Lewis structures
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Introduction to Lewis Structures
Lewis structures are fundamental diagrams in chemistry that illustrate the bonding between atoms and the arrangement of valence electrons in molecules. Named after Gilbert N. Lewis, these structures provide a visual representation of chemical bonding and electron distribution. The introduction video serves as an essential starting point for understanding this crucial concept in chemistry. Lewis structures depict atoms as their elemental symbols surrounded by dots representing valence electrons. These diagrams are invaluable tools for chemists, helping to predict molecular geometry, reactivity, and bonding patterns. By showing how electrons are shared or transferred between atoms, Lewis structures offer insights into the formation of covalent and ionic bonds. They also aid in understanding concepts like resonance structures, formal charges, and octet rule exceptions. Mastering Lewis structures is vital for students and professionals alike, as they form the foundation for more advanced topics in organic and inorganic chemistry.
Drawing Lewis Structures for Covalent Compounds
Lewis structures are essential diagrams in chemistry that illustrate the bonding between atoms in covalent compounds and show the arrangement of valence electrons. Learning how to draw Lewis structures is crucial for understanding molecular geometry and chemical bonding. This guide will walk you through the step-by-step process of creating Lewis structures, using H2O (water) and CO2 (carbon dioxide) as examples.
Step 1: Identifying the Central Atom
The first step in drawing a Lewis structure is to identify the central atom. Generally, the least electronegative atom that isn't hydrogen becomes the central atom. In H2O, oxygen is the central atom, while in CO2, carbon is central.
Step 2: Arranging Atoms
Once you've identified the central atom, arrange the other atoms around it. For H2O, place the two hydrogen atoms on either side of the oxygen. For CO2, position the oxygen atoms on both sides of the carbon.
Step 3: Counting Valence Electrons
Next, count the total number of valence electrons in the molecule. For H2O, oxygen contributes 6 valence electrons, and each hydrogen contributes 1, totaling 8 electrons. In CO2, carbon has 4 valence electrons, and each oxygen has 6, summing to 16 electrons.
Step 4: Placing Electron Pairs
Now, start placing electron pairs around the atoms. Begin by forming single bonds between the central atom and outer atoms. For H2O, place one electron pair between oxygen and each hydrogen. For CO2, put one electron pair between carbon and each oxygen. Then, distribute the remaining electrons as lone pairs, prioritizing outer atoms to satisfy the octet rule.
Step 5: Checking the Octet Rule
The octet rule states that atoms tend to have eight valence electrons in their outer shell (except for hydrogen, which needs two). Check if all atoms in your structure follow this rule. In H2O, oxygen has two bonding pairs and two lone pairs, totaling 8 electrons. Each hydrogen has 2 electrons. For CO2, ensure carbon and both oxygens have 8 electrons each.
Single, Double, and Triple Bonds
Sometimes, single bonds aren't enough to satisfy the octet rule. In these cases, you'll need to form double or triple bonds. Let's use N2 (nitrogen gas) as an example to explain this concept.
When drawing the Lewis structure for N2, start by placing the two nitrogen atoms side by side. Each nitrogen has 5 valence electrons, totaling 10 for the molecule. If we use two electrons to form a single bond between the nitrogens, we're left with 8 electrons to distribute. Placing these as lone pairs still leaves each nitrogen with only 6 electrons, violating the octet rule.
To solve this, we need to create a triple bond between the nitrogen atoms. This uses 6 electrons for bonding, leaving 4 electrons to be distributed as lone pairs (two on each nitrogen). The result is a structure where each nitrogen has 8 electrons, satisfying the octet rule.
Applying the Process to H2O and CO2
For H2O, the final Lewis structure shows oxygen with two single bonds to hydrogen atoms and two lone pairs. This gives oxygen 8 electrons (octet fulfilled) and each hydrogen 2 electrons.
In the case of CO2, after initial placement, we find that carbon only has 4 electrons (violating the octet rule) while each oxygen has 7. To fix this, we need to form double bonds between carbon and each oxygen. The final structure shows O=C=O, where carbon shares 4 electrons with each oxygen, giving all atoms 8 electrons and satisfying the octet rule.
Practice and Common Mistakes
Drawing Lewis structures takes practice. Common mistakes include forgetting to count all valence electrons, misplacing lone pairs, or not forming the correct number of bonds. Understanding molecular geometry can also help in visualizing and verifying your Lewis structures.
Lewis Structures for Ionic Compounds
Drawing Lewis structures for ionic compounds is a fundamental skill in chemistry that helps visualize the electron arrangement and bonding between atoms. Unlike covalent compounds, ionic compounds involve the transfer of electrons between atoms, resulting in oppositely charged ions. Let's explore this process using sodium chloride (NaCl) as our primary example.
To draw the Lewis structure for NaCl, we first need to understand the electron configuration of each atom. Sodium (Na) has one valence electron in its outermost shell, while chlorine (Cl) has seven valence electrons. In ionic bonding, the sodium atom tends to lose its single valence electron, becoming a positively charged sodium ion (Na+). Simultaneously, the chlorine atom gains this electron, becoming a negatively charged chloride ion (Cl-).
The Lewis structure for NaCl would show the sodium ion as a bare atom with no valence electrons, represented by [Na]+. The chloride ion would be shown with a complete octet of eight electrons, represented as [:Cl:]-. The square brackets indicate the ionic nature of the compound, and the superscript shows the charge on each ion.
This electron transfer in ionic bonding differs significantly from covalent bonding. In covalent bonds, atoms share electrons to achieve stable electron configurations. However, in ionic bonds, electrons are completely transferred from one atom to another. This transfer creates strong electrostatic attractions between the oppositely charged ions, forming the ionic bond.
When drawing Lewis structures for ionic compounds, it's crucial to remember that the goal is to achieve a stable electron configuration for each ion. For most main group elements, this means attaining an octet (eight valence electrons) or, in the case of hydrogen and helium, a duet (two electrons).
Polyatomic Ions and Their Lewis Structures
Polyatomic ions add another layer of complexity to Lewis structures for ionic compounds. These are ions composed of multiple atoms bonded together, carrying an overall charge. Examples include the ammonium ion (NH4+) and the carbonate ion (CO3^2-).
When drawing Lewis structures for polyatomic ions, we follow similar principles as for covalent molecules, but we must account for the overall charge. Let's consider the carbonate ion (CO3^2-) as an example:
- Count the total number of valence electrons, including those contributed by the negative charge.
- Arrange the atoms, typically with the central atom (usually the least electronegative) in the middle.
- Connect the atoms with single bonds.
- Distribute the remaining electrons as lone pairs to satisfy the octet rule for each atom.
- Use double bonds if necessary to achieve octets for all atoms.
For CO3^2-, the carbon atom is central, bonded to three oxygen atoms. The structure includes two double bonds and one single bond to the oxygen atoms, with the remaining electrons distributed as lone pairs on the oxygen atoms to complete their octets.
Understanding how to draw Lewis structures for ionic compounds, including those with polyatomic ions, is crucial for visualizing electron arrangements and predicting chemical properties. Whether you're working with simple binary ionic compounds like NaCl or more complex structures involving polyatomic ions, the principles of electron transfer and achieving stable electron configurations remain key to accurately representing these compounds in Lewis dot structures.
Exceptions to the Octet Rule
The octet rule is a fundamental concept in chemistry that states atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with eight valence electrons. However, there are notable exceptions to this rule that challenge our understanding of chemical bonding. Two prime examples of compounds that don't follow the octet rule are boron trifluoride (BF3) and nitric oxide (NO).
Boron trifluoride (BF3) is a classic example of an electron-deficient compound. In its Lewis structure, the central boron atom has only six valence electrons, falling short of the octet. This occurs because boron has only three valence electrons to share, and each fluorine atom contributes one electron to the bonding. The result is a trigonal planar molecule with an incomplete octet for boron. This electron deficiency makes BF3 a strong Lewis acid, readily accepting electron pairs from other molecules.
Nitric oxide (NO) presents another intriguing exception to the octet rule. In its most stable form, NO has an unpaired electron, resulting in an odd number of valence electrons. This unpaired electron gives NO paramagnetic properties and makes it highly reactive. The nitrogen atom in NO has seven valence electrons, while oxygen has seven, totaling 14 electrons an arrangement that doesn't conform to the octet rule for either atom.
These exceptions occur due to various factors, including atomic size, electronegativity differences, and unique electronic configurations. Understanding these exceptions is crucial for grasping the complexity of chemical bonding beyond simple octet-based models. They demonstrate that while the octet rule is a useful guideline, it doesn't apply universally to all compounds.
Expanded Octets
Another category of exceptions involves compounds with expanded octets. Elements in the third period and beyond can accommodate more than eight electrons in their valence shell. Phosphorus pentachloride (PCl5) and sulfur hexafluoride (SF6) are prime examples. In PCl5, the central phosphorus atom is surrounded by five chlorine atoms, resulting in 10 valence electrons around phosphorus. SF6 takes this further, with sulfur bonding to six fluorine atoms, giving it 12 valence electrons. These expanded octets are possible due to the availability of d orbitals in these larger atoms, allowing for additional bonding capacity.
Electron-Deficient Compounds
Electron-deficient compounds, like BF3 mentioned earlier, form another important class of octet rule exceptions. Beryllium compounds, such as BeCl2, also fall into this category. In the Lewis structure for BeCl2, the central beryllium atom has only four valence electrons, significantly less than an octet. This electron deficiency results in a linear molecule with strong Lewis acid properties. Aluminum chloride (AlCl3) is another example, behaving similarly to BF3 in its electron-deficient nature.
The significance of these exceptions lies in their impact on chemical reactivity and bonding theories. They challenge simplistic models of chemical bonding and highlight the need for more sophisticated explanations of molecular structure and behavior. For instance, the electron deficiency in BF3 and BeCl2 explains their high reactivity and their tendency to form adducts with electron-rich species. The unpaired electron in NO accounts for its important role in biological processes and atmospheric chemistry.
Understanding these exceptions is crucial for predicting molecular geometry, reactivity, and properties of compounds. It also paves the way for advanced concepts in inorganic and organometallic chemistry, where compounds frequently deviate from the octet rule. Moreover, these exceptions have practical implications in fields ranging from materials science to pharmaceutical development, where the unique properties of these compounds are exploited for various applications.
In conclusion, while the octet rule remains a valuable tool for understanding basic chemical bonding, the exceptions to this rule provide a deeper insight into the complexity and diversity of molecular structures. From electron-deficient compounds to those with expanded octets, these exceptions enrich our understanding of chemistry and continue to inspire research into the fundamental nature of chemical bonds.
Applications of Lewis Structures
Lewis structures are fundamental tools in chemistry that provide crucial insights into molecular geometry prediction, chemical reactivity, and polarity. These diagrams, developed by Gilbert N. Lewis in 1916, serve as a cornerstone for predicting and understanding various chemical properties and reactions. By representing atoms and their bonds using dots and lines, Lewis structures offer a visual representation of how electrons are distributed within molecules.
One of the primary applications of Lewis structures is in predicting molecular geometry prediction. The arrangement of electron pairs around a central atom, as depicted in a Lewis structure, directly correlates with the three-dimensional shape of a molecule. This concept, known as VSEPR (Valence Shell Electron Pair Repulsion) theory, allows chemists to anticipate the spatial orientation of atoms in a compound. For instance, the Lewis structure of water (H2O) shows two lone pairs on the oxygen atom, leading to its bent shape, which is crucial for understanding its unique properties like high boiling point and surface tension.
Lewis structures also play a vital role in determining molecular polarity. By illustrating the distribution of electrons, these diagrams help identify polar and nonpolar molecules. This information is essential in predicting intermolecular forces, solubility, and other physical properties. For example, the Lewis structure of carbon dioxide (CO2) shows a linear arrangement with equal sharing of electrons, indicating a nonpolar molecule, which explains its gaseous nature at room temperature.
In terms of chemical reactivity, Lewis structures are invaluable for understanding and predicting reaction mechanisms in organic chemistry. They help identify nucleophilic and electrophilic sites within molecules, which is crucial for organic chemistry reactions. For instance, in the synthesis of pharmaceuticals, Lewis structures guide chemists in designing reaction pathways and predicting potential side products.
Real-world applications of Lewis structures are numerous and diverse. In the field of materials science, these diagrams aid in the development of new polymers by predicting how monomers will interact and bond. In environmental chemistry, Lewis structures help explain the behavior of pollutants in different ecosystems. For example, understanding the Lewis structure of chlorofluorocarbons (CFCs) was crucial in recognizing their potential to deplete the ozone layer.
In biochemistry, Lewis structures are essential for understanding enzyme-substrate interactions and drug design. They help predict how molecules will bind to receptor sites, which is fundamental in developing new medications. Additionally, in the field of nanotechnology, Lewis structures guide the design of molecular machines and sensors by providing insights into how atoms can be arranged to create specific functionalities.
In conclusion, Lewis structures are powerful tools that bridge the gap between atomic structure and macroscopic properties. Their ability to predict molecular geometry prediction, polarity, and reactivity makes them indispensable in various fields of chemistry and beyond. From developing new materials to understanding biological processes, Lewis structures continue to be a cornerstone of chemical understanding and innovation.
Advanced Lewis Structures: Resonance and Formal Charge
Lewis structures are fundamental tools in chemistry for representing molecular structures and bonding. However, some molecules and ions require more advanced concepts to accurately depict their electronic arrangements. Two crucial concepts in this realm are resonance structures and formal charge. Understanding these ideas is essential for grasping the true nature of certain chemical species and their behavior.
Resonance structures is a concept used when a single Lewis structure cannot adequately represent a molecule or ion. In such cases, multiple Lewis structures, called resonance structures, are drawn to describe the overall electronic structure. These structures differ only in the placement of electrons, not in the positions of atoms. The actual structure is considered a hybrid of all possible resonance forms, with electrons delocalized across the molecule.
To draw resonance structures, follow these steps:
- Draw the primary Lewis structure
- Identify movable electron pairs or lone pairs
- Shift these electrons to create new structures
- Ensure all structures obey the octet rule (when possible)
- Use double-headed arrows to show the relationship between structures
A classic example of resonance is the carbonate ion (CO32-). In this ion, three equivalent resonance structures can be drawn, each showing a double bond to a different oxygen atom. The actual structure is a hybrid of these three forms, with the negative charge and double bond character distributed equally among all three oxygen atoms.
Formal charge is another important concept in Lewis structures. It helps determine the most likely structure when multiple possibilities exist and can indicate reactive sites in molecules. Formal charge is calculated for each atom in a Lewis structure and represents the difference between the number of valence electrons in the free atom and the number of electrons assigned to that atom in the Lewis structure.
To calculate formal charge, use the following formula:
Formal Charge = [# of valence electrons] - [# of non-bonding electrons] - (1/2)[# of bonding electrons]
When drawing Lewis structures, the goal is to minimize formal charges and place any remaining charges on the most electronegative atoms.
The ozone molecule (O3) serves as an excellent example to illustrate both resonance and formal charge. Ozone has two resonance structures, each with a single bond and a double bond between oxygen atoms. Calculating formal charges helps explain why these two structures are favored over other possible arrangements.
In the first resonance structure of ozone:
- The central O has a formal charge of +1
- The O with the double bond has a formal charge of 0
- The O with the single bond and lone pairs has a formal charge of -1
The second resonance structure simply switches the positions of the single and double bonds, maintaining the same formal charges. This equal distribution of charge contributes to ozone's stability and reactivity.
Understanding resonance and formal charge is crucial for accurately representing molecules and ions, predicting their properties, and explaining their reactivity. These concepts build upon basic Lewis structures to provide a more comprehensive view of molecular bonding and electron distribution. By mastering these ideas, chemists can better interpret and predict the behavior of complex chemical species in various reactions and environments.
Conclusion
Lewis structures are fundamental tools in understanding chemical bonding and molecular geometry. They provide a visual representation of valence electrons and how atoms share or transfer them to form bonds. These structures are crucial for predicting molecular geometry, reactivity, and properties. The introduction video serves as an essential foundation for grasping these concepts, offering clear explanations and examples. However, it's important to recognize the limitations of Lewis structures, such as their inability to accurately depict some complex molecules or account for certain bonding phenomena. To truly master this skill, regular practice in drawing Lewis structures is key. Start with simple molecules and gradually progress to more complex ones. By doing so, you'll develop a deeper understanding of chemical bonding and molecular behavior. Remember, Lewis structures are just the beginning of your journey in chemistry, opening doors to more advanced concepts and theories in the field.
Introduction to Dot and Cross Diagrams
Lewis structures: introduction.
Step 1: Understanding Lewis Structures
Lewis structures, also known as Lewis dot diagrams or electron dot structures, are a way to represent the valence electrons of atoms within a molecule. These diagrams help predict the bonding and structure of a substance, whether it is ionic or covalent. The primary goal is to visualize how atoms share or transfer electrons to achieve stable electron configurations.
Step 2: The Octet Rule
The octet rule is a fundamental concept in chemistry that states atoms tend to form bonds until they are surrounded by eight valence electrons. This rule is derived from the observation that atoms with eight electrons in their outer shell are particularly stable. In Lewis structures, this rule helps determine how atoms will bond and what kind of bonds (single, double, or triple) will form.
Step 3: Constructing Lewis Diagrams
To construct a Lewis diagram, follow these steps:
- Identify the total number of valence electrons in the molecule.
- Write the symbols for the atoms involved and connect them with single bonds initially.
- Distribute the remaining electrons to satisfy the octet rule for each atom, starting with the outer atoms and moving inward.
- If necessary, form double or triple bonds to ensure all atoms achieve an octet configuration.
Step 4: Differentiating Ionic and Covalent Bonds
Lewis structures can illustrate the difference between ionic and covalent bonds. In ionic bonding, electrons are transferred from one atom to another, resulting in the formation of positive and negative ions. In covalent bonding, electrons are shared between atoms, creating a bond through the overlap of electron clouds. The diagrams will show either the transfer or sharing of electrons, helping to visualize the type of bond formed.
Step 5: Exceptions to the Octet Rule
While the octet rule is a useful guideline, there are exceptions. Some molecules have atoms that do not follow the octet rule, either having fewer or more than eight electrons in their valence shell. These exceptions are significant because they highlight the limitations of the octet rule and the need for more advanced theories to explain bonding in certain compounds.
Step 6: Practical Applications of Lewis Structures
Lewis structures are part of valence bond theory, which helps predict the existence, stability, and type of bonds in molecules. For example, they can explain why water is H2O and not H3O, or why nitrogen forms N2 molecules. By constructing these diagrams, chemists can predict molecular shapes, bond angles, and reactivity, making them a valuable tool in both theoretical and practical chemistry.
Step 7: Drawing Lewis Structures for Covalent and Ionic Compounds
To draw Lewis structures for covalent compounds, start by identifying the central atom (usually the least electronegative) and arrange the other atoms around it. For ionic compounds, show the transfer of electrons from the metal to the non-metal, resulting in the formation of ions. Use brackets to indicate the charge on each ion. These visual representations help in understanding the nature of the bonds and the overall structure of the compound.
Step 8: Predicting Molecular Stability
By using Lewis structures, chemists can predict whether a molecule will be stable or not. If a proposed structure does not satisfy the octet rule or results in an unstable arrangement of electrons, it is likely that the molecule will not exist in nature. This predictive power is one of the key benefits of mastering Lewis structures.
FAQs
Here are some frequently asked questions about Lewis structures:
1. How do you draw Lewis structures step by step?
To draw Lewis structures, follow these steps: 1) Count the total valence electrons. 2) Identify the central atom. 3) Connect atoms with single bonds. 4) Place remaining electrons as lone pairs. 5) Check for octets and adjust bonds if necessary.
2. What is the Lewis structure for H2O?
The Lewis structure for H2O (water) shows the oxygen atom in the center with two single bonds to hydrogen atoms. The oxygen also has two lone pairs, giving it a total of 8 electrons (octet). The structure resembles a bent shape.
3. How do you determine formal charges in Lewis structures?
Calculate formal charge using the formula: FC = [# valence e-] - [# non-bonding e-] - ½[# bonding e-]. Assign formal charges to atoms and aim for structures with the lowest possible charges, ideally zero.
4. What are resonance structures?
Resonance structures are multiple valid Lewis structures for a molecule where only the electron positions differ. The actual structure is a hybrid of these forms. They're common in molecules with delocalized electrons, like benzene or carbonate ion.
5. How do Lewis structures help predict molecular geometry?
Lewis structures show the arrangement of electron pairs around atoms, which directly relates to molecular geometry through VSEPR theory. The number of electron domains (bonding and non-bonding) determines the overall shape of the molecule.
Prerequisite Topics for Understanding Lewis Structures
Before diving into the intricacies of Lewis structures, it's crucial to have a solid foundation in related chemical concepts. Two key prerequisite topics that significantly enhance your understanding of Lewis structures are molecular geometry and VSEPR theory, as well as conjugation and resonance structures.
Understanding molecular geometry prediction is essential when working with Lewis structures. This concept helps you visualize the three-dimensional arrangement of atoms in a molecule, which is directly related to how electrons are distributed in Lewis structures. The VSEPR (Valence Shell Electron Pair Repulsion) theory, a key component of molecular geometry, explains how electron pairs around a central atom influence the shape of a molecule. This knowledge is invaluable when drawing Lewis structures, as it helps you predict and represent the correct arrangement of atoms and bonds.
Similarly, familiarity with resonance structures is crucial for a comprehensive understanding of Lewis structures. Resonance occurs when a single Lewis structure cannot adequately represent the electron distribution in a molecule. By studying conjugation and resonance, you'll learn how to draw multiple Lewis structures for a single molecule and understand how these structures contribute to the overall stability and properties of the compound.
The interplay between Lewis structures and these prerequisite topics is significant. When you draw a Lewis structure, you're essentially creating a two-dimensional representation of a molecule's electron distribution. Your knowledge of molecular geometry allows you to mentally translate this 2D structure into a 3D model, giving you a more accurate picture of the molecule's true form. Additionally, understanding resonance helps you recognize when a single Lewis structure is insufficient and how to represent more complex electron distributions.
Moreover, these prerequisite topics enhance your ability to predict molecular properties based on Lewis structures. For instance, your understanding of molecular geometry can help you deduce polarity and intermolecular forces, while knowledge of resonance structures aids in predicting reactivity and stability. This interconnectedness demonstrates why a strong grasp of these foundational concepts is vital for mastering Lewis structures and, by extension, more advanced topics in chemistry.
In conclusion, investing time in understanding molecular geometry and resonance structures will significantly enhance your ability to work with Lewis structures. These prerequisite topics provide the necessary context and tools to not only draw accurate Lewis structures but also to interpret and apply them in various chemical scenarios. By building this strong foundation, you'll be better equipped to tackle more complex chemical concepts and problem-solving in your studies and beyond.
- To construct Lewis (dot and cross) structures using knowledge of electrons shells and the octet rule.
- To use Lewis diagrams to explain the difference between ionic and covalent bonding.
- Some of the exceptions to the octet rule and the significance of this to theories of bonding.
Notes:
- In C11.4.2: Ionic and covalent bonding we saw the two main types of chemical bond: ionic and covalent bonding. These are attractive forces that hold individual atoms together allowing them to form molecules and larger, more complex structures.
We can obviously see that most atoms are just more stable existing in compounds or molecules, such as CH4 or O2, compared to being single isolated atoms, like H or O.
In C11.4.1: Introduction to bonding we mentioned that theories of bonding exist to try to explain why this is true. One theory of bonding is valence bond theory, which uses Lewis structures to show the bonding in simple molecules. - In Lewis structures, dots and crosses are used to show the valence electrons, where dots or crosses are used for the central atom and the other is used for the atom(s) it is bonding to. They can be used to show covalent and ionic bonding between atoms in molecules and compounds, where electrons are shared or donated/accepted to achieve a full outer shell. Valence means outer shell or highest energy electrons; dont draw core electrons in Lewis structures!
- When drawing these diagrams, look at our lesson Electronic structure: 288 rule for help when finding the number of outer shell electrons. Remember the octet rule, where electrons will try and fill their outer shell with an octet of (eight) electrons.
- Lewis structures for covalent compounds can be drawn step-by-step as below, with CO and HO): as examples:
- Step 1: Write the chemical symbol for the central atom in the molecule and draw a ring around it (this will be the outer electron shell). If the molecule is diatomic, draw either one atom.
- Step 2: Write the symbols of the other atoms equally spaced around it. Draw a ring around each of these so that their rings each have an overlap into the central ring. These rings are the outer electron shells.
- Step 3: For the non-central atoms, fill in their outer shells of electrons using the 288 rule fromElectronic structure: 288 rule (288 rule). Remember to fill up to four electrons individually then start pairing them up. Mark all these electrons using either a dot or a cross. Be consistent with the dots and crosses; if you use crosses for one H atom, use it for all H atoms.
- Step 4: LOOK CAREFULLY AT YOUR DIAGRAM. Now make sure any unpaired electrons are placed in the overlap with the central atom.
This is a clue to show how many pairs of electrons (how many covalent bonds) there will be with the central atom. Each of these will be paired with an electron from the central atom too. - Step 5: Now fill the outer shell of electrons in the central atom. Use the other sign to the one you used with the non-central atom. Start by making pairs of dots and crosses with the electrons of the non-central atoms in the overlapping area. You can only pair dots and crosses in a 1:1 ratio. Now, any extra electrons can be placed outside the overlaps on the outer shell. Remember that if there are more than four electrons in the shell they must now be paired. Electrons in the overlapping area count for both overlapping atoms!
- Step 6: Look carefully at your diagram,. If you have followed these steps, all of your atoms should have exactly 8 electrons (2 if it is hydrogen) in the outer shell. Remember that electrons in the overlaps count for both atoms.
- Dot and cross diagrams for ionic compounds can be drawn like this:
- Step 1: Write the chemical symbols of the atoms side by side and draw a ring (the outer shell) around them but keep them separate. They must not overlap like in covalent compounds.
- Step 2: Fill the outer shell with the correct number of electrons for each atom.
- Step 3: Move the electrons from the metal atom(s) to the non-metal atom(s). This should complete the outer shell of the non-metal atom(s) while leaving empty the outer shell of the metal atom(s). The atoms are now ions.
- Step 4: Ions should have square brackets surrounding them with their charge on the top-right corner. This shows an ion has been formed. In order to show an electron was transferred and not shared, keep the original dot/cross signage you used. See below, where the Na outer electron is still a dot on Cl.
- These Lewis diagrams should help to show the difference between bonding in covalent and ionic bonding in terms of electrons:
- Covalent bonding involves a sharing of electrons to complete outer shells in both atoms.
- Ionic bonding involves a discrete transfer of electrons usually from a metal atom to a non-metal atom, which also leads to complete outer shells.
- Building Lewis structures using the octet rule is applying a theory of chemical bonding. These theories are made to try and explain what we observe and to make predictions that can be tested, but there are exceptions to the octet rule.
- Nitric oxide (NO) would have a Lewis structure beginning with six electrons around oxygen, and five electrons around nitrogen. This means nitrogen needs three more electrons and oxygen two to satisfy the octet rule. This cannot be satisfied by:
- A single bond (one pair of electrons) between the two atoms; it leaves only six electrons on nitrogen and seven on oxygen.
- A double bond (or two pairs of electrons) between the two atoms; it leaves only seven electrons on nitrogen and eight on oxygen.
- A triple bond between the two atoms; oxygen now has nine electrons in the outer shell, violating the octet rule.
- Boron trifluoride (BF3) would have a Lewis structure beginning with three electrons around boron and seven around all three fluorine atoms each. The octet rule cannot be satisfied for boron. It has only three electrons, and with one covalent bond each to the three fluorine atoms, it is still has only six electrons when it should have eight.
For example, the molecules NO (nitric oxide) and BF3 (boron trifluoride):
If the theory behind Lewis structures was complete and totally accurate, neither of these two molecules should exist. The fact that they do shows the theory is not totally accurate!