Group I PropertiesCambridge IGCSE Chemistry: Revision notes
Section 1
What are the Group I alkali metals?
Group I elements are the alkali metals found in the first column of the periodic table. The three most commonly studied are lithium (Li), sodium (Na), and potassium (K). These metals are highly reactive and naturally occur as compounds rather than pure elements in nature. They are called 'alkali metals' because they form strongly alkaline hydroxides when they react with water.
Examiners expect you to recall that Group I elements are stored under oil in the laboratory—mention this if asked why, as it prevents reaction with air and moisture.
Section 2
Why are Group I metals soft and what does this tell us?
Group I metals are relatively soft—they can be cut easily with a knife, unlike harder metals such as iron or copper. This softness reflects their metallic structure and the strength of metallic bonding within them. The ability to cut them easily is a key characteristic that allows chemists to quickly identify Group I metals experimentally. Softness increases as you move down the group: lithium is harder than sodium, which is harder than potassium.
Think of Group I metals as being like butter—the further down the group you go, the softer and more spreadable they become.
Students often confuse 'soft' with 'weak'—soft refers to physical deformability, not structural strength. These metals are still metals and conduct electricity effectively.
Section 3
What are the key trends in Group I as you move down the group?
As you move down Group I from lithium to potassium, three important trends occur:
| Property | Trend Down Group I | Reason |
|---|---|---|
| Melting point | Decreases | Weaker metallic bonding due to increased atomic size and more distant electrons |
| Density | Increases | Atomic mass increases faster than atomic radius |
| Reactivity | Increases | Outermost electron is further from nucleus, easier to lose |
These trends are predictable and follow the periodic patterns caused by atomic structure changes down a group.
When answering trend questions, always explain why the trend occurs by referring to atomic structure and electron configuration—examiners mark explanation marks separately.
Sodium has a higher melting point than potassium because potassium's valence electron is in the fourth shell, further from the nucleus, so it experiences weaker attraction and the metallic bonding is weaker overall.
Section 4
How does reactivity increase down Group I?
Reactivity with water is the most obvious way to observe trends in Group I. The order is:
Lithium < Sodium < Potassium (increasing reactivity)
This occurs because:
- The outermost electron (valence electron) becomes further from the nucleus as you go down the group
- The increased atomic radius means the valence electron experiences greater shielding from inner electrons
- This makes the valence electron easier to lose, so the atom reacts more readily to form the alkali metal ion (M⁺)
- Elements that lose electrons more easily are more reactive
When Group I metals react with water, they produce a hydroxide and hydrogen gas: for example, 2Na + 2H₂O → 2NaOH + H₂↑
Potassium is more reactive than sodium with water because potassium's valence electron is in the 4th shell compared to sodium's 3rd shell, making it further away and more easily removed.
Do not say Group I metals are reactive 'because they want to gain electrons'—they are metals and they lose electrons, not gain them, to form positive ions.
Section 5
How can you predict properties of other Group I elements?
Using the trends established by lithium, sodium, and potassium, you can predict the properties of other Group I elements such as rubidium (Rb) and caesium (Cs):
For any unknown Group I element:
- Identify its position in the group (how many elements are below or above it)
- Apply the known trends: melting point continues to decrease, density continues to increase, reactivity continues to increase
- Estimate numerical values by continuing the pattern shown by Li, Na, and K
Example: Rubidium is below potassium, so you would predict that rubidium:
- Has a lower melting point than potassium
- Has a higher density than potassium
- Is more reactive with water than potassium
- Has a softer appearance than potassium
This predictive skill is essential in chemistry and demonstrates understanding of periodic patterns.
Examiners often test prediction by asking 'predict the melting point of rubidium' or 'suggest why caesium would be even more reactive'—always reference the trend and explain your reasoning.
If asked to predict the reactivity of rubidium: 'Rubidium is below potassium in Group I, so its valence electron is in a higher shell and experiences greater shielding. This makes it easier to lose, so rubidium is more reactive than potassium.'
Must Know
- Group I elements (lithium, sodium, potassium) are relatively soft metals that can be cut with a knife; they become softer down the group
- Melting point decreases down Group I because metallic bonding weakens as atoms get larger and electrons become more distant from the nucleus
- Density increases down Group I because atomic mass increases faster than atomic radius increases
- Reactivity increases down Group I because the valence electron is further from the nucleus and more shielded, making it easier to lose
- Group I metals react with water to produce metal hydroxides and hydrogen gas; reactivity with water demonstrates the group trend (K > Na > Li)
- Periodic trends can be used to predict properties of any Group I element—extend the pattern shown by known elements to estimate properties of unknown ones
That's the notes covered.
Carry on to the next subtopic.