Electron pair geometry (EPG) describes the spatial arrangement of electron clouds surrounding a molecule’s central atom. This arrangement includes both bonding electrons and non-bonding electrons (lone pairs). Understanding EPG is fundamental because it dictates the three-dimensional structure that the atoms themselves will adopt. This predictable shape affects a molecule’s properties, such as its polarity and how it interacts with other substances.
Essential First Step: Drawing the Lewis Structure
The first step to determine the electron pair geometry is to accurately map the molecule’s electron distribution using a Lewis structure. This two-dimensional diagram shows the connectivity between atoms and the location of all valence electrons. The process begins with calculating the total number of valence electrons, adjusting for any overall ionic charge.
Next, arrange the atoms into a skeletal structure, typically placing the least electronegative atom in the center. Draw single bonds between the central atom and the surrounding atoms. The remaining valence electrons are then distributed to satisfy the octet rule for all outer atoms (eight electrons, or two for hydrogen). Finally, any leftover electrons are placed on the central atom as lone pairs, and multiple bonds are formed if the central atom still lacks a complete octet.
Identifying and Counting Electron Groups
Once the Lewis structure is complete, focus shifts to the central atom to count its surrounding “electron groups,” also referred to as electron domains. This total count is the sole factor determining the electron pair geometry. An electron group is any region of high electron density around the central atom. These clouds of negative charge repel one another, pushing as far apart as possible in three-dimensional space to minimize repulsion.
A crucial rule for counting is that every lone pair on the central atom counts as a single electron group. Furthermore, any bond connecting the central atom to another atom—whether single, double, or triple—is also counted as only one electron group. For example, the carbon atom in carbon dioxide (\(\text{CO}_2\)) has two double bonds and no lone pairs, resulting in two electron groups. In contrast, the oxygen atom in water (\(\text{H}_2\text{O}\)) has two single bonds and two lone pairs, resulting in four electron groups.
This total count of electron groups is called the steric number. This number dictates the fundamental spatial arrangement and is the direct link between the two-dimensional drawing and the three-dimensional prediction of the electron cloud arrangement.
The Standard Electron Pair Geometries
The total count of electron groups around the central atom dictates one of the five primary electron pair geometries.
When the count is two, the two electron domains arrange themselves on opposite sides of the central atom for maximum separation. This arrangement creates a Linear geometry, with an angle of \(\text{180}^\circ\) between the electron groups. Carbon dioxide is an example of this linear arrangement.
A central atom with three electron groups adopts a Trigonal Planar geometry. The groups spread out into a single plane, forming the corners of an equilateral triangle. This planar arrangement results in ideal angles of \(\text{120}^\circ\). Boron trifluoride (\(\text{BF}_3\)), where all three groups are bonding pairs, is characteristic of this structure.
When the electron group count is four, the geometry established is Tetrahedral. This structure expands into three dimensions, with the four domains pointing towards the corners of a tetrahedron. The bond angles in a perfect tetrahedral structure are \(\text{109.5}^\circ\), representing the greatest separation possible. Methane (\(\text{CH}_4\)) is the primary example of this geometry.
If the central atom is surrounded by five electron groups, the geometry is Trigonal Bipyramidal. This structure contains two non-equivalent spatial positions: three groups lie in an equatorial plane, and two groups occupy axial positions. The equatorial groups are separated by \(\text{120}^\circ\) angles, while the axial groups are positioned at \(\text{90}^\circ\) angles relative to the equatorial plane.
Finally, a count of six electron groups yields the Octahedral geometry. This arrangement features six corners, with all six electron domains positioned at \(\text{90}^\circ\) angles relative to their neighbors. Sulfur hexafluoride (\(\text{SF}_6\)) is a classic example of this highly symmetric structure.
Understanding the Difference: EPG vs. Molecular Geometry
A frequent point of confusion is the distinction between electron pair geometry and molecular geometry. EPG is the arrangement of all electron groups, including bonding electrons and lone pairs, determined by the total count. Molecular geometry, however, describes only the spatial arrangement of the atoms in the molecule. The lone pairs are ignored when naming the final shape.
When a central atom has no lone pairs, the EPG and the molecular geometry are identical, such as the tetrahedral EPG of methane. The presence of one or more lone pairs causes the two geometries to diverge. These non-bonding pairs still occupy space and exert a stronger repulsive force on the bonding electrons.
Consider the ammonia molecule (\(\text{NH}_3\)), which has three bonding pairs and one lone pair, giving it a tetrahedral EPG. The lone pair forces the three hydrogen atoms closer together, resulting in a trigonal pyramidal molecular geometry. Similarly, the oxygen atom in water (\(\text{H}_2\text{O}\)) has two bonding pairs and two lone pairs, maintaining a tetrahedral EPG. The two lone pairs push the hydrogen atoms into a non-linear, bent shape, which is its molecular geometry.