Covalent Bonding, Lewis Structures, and Molecular Geometry

I. Covalent Bonds & The Periodic Table

Learning Objectives:

• Predict the number of covalent bonds an atom forms based on its periodic table position and valence electron count.

• Memorize specific exceptions for Hydrogen (H) and Boron (B) that do not follow the octet rule.

• Use the octet rule as a guiding principle to determine when multiple covalent bonds (double and triple) are necessary in a molecule.

• Draw accurate Lewis structures from given molecular formulas, effectively applying the octet rule and considering formal charges.

• Draw accurate 3-D wedge/dash representations of molecules to depict their spatial arrangement.

• Utilize Lewis structures to accurately predict and describe molecular geometry based on the VSEPR model.

• (Note: Expanded octets will be named, but their 3-D representations will not be drawn in detail.)

General Principles:

• Covalent bonds primarily form between two non-metal atoms (or between a non-metal and a metalloid like Boron or Silicon). This occurs because non-metals have high electronegativity and tend to gain electrons, and by sharing, they can mutually achieve a stable electron configuration.

• Atoms share valence electrons to achieve a stable electron configuration, typically an octet (8 valence electrons) in their outermost shell, mimicking the electron configuration of noble gases. This sharing results in a lower energy state for the molecule compared to isolated atoms.

Exceptions to the Octet Rule:

• Boron (B):

  • Possesses only 3 valence electrons to share (electron configuration [He]2s^22p^1).

  • In most stable compounds, Boron forms only 3 covalent bonds, such as in BF_3.

  • It achieves stability with only 6 electrons in its valence shell, rather than the typical 8. It is considered an electron-deficient atom.

• Hydrogen (H):

  • Possesses only 1 valence electron to share (electron configuration 1s^1).

  • Forms compounds with only 1 covalent bond, for example, in H_2 or HCl.

  • Achieves stability by having only 2 electrons in its valence shell, mimicking Helium's electron configuration. This is often referred to as the duet rule.

• Expanded Octet (Elements in the third row and below on the periodic table):

  • These elements (e.g., Phosphorus (P), Sulfur (S), Chlorine (Cl), Bromine (Br), Iodine (I)) possess vacant d orbitals in their valence shell.

  • The availability of these vacant d orbitals allows them to accommodate more than 8 valence electrons, enabling them to form more than 4 bonds.

  • This ability to make more than 4 bonds and hold more than 8 electrons is known as an expanded octet.

  • (These expanded octets will be identified in Lewis structures but will not be depicted with detailed 3-D geometry focuses on their specific arrangement due to the complexity of d orbital involvement.)

Typical Bonding Patterns (HONC 1234 mnemonic):

This mnemonic helps quickly recall the most common number of bonds formed by important non-metal elements to satisfy the octet rule (or duet rule for H):

• Hydrogen: Typically forms 1 bond (due to its need for a duet).

• Oxygen: Typically forms 2 bonds (e.g., in water H2O with two single bonds, or in oxygen gas O2 with one double bond).

• Nitrogen: Typically forms 3 bonds (e.g., in ammonia NH3 with three single bonds, or in nitrogen gas N2 with one triple bond).

• Carbon: Typically forms 4 bonds (always, to satisfy its octet, e.g., in methane CH_4 with four single bonds).

• Boron (B): Typically forms 3 bonds (as it satisfies its electron deficiency with a sextet).

• Halogens (F, Cl, Br, I): Typically form 1 bond when acting as terminal atoms, as they have 7 valence electrons and need only one more to complete an octet. They commonly form single bonds with the central atom.

• Note: While these are typical patterns, atoms can deviate from them (e.g., oxygen can form three bonds in a hydronium ion, H_3O^+), but they will always bond in a way that meets the octet rule (with the established Hydrogen and Boron exceptions) or satisfy a formal charge requirement.

II. Multiple Covalent Bonds

Definition:

• Multiple covalent bonds occur when two atoms need to share more than two electrons (i.e., more than one electron pair) to achieve stable outer-shell electron octets. This involves enhanced sharing of electrons between the bonding atoms.

• These bonds are generally stronger and shorter than single bonds due to increased electron density between the nuclei.

Types of Multiple Covalent Bonds:

• Single Bond: Involves the sharing of one pair of electrons (total of 2 electrons) between two atoms. It is represented by a single line (—).

  • Example: Hydrogen molecule (H—H). This is the longest and weakest type of covalent bond between two given atoms.

• Double Bond: Involves the sharing of two pairs of electrons (total of 4 electrons) between two atoms. It is represented by a double line (=\).

  • Example: Oxygen molecule (O=O). Double bonds are shorter and stronger than single bonds between the same atoms.

• Triple Bond: Involves the sharing of three pairs of electrons (total of 6 electrons) between two atoms. It is represented by a triple line ($\equiv$).

  • Example: Nitrogen molecule (N\equiv N). Triple bonds are the shortest and strongest type of covalent bond between two given atoms.

Common Elements Forming Multiple Bonds:

• Carbon (C): Highly versatile, can form double bonds (e.g., in CO_2 or alkenes) and triple bonds (e.g., in CO or alkynes).

• Nitrogen (N): Frequently forms double bonds (e.g., in NO2^-) and triple bonds (e.g., in N2 or nitriles).

• Oxygen (O): Can only form double bonds (e.g., in O2 or CO2) because forming a triple bond would require it to exceed its typical bonding capacity and result in an unstable electronic configuration.

• Sulfur (S): Also can form double bonds (e.g., in SO_2).

III. Lewis Structures

Definitions:

• Molecular formula: A shorthand notation that shows the exact numbers and types of atoms in one molecule (e.g., H2O, NH3, CH_4). It does not convey structural information.

• Lewis structure: A two-dimensional molecular representation depicting the connections between atoms using lines for shared electron pairs and showing the locations of all non-bonding (lone-pair) valence electrons as dots. It helps visualize electron distribution and bonding.

• Lone pair: Also known as a non-bonding pair, it is a pair of valence electrons that are not involved in covalent bonding and belong exclusively to one atom.

Steps for Drawing Lewis Structures:

1. Decide the bonding pattern (Central Atom):

   • Typically, the least electronegative atom (excluding Hydrogen) is the central atom. Hydrogen can never be a central atom as it only forms one bond.

   • For elements in the same group, the element farther down the group typically goes in the center because it is larger and can accommodate more bonds or an expanded octet.

   • Often, the atom that appears only once in the formula is the central atom (e.g., C in CH4, S in SO2).

2. Calculate Total Valence Electrons:

   • Sum the valence electrons contributed by all atoms in the molecule based on their group number in the periodic table.

   • If the molecule is a cation (has a positive charge, +, +2), subtract that many electrons from the total, as electrons have been lost.

   • If the molecule is an anion (has a negative charge, - ,-2), add that many electrons to the total, as electrons have been gained.

3. Form Single Bonds:

   • Draw a single line (representing 2 shared electrons) between the central atom and each peripheral (surrounding) atom to form the basic skeletal structure.

   • Subtract the electrons used for these single bonds from the total valence electrons calculated in Step 2.

4. Add Lone Pairs to Peripheral Atoms:

   • Distribute the remaining electrons as lone pairs onto each peripheral atom first until each one achieves a full octet (8 electrons, including shared bonding electrons).

   • Remember the exceptions: Hydrogen (H) needs only 2 electrons, and Boron (B) needs only 6 electrons to achieve stability. Halogens (Group 17) typically complete their octet with three lone pairs.

5. Place Remaining Electrons on Central Atom:

   • Any electrons left after satisfying the peripheral atoms' octets in Step 4 are placed as lone pairs on the central atom. This is especially relevant for expanded octets.

6. Form Multiple Bonds (If Necessary):

   • Only if the central atom still lacks a full octet (8 electrons) after all lone pairs have been placed, move one or more lone pairs from surrounding atoms to form double or triple bonds with the central atom.

   • Do this until the central atom achieves an octet. Avoid giving peripheral atoms (especially halogens) more than an octet unless they are in the third period or below and the central atom is already octet-satisfied (which is rare).

Examples for Practice:

• O_2

• Cl_2O

• OCS

• CO_3^{2-}(carbonate ion)

• CF_4

• SO_2

• PCl_3

• CHF_3

IV. Molecular Shapes (VSEPR Model)

Valence-Shell Electron-Pair Repulsion (VSEPR) Model:

• The VSEPR model is a fundamental method for predicting the three-dimensional geometry (shape) of molecules based on the repulsion between electron charge clouds around a central atom.

• It operates on the principle that electron charge clouds (whether bonding pairs or lone pairs) are negatively charged and will orient themselves in space as far away from each other as possible to minimize electrostatic repulsion, thereby achieving the lowest energy and most stable molecular structure.

• Electron charge cloud: Defined as a region of electron density. This includes:

  • A single bond (e.g., C—H)

  • A double bond (e.g., C=O)

  • A triple bond (e.g., C\equiv N)

  • A lone pair of electrons

• Each of these counts as one electron charge cloud or electron domain for VSEPR purposes.

• Valence electrons in bonds and lone pairs exert repulsive forces on one another, which is the primary factor dictating the molecule's specific shape around the central atom.

Steps for Predicting Molecular Shape using VSEPR:

1. Draw Lewis Structure: Accurately draw the Lewis structure of the molecule. Identify the central atom whose geometry is of interest. If there are multiple central atoms, determine the geometry for each independently.

2. Count Electron Charge Clouds: Count the total number of electron charge clouds (also called electron domains or steric number) surrounding the central atom of interest. Remember that single, double, and triple bonds each count as one charge cloud, as does each lone pair of electrons.

3. Determine Electron Geometry: Based on the total number of electron charge clouds, predict the electron geometry. This is the arrangement of all electron charge clouds (bonding pairs and lone pairs) around the central atom.

4. Predict Molecular Shape: Based on the number of bonded atoms and lone pairs within that electron geometry, predict the specific molecular geometry. This describes only the arrangement of the atoms themselves.

   • Important Note: Lone pairs typically exert a greater repulsive force than bonding pairs because their electron density is spread out more broadly around the central atom compared to bonding pairs, which are held between two nuclei. This stronger repulsion by lone pairs often causes slight distortions in bond angles, making them smaller than ideal VSEPR angles.

Molecular Geometry Table (Based on VSEPR):

V. Representing 3D on Paper

Conventions for Wedge/Dash Representations:

To accurately depict the three-dimensional structure of molecules on a two-dimensional surface, specific conventions for drawing bonds are used:

• The central atom and as many other atoms as possible are conventionally drawn as if they lie in the plane of the paper. These bonds are indicated by a straight line (—).

• Atoms or groups of atoms that project in front of the plane of the paper (coming out towards the viewer) are indicated with a thick, solid wedge ($\blacktriangle$). The broad end of the wedge is towards the atom closer to the viewer.

• Atoms or groups of atoms that project behind the plane of the paper (going away from the viewer) are indicated with a hatched wedge ($\triangledown$ dotted or striped). The thick end of the hatched wedge is towards the atom farther from the viewer.

• Example: In methane (CH_4), a central carbon atom in the plane of the paper would be bonded to two hydrogen atoms also in the plane (straight lines), one hydrogen atom coming forward (solid wedge), and one hydrogen atom going backward (hatched wedge), illustrating its tetrahedral geometry.

VI. Assignments

For Friday (9/19):

• Read Chapter 4.5 in your textbook, which covers covalent bonding and molecular geometry in more detail. Complete Reading Questions 4.5.3$$ to test your understanding.

• Begin working diligently on Problem-Set #4 (available as a D2L Quiz), which will be due by Monday. This problem set will provide practice in drawing Lewis structures and predicting molecular shapes.