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Representing Valence Electrons in Lewis Symbols

Get the free 'Lewis structure' widget for your website, blog, Wordpress, Blogger, or iGoogle. Find more Chemistry widgets in Wolfram Alpha. Is it possible to make a Lewis structure for metallic bonds.? For our project on bonds, the teacher wants us to draw at least two examples of each bonds, including metallic bonds. I have searched everywhere on how to do this, but there is not an answer anywhere. Lewis theory (Gilbert Newton Lewis, 1875-1946) focuses on the valence electrons, since the outermost electrons are the ones that are highest in energy and farthest from the nucleus, and are therefore the ones that are most exposed to other atoms when bonds form. Lewis dot diagrams for elements are a handy way of picturing valence electrons,. Given descriptions, diagrams, scenarios, or chemical symbols, students will model covalent bonds using electron dot formula (Lewis structures).

Lewis symbols use dots to visually represent the valence electrons of an atom.

Learning Objectives

Recall the Lewis structure formalism for representing valance electrons

Key Takeaways

Key Points

• Electrons exist outside of an atom ‘s nucleus and are found in principal energy levels that contain only up to a specific number of electrons.
• The outermost principal energy level that contains electrons is called the valence level and contains valence electrons.
• Lewis symbols are diagrams that show the number of valence electrons of a particular element with dots that represent lone pairs.
• Lewis symbols do not visualize the electrons in the inner principal energy levels.

Key Terms

• principal energy levels: The different levels where electrons can be found and that occur at specific distances from the atom’s nucleus. Each level is associated with a particular energy value that electrons within it have.
• valence level: The outermost principal energy level, which is the level furthest away from the nucleus that still contains electrons.
• valence electrons: The electrons of atoms that participate in the formation of chemical bonds.
• Lewis symbols: Symbols of the elements with their number of valence electrons represented as dots

Lewis symbols (also known as Lewis dot diagrams or electron dot diagrams) are diagrams that represent the valence electrons of an atom. Lewis structures (also known as Lewis dot structures or electron dot structures) are diagrams that represent the valence electrons of atoms within a molecule. These Lewis symbols and Lewis structures help visualize the valence electrons of atoms and molecules, whether they exist as lone pairs or within bonds.

Principal Energy Levels

An atom consists of a positively charged nucleus and negatively charged electrons. The electrostatic attraction between them keeps electrons ‘bound’ to the nucleus so they stay within a certain distance of it. Careful investigations have shown that not all electrons within an atom have the same average position or energy. We say the electrons ‘reside’ in different principal energy levels, and these levels exist at different radii from the nucleus and have rules regarding how many electrons they can accommodate.

Principal energy levels of gold (Au): The figure shows the organization of the electrons around the nucleus of a gold (Au) atom. Notice that the first energy level (closest to the nucleus) can have only two electrons, while more electrons can ‘fit’ within a given level further out. The number of electrons in each level is listed on the upper right corner of the figure. Notice that the outermost level has only one electron.

As an example, a neutral atom of gold (Au) contains 79 protons in its nucleus and 79 electrons. The first principal energy level, which is the one closest to the nucleus, can hold a maximum of two electrons. The second principal energy level can have 8, the third can have 18, and so on, until all 79 electrons have been distributed.

The outermost principal energy level is of great interest in chemistry because the electrons it holds are the furthest away from the nucleus, and therefore are the ones most loosely held by its attractive force; the larger the distance between two charged objects, the smaller the force they exert on each other. Chemical reactivity of all of the different elements in the periodic table depends on the number of electrons in that last, outermost level, called the valence level or valence shell. In the case of gold, there is only one valence electron in its valence level.

Octet of Valence Electrons

Atoms gain, lose, or share electrons in their valence level in order to achieve greater stability, or a lower energy state. From this perspective, bonds between atoms form so that the bonded atoms are in a lower energy state compared to when they were by themselves. Atoms can achieve this more stable state by having a valence level which contains as many electrons as it can hold. For the first principal energy level, having two electrons in it is the most stable arrangement, while for all other levels outside of the first, eight electrons are necessary to achieve the most stable state.

Lewis Symbols

In the Lewis symbol for an atom, the chemical symbol of the element (as found on the periodic table) is written, and the valence electrons are represented as dots surrounding it. Only the electrons in the valence level are shown using this notation. For example, the Lewis symbol of carbon depicts a “C’ surrounded by 4 valence electrons because carbon has an electron configuration of 1s²2s²2p².

The Lewis symbol for carbon: Each of the four valence electrons is represented as a dot.

Electrons that are not in the valence level are not shown in the Lewis symbol. The reason for this is that the chemical reactivity of an atom of the element is solely determined by the number of its valence electrons, and not its inner electrons. Lewis symbols for atoms are combined to write Lewis structures for compounds or molecules with bonds between atoms.

Writing Lewis Symbols for Atoms

The Lewis symbol for an atom depicts its valence electrons as dots around the symbol for the element.

Key Takeaways

Key Points

• The columns, or groups, in the periodic table are used to determine the number of valence electrons for each element.
• The noble/ inert gases are chemically stable and have a full valence level of electrons.
• Other elements react in order to achieve the same stability as the noble gases.
• Lewis symbols represent the valence electrons as dots surrounding the elemental symbol for the atom.

Key Terms

• group: A column in the periodic table that consists of elements with similar chemical reactivity, because they have the same number of valence electrons.
• Noble Gases: Inert, or unreactive, elements in the last group in the periodic table which are typically found in the gaseous form.
• Lewis symbol: Formalism in which the valence electrons of an atom are represented as dots.

Determining the Number of Valence Electrons

In order to write the Lewis symbol for an atom, you must first determine the number of valence electrons for that element. The arrangement of the periodic table can help you figure out this information. Since we have established that the number of valence electrons determines the chemical reactivity of an element, the table orders the elements by number of valence electrons.

Each column (or group) of the periodic table contains elements that have the same number of valence electrons. Furthermore, the number of columns (or groups) from the left edge of the table tells us the exact number of valence electrons for that element. Recall that any valence level can have up to eight electrons, except for the first principal energy level, which can only have two.

Periodic table of the elements: Group numbers shown by Roman numerals (above the table) tell us how many valence electrons there are for each element.

Some periodic tables list the group numbers in Arabic numbers instead of Roman numerals. In that case, the transition metal groups are included in the counting and the groups indicated at the top of the periodic table have numbers 1, 2, 13, 14, 15, 16, 17, 18. The corresponding roman numerals used are I, II, III, IV, V, VI, VII, VIII.

Survey of the Groups in the Periodic Table

Take the first column or group of the periodic table (labeled ‘I’): hydrogen (H), lithium (Li), sodium (Na), potassium (K), etc. Each of these elements has one valence electron. The second column or group (labeled ‘II’) means that beryllium (Be), magnesium (Mg), calcium (Ca), etc., all have two valence electrons.

The middle part of the periodic table that contains the transition metals is skipped in this process for reasons having to do with the electronic configuration of these elements.

Proceeding to the column labeled ‘III’, we find that those elements (B, Al, Ga, In,…) have three valence electrons in their outermost or valence level.

We can continue this inspection of the groups until we reach the eighth and final column, in which the most stable elements are listed. These are all gaseous under normal conditions of temperature and pressure, and are called ‘noble gases.’ Neon (Ne), argon (Ar), krypton (Kr), etc., each contain eight electrons in their valence level. Therefore, these elements have a full valence level that has the maximum number of electrons possible. Helium (He), at the very top of this column is an exception because it has two valence electrons; its valence level is the first principal energy level which can only have two electrons, so it has the maximum number of electrons in its valence level as well.

The Lewis symbol for helium: Helium is one of the noble gases and contains a full valence shell. Unlike the other noble gases in Group 8, Helium only contains two valence electrons. In the Lewis symbol, the electrons are depicted as two lone pair dots.

The noble gases represent elements of such stability that they are not chemically reactive, so they can be called inert. In other words, they don’t need to bond with any other elements in order to attain a lower energy configuration. We explain this phenomenon by attributing their stability to having a ‘full’ valence level.

The significance in understanding the nature of the stability of noble gases is that it guides us in predicting how other elements will react in order to achieve the same electronic configuration as the noble gases by having a full valence level.

Writing Lewis Symbols for Atoms

Lewis symbols for the elements depict the number of valence electrons as dots. In accordance with what we discussed above, here are the Lewis symbols for the first twenty elements in the periodic table. The heavier elements will follow the same trends depending on their group.

Once you can draw a Lewis symbol for an atom, you can use the knowledge of Lewis symbols to create Lewis structures for molecules.

Valence Electrons and the Periodic Table: Electrons can inhabit a number of energy shells. Different shells are different distances from the nucleus. The electrons in the outermost electron shell are called valence electrons, and are responsible for many of the chemical properties of an atom. This video will look at how to find the number of valence electrons in an atom depending on its column in the periodic table.

Introduction to Lewis Structures for Covalent Molecules

In covalent molecules, atoms share pairs of electrons in order to achieve a full valence level.

Learning Objectives

Predict and draw the Lewis structure of simple covalent molecules and compounds

Key Takeaways

Key Points

• The octet rule says that the noble gas electronic configuration is a particularly favorable one that can be achieved through formation of electron pair bonds between atoms.
• In many atoms, not all of the electron pairs comprising the octet are shared between atoms. These unshared, non-bonding electrons are called ‘ lone pairs ‘ of electrons.
• Although lone pairs are not directly involved in bond formation, they should always be shown in Lewis structures.
• There is a logical procedure that can be followed to draw the Lewis structure of a molecule or compound.

Key Terms

• octet rule: Atoms try to achieve the electronic configuration of the noble gas nearest to them in the periodic table by achieving a full valence level with eight electrons.
• exceptions to the octet rule: Hydrogen (H) and helium (He) only need two electrons to have a full valence level.
• covalent bond: Two atoms share valence electrons in order to achieve a noble gas electronic configuration.
• Lewis structure: Formalism used to show the structure of a molecule or compound, in which shared electrons pairs between atoms are indicated by dashes. Non-bonding, lone pairs of electrons must also be shown.

The Octet Rule

Noble gases like He, Ne, Ar, Kr, etc., are stable because their valence level is filled with as many electrons as possible. Eight electrons fill the valence level for all noble gases, except helium, which has two electrons in its full valence level. Other elements in the periodic table react to form bonds in which valence electrons are exchanged or shared in order to achieve a valence level which is filled, just like in the noble gases. We refer to this chemical tendency of atoms as ‘the octet rule,’ and it guides us in predicting how atoms combine to form molecules and compounds.

Covalent Bonds and Lewis Diagrams of Simple Molecules

The simplest example to consider is hydrogen (H), which is the smallest element in the periodic table with one proton and one electron. Hydrogen can become stable if it achieves a full valence level like the noble gas that is closest to it in the periodic table, helium (He). These are exceptions to the octet rule because they only require 2 electrons to have a full valence level.

Two H atoms can come together and share each of their electrons to create a ‘ covalent bond.’ The shared pair of electrons can be thought of as belonging to either atom, and thus each atom now has two electrons in its valence level, like He. The molecule that results is H₂, and it is the most abundant molecule in the universe.

Lewis structure of diatomic hydrogen: This is the process through which the H₂ molecule is formed. Two H atoms, each contributing an electron, share a pair of electrons. This is known as a ‘single covalent bond.’ Notice how the two electrons can be found in a region of space between the two atomic nuclei.

The Lewis formalism used for the H₂ molecule is H:H or H—H. The former, known as a ‘Lewis dot diagram,’ indicates a pair of shared electrons between the atomic symbols, while the latter, known as a ‘Lewis structure,’ uses a dash to indicate the pair of shared electrons that form a covalent bond. More complicated molecules are depicted this way as well.

Lewis dot dragram for methane: Methane, with molecular formula CH₄, is shown. The electrons are color-coded to indicate which atoms they belonged to before the covalent bonds formed, with red representing hydrogen and blue representing carbon. Four covalent bonds are formed so that C has an octet of valence electrons, and each H has two valence electrons—one from the carbon atom and one from one of the hydrogen atoms.

Lewis Dot Structure Practice

Now consider the case of fluorine (F), which is found in group VII (or 17) of the periodic table. It therefore has 7 valence electrons and only needs 1 more in order to have an octet. One way that this can happen is if two F atoms make a bond, in which each atom provides one electron that can be shared between the two atoms. The resulting molecule that is formed is F₂, and its Lewis structure is F—F.

Achieving an octet of valence electrons: Two fluorine atoms are able to share an electron pair, which becomes a covalent bond. Notice that only the outer (valence level) electrons are involved, and that in each F atom, 6 valence electrons do not participate in bonding. These are ‘lone pairs’ of electrons.

After a bond has formed, each F atom has 6 electrons in its valence level which are not used to form a bond. These non-bonding valence electrons are called ‘lone pairs’ of electrons and should always be indicated in Lewis diagrams.

Lewis structure of acetic acid: Acetic acid, CH₃COOH, can be written out with dots indicating the shared electrons, or, preferably, with dashes representing covalent bonds. Notice the lone pairs of electrons on the oxygen atoms are still shown. The methyl group carbon atom has six valence electrons from its bonds to the hydrogen atoms because carbon is more electronegative than hydrogen. Also, one electron is gained from its bond with the other carbon atom because the electron pair in the C−C bond is split equally.

Procedure for Drawing Simple Lewis Structures

We have looked at how to determine Lewis structures for simple molecules. The procedure is as follows:

1. Write a structural diagram of the molecule to clearly show which atom is connected to which (although many possibilities exist, we usually pick the element with the most number of possible bonds to be the central atom).
2. Draw Lewis symbols of the individual atoms in the molecule.
3. Bring the atoms together in a way that places eight electrons around each atom (or two electrons for H, hydrogen) wherever possible.
4. Each pair of shared electrons is a covalent bond which can be represented by a dash.

Alternate view of lewis dot structure of water: This arrangement of shared electrons between O and H results in the oxygen atom having an octet of electrons, and each H atom having two valence electrons.

Multiple bonds can also form between elements when two or three pairs of electrons are shared to produce double or triple bonds, respectively. The Lewis structure for carbon dioxide, CO₂, is a good example of this.

Lewis structure of carbon dioxide: This figure explains the bonding in a CO₂ molecule. Each O atom starts out with six (red) electrons and C with four (black) electrons, and each bond behind an O atom and the C atom consists of two electrons from the O and two of the four electrons from the C.

In order to achieve an octet for all three atoms in CO₂, two pairs of electrons must be shared between the carbon and each oxygen. Since four electrons are involved in each bond, a double covalent bond is formed. You can see that this is how the octet rule is satisfied for all atoms in this case. When a double bond is formed, you still need to show all electrons, so double dashes between the atoms show that four electrons are shared.

Final Lewis structure for carbon dioxide: Covalent bonds are indicated as dashes and lone pairs of electrons are shown as pairs of dots. in carbon dioxide, each oxygen atom has two lone pairs of electrons remaining; the covalent bonds between the oxygen and carbon atoms each use two electrons from the oxygen atom and two from the carbon.

Lewis Structures for Polyatomic Ions

The Lewis structure of an ion is placed in brackets and its charge is written as a superscript outside of the brackets, on the upper right.

Learning Objectives

Apply the rules for drawing Lewis structures to polyatomic ions

Key Takeaways

Key Points

• Ions are treated almost the same way as a molecule with no charge. However, the number of electrons must be adjusted to account for the net electric charge of the ion.
• When counting electrons, negative ions should have extra electrons placed in their Lewis structures, while positive ions should have fewer electrons than an uncharged molecule.

Key Terms

Lewis Dot Structure Finder

• polyatomic ion: A charged species composed of two or more atoms covalently bonded, or of a metal complex that acts as a single unit in acid-base chemistry or in the formation of salts. Also known as a molecular ion.

The total number of electrons represented in a Lewis structure is equal to the sum of the numbers of valence electrons in each individual atom. Non-valence electrons are not represented in Lewis structures. After the total number of available electrons has been determined, electrons must be placed into the structure.

Lewis structures for polyatomic ions are drawn by the same methods that we have already learned. 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 of the brackets. For example, consider the ammonium ion, NH₄⁺, which contains 9 (5 from N and 1 from each of the four H atoms) –1 = 8 electrons. One electron is subtracted because the entire molecule has a +1 charge.

Coordinate covalent bonding: The ammonium ion, NH4⁺, contains 9–1 = 8 electrons.

Negative ions follow the same procedure. The chlorite ion, ClO₂^–, contains 19 (7 from the Cl and 6 from each of the two O atoms) +1 = 20 electrons. One electron is added because the entire molecule has a -1 charge.

Hypochlorite ion Lewis structure: The hypochlorite ion, ClO⁻, contains 13 + 1 = 14 electrons.

The sharing of pair of electrons between two atoms is referred to as a covalent bond. Normally, each atom that is participating in the covalent bond formation, contributes equal number of electrons to form pair(s) of electrons. The pair of electrons shared between the atoms is also known as bond pair.

The bond pair is strongly attracted by the nuclei of two atoms and thus by reducing the potential energy of them. This is the driving force of formation of covalent bond, which stabilizes the two atoms.

The distance between the nuclei of two atoms when their potential energy reaches a minima is also known as bond length. The forces of attraction between these two atoms are maximum at this point. However, the repulsion forces dominate over the attraction forces when these atoms are bring further closer together and thus by increasing the potential energy.

If two atoms share only one bond pair, that bond is referred to as a single bond. If two bond pairs are shared, that is known as a double bond. Likewise, a triple bond is formed when the atoms share three bond pairs.

A covalent bond is formed between two atoms when their electronegativity difference is less than 1.7 on Pauling's scale. Usually it is formed between two nonmetals.

Polarity of covalent bond: The bond pair is equally shared in between two atoms when the electronegativity difference between them is zero or nearer to zero. In this case, neither of the atoms gets excess of electron density and hence carry no charge. This is called non polar covalent bond.

However, when there is a considerable difference in the electronegativity, the bond pair is no longer shared equally between the atoms. It is shifted slightly towards the atom with higher electronegativity by creating partial negative charge (represented by δ-) over it. Whereas, the atom with less electronegativity gets partial positive charge (represented by δ+). This type of bond is also referred to as polar covalent bond.

Coordinate covalent bond: Some times, during the formation of covalent bond, the shared pair is entirely contributed by only one atom This is called as coordinate covalent bond or dative bond.

LEWIS DOT MODEL

To explain the formation of covalent bond, a simple qualitative model was developed by Gilbert Newton Lewis in 1916.

According to this model:

* Octet rule: The inert gas atoms with 8 electrons in their outer shell (also known as valence shell) are highly stable. The Helium atom with 2 electrons in its outer shell is also stable.

Hence every atom tries to get nearest inert gas configuration by sharing electrons. The bond formed due to sharing of electrons is otherwise known as a covalent bond.

* Only the electrons in the valence shell are contributed for sharing. The inner electrons, which are also known as core electrons do not participate in the bond formation.

* In the formation of covalent bond between two atoms, each atom contributes its valence electrons to form pair(s) of electrons, which in turn is/are shared by both of them.

Due to sharing of electrons, each atom gets nearest inert gas configuration.

Covalency : The number of electrons contributed by the atom of an element in the formation of covalent compound is known as covalency of that element.

In Lewis dot model, the electrons in the valence shell of the atom are shown as dots around it.

ILLUSTRATIONS OF LEWIS DOT STRUCTURES

1) H₂ molecule:

* The electronic configuration of hydrogen is 1s¹. It requires one electron to get the configuration of Helium.

* Therefore, during the formation of H₂ molecule, each hydrogen atom contributes one electron to form a pair of electrons. This in turn is shared between the two hydrogen atoms to form a covalent bond.

* Thus in H₂ molecule, each hydrogen atom gets its nearest inert gas: Helium's configuration, 1s².

* The covalency of hydrogen is 1.

* The Lewis dot structure for H₂ molecule is shown below. Note that each hydrogen gets two electrons after forming the bond. The bond between two hydrogen atoms can be shown as a line, which represents a bond pair of electrons.

Note: The bond between two hydrogen atoms is non polar since the electronegativity difference is zero.

2) Cl₂ molecule:

* The electronic configuration of Cl is [Ne]3s²3p⁵. There are 7 electrons in the outer shell.

* Therefore, in order to get the nearest inert gas: Argon's configuration, [Ne]3s²3p⁶, each chlorine atom contributes one electron for the bond formation and form a bond pair, which is shared between the two chlorine atoms.

* Thus Cl₂ molecule is formed with a covalent bond between two chlorine atoms.

* In Cl₂ molecule, each Cl atom gets 8 electrons in its outer shell. See the diagram below.

The bond pair is also shown as a line. The electron pairs which do not participate in the bonding are known as lone pairs. There are three such lone pairs on each of the chlorine atom.

* Hence the covalency of chlorine is 1.

* Remember that, in the Lewis dot structures, only the electrons in the valence shell are shown.

Note: The bond between two chlorine atoms is non polar since the electronegativity difference is zero.

3) Hydrogen chloride (HCl):

* The electronic configuration of hydrogen is 1s¹. It requires one electron to get the configuration of He.

* The electronic configuration of chlorine is [Ne]3s²3p⁵. It also requires one electron to get the octet configuration.

* Hence, in the formation of HCl molecule, the hydrogen and chlorine atoms contribute one electron each for the bond formation.

Note: The bond between hydrogen and chlorine atoms is considerably polar, since the electronegativity difference between them is 3.5 - 2.1 = 1.4. However they do not form ionic bond since the e.n. difference is less than 1.7.

The hydrogen atom gets partial positive charge, whereas the chlorine atom gets partial negative charge. However, they are not completely ionized.

4) Methane (CH₄):

* The electronic configuration of carbon is [He]2s²2p².

* The electronic configuration of hydrogen is 1s¹. Hence it contributes this electron to get the nearest inert gas, Helium's configuration.

* The carbon atom forms four covalent bonds by contributing four of its valence electrons. It forms 4 bonds with four hydrogen atoms. Thus it gets octet configuration.

* Covalency of carbon is 4.

* There are no lone pairs on carbon in methane.

Note: The electronegativity difference between carbon (e.n. = 2.5) and hydrogen (e.n. = 2.1) is not considerable. Hence the covalent bond formed is practically considered as non polar.

5) Ammonia (NH₃) :

* The electronic configuration of nitrogen is [He]2s²2p³. Nitrogen require three electrons to complete the octet.

* The electronic configuration of hydrogen is 1s¹ and hence in need of one electron to complete the shell.

* In the formation of Ammonia molecule, the nitrogen atom contributes 3 of its valence electrons to form three bond pairs which are shared with hydrogen atoms. Thus nitrogen forms 3 single bonds with three hydrogen atoms and gets the configuration of Neon.

* There is also one lone pair on nitrogen atom.

* Covalency of nitrogen is 3.

Note: There is considerable polarity in N-H bond since the electronegativity difference between nitrogen (e.n. = 3.0) and hydrogen (e.n. = 2.1) is 0.9.

Being highly electronegative than hydrogen, the nitrogen atom gets partial negative charge, whereas the hydrogen atom gets partial positive charge.

6) H₂O molecule :

* The electronic configuration of oxygen is [He]2s²2p⁴. Hence it requires two more electrons to complete its octet.

* In the formation of water molecule, the oxygen atom contributes two of its valence electrons to form two bond pairs that are shared with two hydrogen atoms separately. Thus two bonds are formed by oxygen atom to get the configuration of Neon.

* There are also two lone pairs on oxygen atom.

* Covalency of oxygen is 2.

Note: The O-H bond is also considerably polar since the electronegativity difference between oxygen (e.n. = 3.5) and hydrogen (e.n. = 2.1) is 1.4.

Since the oxygen atom is more electronegative, it gets partial negative charge.

7) Dioxygen molecule (O₂):

* In the formation of dioxygen molecule, each oxygen atom contributes 2 electrons to form 2 bond pairs, which are shared by the two oxygen atoms. Thus each oxygen atom gets Neon's configuration.

* There is a double bond between two oxygen atoms.

* However there are two lone pairs on each oxygen atom.

8) Dinitrogen molecule (N₂):

* In the formation of Dinitrogen molecule, each nitrogen atom contributes 3 electrons to form 3 bond pairs, which in turn are shared by two nitrogen atoms. Each nitrogen atom gets Neon's configuration.

* Thus a triple bond between nitrogen atoms is formed.

* Each nitrogen also contains one lone pair.

9) Carbon dioxide molecule (CO₂):

* The carbon atom contributes four of its valence electrons, whereas each oxygen atom contributes two electrons for the bond formation to complete their octets.

* There are two electron pairs shared between carbon and one of the oxygen atom i.e., a double bond, C=O is formed. There are two such C=O bonds in CO₂ molecule.

Note: The C=O bond is polar due to electronegativity difference between carbon (e.n. = 2.1) and oxygen (e.n. = 3.5) is 1.4.

Question to ponder: What is the overall polarity of the carbon dioxide molecule? Is it polar or non polar?

Molecules violating octet rule:

1) BeCl₂ (Beryllium chloride):

* The electronic configuration of Beryllium is 1s²2s².

* The electronic configuration of Chlorine is [Ne]3s²3p⁵.

* During the formation of Beryllium chloride, the beryllium atom contributes its two valence electrons and forms two bond pairs. These are shared with chlorine atoms.

* In BeCl₂ there are only four electrons in the valence shell of Be. However the Beryllium chloride is a relatively stable molecule. This is the violation of octet rule.

* Covalency of beryllium is 2.

Note: The beryllium atom may lose two of its valence electrons to get the nearest inert gas: Helium's configuration and form Be²⁺ ion. However, the electronegativity difference between Beryllium (e.n. = 1.5) and that of chlorine (e.n. = 3.0) is less than 1.7. Hence the bond between them has considerable covalent character rather than the ionic nature.

2) BCl₃ (Boron trichloride):

* The electronic configuration of Boron is 1s²2s²2p¹.

* During the formation of BCl₃ molecule, the boron atom contributes 3 of its valence electrons to form three bond pairs with chlorine atoms.

* Thus there are only six electrons in the valence shell of boron atom in BCl₃. However, still this molecule is stable. It is an electron deficient compound. It is also the violation of the octet rule.

* Covalency of boron is 3.

Note: The bond between boron (e.n. = 2.0) and chlorine (e.n. = 3.0) is covalent since the electronegativity difference between them is less than 1.7.

3) PCl₅ (Phosphorous pentachloride):

* Electronic configuration of phosphorus is [Ne]3s²3p³.

* Electronic configuration of chlorine is [Ne]3s²3p⁵.

* In the formation of PCl₅ molecule, phosphorus contributes five electrons in it's valence shell and forms five bonds with chlorine atoms.

* Thus there are 10 electrons in the valence shell of phosphorus in this molecule. It is a stable molecule, however, violating the octet rule.

* Covalency of phosphorus in this molecule is 5.

4) SF₆ (Sulphur hexafluoride) :

* The electronic configuration of Sulphur is [Ne]3s²3p⁴.

* The electronic configuration of Fluorine is [He]2s²2p⁵.

* The fluorine atom requires one electron to complete its octet. Hence it contributes one electron for bonding.

* The Sulphur atom contributes six of its valence electrons to form 6 bonds with six fluorine atoms. Thus there are 12 electrons in the valence shell in sulphur atom.

* The SF₆ molecule is quite stable even though it violates the octet rule.

* Covalency of sulphur in this molecule is 6.

DRAWBACKS OF LEWIS THEORY

* Lewis theory could not explain the geometry of molecules and bond angle in them.

* It could not explain why some molecules are violating the octet rule.

* This is a qualitative explanation for covalent bond only.

To fulfill these gaps and to explain the covalent bond formation quantitatively, the Valence bond theory (VBT) was put forwarded. You can find more insight into this theory in next section.

However it is important to learn VSEPR theory, another qualitative model, which was put forwarded to explain the shapes of molecules, before moving on to Valence bond theory.

< Ionic bond

Chemical bonding: TOC

VSEPR Theory >

By Aditya vardhan Vutturi