Important Groups
1. Group 1: Alkali Metals
Learning outcomes
- I can identify the alkali metals on the periodic table.
- I can describe the physical and chemical properties of alkali metals.
- I can explain why alkali metals form ions.
- I can describe how the reactivity of alkali metals changes down the group.
- I can explain observations from reactions involving alkali metals.
What Are the Alkali Metals?
The alkali metals are the metallic elements found in Group 1 of the periodic table.
They include:
- lithium (Li)
- sodium (Na)
- potassium (K)
- rubidium (Rb)
- caesium (Cs)
- francium (Fr)
Hydrogen is also located in Group 1 because it has one valence electron, but it is not classified as an alkali metal.
The alkali metals share many similar chemical properties because they all have one electron in their outer shell.
Finding Group 1 on the Periodic Table
Group 1 is the first vertical column on the left side of the periodic table.
Moving down the group:
Li → Na → K → Rb → Cs → Fr
These elements show clear patterns in their physical and chemical properties.
This is an example of periodicity: repeating patterns of properties across the periodic table.
Electron Arrangements of Alkali Metals
All alkali metals have one valence electron.
For example:
Lithium:
2,1
Sodium:
2,8,1
Potassium:
2,8,8,1
Although these atoms have different numbers of occupied electron shells, they all have the same basic outer-shell pattern:
one outer electron
This explains many of their similar chemical properties.
Why Elements in the Same Group Behave Similarly
Chemical reactions usually involve an atom's outer electrons.
Because every alkali metal has one valence electron, they tend to undergo similar reactions.
In particular, Group 1 metals tend to:
lose one electron
and form:
+1 ions
For example:
Li → Li⁺ + e⁻
Na → Na⁺ + e⁻
K → K⁺ + e⁻
This is one of the defining chemical behaviours of Group 1.
Why Alkali Metals Form +1 Ions
Consider sodium.
Its electron arrangement is:
2,8,1
Sodium can lose its single outer electron.
After losing that electron, its arrangement becomes:
2,8
The sodium atom originally has equal numbers of protons and electrons, so it is neutral.
When it loses one negatively charged electron, it has one more proton than electrons.
Therefore, it has a charge of:
+1
and becomes:
Na⁺
Group 1 Ion Formation
The general process can be written:
M → M⁺ + e⁻
where M represents a Group 1 metal.
For example:
Li → Li⁺ + e⁻
Na → Na⁺ + e⁻
K → K⁺ + e⁻
All alkali metals therefore typically form ions with a:
+1 charge
These positive ions are called cations.
Physical Properties of Alkali Metals
Alkali metals share several physical properties.
They are generally:
- soft
- relatively low in density compared with many other metals
- shiny when freshly cut
- good conductors of electricity
- good conductors of heat
- solids at room temperature
- relatively low-melting compared with many other metals
Although they are metals, they are very different from strong structural metals such as iron.
Alkali Metals Are Soft
Many metals are hard and strong.
Alkali metals are unusual because they are very soft.
Lithium, sodium, and potassium can be cut relatively easily.
A freshly cut surface appears shiny because the newly exposed metal reflects light.
However, this shiny surface does not usually remain visible for long.
Why the Surface Becomes Dull
Alkali metals react with substances in the air.
For example, they can react with oxygen and moisture.
As reaction products form on the surface, the metal loses its shiny appearance and becomes dull or tarnished.
This means an observation such as:
"The freshly cut surface quickly became dull."
provides evidence that the metal is reacting with substances in the air.
Storage of Alkali Metals
Because alkali metals react readily with oxygen and water, laboratory samples are often stored under a suitable hydrocarbon oil.
The oil helps prevent contact with:
- oxygen
- water vapour
- moisture
This slows unwanted reactions during storage.
Density of Alkali Metals
Alkali metals have relatively low densities compared with many familiar metals.
Lithium, sodium, and potassium are particularly low-density metals.
This affects observations during reactions with water.
For example, sodium and potassium can float on water.
Floating is an observation related to density rather than evidence by itself of chemical reactivity.
Melting Points
Alkali metals have relatively low melting points compared with many other metals.
The melting points generally decrease down Group 1.
This can affect what we observe during reactions.
For example, sodium may melt into a small ball during its reaction with water because the reaction releases heat and sodium has a relatively low melting point.
Chemical Properties of Alkali Metals
Alkali metals are reactive metals.
They readily participate in reactions in which they lose their outer electron.
Important reactions include reactions with:
- water
- oxygen
- halogens
These reactions usually produce compounds containing M⁺ ions.
Reaction with Water
One of the most characteristic reactions of Group 1 metals is their reaction with water.
The general word equation is:
alkali metal + water → metal hydroxide + hydrogen
For lithium:
lithium + water → lithium hydroxide + hydrogen
For sodium:
sodium + water → sodium hydroxide + hydrogen
For potassium:
potassium + water → potassium hydroxide + hydrogen
General Equation for the Reaction with Water
Using M to represent a Group 1 metal:
2M + 2H₂O → 2MOH + H₂
For sodium:
2Na + 2H₂O → 2NaOH + H₂
The products are:
- sodium hydroxide
- hydrogen gas
Why the Solution Becomes Alkaline
The metal hydroxide produced during the reaction dissolves in the water.
For example, sodium produces:
sodium hydroxide, NaOH
Sodium hydroxide forms an alkaline solution.
This is where the name:
alkali metals
comes from.
They produce alkaline hydroxides when they react with water.
Testing the Solution
An indicator can be used to show that the solution produced is alkaline.
For example, universal indicator would show a colour associated with high pH.
The observation provides evidence that an alkaline metal hydroxide has formed.
Lithium and Water
Lithium reacts with water.
Typical observations include:
- lithium floats
- it moves slowly across the surface
- bubbles of gas are produced
- the metal gradually becomes smaller
- it eventually disappears
The bubbles are caused by the production of:
hydrogen gas
The lithium disappears because lithium atoms are being converted into lithium ions in solution as lithium hydroxide forms.
Sodium and Water
Sodium reacts more vigorously with water than lithium.
Typical observations include:
- sodium floats
- rapid fizzing occurs
- sodium moves across the surface
- the metal often forms a small ball
- the metal eventually disappears
Why does sodium form a ball?
The reaction releases heat.
Sodium has a relatively low melting point.
The heat can therefore melt the sodium.
Surface tension then causes the liquid metal to form a roughly spherical shape.
Potassium and Water
Potassium reacts even more vigorously with water.
Typical observations can include:
- very rapid movement
- vigorous fizzing
- enough heat for the hydrogen or potassium to ignite
- a characteristic lilac flame may be observed
These observations show that potassium reacts more vigorously than sodium or lithium.
Explaining the Bubbles
During the reaction:
metal + water → metal hydroxide + hydrogen
Therefore, the bubbles observed are:
hydrogen gas
For sodium:
2Na + 2H₂O → 2NaOH + H₂
Gas production causes the visible fizzing.
More vigorous gas production generally indicates a faster reaction.
Testing for Hydrogen
Hydrogen gas can be identified using a burning splint.
Hydrogen burns rapidly in oxygen and can produce a characteristic squeaky pop under appropriate small-scale test conditions.
Therefore, collecting and identifying the gas provides evidence that hydrogen is one of the reaction products.
Reactivity Down Group 1
The most important chemical trend is:
reactivity increases down Group 1.
Therefore:
Li < Na < K < Rb < Cs
Lithium is less reactive than sodium.
Sodium is less reactive than potassium.
Potassium is less reactive than rubidium.
The trend continues down the group.
Why Reactivity Increases Down Group 1
All alkali metals react by losing their single outer electron.
Moving down Group 1:
- the number of occupied electron shells increases
- atomic radius increases
- shielding increases
- the outer electron becomes farther from the nucleus
- the effective attraction between the nucleus and outer electron decreases
Therefore, the outer electron becomes easier to remove.
The easier the electron is to remove, the more readily the metal reacts.
Visualizing the Periodic Trend
The change in atomic radius and first ionization energy helps explain the Group 1 reactivity pattern.
Down Group 1, atomic radius generally increases while first ionization energy decreases.
A lower first ionization energy means less energy is required to remove the outer electron.
This helps explain why the metals become increasingly reactive.
The Reactivity Explanation Chain
A strong explanation can be organized as:
Down Group 1
↓
More occupied electron shells
↓
Greater atomic radius
↓
Greater shielding
↓
Weaker effective attraction for the outer electron
↓
Outer electron is easier to remove
↓
+1 ion forms more readily
↓
Reactivity increases
This is much stronger than simply memorizing:
"Group 1 becomes more reactive down the group."
Comparing Lithium, Sodium, and Potassium
The reaction with water makes the trend easy to observe.
Lithium
Relatively gentle fizzing and slower movement.
Sodium
More rapid fizzing and movement; often melts into a ball.
Potassium
Very vigorous reaction; ignition and a lilac flame may occur.
These observations provide experimental evidence that:
Li < Na < K
in reactivity.
Reaction with Oxygen
Alkali metals also react readily with oxygen.
For example, lithium reacts with oxygen to form lithium oxide:
4Li + O₂ → 2Li₂O
The exact oxygen-containing products can differ among the Group 1 metals under different conditions, especially farther down the group.
The important general idea is that alkali metals readily react with oxygen and form compounds containing +1 metal ions.
Reaction with Halogens
Alkali metals react with halogens to produce ionic compounds.
For example:
2Na + Cl₂ → 2NaCl
Sodium loses electrons:
Na → Na⁺ + e⁻
Chlorine gains electrons.
The resulting ions attract each other and form sodium chloride.
Group 1 metals therefore commonly form compounds with formulas such as:
NaCl
KBr
LiF
because the metal ion has a charge of +1 and the halide ion has a charge of −1.
Flame Colours
Some alkali metal ions produce characteristic colours when heated in a flame.
Common examples include:
Lithium → crimson red
Sodium → yellow
Potassium → lilac
These colours can help identify certain metal ions.
They occur because electrons absorb energy and move to higher energy levels. When they return to lower energy levels, energy is released as light of characteristic wavelengths.
Explaining Experimental Observations
Chemistry requires us to do more than simply record what happened.
We should connect observations to scientific explanations.
For example:
Observation: Bubbles appear.
Explanation: Hydrogen gas is being produced.
Observation: Sodium becomes a ball.
Explanation: Heat from the reaction melts the sodium.
Observation: Potassium reacts more vigorously than sodium.
Explanation: Potassium's outer electron is farther from the nucleus and more shielded, so it is easier to remove.
Observation: The solution becomes alkaline.
Explanation: A soluble metal hydroxide has formed.
Observation vs Explanation
It is important to distinguish an observation from an explanation.
An observation is something that can be detected or measured.
Examples:
- bubbles formed
- the metal moved
- a flame appeared
- the metal became smaller
- the solution changed indicator colour
An explanation uses scientific ideas to explain why the observation occurred.
Examples:
- hydrogen gas was produced
- the reaction released heat
- a metal hydroxide formed
- electrons were transferred
Good experimental writing often includes both.
Predicting an Unfamiliar Alkali Metal
Suppose element X is located below potassium in Group 1.
Even if we have never observed X, we can predict that it:
- has one valence electron
- forms X⁺ ions
- is metallic
- should lose its outer electron relatively easily
- should be more reactive than potassium
- would be expected to react vigorously with water
This demonstrates one of the most powerful uses of the periodic table: predicting properties from position.
Worked Example 1
Which element is more reactive?
Lithium or sodium?
Sodium is below lithium in Group 1.
Moving down the group:
- atomic radius increases
- shielding increases
- the outer electron becomes easier to remove
Therefore:
sodium is more reactive than lithium.
Worked Example 2
Explain why potassium reacts more vigorously with water than lithium.
Potassium has more occupied electron shells than lithium.
Therefore:
- potassium has greater shielding
- its outer electron is farther from the nucleus
- the effective attraction holding that electron is weaker
- the electron is easier to remove
Therefore:
potassium reacts more readily than lithium.
Worked Example 3
A Group 1 metal reacts with water and produces a gas.
What is the gas?
The general reaction is:
metal + water → metal hydroxide + hydrogen
Therefore:
the gas is hydrogen.
Worked Example 4
A freshly cut piece of sodium is shiny but quickly becomes dull.
Explain the observation.
The freshly exposed sodium surface reflects light and appears shiny.
Sodium is reactive and quickly reacts with substances in the air.
Reaction products form on the surface.
Therefore, the shiny surface becomes:
dull or tarnished.
Worked Example 5
Why does sodium form Na⁺ rather than Na²⁺ in its usual compounds?
Sodium has the electron arrangement:
2,8,1
It can achieve a stable outer electron arrangement by losing one electron.
This produces:
Na⁺
Removing a second electron would require removing an electron from the now-filled inner shell and requires far more energy.
Therefore, sodium normally forms:
Na⁺ rather than Na²⁺.
Alkali Metals and Ionic Compounds
Because Group 1 metals form +1 ions, their ionic formulas can often be predicted easily.
With chloride:
Na⁺ + Cl⁻ → NaCl
With oxide:
Two Na⁺ ions are needed to balance one O²⁻ ion:
Na₂O
With sulfide:
K₂S
Knowing the charge of Group 1 ions therefore helps us predict chemical formulas.
Group 1 and Periodicity
Group 1 demonstrates several important periodic trends.
Moving down the group:
Atomic radius: generally increases
Shielding: increases
First ionization energy: generally decreases
Reactivity: increases
Metallic character: increases
Melting point: generally decreases
These trends are connected rather than being unrelated facts.
Atomic structure helps explain many of them.
Safety and Alkali Metals
Alkali metals are important laboratory substances, but their reactions can become vigorous, especially farther down the group.
For this reason, reactions involving reactive alkali metals should be conducted only using appropriate laboratory procedures and supervision.
The dramatic nature of the reactions is also useful evidence of the increasing reactivity down the group.
Real-World Uses
Despite their reactivity, Group 1 elements and their compounds have many important applications.
Examples include:
Lithium compounds
Used in rechargeable batteries and other technologies.
Sodium compounds
Common in products such as sodium chloride and sodium hydroxide.
Potassium compounds
Important in fertilizers because potassium is an essential plant nutrient.
The pure metals themselves are highly reactive, so many everyday applications involve their compounds rather than the elemental metals.
A Strategy for Group 1 Questions
When solving a Group 1 problem, ask:
1. Which element is being discussed?
2. Where is it located within Group 1?
3. How many valence electrons does it have?
Answer:
one
4. What ion will it form?
Answer:
+1
5. If comparing reactivity, which element is farther down the group?
The lower element will generally be more reactive.
6. Can the trend be explained using atomic structure?
Think about:
- number of electron shells
- atomic radius
- shielding
- attraction to the outer electron
- ease of electron loss
Common Mistakes
Mistake 1: Including hydrogen as an alkali metal
Hydrogen is located in Group 1 but is not an alkali metal.
Mistake 2: Saying Group 1 metals gain one electron
Group 1 metals lose one electron.
Mistake 3: Saying they form −1 ions
Losing one negatively charged electron produces:
+1 ions
Mistake 4: Saying reactivity decreases down Group 1
The correct trend is:
reactivity increases down Group 1.
Mistake 5: Saying the nucleus itself becomes weaker
A better explanation is that increased atomic radius and shielding reduce the effective attraction experienced by the outer electron.
Mistake 6: Saying bubbles are an explanation
"Bubbles formed" is an observation.
"The bubbles are hydrogen gas produced by the reaction" is an explanation.
Mistake 7: Assuming the metal hydroxide is the gas
The metal hydroxide remains in solution.
The gas produced is:
hydrogen.
Did You Know?
The word alkali refers to a soluble base.
Group 1 metals are called alkali metals because their reactions with water produce soluble metal hydroxides, which form alkaline solutions.
For example:
sodium + water → sodium hydroxide + hydrogen
The sodium hydroxide makes the resulting solution strongly alkaline.
Key Terms
Alkali metal: A metallic element in Group 1 of the periodic table, excluding hydrogen.
Valence electron: An electron in the outermost occupied shell of an atom.
Reactivity: How readily a substance undergoes a chemical reaction.
Cation: A positively charged ion.
Shielding: The reduction in effective nuclear attraction caused by inner electrons.
Atomic radius: A measure of the size of an atom.
Ionization energy: The energy required to remove an electron from an isolated gaseous atom.
Metal hydroxide: A compound containing a metal ion and hydroxide ions.
Alkaline: Having a pH greater than 7 in aqueous solution.
Observation: Something detected or measured during an experiment.
Explanation: A scientific reason for an observation.
Key Equations and Patterns
Ion formation:
M → M⁺ + e⁻
Reaction with water:
2M + 2H₂O → 2MOH + H₂
Example:
2Na + 2H₂O → 2NaOH + H₂
Reaction with chlorine:
2M + Cl₂ → 2MCl
Example:
2Na + Cl₂ → 2NaCl
Reactivity trend:
Li < Na < K < Rb < Cs
Reactivity:
increases down Group 1 ↓
Key Takeaways
- The alkali metals are found in Group 1 of the periodic table.
- Lithium, sodium, potassium, rubidium, caesium, and francium are alkali metals.
- Hydrogen is located in Group 1 but is not classified as an alkali metal.
- All alkali metals have one valence electron.
- They tend to lose this electron during chemical reactions.
- Losing one electron produces a +1 ion.
- Alkali metals are generally soft, conductive, and relatively low-density compared with many other metals.
- Freshly cut alkali metals can appear shiny but quickly become dull because they react with substances in the air.
- Alkali metals react with water to produce a metal hydroxide and hydrogen gas.
- The metal hydroxide makes the resulting solution alkaline.
- Lithium reacts less vigorously with water than sodium.
- Sodium reacts less vigorously than potassium.
- Reactivity increases down Group 1.
- Down the group, atomic radius and shielding increase.
- The outer electron therefore experiences weaker effective nuclear attraction and becomes easier to remove.
- Easier electron loss explains the increasing reactivity down Group 1.
- Observations such as fizzing, movement, melting, flames, and disappearing metal can be explained using the products and energy changes of the reaction.
- Group 1 metals also react with oxygen and halogens.
- Their +1 ion charge allows us to predict formulas of many ionic compounds.
- The periodic table allows us to predict the properties and reactivity of unfamiliar Group 1 elements from their position.