¿CÓMO dibujar estructuras de LEWIS? 1º parte.
How to Create Lewis Structures for Ionic and Covalent Compounds
Introduction to Lewis Structures
- The video introduces the concept of Lewis structures, which represent valence electrons in atoms. Valence electrons are those in the outermost energy level, farthest from the nucleus.
- Sodium (Na), with one valence electron, is represented by its symbol and a single dot. In contrast, neon (Ne), with eight valence electrons, is depicted with its symbol surrounded by eight dots.
Understanding Groups in the Periodic Table
- Elements in Group 1 have one valence electron; Group 2 has two; Group 3 (Boron group) has three; and so forth up to noble gases that typically have eight.
- An exception is helium (He), which has only two valence electrons but does not form bonds like other noble gases.
Example of Ionic Compound: Potassium Iodide
- The example begins with potassium iodide (KI). Potassium has one valence electron while iodine has seven.
- To form a bond, potassium transfers its single electron to iodine so that iodine can complete its octet. The resulting charges are noted: potassium becomes positive (+1), and iodine becomes negative (-1).
Example of Ionic Compound: Lithium Oxide
- Next, lithium oxide (Li₂O) is discussed. Lithium has one valence electron while oxygen has six.
- Oxygen acts as the central atom since it appears once compared to lithium's two atoms. Each lithium donates an electron to oxygen, allowing it to complete its octet.
Example of Ionic Compound: Aluminum Chloride
- The discussion moves on to aluminum chloride (AlCl₃). Aluminum possesses three valence electrons while each chlorine atom holds seven.
- Aluminum serves as the central atom again due to being singular. Each chlorine takes an electron from aluminum until all chlorines achieve their octets.
Transitioning to Covalent Compounds
- The video shifts focus towards covalent compounds where sharing of electrons occurs rather than transferring them.
Example of Covalent Compound: Diatomic Chlorine
- For diatomic chlorine (Cl₂), both atoms possess seven valence electrons.
- Since there are only two identical atoms, no central atom needs identification; they directly share a pair of electrons forming a single covalent bond.
Ensuring Octet Completion in Covalent Bonds
- It’s emphasized that each chlorine must complete its octet by sharing one pair of electrons between them—resulting in each having access to eight total shared or unshared electrons.
Understanding Lewis Structures: Bromine Monoxide Example
Key Concepts in Lewis Structure Formation
- The formation of a complete octet is essential for stability; in this case, 8 electrons are needed to satisfy the octet rule for both chlorine atoms.
- Oxygen has 6 valence electrons while bromine has 7; oxygen is chosen as the central atom due to its lower electron count and need for bonding.
- A single bond is formed between bromine and oxygen by sharing one electron from each atom, establishing a basic connection.
- To confirm the structure's accuracy, it’s important to check that all atoms fulfill the octet rule; here, oxygen needs two bonds to reach 8 electrons.
- Each bromine contributes enough electrons through bonding (2 from each bond), ensuring that all involved atoms meet their octet requirements.
Steps in Constructing Lewis Structures
- Begin by identifying total valence electrons available from all atoms involved in the molecule.
- Select an appropriate central atom based on electron counts and bonding potential; typically, this will be the least electronegative element.
- Form initial bonds between the central atom and surrounding atoms while ensuring shared electrons contribute towards completing their respective octets.
- Verify that every atom achieves an octet by counting shared and unshared electrons after constructing initial bonds.