Dot and Cross Diagrams: A Guide to Electron Sharing and Transfer
The Building Blocks: Atoms and Electrons
Before we can draw bonds, we need to understand the players. Everything around us is made of atoms. Each atom has a tiny nucleus at its center, surrounded by even smaller particles called electrons. These electrons zoom around the nucleus in regions called shells or energy levels.
The electrons in the outermost shell are the most important for bonding and are called valence electrons. Atoms are generally most stable when they have a full outer shell. For many common elements, this stable configuration resembles that of a noble gas[1], with 8 electrons in its outer shell (or 2 for the very first shell). This desire for a full outer shell is the driving force behind all chemical bonding.
Two Ways to Bond: An Overview
Atoms can achieve stable electron configurations in two primary ways, leading to two main types of bonds, each represented differently in dot and cross diagrams.
| Feature | Ionic Bonding | Covalent Bonding |
|---|---|---|
| Process | Transfer of electrons from one atom to another. | Sharing of electrons between atoms. |
| Between | A metal and a non-metal. | Two or more non-metals. |
| Resulting Particles | Positive and negative ions[2] that attract. | Molecules. |
| Diagram Focus | Shows electron transfer and resulting charges. | Shows shared pairs of electrons between atomic nuclei. |
Drawing Covalent Bonds: Sharing is Caring
In covalent bonding, atoms share pairs of valence electrons to fill their outer shells. In a dot and cross diagram, we use different symbols (dots and crosses) for the electrons coming from different atoms. This makes it easy to see where each electron originated.
Example 1: Hydrogen Molecule (H$_2$)
A hydrogen atom has 1 valence electron. It needs one more to fill its first shell (which holds a maximum of 2 electrons).
- Each H atom is represented by its symbol: H.
- One H atom contributes an electron (shown as a dot), and the other contributes an electron (shown as a cross).
- These two electrons are shared between the atoms, forming a single covalent bond.
The diagram shows: H· + ×H → H·×H
This shared pair ($\centerdot$$\times$) means each hydrogen atom now has a share of 2 electrons, achieving a stable arrangement.
Example 2: Water Molecule (H$_2$O)
An oxygen atom has 6 valence electrons. It needs 2 more to achieve a full shell of 8.
- The oxygen atom is the central atom.
- It shares one pair of electrons with one hydrogen atom and another pair with a second hydrogen atom.
- Oxygen's electrons might be shown as dots, while the electrons from the hydrogen atoms are shown as crosses.
This results in two O-H covalent bonds. The final diagram shows oxygen surrounded by 8 electrons (both its own and the shared ones), and each hydrogen is surrounded by 2 electrons.
Drawing Ionic Bonds: The Great Electron Handover
In ionic bonding, one atom completely gives up one or more electrons to another atom. This transfer creates particles with a net electrical charge called ions. The atom that loses electrons becomes a positively charged cation, and the atom that gains electrons becomes a negatively charged anion. The opposite charges attract, forming the ionic bond.
Example: Sodium Chloride (NaCl)
Sodium (Na) is a metal with 1 valence electron. Chlorine (Cl) is a non-metal with 7 valence electrons.
- Sodium wants to lose its 1 valence electron to expose its full inner shell.
- Chlorine wants to gain 1 electron to complete its outer shell with 8 electrons.
The process is simple: Sodium donates its single valence electron (shown as a cross) to chlorine.
Before: Na (with ×) and Cl (with 7 dots).
After: The sodium ion (Na$^+$) has lost the cross, and the chloride ion (Cl$^-$) has gained it, giving it 8 electrons in its outer shell (7 dots and 1 cross). The diagram shows the ions as [Na]$^+$ and [Cl······×]$–. The + and - charges are crucial as they show the electrostatic attraction that is the ionic bond.
From Simple Molecules to Complex Structures
Dot and cross diagrams can also be used for more complex molecules and giant structures.
Carbon Dioxide (CO$_2$): Carbon has 4 valence electrons and needs 4 more. Each oxygen has 6 and needs 2 more. The carbon atom forms two double bonds, sharing two pairs of electrons with each oxygen atom. The diagram shows O::C::O, with dots and crosses distinguishing the electrons from carbon and oxygen.
Giant Covalent Structures: For substances like diamond (a form of carbon), each carbon atom is bonded to four others in a giant network. A dot and cross diagram can only show a tiny part of this vast structure, focusing on one central atom and its four immediate neighbors.
Giant Ionic Lattices: Similarly, ionic compounds like sodium chloride form a 3D lattice where each Na$^+$ ion is surrounded by Cl$^-$ ions and vice versa. A dot and cross diagram typically illustrates the formation of a single formula unit, like one Na$^+$ and one Cl$^-$, but it's important to remember that this repeats billions of times in all directions in a crystal.
Common Mistakes and Important Questions
Q: Why do we use dots and crosses instead of just dots for all electrons?
A: Using dots and crosses helps us track where the electrons came from. This is especially important for understanding how the bond formed and for correctly drawing the diagram without accidentally creating or destroying electrons. It clarifies the process of sharing or transfer.
Q: What is the most common mistake when drawing these diagrams?
A: The most common mistake is showing the wrong number of electrons in the final diagram. For covalent bonds, ensure that each atom (especially the central one) has the correct number of electrons around it (often 8). For ionic bonds, a frequent error is forgetting to include the charges on the ions or drawing the brackets incorrectly. The charges are what hold the compound together.
Q: Can a dot and cross diagram predict the shape of a molecule?
A: No, a basic dot and cross diagram only shows the connectivity (which atoms are bonded) and the origin of the electrons. It is a 2D representation. The 3D shape of a molecule is determined by the repulsion between electron pairs in the outer shell and requires different models, like VSEPR[3] theory, to predict.
Footnote
[1] Noble Gas: Elements in Group 8/18 of the periodic table, such as Helium (He) and Neon (Ne). They are very stable and unreactive because they already have full outer electron shells.
[2] Ion: An atom or molecule that has a net electrical charge because it has lost or gained one or more electrons.
[3] VSEPR: Valence Shell Electron Pair Repulsion. A model used in chemistry to predict the geometry of individual molecules based on the repulsion between electron pairs in the valence shell of the central atom.