What do we want our students to say when we ask them to describe ionic bonding? Ideally, they would be talking about the electrostatic attraction between oppositely charged ions within a giant lattice of alternating positive and negative ions. However, the way we have traditionally taught ionic bonding in pre-16 courses sometimes leads to confusion. Students often discuss the electron transfer process that creates ions from the reaction between metal and nonmetal atoms. This misconception can hinder their progress in more advanced courses. So, how can we help students better understand and discern these ideas?
Dot and cross diagrams are regularly used to explain ionic bonding. These show the transfer of the outer electrons from the metal atoms to fill the outer shell of the nonmetal atoms. However, this transfer of electrons is what happens during the redox reaction between a metal and a non-metal, and is not a type of bond.
The emphasis on dot and cross diagrams also leads to an appeal to the octet rule to explain why binding occurs. This rule can lead to the idea that atoms are always trying to get a stable full outer shell of electrons. The full outer shell rule can be helpful, but it can also hinder students from developing more advanced understanding if they see the rule as the driving force behind reaction.
- Read Beyond appearances by Vanessa Kind, starting on page 56, for an in-depth analysis of the research that identifies misconceptions and offers suggestions for how to address them.
- When teaching younger students, emphasize the importance of electrostatic attractions with a classroom activity on modeling ionic bonds with plasticine.
- Use this ionic connection probe to find out what your students are thinking and facilitate discussions about their understanding.
- Discover and challenge misconceptions about binding with a toolkit of pre-prepared diagnostic and response activities about particles and structure.
- Read Beyond appearances by Vanessa Kind, starting on page 56, for an in-depth analysis of the research identifying misconceptions, as well as suggestions on how to address them: rsc.li/3iCrrSp
- When teaching younger students, emphasize the importance of electrostatic attractions with a classroom activity on modeling ionic bonds with plasticine: rsc.li/3UC3mIY
- Use an ionic binding diagnostic probe to find out what your students are thinking and facilitate discussion of their misconceptions: rsc.li/3iz8jVo
- Discover and challenge misconceptions about binding with a toolkit of pre-prepared particle and structure diagnostic and response activities: bit.ly/3FzxiRN
What you need to know
Students progressing to post-16 courses will study the energy transfers involved in ion formation. They will find that the first enthalpy of ionization of a gaseous metal atom, such as sodium, is very endothermic. Atoms certainly do not ‘want’ to lose their outermost electrons. An isolated metal atom is therefore much more stable with its outermost electrons than as a positive ion. However, it is common for students to hold on to the idea that metal atoms lose their outermost electrons as they try to get a full outer shell.
What about adding an electron to an isolated nonmetal atom like chlorine? This process, called first electron affinity, is exothermic. However, the magnitude of this enthalpy change does not compensate for the highly endothermic removal of the electrons from the metal. Also, the transfer of a second electron to an outer shell is an endothermic process, due to the repulsion between the second electron and the already negative ion.
However you look at it, the electron transfer process represented by the dot and cross diagrams on pre-16 is not energetically viable. Explaining that ionic bonding occurs so that atoms get stable full outer shells just doesn’t fit energetics. This begs the question, why do these ions form at all?
After 16, students move on to look at Born-Haber enthalpy cycles, which describe the formation of ionic compounds from elements. Born-Haber cycles bring together all the energy transfer processes involved. They show that to get to the isolated atoms depicted in the dot and cross diagrams, you need to break the metallic bond in the solid metal and break the covalent bond in the diatomic non-metal molecules. When this is added to the endothermic electron transfer process, it means that a large amount of energy must be transferred to the reacting elements to form the gaseous ions.
Suggestions for your teaching
With current specifications and textbooks, it is not possible to completely separate the dot and cross diagrams from the description of ionic bonding itself. However, here are a few suggestions for a learning sequence that can help students avoid potential misconceptions.
- Introduce ions and their formulas to students before they study bonding, ideally when they first encounter atoms and molecules. Point out that some atoms and molecules can become charged and these charged particles are then called ions. Show them labels of mineral water bottles for interesting context.
- After you’ve learned atomic structure, you’ll move on to introduce dot and cross diagrams to show how atoms form ions and molecules. At this stage, separate the loss of electrons from metal atoms from the gain of electrons from non-metal atoms and do not call it an ionic bond. Point out that atoms often acquire a complete outer shell after forming molecules or ions. Emphasize that we can use this rule to predict the number of covalent bonds or charges, but emphasize that this is not the whole story.
- Focus on comparing the origin and strength of the electrostatic forces that hold different types of structures together – giant metals, giant ionic, giant covalent, and simple molecular covalent. Do not use dot and cross charts at this point. Cement the ideas of bonding and electrostatic forces in students’ minds.
- Use these ideas to position students to understand more topics related to ions in redox reactions, such as reactions of group 1 and 7 elements, metal acid reactions, and electrolysis. That’s the time to show the transfer of electrons from metals to non-metal atoms and the half equations that go with it.