Problem 31 - Chemistry: Principles and Reactions

Question

Predict the geometry of the following species: (a) \(\mathrm{SCO}\) (b) \(\mathrm{IBr}_{2}^{-}\) (c) \(\mathrm{NO}_{3}^{-}\) (d) \(\mathrm{RnF}_{4}\)

Step-by-Step Solution

Answer

Answer: The predicted molecular geometries are as follows: - SCO: Trigonal planar - IBr2-: Bent (V-shaped) - NO3-: Trigonal planar - RnF4: Tetrahedral

Step 1: Predicting the geometry of \(\mathrm{SCO}\)

First, identify the central atom, which is S. Draw the Lewis structure of the molecule. Sulfur has 6 valence electrons, and oxygen and carbon each have 4. After connecting them to make SCO (with carbon in the middle), oxygen and sulfur will have single bonds to carbon. Carbon will have a double bond to oxygen. Therefore, carbon has 3 electron groups surrounding it: a single bond to sulfur, a double bond to oxygen, and another single bond to the other oxygen atom. These 3 electron groups will result in a trigonal planar geometry.

Step 2: Predicting the geometry of \(\mathrm{IBr}_{2}^{-}\)

First, identify the central atom, which is I. Draw the Lewis structure of the ion. Iodine has 7 valence electrons, and each bromine atom has 7 valence electrons. There is also an additional electron from the negative charge. Connecting iodine to each bromine atom with a single bond leaves one lone pair on iodine. Now, iodine is surrounded by 3 electron groups: 2 single bonds with bromine, and one lone pair. These 3 electron groups result in a trigonal planar electron-group geometry, but because of the lone pair, the molecular geometry will be bent (or V-shaped) instead.

Step 3: Predicting the geometry of \(\mathrm{NO}_{3}^{-}\)

First, identify the central atom, which is N. Draw the Lewis structure of the ion. Nitrogen has 5 valence electrons, and each oxygen atom has 6 valence electrons. There is also an additional electron from the negative charge. Nitrogen will form a single bond to two of the oxygen atoms and a double bond to the third oxygen atoms. This allows all atoms to fulfill the octet rule. Nitrogen has 3 electron groups surrounding it: 2 single bonds to oxygen and 1 double bond to another oxygen. These 3 electron groups will result in a trigonal planar geometry.

Step 4: Predicting the geometry of \(\mathrm{RnF}_{4}\)

First, identify the central atom, which is Rn. Draw the Lewis structure of the molecule. Radon has 8 valence electrons, and each fluorine atom has 7 valence electrons. Connecting radon to each fluorine atom with a single bond leaves no lone pairs on radon. Radon has 4 electron groups surrounding it: 4 single bonds with fluorine atoms. These 4 electron groups will result in a tetrahedral geometry.

Key Concepts

Chemical Bonding

Chemical bonding is the force that holds atoms together to form molecules and compounds. At the foundation of chemical bonding lies the electron configuration of atoms. An atom will typically bond with another atom in order to reach a more stable configuration, generally by filling its valence shell, which is the outermost shell of electrons.

There are three primary types of chemical bonds: ionic, covalent, and metallic bonds.

• Ionic bonds form when there is a transfer of electrons from one atom to another, typically between metals and non-metals, resulting in positively and negatively charged ions that attract each other.
• Covalent bonds form when two atoms share electrons to fill their valence shells, often seen in compounds consisting of non-metals.
• Metallic bonds involve a 'sea' of delocalized electrons shared by metal cations.

Understanding the type of bonding and the electron arrangement allows us to predict the properties of the substance and its molecular geometry.

Lewis Structures

Lewis structures, also known as Lewis dot diagrams, are a visual representation of the bonding between atoms and the arrangement of electrons within a molecule. Drawing a Lewis structure involves several steps:

• Determining the total number of valence electrons available for bonding.
• Arranging the atoms with the least electronegative atom typically in the center (except for hydrogen) and connecting them with single bonds.
• Distributing the remaining valence electrons to complete the octets (or duets for hydrogen) while giving preference to more electronegative atoms first.
• Making double or triple bonds if necessary to fulfill the octet rule.

Lewis structures are essential for predicting the geometry of molecules because they show how many electron groups surround a central atom, a critical factor in molecular shape.

VSEPR Theory

VSEPR theory, which stands for Valence Shell Electron Pair Repulsion theory, is a model used to predict the geometry of individual molecules based on the number of electron pairs surrounding their central atoms. The premise of the VSEPR theory is that electron pairs, including bonding pairs and lone pairs, repel each other and will therefore arrange themselves as far apart as possible.

This repulsion causes the molecule to adopt a specific shape that minimizes the repulsion. Important shapes include linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral. The presence of lone pairs can also alter the observable shape of a molecule because lone pairs exert greater repulsion than bonding pairs.

Electron Group Geometry

Electron group geometry refers to the spatial arrangement of electron groups (bonding and lone pairs) around the central atom in a molecule. Unlike molecular geometry, which describes the layout of atoms, electron group geometry only considers the electron groups.

The main electron group geometries are:

• Linear: 180-degree bond angles, with two electron groups.
• Trigonal planar: 120-degree bond angles, with three electron groups.
• Tetrahedral: 109.5-degree bond angles, with four electron groups.
• Trigonal bipyramidal: a combination of 90-degree and 120-degree bond angles, with five electron groups.
• Octahedral: 90-degree bond angles, with six electron groups.

This geometry can be predicted after drawing the Lewis structure by counting bonding pairs and lone pairs of electrons. The molecular geometry, accounting for lone pairs, may differ from the electron group geometry depending on how the electron groups are distributed.