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Lewis Dot Diagram For N

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Last Updated: 27 October 2020

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General | Latest Info

Why are some substances chemically bond molecules and others are association of ions? The answer to this question depends upon the electronic structures of atoms and the nature of chemical forces within compounds. Although there are no sharply defined boundaries, chemical bonds are typically classified into three main types: ionic bonds, Covalent bonds, and Metallic bonds. In this chapter, each type of bond will be discussed and general properties found in typical substances in which bond type occur ionic bonds result from electrostatic forces that exist between ions of opposite charge. These bonds typically involve metal with nonmetal Covalent bonds result from sharing of electrons between two atoms. Bonds typically involve one nonmetallic element with another metallic bond. These bonds are found in solid metals with each metal bond to several neighboring groups and bonding electrons free to move throughout the 3-dimensional structure. Each bond classification is discussed in detail in subsequent sections of the chapter. Let's look at preferred arrangements of electrons in atoms when they form chemical compounds. Figure: G. N. Lewis and Octet Rule. Lewis is working in a laboratory. In Lewiss original sketch for Octet Rule, he initially placed electrons at corners of the cube rather than placing them as we do now. Normally, two electrons pair up and form a bond, eg, {matheq}H_2{endmatheq} for most atoms there will be a maximum of eight electrons in valence shell, eg, {matheq}CH_4{endmatheq} figure: Bonding in {matheq}H_2{endmatheq} and methane {matheq}CH_4{endmatheq} other tendency of atoms is to maintain neutral charge. Only noble gases have zero charge with filled valence octets. All of the other elements have a charge when they have eight electrons all to themselves. The result of these two guiding principles is an explanation for much of the reactivity and bonding that is observed within atoms: atoms seek to share electrons in a way that minimizes charge while fulfilling Octet in valence shell. The formula for table salt is NaCl. It is the result of Na + ions and Cl-ions bonding together. If sodium metal and chlorine gas mix under the right conditions, they will form salt. Sodium loses electron, and chlorine gains that electron. In the process, great amount of light and heat is release. The resulting salt is mostly unreactive it is stable. It will not undergo any explosive reactions, unlike sodium and chlorine that it is made of. Why? Referring to the Octet Rule, atoms attempt to get noble gas electron configuration, which is eight valence electrons. Sodium has one valence electron, so giving it up would result in the same electron configuration as neon. Chlorine has seven valence electrons, so if it takes one it will have eight. Chlorine has electron configuration of argon when it gains electron. The Octet Rule could have been satisfied if chlorine gave up all seven of its valence electrons and sodium took them.

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Electron Dot Diagrams

Table

lithium1 s 2 2 s 11 valence electron
beryllium1 s 2 2 s 22 valence electrons
nitrogen1 s 2 2 s 2 2 p 35 valence electrons
neon1 s 2 2 s 2 2 p 68 valence electrons

Recall that valence electrons of an atom are electrons located at the highest occupied principal energy level. Valence electrons are primarily responsible for the chemical properties of elements. The number of valence electrons can be easily determined from electron configuration. Several examples of second period elements are shown in the table below. In each case, valence electrons are those at second principal energy level. As one proceeds leave to right across period, number of valence electrons increases by one. In s block, Group 1 elements have one valence electron, while Group 2 elements have two valence electrons. In p block, number of valence electrons is equal to Group number minus ten. Group 13 has three valence electrons, Group 14 has four, up through Group 18 with eight. Eight valence electrons, full outer s and p sublevel, give noble gases their special stability. When examining chemical bonding, it is necessary to keep track of the valence of electrons of each atom. Electron dot diagrams are diagrams in which valence electrons of atom are shown as dots distributed around elements symbol. A Beryllium atom, with two valence electrons, would have an electron dot diagram below. Since electrons repel each other, dots for give atoms are distributed evenly around symbol before they are pair. The table below shows electron dot diagrams for the entire second period.

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Lewis Symbols of Monoatomic Elements

In almost all cases, chemical bonds are formed by interactions of valence electrons in atoms. To facilitate our understanding of how valence electrons interact, simple way of representing those valence electrons would be useful. The Lewis electron Dot diagram is a representation of valence electrons of an atom that uses dots around the symbol of element. The number of dots equals the number of valence electrons in an atom. These dots are arranged to right and left and above and below the symbol, with no more than two dots on side. For example, Lewis electron Dot diagram for calcium is simply Figure 1 shows Lewis Symbols for Elements of third Period of Periodic Table.


Example 1: Writing Lewis Structures

Writing Lewis Structures NASAs Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide on Titan, one of Saturn's moons. Titan also contains ethane, acetylene, and ammonia. What are the Lewis structures of these molecules? Calculate the number of valence electrons.: + = 10: + = 14: + = 10: + = 8 Draw skeleton and connect atoms with single bonds. Remember that H is never a central atom: Where needed distribute electrons to terminal atoms: six electrons place on: no electrons remain: no terminal atoms capable of accepting electrons: no terminal atoms capable of accepting electrons Where needed place remaining electrons on the central atom: no electrons remain: no electrons remain: four electrons place on carbon: two electrons place on nitrogen Where need, rearrange electrons to form multiple bonds in order to obtain octet on each atom: form two more C-N bonds: all atoms have correct number of electrons: form triple bond between two carbon atoms: all atoms have correct number of electrons check Your Learning Both carbon monoxide, and carbon dioxide, are products of combustion of fossil fuels. Both of these gases also cause problems: are toxic and have been implicated in global climate change. What are the Lewis structures of these two molecules?


Introduction to Lewis Structures for Covalent Molecules

We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in Lewis structures, drawings that describe bonding in molecules and polyatomic ions. For example, when two chlorine atoms form chlorine molecule, they share one pair of electrons: Lewis structure indicates that each atom has three pairs of electrons that are not used in bonding and one share pair of electrons. Dash is sometimes used to indicate shared pair of electrons: single shared pair of electrons is called single bond. Each atom interacts with eight valence electrons: six in lone pairs and two in single bond.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions

The Octet Rule

For very simple molecules and molecular ions, we can write Lewis structures by merely pairing up unpaired electrons on constituent atoms. See these examples: For more complicated molecules and molecular ions, it is helpful to follow the Step-by-Step procedure outline here: Determine Total number OF valence electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge. Draw skeleton structure of a molecule or ion, arranging atoms around the central atom. Connect each atom to the central atom with a single bond. Distribute remaining electrons as lone pairs on terminal atoms, completing Octet around each atom. Place all remaining electrons on the central atom. Rearrange electrons OF outer atoms to make multiple bonds with central atom in order to obtain octets wherever possible. Let us determine Lewis structures OF SiH 4, CHO 2, NO +, and OF 2 as examples in following this procedure: Determine total number OF valence electrons in molecule or ion. For molecule, we add the number OF valence electrons on each atom in molecule: {matheq}\begin{array}{r r l} \text{SiH}_4 & & \ {matheq}1em] & \text{Si: 4 valence electrons/atom} \times 1 \;\text{atom} & = 4 \ {matheq}1em] \rule[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & \text{H: 1 valence electron/atom} \times 4 \;\text{atoms} & = 4 \ {matheq}1em] & & = 8 \;\text{valence electrons} \end{array}{endmatheq} For negative ion, such as CHO 2 −, we add the number OF valence electrons on atoms to the number of negative charges on ion: {matheq}\begin{array}{r r l} {\text{CHO}_2}^{-} & & \ {matheq}1em] & \text{C: 4 valence electrons/atom} \times 1 \;\text{atom} & = 4 \ {matheq}1em] & \text{H: 1 valence electron/atom} \times 1 \;\text{atom} & = 1 \ {matheq}1em] & \text{O: 6 valence electrons/atom} \times 2 \;\text{atoms} & = 12 \ {matheq}1em] \rule[-0.5ex]{21.5em}{0.1ex}\hspace{-21.5em} + & 1\;\text{additional electron} & = 1 \ {matheq}1em] & & = 18 \;\text{valence electrons} \end{array}{endmatheq} For positive ion, such as NO +, we add the number OF valence electrons on atoms in ion and then subtract number OF positive charges on ion from Total number OF valence electrons: {matheq}\begin{array}{r r l} \text{NO}^{+} & & \ {matheq}1em] & \text{N: 5 valence electrons/atom} \times 1 \;\text{atom} & = 5 \ {matheq}1em] & \text{O: 6 valence electrons/atom} \times 1 \;\text{atom} & = 6 \ {matheq}1em] \rule[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & -1 \;\text{electron (positive charge)} & = -1 \ {matheq}1em] & & = 10 \;\text{valence electrons} \end{array}{endmatheq} since OF 2 is neutral molecule, We simply add number OF valence electrons: {matheq}\begin{array}{r r l} \text{OF}_{2} & & \ {matheq}1em] & \text{O: 6 valence electrons/atom} \times 1 \;\text{atom} & = 6 \ {matheq}1em] \rule[-0.5ex]{21em}{0.1ex}\hspace{-21em} + & \text{F: 7 valence electrons/atom} \times 2 \;\text{atoms} & = 14 \ {matheq}1em] & & = 20 \;\text{valence electrons} \end{array}{endmatheq} Draw skeleton structure OF molecule or ion, arranging atoms around central atom and connecting each atom to central atom with single bond. When several arrangements OF atoms are possible, as for CHO 2 −, we must use experimental evidence to choose the correct one. In general, less electronegative elements are more likely to be central atoms. In CHO 2 −, less electronegative carbon atoms occupy central position with oxygen and hydrogen atoms surrounding them. Other examples include P in POCl 3, S in SO 2, and Cl in ClO 4 −. The exception is that hydrogen is almost never the central atom. As most electronegative element, fluorine also cannot be central atom. Distribute remaining electrons as lone pairs on terminal atoms to complete their valence shells with Octet OF electrons. There are NO remaining electrons on SiH 4, SO it is unchanged: Place all remaining electrons on the central atom. For SiH 4, CHO 2 −, and NO +, there are NO remaining electrons; We already Place all OF electrons Determine in Step 1.


Limitations of Lewis Structures

Lewis structures first came into use early in the twentieth century when chemical bonding was poorly understood. Electron dot diagrams help illustrate the electronic structure of molecules and chemical reactivity. Their use remains popular with chemistry educators introducing valence-bond models of chemical bonds and they are often used in organic chemistry, where valence-bond model is most appropriate. However, in the fields of inorganic chemistry and organometallic chemistry, delocalize molecular orbitals are common and Lewis structures don't accurately predict behavior. While it's possible to draw Lewis structure for molecule known empirically to contain unpaired electrons, use of such structures leads to errors in estimating bond length, magnetic properties, and aromaticity. Examples of these molecules include molecular oxygen, nitric oxide, and chlorine dioxide. While Lewis structures have some value, reader is advise valence bond theory and molecular orbital theory do a better job describing behavior of valence shell electrons.

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Summary

Principal energy levels of gold: figure shows the organization of electrons around the nucleus of gold atom. Notice that the first energy level can have only two electrons, while more electrons can fit within give level further out. The number of electrons in each level is listed in upper right corner of the figure. Notice that the outermost level has only one electron. Lewis dot dragram for Methane: Methane, with molecular formula CH 4, is show. Electrons are color-cod to indicate which atoms they belong to before covalent bonds form, 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 electronsone, from the carbon atom and one from one of the hydrogen atoms. Lewis structure of Acetic acid: Acetic acid, CH 3 COOH, can be written out with dots indicating shared electrons, or, preferably, with dashes representing covalent bonds. Notice lone pairs of electrons on oxygen atoms are still on show. The Methyl group carbon atom has six valence electrons from its bonds to hydrogen atoms because carbon is more electronegative than hydrogen. Also, one electron is gained from its bond with other carbon atom because the electron pair in CC bond is split equally.


Lewis Structures for Polyatomic Ions

We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in Lewis structures, drawings that describe bonding in molecules and polyatomic ions. For example, when two chlorine atoms form chlorine molecule, they share one pair of electrons: Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding and one share pair of electrons. Dash is sometimes used to indicate shared pair of electrons: single shared pair of electrons is called single bond. Each Cl atom interacts with eight valence electrons: six in lone pairs and two in single bond.


Representing Valence Electrons in Lewis Symbols

We use Lewis symbols to describe valence electron configurations of atoms and monatomic ions. Lewis symbols consist of an elemental symbol surrounded by one dot for each of its valence electrons: table below shows Lewis symbols for elements of the third period of the periodic table. Lewis symbols can also be used to illustrate formation of cations from atoms, as shown here for sodium and calcium: likewise, they can be used to show formation of anions from atoms, as shown here for chlorine and sulfur: following table demonstrates use of Lewis symbols to show transfer of electrons during formation of ionic compounds.


The Octet Rule

Other halogen molecules form bonds like those in chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the octet rule. The number of bonds that atom can form can often be predicted from the number of electrons needed to reach octet; this is especially true of nonmetals OF second period OF periodic table. For example, each atom OF group 14 element has four electrons in its outermost shell and therefore requires four more electrons to reach the octet. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CCl 4 and silicon in SiH 4. Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. Transition elements and inner transition elements also do not follow the octet rule: group 15 elements such as nitrogen have five valence electrons in atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain octet, these atoms form three covalent bonds, as in NH 3. Oxygen and other atoms in group 16 obtain octets by forming two covalent bonds:

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

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How to draw Lewis Diagrams

Table

N5
O (x 3)18
charge1
24

The following is an example of how to draw the best Lewis structure for NO 3-. First, determine the total number of valence electrons in the molecule. This will be the sum of the group number of all atoms plus charge. Draw skeletal structure for molecule which connect all atoms using only single bonds. Central atom will be one that can form the greatest number of bonds and / or expand its octet. This usually means atom lower and / or to the right in Periodic Table, N in this case. Now we need to add lone pairs of electrons. Of 24 valence electrons available in NO 3-, 6 were used to make skeletal structure. Add lone pairs of electrons on terminal atoms until their octet is complete or you run out of electrons. If there are remaining electrons, they can be used to complete the octet of the central atom. If you have run out of electrons, you are required to use lone pairs of electrons from the terminal atom to complete octet on the central atom by forming multiple bond. In this case, N is short for 2 electrons, so we can use lone pair from leave most O atom to form a double bond and complete octet on N atom. Now you need to determine FORMAL CHARGES for all of the atoms. FORMAL charge is calculated by:-, ie see figure below. NO Lewis structure is complete without FORMAL CHARGES. In general, you want: fewest number of FORMAL CHARGES possible, ie FORMAL CHARGES of 0 for as many atoms in structure as possible. FORMAL CHARGES should match the electronegativity of the atom, that is negative CHARGES should be on more electronegative atoms and positive CHARGES on least electronegative atoms if possible. CHARGES of-1 and + 1 on adjacent atoms can usually be removed by using lone pair of electrons from-1 atom to form a double bond to an atom with + 1 charge. Note: octet can be expanded beyond 8 electrons but only for atoms in period 3 or below in Periodic Table. In our present example, N can not expand beyond 8 electrons so retain FORMAL charge of + 1, but S atom below can expand its octet. You have to determine the best Lewis structure for NO 3-, but there are a number of ways to show this structure. Although it is most common to use line to indicate bonding pair of electrons, they can be shown as electrons, see leave most image below. It is also common to show only net charge on ion rather than all of FORMAL CHARGES, ie see right most figure below. Why are there different ways of the same Lewis structure? It depends what you want to show.

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Double and Triple Bonds

Some molecules are not able to satisfy the octet rule by making only single Covalent bonds between atoms. Consider compound ethene, which has molecular formula of C 2 H 4. Carbon atoms are Bond together, with each carbon also being Bond to two hydrogen atoms. If the Lewis electron dot structure was drawn with a single Bond between carbon atoms and with octet rule follow, it would look like this: this Lewis structure is incorrect because it contains a total of 14 electrons. However, Lewis ' structure can be changed by eliminating lone pairs on carbon atoms and having share two pairs instead of only one pair. A double Covalent Bond is a covalent bond formed by atoms that share two pairs of electrons. The Double Covalent Bond that occurs between two carbon atoms in ethane can also be represented by structural formula and with molecular model as shown below. A Triple Covalent Bond is a covalent bond formed by atoms that share three pairs of electrons. Element nitrogen is a gas that composes the majority of Earth's atmosphere. The nitrogen atom has five valence electrons, which can be show as one pair and three single electrons. When combined with another nitrogen atom to form a diatomic molecule, three single electrons on each atom combine to form three share pairs of electrons. Each nitrogen atom follows the octet rule with one lone pair of electrons and six electrons that are shared between atoms.

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Fullerene Chemistry

Thus far in this chapter, we have discussed various types of bonds that form between atoms and / or ions. In all cases, these bonds involve sharing or transfer of valence shell electrons between atoms. In this section, we will explore typical method for depicting valence shell electrons and chemical bonds, namely Lewis symbols and Lewis structures. Step 1: Calculate the number of valence electrons. HCN: + = 10 H 3 CCH 3: + = 14 HCCH: + = 10 NH 3: + = 8 step 2. Draw skeleton and connect atoms with single bonds. Remember that H is never central atom: step 3: Where needed to distribute electrons to terminal atoms: step 4: Where needed to place remaining electrons on central atom: carbon soot has been known to man since prehistoric times, but it was not until fairly recently that the molecular structure of the main component of soot was discover. In 1996, Nobel Prize in Chemistry was awarded to Richard Smalley, Robert Curl, and Harold Kroto for their work in discovering a new form of carbon, C 60 buckminsterfullerene molecule. An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C 60. This type of molecule, call fullerene, shows promise in a variety of applications. Because of their size and shape, fullerenes can encapsulate other molecules, so they have shown potential in various applications from hydrogen storage to target drug delivery systems. They also possess unique electronic and optical properties that have been put to good use in solar powered devices and chemical sensors. Richard Smalley, professor of physics, Chemistry, and astronomy at Rice University, was one of the leading advocates for fullerene Chemistry. Upon his death in 2005, US Senate honored him as Father of Nanotechnology. We can draw the Lewis structure of any covalent molecule by following six steps discussed earlier. In this case, we can condense the last few steps, since not all of them apply. 1. Calculate the number of valence electrons: XeF 2: 8 + = 22XeF 6: 8 + = 50 2. Draw skeleton joining atoms by single bonds. Xenon will be central atom because fluorine cannot be central atom: 3. Distribute remaining electrons. XeF 2: We place three lone pairs of electrons around each F atom, accounting for 12 electrons and giving each F atom 8 electrons. Thus, six electrons remain. These lone pairs must be placed on Xe atom. This is acceptable because Xe atoms have empty valence shell d orbitals and can accommodate more than eight electrons. The Lewis structure of XeF 2 shows two bonding pairs and three lone pairs of electrons around Xe atom: 4. XeF 6: We place three lone pairs of electrons around each F atom, accounting for 36 electrons.


Lewis Structures

For very simple molecules and molecular ions, we can write Lewis structures by merely pairing up unpaired electrons on constituent atoms. See these examples: For more complicated molecules and molecular ions, it is helpful to follow the Step-by-Step procedure outline here: determine total number of valence electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge. Draw skeleton structure of a molecule or ion, arranging atoms around the central atom. Connect each atom to the central atom with a single bond. Distribute remaining electrons as lone pairs on terminal atoms, completing octet around each atom. Place all remaining electrons on the central atom. Rearrange electrons OF outer atoms to make multiple bonds with central atom in order to obtain octets wherever possible. Let us determine Lewis structures OF, and as example in following this procedure: determine the total number OF valence electrons in molecule or ion. For molecule, we add the number OF valence electrons on each atom in the molecule: SiH 4 Si: 4 valence electrons / atom 1 atom = 4 + H: 1 valence electron / atom 4 atoms = 4 = 8 valence electrons For negative ion, such as, we add the number OF valence electrons on atoms to number OF negative charges on ion: CHO 2-C: 4 valence electrons / atom 1 atom = 4 H: 1 valence electrons / atom 1 atom = 1 O: 6 valence electrons / atom 2 & atoms = 12 + 1 additional electron = 18 valence electrons For positive ion, such as, We add number OF valence electrons on atoms in ion and then subtract number OF positive charges on ion from total number OF valence electrons: NO + N: 5 valence electrons / atom 1 atom = 5 O: 6 valence electrons / atom 1 atom = 6 +-1 electron = 10 valence electrons Since is neutral molecule, We simply add number OF valence electrons: OF 2 O: 6 valence electrons / atom 1 atom = 6 + F: 7 valence electrons / atom 2 atoms = 14 = 20 valence electrons Draw skeleton structure OF molecule or ion, arranging atoms around central atom and connecting each atom to central atom with single bond. When several arrangements OF atoms are possible, as For, we must use experimental evidence to choose the correct one. In general, less electronegative elements are more likely to be central atoms. In, less electronegative carbon atoms occupy central position with oxygen and hydrogen atoms surrounding them. Other examples include In, In, and In. The exception is that hydrogen is almost never the central atom. As most electronegative element, fluorine also cannot be central atom. Distribute remaining electrons as lone pairs on terminal atoms to complete their valence shells with octet OF electrons.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

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Solutions

In this case, Lewis structure is inadequate to depict the fact that experimental studies have shown two unpaired electrons in each oxygen molecule. 11. Two valence electrons per Pb atom are transferred to cl atoms; resulting Pb 2 + ion has a 6 s 2 valence shell configuration. Two of the valence electrons in HCl molecule are share, and the other six are located on Cl atom as lone pairs of electrons. 21. Each bond includes sharing of electrons between atoms. Two electrons are shared in single bond; four electrons are shared in double bond; and six electrons are shared in triple bond.


Lewis Structures

For very simple molecules and molecular ions, we can write Lewis structures by merely pairing up unpaired electrons on constituent atoms. See these examples: For more complicated molecules and molecular ions, it is helpful to follow the Step-by-Step procedure outline here: determine total number of valence electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge. Draw skeleton structure of a molecule or ion, arranging atoms around the central atom. Connect each atom to the central atom with a single bond. Distribute remaining electrons as lone pairs on terminal atoms, completing octet around each atom. Place all remaining electrons on the central atom. Rearrange electrons OF outer atoms to make multiple bonds with central atom in order to obtain octets wherever possible. Let us determine Lewis structures OF SiH 4, {matheq}{endmatheq} NO +, and OF 2 as examples in following this procedure: determine total number OF valence electrons in molecule or ion. For molecule, we add the number OF valence electrons on each atom in molecule: {matheq}SiH4 Si: 4 valence electrons/atomA1 atom=4 A =8 valence electrons{endmatheq} For negative ion, such as {matheq}{endmatheq} we add the number OF valence electrons on atoms to the number of negative charges on ion: {matheq}CHO2aC: 4 valence electrons/atomA1 atom=4H: 1 valence electron/atomA1 atom=1O: 6 valence electrons/atomA2 atoms=12A=18 valence electrons{endmatheq} For positive ion, such as NO +, we add the number OF valence electrons on atoms in ion and then subtract number OF positive charges on ion from total number OF valence electrons: {matheq}NO+N: 5 valence electrons/atomA1 atom=5O: 6 valence electron/atomA1 atom=6A=10 valence electrons{endmatheq} since OF 2 is neutral molecule, we simply add number OF valence electrons: {matheq}OF2O: 6 valence electrons/atomA1 atom=6A=20 valence electrons{endmatheq} draw skeleton structure OF molecule or ion, arranging atoms around central atom and connecting each atom to central atom with single bond. When several arrangements OF atoms are possible, as For {matheq}{endmatheq} we must use experimental evidence to choose the correct one. In general, less electronegative elements are more likely to be central atoms. In {matheq}{endmatheq} less electronegative carbon atom occupy central position with oxygen and hydrogen atoms surrounding it. Other examples include P in POCl 3, S in SO 2, and Cl. In {matheq}{endmatheq} exception is that hydrogen is almost never the central atom. As most electronegative element, fluorine also cannot be central atom. Distribute remaining electrons as lone pairs on terminal atoms to complete their valence shells with octet OF electrons. There are NO remaining electrons on SiH 4, SO it is unchanged: Place all remaining electrons on the central atom. For SiH 4, {matheq}{endmatheq} and NO +, there are NO remaining electrons; we already Place all OF electrons determined in Step 1. For OF 2, we had 16 electrons remaining in Step 3, and we Place 12, leaving 4 to be Place on the central atom: Rearrange electrons OF outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions

Sources

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions

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