).

Key Takeaway: Having only 1 outer electron makes Group I metals extremely eager to react, giving that electron away easily.

2. Physical Properties and Trends

Group I metals are very different from the common metals you are used to, like iron or copper. They exhibit clear trends in their physical properties as you move down the group (from Li to Na to K).

2.1 General Physical Characteristics

Group I metals are described as relatively soft metals.

• They are shiny when freshly cut, but quickly tarnish (dull) in air due to reaction with oxygen.
• They can be easily cut with a knife (think of cutting a stick of butter—that's how soft they are!)

2.2 Trends Moving Down the Group (Li → Na → K)

The syllabus requires you to know three specific trends:

Trend 1: Melting Point (Decreasing)

As you move down Group I, the melting points decrease.

• Lithium melts at 181 °C.
• Sodium melts at 98 °C.
• Potassium melts at 63 °C.

Did you know? If you went further down to Caesium (Cs), it would melt in your hand (28 °C)!

Why? As the atoms get larger, the metallic bonds holding the lattice together get weaker because the distance between the positive ions and the 'sea' of delocalised electrons increases. Less energy (heat) is needed to break these weaker bonds.

Trend 2: Density (Increasing)

As you move down Group I, the densities increase.

Important Check: Lithium, Sodium, and Potassium are all less dense than typical metals like iron. In fact, Lithium, Sodium, and Potassium all float on water!

Trend 3: Reactivity (Increasing)

This is the most important chemical trend: reactivity increases dramatically down the group. This is best seen in their reaction with water.

3. Chemical Properties: Increasing Reactivity

Group I metals are stored under oil (like paraffin) in the laboratory to prevent them from reacting instantly with oxygen and water vapour in the air.

3.1 Reaction with Water

All alkali metals react vigorously with cold water to produce a metal hydroxide (an alkali solution) and hydrogen gas.

General Equation:
\(2M(s) + 2H_2O(l) \rightarrow 2MOH(aq) + H_2(g)\)

Let’s compare the reactions for the top three elements:

Lithium (Li):

• Reaction is brisk (fast), but the metal stays solid and moves on the water surface.
• Hydrogen gas is produced, but the reaction is generally not hot enough to ignite the hydrogen.

Sodium (Na):

• Reaction is much more vigorous.
• The metal melts (due to the heat released) and forms a shiny ball, skittering across the surface.
• Hydrogen gas often catches fire, burning with a yellow/orange flame.

Potassium (K):

• Reaction is extremely vigorous and almost instantaneous.
• The heat produced is so great that the hydrogen gas ignites immediately.
• Potassium burns with a characteristic lilac (pink/purple) flame.

3.2 The Trend Explained: Why Reactivity Increases

Don't worry if this seems tricky at first, we can break it down! The reactivity depends entirely on how easily the atom loses its outer electron.

1. More Shells: Moving down the group (Li to K), the atoms have more occupied electron shells (they get physically larger).

2. Distance: The single outer electron is getting further away from the positively charged nucleus. The attraction between the nucleus and this valence electron weakens.

3. Shielding: The inner shells of electrons act like a 'shield' or blanket, reducing the attractive force felt by the outer electron from the nucleus. This effect is called electron shielding.

4. Result: Because the outer electron is held less strongly in Potassium than in Lithium, it is lost more easily.
Therefore, reactivity increases as you go down Group I.

4. Predicting Properties in Group I

A key skill in IGCSE Chemistry is using observed trends (like Li, Na, K) to predict the properties of elements further down the group, such as Rubidium (Rb) or Caesium (Cs).

If you were given information about Li, Na, and K, and asked to predict Rb (which is below K):

1. Reactivity: Rb would be more reactive than K. Its reaction with water would be even more explosive.
2. Softness: Rb would be softer than K.
3. Melting Point: Rb would have a lower melting point than K (which is 63 °C).
4. Density: Rb would have a higher density than K.

Memory Aid: For Group I metals (the ones on the left):
Go Down the group, things get BIGGER, which means they are BETTER at reacting (more reactive)!

Common Mistake to Avoid: Don't confuse the trend in Group I (Reactivity INCREASES down the group) with the trend in Group VII (Halogens), where reactivity DECREASES down the group!