||Why do atoms form ions?
In this activity, students, in small groups, decide whether statements they are given about the formation of sodium and chloride ions are true or false. By thinking about and discussing the statements, students check their understanding about ionic bonding.
This activity could be used to assess prior learning.
- understand how and why ionic compounds such as sodium chloride are formed.
Sequence of activities
Demonstrate the reaction of sodium with chlorine (Classic chemistry demonstration 49) and the subsequent test for chloride ions using silver nitrate solution.
Issue a mini whiteboard to each student and ask them to write down:
- what particles are present in sodium and chlorine before the reaction
- what particles are present in sodium chloride after the reaction.
Pose questions, for example:
What has changed during the chemical reaction to cause such an enormous change in the appearance of reactants and products?
to stimulate discussion about the use of the words atoms, molecules and ions before sharing the purpose of the session.
Organise students into groups of three. Tell each group to number its members, one, two or three. Ask:
- number one’s to explain what they mean by an atom
- number two’s to explain what they mean by a molecule
- number three’s what they mean by an ion
to the other two students in the group.
Allow time for discussions to come to agreed explanations. Support individual groups by using questions to focus their ideas.
Give each student a Why do atoms form ions? sheet.
Ask them to:
- work individually
- decide whether they think that the 10 statements are true or false
- record this in the table.
Invite students to:
- return to their groups of three
- share their ideas
- agree group responses.
||If a data projector is available, show an electronic animation of the formation of the sodium chloride ionic lattice.
In a plenary:
- select members of different groups to describe and explain their responses
- encourage other groups to challenge or add to explanations.
Provide an opportunity for students to write down in what ways they have changed their original ideas, as a result of listening to others during the session.
Take in the sheets and comment on how their ideas have developed. Draw attention to ways in which individuals might need to develop their explanations further.
Assessment for learning commentary
The initial demonstration and discussion about the meaning of words such as atom, molecule and ion draw attention to the learning objectives.
Students compare their responses to statements about the formation of ions and have the opportunity to articulate their ideas. The electronic animation showing the formation of a sodium chloride lattice helps students recognise the standard they are aiming for.
They examine how their ideas have developed during the session and written feedback supports students should they need to develop their understanding further.
For each student
||Why do atoms form ions?
For the demonstration of the reaction of sodium with chlorine
- Chlorine generator
- Gas jar with lid or 500 cm3 conical flask with bung to fit
- Small piece of ceramic material, such as a piece of evaporating basin or heat-proof mat, to fit inside jar or flask
- Dropping pipette or wash bottle
- Knife to cut sodium
- Filter paper or paper towels to wipe the oil from the sodium
- Access to a fume cupboard
- 10 g of potassium manganate(VII) and about 50 cm3 of concentrated hydrochloric acid
- Piece of sodium about half the size of a pea
- A little silver nitrate solution
- A few cm3 of a hydrocarbon solvent such as hexane to clean the oil from the sodium
- Vaseline for the gas jar lid
For the electronic animation
- Animation showing the formation of the sodium chloride lattice
- Data projector.
Notes on responses to statements
- A sodium atom spontaneously loses an electron to get a full shell of electrons.
The ionisation energy has to be supplied to remove an electron. It is not a spontaneous process.
- A Na7- ion is more stable than a sodium atom because it has a full shell of electrons.
The Na7- ion would be highly unstable because the repulsions between electrons would far out balance the attractions between protons and electrons. Having a full shell of electrons does not automatically mean that an ion is stable.
- A Cl7+ ion is just as stable as a Cl- ion because they both have a full shell of electrons.
A great deal of energy would be required to ionise a chlorine atom to give a 7+ cation. Having a full shell of electrons does not automatically mean that an ion is stable.
- Each proton in the nucleus of an atom attracts one specific electron.
All the protons in the nucleus attract all of the electrons (and vice versa).
- Energy is required to remove an electron from an atom.
This is the ionisation enthalpy.
- When an atom is ionised, it then requires even more energy to remove a second electron.
This is because there is less repulsion from other electrons counteracting the attraction from the nucleus and because the ionic radius decreases so that the second electron is closer to the nucleus.
- Once you’ve removed an electron from a sodium atom you can never put it back.
The separated cation and electron are less stable which is why energy is needed to ionise an atom.
- Once you have removed one electron from a sodium atom you can’t remove another because that would mean it no longer had a full electron shell.
It is perfectly possible to remove another electron from the sodium ion, although the ionisation energy needed will be significantly higher than that needed to remove the first electron.
- Solid sodium chloride contains pairs of sodium and chloride ions which are kept together by their opposite charges.
Solid sodium chloride is made up of a three dimensional lattice of sodium and chloride ions.
- When sodium chloride dissolves in water the solution contains molecules of sodium chloride.
When solid sodium chloride dissolves in water the sodium and chloride ions become hydrated and are dispersed through the solution.
T. Lister, Classic Chemistry Demonstrations. London: Royal Society of Chemistry, 1995.
K. Taber, Chemical misconceptions – prevention, diagnosis and cure Volume 2: classroom resources. London: Royal Society of Chemistry, 2002.