Lewis Model Chemistry

[Favio Cassidy]

Lewisstructures show each atom and its position in the structure of the molecule using its chemical symbol. Lines are drawn between atoms that are bonded to one another (pairs of dots can be used instead of lines). Excess electrons that form lone pairs are represented as pairs of dots, and are placed next to the atoms.

Although main group elements of the second period and beyond usually react by gaining, losing, or sharing electrons until they have achieved a valence shell electron configuration with a full octet of (8) electrons, hydrogen (H) can only form bonds which share just two electrons.

The total number of electrons represented in a Lewis structure is equal to the sum of the numbers of valence electrons on each individual atom. Non-valence electrons are not represented in Lewis structures.

Lewis structures for polyatomic ions may be drawn by the same method. When counting electrons, negative ions should have extra electrons placed in their Lewis structures; positive ions should have fewer electrons than an uncharged molecule. When the Lewis structure of an ion is written, the entire structure is placed in brackets, and the charge is written as a superscript on the upper right, outside the brackets.

A simpler method has been proposed for constructing Lewis structures, eliminating the need for electron counting: the atoms are drawn showing the valence electrons; bonds are then formed by pairing up valence electrons of the atoms involved in the bond-making process, and anions and cations are formed by adding or removing electrons to/from the appropriate atoms.[5]

A trick is to count up valence electrons, then count up the number of electrons needed to complete the octet rule (or with hydrogen just 2 electrons), then take the difference of these two numbers. The answer is the number of electrons that make up the bonds. The rest of the electrons just go to fill all the other atoms' octets.

In the column titled "Molecular Formula," write each individual atom of the molecule in individual rows. Then, in the column titled "Octect electrons," write the number of electrons each atom requires to achieve an octect (this will be 8 for every element except for Hydrogen, which can only hold 2 valence electrons). Finally, in the column titled "Total Valence electrons," write the number of valence electrons each atom has when unbonded. This information comes from the periodic table. For Group 1-8 elements (everything excluding transition metals, lanthanides, and actinides), the number of valence electrons is equal to their Group number. Now that the table has been made, calculating number of bonds and lone pairs is possible.

In terms of Lewis structures, formal charge is used in the description, comparison, and assessment of likely topological and resonance structures[7] by determining the apparent electronic charge of each atom within, based upon its electron dot structure, assuming exclusive covalency or non-polar bonding. It has uses in determining possible electron re-configuration when referring to reaction mechanisms, and often results in the same sign as the partial charge of the atom, with exceptions. In general, the formal charge of an atom can be calculated using the following formula, assuming non-standard definitions for the markup used:

The formal charge of an atom is computed as the difference between the number of valence electrons that a neutral atom would have and the number of electrons that belong to it in the Lewis structure. Electrons in covalent bonds are split equally between the atoms involved in the bond. The total of the formal charges on an ion should be equal to the charge on the ion, and the total of the formal charges on a neutral molecule should be equal to zero.

For some molecules and ions, it is difficult to determine which lone pairs should be moved to form double or triple bonds, and two or more different resonance structures may be written for the same molecule or ion. In such cases it is usual to write all of them with two-way arrows in between .mw-parser-output div.crossreferencepadding-left:0(see Example below). This is sometimes the case when multiple atoms of the same type surround the central atom, and is especially common for polyatomic ions.

When this situation occurs, the molecule's Lewis structure is said to be a resonance structure, and the molecule exists as a resonance hybrid. Each of the different possibilities is superimposed on the others, and the molecule is considered to have a Lewis structure equivalent to some combination of these states.

When comparing resonance structures for the same molecule, usually those with the fewest formal charges contribute more to the overall resonance hybrid. When formal charges are necessary, resonance structures that have negative charges on the more electronegative elements and positive charges on the less electronegative elements are favored.

Single bonds can also be moved in the same way to create resonance structures for hypervalent molecules such as sulfur hexafluoride, which is the correct description according to quantum chemical calculations instead of the common expanded octet model.

Despite their simplicity and development in the early twentieth century, when understanding of chemical bonding was still rudimentary, Lewis structures capture many of the key features of the electronic structure of a range of molecular systems, including those of relevance to chemical reactivity. Thus, they continue to enjoy widespread use by chemists and chemistry educators. This is especially true in the field of organic chemistry, where the traditional valence-bond model of bonding still dominates, and mechanisms are often understood in terms of curve-arrow notation superimposed upon skeletal formulae, which are shorthand versions of Lewis structures. Due to the greater variety of bonding schemes encountered in inorganic and organometallic chemistry, many of the molecules encountered require the use of fully delocalized molecular orbitals to adequately describe their bonding, making Lewis structures comparatively less important (although they are still common).

There are simple and archetypal molecular systems for which a Lewis description, at least in unmodified form, is misleading or inaccurate. Notably, the naive drawing of Lewis structures for molecules known experimentally to contain unpaired electrons (e.g., O2, NO, and ClO2) leads to incorrect inferences of bond orders, bond lengths, and/or magnetic properties. A simple Lewis model also does not account for the phenomenon of aromaticity. For instance, Lewis structures do not offer an explanation for why cyclic C6H6 (benzene) experiences special stabilization beyond normal delocalization effects, while C4H4 (cyclobutadiene) actually experiences a special destabilization.[citation needed] Molecular orbital theory provides the most straightforward explanation for these phenomena.[original research?]

After 29 years of teaching high school chemistry and teaching Lewis Structure, I am curious to see how this helps my students next year. They want to do everything on computers these days instead of building the models.

Typically most college curricula include three acid base models: Arrhenius', Bronsted-Lowry's, and Lewis'. Although Lewis' acid base model is generally thought to be the most sophisticated among these three models, and can be further applied in reaction mechanisms, most general chemistry curricula either do not include Lewis' acid base model, or quickly mention it at the end of the acid base chapter, because of the concern that Lewis' model may confuse general chemistry students (Shaffer 2006). While such a disconnection in curriculum might put students to disadvantage as they try to construct solid and coherent acid base mental models, there has not been any research data to favor one curriculum over another. The large sizes of general chemistry courses at most universities (from one hundred to several hundred students per lecture section) pose further challenges to the comparison of different general chemistry curricula on their effectiveness in helping students construct acid base mental models. In light of these challenges, the research questions I focused on were: 1) What are the important characteristics of activities that effectively promote and retain argumentation skills among college students? 2) In what ways is argumentation an effective assessment method for student understanding of acid base models? 3) How do different curricula affect students' acid base models? This dissertation presents promising results from using BeSocratic activities in promoting argumentation skills among college students and at the same time using their responses in the activities to understand aspects of their acid base mental models, and compare how two different general chemistry curricula affected students' acid base mental models.

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It has been accepted for decades now that students entering college present significant deficiencies in their conceptual understanding of chemistry and their ability to express themselves symbolically (Kozma & Russell, 1997). Recent evidence confirms the difficulties experienced by students to simply construct representations, let alone use them effectively (Grove, Cooper, & Rush, 2012). To complicate the matter, current evidence suggests advanced undergraduates and graduate students alike are not necessarily representationally competent (Cooper, Grove, Underwood, & Klymkowsky, 2010; Strickland, Kraft, & Bhattacharyya, 2010). Less than satisfactory learning is product of ineffective learning experiences, thus we have become interested in investigating the challenges instructors encounter when they use symbolic language to teach chemistry.

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