Fundamentals of Atomic Bonding and Lewis Structures

Magnetic Forces and the Radius of Atomic Attraction

• The relationship between magnetic intensity and physical separation provides a fundamental metaphor for electrostatic forces in atoms.

• The primary variable affecting the force of attraction is Distance.

  • In atomic chemistry, distance is referred to as the Radius of the atom.

  • Atoms consist of approximately $99\%$ empty space; therefore, the centers (nuclei) of two atoms in an ionic lattice are not literally touching even if they are "bonded."

  • The separation between these centers is determined by the atomic radius. A larger radius results in a greater physical separation between charges, which weakens the force of attraction.

• The Electrostatic Force (also known as the electrostatic attraction) operates similarly to magnetism.

  • A cation (positive ion) can be viewed as a positive magnet.

  • An anion (negative ion) can be viewed as a negative magnet.

  • The strength of the bond between them is determined by the interplay of their radius and their respective charges.

Determinants of Bond Strength: Charge and Radius

• Bond strength is influenced by two main factors:

  • Radius: The larger the atom, the more separate the charges become, leading to a weaker attraction.

  • Charge: A higher charge magnitude (e.g., $+2$ vs. $+1$) results in a stronger "magnetic" pull.

• Comparison of Sodium ($Na^+$) and Magnesium ($Mg^{2+})$:

  • Magnesium has a higher positive charge ($+2$) than Sodium ($+1$). This acts as a "stronger magnet," theoretically creating a stronger bond.

  • Magnesium is also physically smaller in radius than Sodium, further increasing the force of attraction due to the decrease in distance.

• Data Comparison regarding Melting Points (MP):

  • Melting points offer a practical measure of bond strength, as higher melting points typically indicate stronger atomic bonds.

  • Sodium ($Na$) in its metallic form has a melting point of approximately $100\,^{\circ}\text{C}$.

  • Magnesium ($Mg$) in its metallic form has a melting point of approximately $650\,^{\circ}\text{C}$.

  • Note: In specific ionic compounds like Sodium Chloride ($NaCl$) and Magnesium Chloride ($MgCl_2$), the data can be unusual. The MP of $NaCl$ is $801\,^{\circ}\text{C}$ while the MP of $MgCl_2$ is $714\,^{\circ}\text{C}$. This suggests that factors other than just charge and radius, such as the specific internal crystal structure of the lattice, influence melting points.

• Neutral Elements: Noble gases like Argon ($Ar$) are neutral and do not participate in ionic bonding because they lack a net charge to facilitate attraction.

Valence Electrons and the Octet Rule

• Valence Shell: The outermost shell of electrons that participates in chemical reactions.

  • The word "valence" implies "reactive."

  • Mentally modeled like an onion: you cannot interact with the inner layers without first "peeling" the outer (valence) layer.

  • Inner shells are protected and stable; they generally do not involve themselves in interactions unless the atom is broken.

• The Octet Rule: A principle stating that atoms are most stable when they possess eight electrons in their valence shell.

  • This mirrors the electronic configuration of noble gases.

  • Atoms achieve this stability by gaining, losing, or sharing electrons.

  • Hydrogen Exception: Hydrogen does not follow the octet rule; it only requires $2$ electrons (a full outer shell of $1s^2$) to be stable.

• Electron Configuration Examples:

  • Oxygen: $1s^2 2s^2 2p^4$. It has $6$ valence electrons (the $2s$ and $2p$ electrons) and needs $2$ more to reach eight.

  • Chlorine: $1s^2 2s^2 2p^6 3s^2 3p^5$. It has $7$ valence electrons and needs $1$ more.

  • Sodium: Ends in $3s^1$. It is easier to lose one electron to revert to a full shell underneath ($2p^6$).

• Periodic Table Trick: For elements in the main groups ($13-18$), the number of valence electrons is the group number minus $10$ (e.g., Nitrogen in Group $15$ has $5$ valence electrons).

Procedure for Drawing Lewis Structures

1. Place the Atoms: Identify a central atom (usually the single atom of its type or the larger atom further down the periodic table) and arrange surrounding atoms.

2. Count Total Valence Electrons: Sum the valence electrons for all atoms in the molecule to determine the "Lego stack" of available components.

3. Draw Single Bonds: Connect the central atom to each surrounding atom with a single line. Each line (bond) represents a pair of shared electrons ($2$ electrons).

4. Fill Outer Octets: Distribute the remaining electrons as lone pairs to the surrounding atoms to satisfy their octets (or duets for Hydrogen).

5. Place Remaining Electrons on the Central Atom: Any leftover electrons from the initial count must be placed on the central atom as lone pairs.

6. Calculate Formal Charge: If the central atom does not have a formal charge of zero, adjust by moving lone pairs from surrounding atoms into multiple bonds (double or triple bonds).

Examples of Geometric and Lewis Representations

• Ammonia ($NH_3$):

  • Nitrogen: $5$ valence; Hydrogen: $3 \times 1 = 3$. Total: $8$ electrons.

  • Three single bonds to Hydrogen use $6$ electrons, leaving $2$ for a lone pair on Nitrogen.

  • The final structure has Nitrogen with one lone pair and three bonds ($2 + 6 = 8$), satisfying the octet.

• Methane ($CH_4$):

  • Carbon: $4$ valence; Hydrogen: $4 \times 1 = 4$. Total: $8$ electrons.

  • Four single bonds to Hydrogen use all $8$ electrons. Carbon is stable with four bonds.

• Phosphorus Tribromide ($PBr_3$):

  • Phosphorus: $5$ valence; Bromine: $3 \times 7 = 21$. Total: $26$ electrons.

  • Three single bonds use $6$ electrons, leaving $20$.

  • Filling three Bromine octets uses $18$ electrons ($6 \times 3$), leaving $2$ electrons as a lone pair on Phosphorus.

• Thionyl Chloride ($SOCl_2$):

  • Sulfur: $6$ valence; Oxygen: $6$ valence; Chlorine: $2 \times 7 = 14$. Total: $26$ electrons.

Formal Charge and Expanded Octets

• Formal Charge Calculation:

  • $\text{Formal Charge} = \text{Valence Electrons} - (\text{Bonds} + \text{Unbonded Electrons})$

  • A stable Lewis structure minimizes the formal charge of each atom to as close to zero as possible.

• Expanded Octet: Elements in Period $3$ and below (e.g., Sulfur, Phosphorus) can hold more than $8$ electrons in their valence shell.

  • This is possible because they have access to the $3d$ orbital subshell, which can accommodate extra electrons.

  • Example: In $SOCl_2$, Sulfur may form a double bond with Oxygen (calculated via formal charge) resulting in Sulfur having $10$ valence electrons. This is acceptable for Sulfur because it is a Period $3$ element.

• Three-Dimensional Molecular Symbols:

  • Normal Line: The bond is in the same plane as the page/screen.

  • Shaded Wedge: The bond is "out of the page," pointing toward the viewer.

  • Dashed Line (or thin lines): The bond is "into the page," pointing away from the viewer.

Questions & Discussion

• Student: "The melting point of sodium chloride is $801\,^{\circ}\text{C}$ while magnesium chloride is $714\,^{\circ}\text{C}$. So magnesium chloride is actually less strong?"

• Teacher: "That's interesting… that's a bit unusual. This might just be a special case. Structure might also have other things to do with it as well. We can look at their normal melting points in metallic form where Sodium is about $100\,^{\circ}\text{C}$ and Magnesium is about $650\,^{\circ}\text{C}$. There, Magnesium is much stronger because of the higher charge and smaller radius."

• Student: "What does valence mean in English?"

• Teacher: "Valence means reactive. It means the shell of electrons that's gonna get involved in reactions."

• Student: "Which one do I put in the middle? In $SOCl_2$?"

• Teacher: "You put the atom that is lower in the periodic table or the single atom. In this case, Sulfur. Larger atoms usually go in the center because they are able to form more bonds."

• Student: "What is ATP?"

• Teacher: "ATP is like the battery between your cells. It is the energy-providing protein. Glucose is the primary sugar in metabolism used to generate ATP."