# Lewis Diagram Generator

### General | Latest Info

During chemical bonding, it is valence electrons which move amongst different atoms. In order to keep track of the valence of electrons for each atom and how they may be shared in bonding, we use the Lewis Dot Structure for Atoms and molecules. In this approach, we represent valence electrons as dots around the element symbol. For example, oxygen has 6 valence electrons, so we write the symbol O for oxygen and surround it with 6 dots: unpaired electrons are represented as single dots, and paired electrons as double dots. Placement of single or double dots around symbol is not critical. Alternatively, we can represent paired electrons as line. That is, we replace double dots as shown below: let's consider other examples. Sodium atom has 11 electrons, but only one is valence electron. The other 10 are inside close shell with Neon electron configuration. Thus, we draw the Lewis Structure for Sodium atom as symbol Na with a single dot: chlorine atom has 17 electrons, but only 7 of these are valence electrons. Thus, we draw the Lewis Structure as: in Ionic Bonds valence electrons are completely transfer. Thus, we write Lewis Structure for NaCl as: as you can see, Chlorine is now surrounded by 8 electrons in N = 3 shell and Sodium has lost its one valence electron in N = 3 shell. Of course, sodium, is still surrounded by 8 electrons of N = 2 shell, but we do not show electrons in inner close shells. For period 2 elements, where all valence electrons of an atom are in s and p orbitals, we find that the Lewis Dot Structure of molecules will often follow the Octet Rule: Octet Rule - Atoms tend to gain, lose, or share electrons until they are surrounded by eight electrons. Using Lewis Dot structures and Octet Rule, we can predict and represent the electronic structure of covalently bonded molecules. For example, when two Chlorine atoms, each with 7 valence electrons, come together to form a diatomic Chlorine molecule, Lewis Structure shows that there will be sharing of two electrons between two Chlorine atoms, which allows both Chlorine to be surrounded by 8 electrons. Of course, hydrogen is period 1 element, with only 1s Orbital, so it has a maximum of two electrons allowed in its valence shell. When two hydrogen atoms come together into a diatomic H 2 molecule, the Lewis Structure shows that there will be sharing of two electrons between two hydrogen, allowing both hydrogen to be surrounded by a closed N = 1 shell of 2 electrons: we can represent the electronic structure and reaction of hydrogen and Chlorine Atoms to form HCl with Lewis structures: for diatomic oxygen, Lewis Dot Structure predict double bond. While the Lewis diagram correctly predicts that there is a double bond between o atoms, it incorrectly predicts that all valence electrons are pair.

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### Representation

#### Table

N5
O (x 3)18
charge1
24

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 the third period of the Periodic Table.

##### 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: figure 1 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: figure 2 demonstrates use of Lewis symbols to show transfer of electrons during formation of ionic compounds.

##### 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 outlined here: determining 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 the 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. Like most electronegative element,ss 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 determined in Step 1.

* 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.

#### Table : Common Compounds

ElementAtomic WeightHydrogen CompoundsOxygen CompoundsChlorine Compounds
Hydrogen1.01H 2H 2 O , H 2 O 2HCl
Helium4.00None formedNone formedNone formed
Lithium6.94LiHLi 2 O , Li 2 O 2LiCl
Beryllium9.01BeH 2BeOBeCl 2
Boron10.81B 2 H 6B 2 O 3BCl 3
Carbon12.01CH 4 , C 2 H 6 , C 3 H 8CO 2 , CO, C 2 O 3CCl 4 , C 2 Cl 6
Nitrogen14.01NH 3 , N 2 H 4 , HN 3N 2 O, NO, NO 2 , N 2 O 5NCl 3
Oxygen16.00H 2 O , H 2 O 2O 2 , O 3<Cl 2 O , ClO 2 , Cl 2 O 7
Fluorine19.00HFOF 2 , O 2 F 2ClF , ClF 3 , ClF 5
Neon20.18None formedNone formedNone formed
Sodium22.99NaHNa 2 O , Na 2 O 2NaCl
Magnesium24.31MgH 2MgOMgCl 2
Aluminum26.98AlH 3Al 2 O 3AlCl 3
Silicon28.09SiH 4 , Si 2 H 6SiO 2SiCl 4 , Si 2 Cl 6
Phosphorus30.97PH 3 , P 2 H 4P 4 O 10 , P 4 O 6PCl 3 , PCl 5 , P 2 Cl 4
Sulfur32.06H 2 S , H 2 S 2SO 2 , SO 3S 2 Cl 2 , SCl 2 , SCl 4
Chlorine35.45HClCl 2 O , ClO 2 , Cl 2 O 7Cl 2
Potassium39.10KHK 2 , K 2 O 2 , KO 2KCl
Argon39.95None formedNone formedNone formed
Calcium40.08CaH 2CaO , CaO 2CaCl 2
Scandium44.96Relatively UnstableSc 2 O 3ScCl 3
Titanium47.90TiH 2TiO 2 , Ti 2 O 3 , TiOTiCl 4 , TiCl 3 , TiCl 2
Vanadium50.94VH 2V 2 O 5 , V 2 O 3 , VO 2 , VOVCl 4 , VCl 3 , VCl 2
Chromium52.00CrH 2Cr 2 O 3 , CrO 2 , CrO 3CrCl 3 , CrCl 2
###### * 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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