Name: drawing lewis structures with valence electron counts
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Lewis Structures of Molecular Compounds and Polyatomic Ions

Lewis structures are simplified representations of chemical bonds between atoms. Writing Lewis structures for neutral compounds like ethene or ammonia, or for polyatomic ions such as hydronium or phosphate involves the following sequence of steps.

First, calculate the total number of valence electrons in the molecule. Consider ethene, for example. Each carbon atom has 4 valence electrons, and each hydrogen atom has 1 valence electron; with 2 carbon and 4 hydrogen atoms, ethene has a total of 12 valence electrons.

Next, draw the skeletal structure of the molecule. Generally, the first atom in the chemical formula or the less electronegative atom is placed centrally. This atom is also called the central atom and must be able to form at least two bonds. 

Hydrogen atoms or more electronegative atoms are positioned terminally around the central atom. Here, the carbon atoms occupy the central positions with hydrogen atoms surrounding them.

The next step is octet completion. First, place a pair of electrons between every two atoms to represent bonding pairs. The remaining electrons are distributed as lone pairs to terminal and central atoms to satisfy the duet or octet configurations. In this case, since hydrogen has reached a duet, the remaining electrons are transferred to the carbon atoms. 

Now, use the lone pairs to form multiple bonds between atoms that lack an octet. In ethene, the carbon atoms haven&#8217;t reached an octet yet. Therefore, the lone pair is moved to the bonding region to form a double bond allowing all atoms to reach a stable electron configuration. 

To write the Lewis structure for charged polyatomic ions like hydronium or phosphate, the same approach is followed, with a modification to accommodate the ionic charge. 

To draw the Lewis structure of hydronium, first, calculate the number of valence electrons on the ion. Since the hydronium ion has a positive charge, remove an electron from the total valence electrons, reducing the count to 8. Next, illustrate the molecule by its skeletal structure. 

Now place an electron pair between every atom, followed by the lone pair of electrons on the central atom. This way, hydrogen and oxygen atoms satisfy the duet and octet, respectively. Finally, write the Lewis structure in brackets with a positive charge on its upper-right corner.

lewis structures of molecular compounds and polyatomic ions

To draw Lewis structures for complicated molecules and molecular ions, it is helpful to follow a step-by-step procedure as outlined:

<ol>
	<li>Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.</li>
	<li>Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).</li>
	<li>Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.</li>
	<li>Place all remaining electrons on the central atom.</li>
	<li>Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.</li>
</ol>

For instance, consider SiH4, CHO2&#8211;, NO+, and OF2 as examples for which this general guideline can be applied to determine their Lewis structures.

<ol>
	<li>Determine the total number of valence (outer shell) electrons in the molecule or ion. </li>
</ol>

For a molecule like SiH4, the number of valence electrons on each atom in the molecule is added: 

= [4 valence e&#8722;/Si atom &#215; 1 Si atom] + [ 1 valence e&#8722;/H atom &#215; 4 H atoms] = 8 valence e&#8722;

<h4>CHO2&#8211;</h4>

For a negative ion, such as CHO2&#8211;, the numbers of valence electrons on the atoms are added to the number of negative charges on the ion (one electron is gained for each single negative charge):

= [4 valence e&#8722;/C atom &#215; 1 C atom] + [1 valence e&#8722;/H atom &#215; 1 H atom] + [6 valence e&#8722;/O atom &#215; 2 O atoms] + [1 additional e&#8722;] = 18 valence e&#8722;

<h4>NO+</h4>

For a positive ion, such as NO+, the numbers of valence electrons on the atoms in the ion are added, followed by subtraction of the number of positive charges on the ion (one electron is lost for every single positive charge) from the total number of valence electrons:

= [5 valence e&#8722;/N atom &#215; 1 N atom] + [6 valence e&#8722;/O atom &#215; 1 O atom] + [&#8722;1 e&#8722;] = 10 valence e&#8722;

<h4>OF2 </h4>

OF2 being a neutral molecule, the number of valence electrons are simply added:

= [6 valence e&#8722;/O atom &#215; 1 O atom] + [7 valence e&#8722;/F atom &#215; 2 F atoms] = 20 valence e&#8722;

<ol start="2">
	<li>Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom and connecting each atom to the central atom with a single (one electron pair) bond. (Note that ions are denoted with brackets around the structure, and the ionic charge outside the brackets:) 
	<img alt="Chemical structures diagram showing silane, formate, nitrite, and difluorine molecules." src="/files/core/CDNSource/lessons/1063/1063Image_1.svg" style="height: 80px;" data-altupdatedat="1754904545000" data-alt="Chemical structure diagram: SiH4, HCO3-, NO+, FOF; molecular geometry comparison." /> 
	In cases where several arrangements of atoms are possible, as for CHO2&#8722;, experimental evidence is used to choose the correct one. In general, the less electronegative elements are more likely to be central atoms. In CHO2&#8722;, the less electronegative carbon atom occupies the central position with the oxygen and hydrogen atoms surrounding it. Other examples include P in POCl3, S in SO2, and Cl in ClO4&#8722;. An exception is that hydrogen is almost never a central atom. As the most electronegative element, fluorine also cannot be a central atom.</li>
	<li>Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen) to complete their valence shells with an octet of electrons. (With no remaining electrons on SiH4, its structure is unchanged.) 
	<img alt="Lewis structures of H&#60;sub&#62;2&#60;/sub&#62;SiO, HCO&#60;sub&#62;3&#60;/sub&#62;&#60;sup&#62;&#8722;&#60;/sup&#62;, NO&#60;sup&#62;&#8722;&#60;/sup&#62;, F&#60;sub&#62;2&#60;/sub&#62;O; molecular diagrams." src="/files/core/CDNSource/lessons/1063/1063image_2.svg" style="height: 80px;" data-altupdatedat="1754904545000" data-alt="Lewis structures of molecules: SiH4, HCO3^-, NO^+, and F2O; chemical bonding diagram." /></li>
	<li>Place all remaining electrons on the central atom. 
	<ul>
		<li>For SiH4, CHO2&#8722;, and NO+, there are no remaining electrons. For OF2, of the 16 electrons remaining, 12 are placed, thereby leaving 4 electrons to be placed on the central atom: 
		<img alt="Lewis structure of oxygen difluoride; F-O-F molecule showing bond and electron pairs." src="/files/core/CDNSource/lessons/1063/1063Image_3.svg" style="height: 40px;" data-altupdatedat="1754904545000" data-alt="Fluorine dioxide Lewis structure, diagram showing bond and electron pairs in FO2 molecule." /></li>
	</ul>
	</li>
	<li>Rearrange the electrons of the outer atoms to make multiple bonds with the central atom, to obtain octets wherever possible. 
	<ul>
		<li>SiH4: Si already has an octet, so nothing needs to be done. 
		<img alt="Silane molecular structure diagram, SiH&#8324;, illustrating chemical bonding in inorganic chemistry." src="/files/core/CDNSource/lessons/1063/1063Image_4.svg" style="height: 70px;" data-altupdatedat="1754904545000" data-alt="Silane (SiH4) molecular structure, chemical formula, bond diagram for organic chemistry study." /></li>
		<li>CHO2&#8722;: The valence electrons are distributed as lone pairs on the oxygen atoms, but the carbon atom lacks an octet. 
		<img alt="Resonance structure diagram, HCOO- to formate ion conversion, electron movement shown with arrows." src="/files/core/CDNSource/lessons/1063/1063Image_5.svg" style="height: 80px;" data-altupdatedat="1754904545000" data-alt="Lewis structure, resonance form conversion diagram for formate ion, highlighting electron shift." /></li>
		<li>Hence, one lone pair of electrons is donated from one of the oxygen to the carbon atom forming a double bond. Depending on which oxygen atom donated the electrons, there can be two possible structures, otherwise called resonance structures. </li>
		<li>NO+: For this ion, eight valence electrons are added, but neither atom has an octet. Additional electrons cannot be added since the total electrons are already used up. In this scenario, electrons must be moved to form multiple bonds. The nitrogen atom has two lone pairs of electrons and the oxygen atom has one. 
		<img alt="Nitrosonium ion resonance structures, chemical equation diagram showing electron movement, &#9127;N=O&#9131;&#8314;." src="/files/core/CDNSource/lessons/1063/1063Image_6.svg" style="height: 55px;" data-altupdatedat="1754904545000" data-alt="Resonance structure transition of nitric oxide ion, diagram with electron movement arrow." /></li>
		<li>This still does not produce an octet, so another pair must be moved to form a triple bond. 
		 <img alt="Nitric oxide ion Lewis structure, chemical bonding diagram, valence electrons visualization." src="/files/core/CDNSource/lessons/1063/1063Image_7.svg" style="height: 50px;" data-altupdatedat="1754904545000" data-alt="Nitric oxide chemical structure, NO&#8314; ion, linear molecule, showing triple bond and electron dot diagram." /></li>
		<li>In OF2, nothing changes as each atom already has an octet. 
		<img alt="Molecular structure of oxygen difluoride (OF2) showing Lewis dot structure and bonding diagram." src="/files/core/CDNSource/lessons/1063/1063Image_8.svg" style="height: 40px;" data-altupdatedat="1754904545000" data-alt="Lewis structure of FOF molecule diagram, showing electron dot representation and bonds." /></li>
	</ul>
	</li>
</ol>

This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chemistry 2e, Section 7.3: Lewis Symbols and Structures.</a>

To draw Lewis structures for complicated molecules and molecular ions, it is helpful to follow a step-by-step procedure as outlined:


	Determine the ...

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drawing lewis structures with valence electron counts

Drawing Lewis structures starts with counting valence electrons and building the atom arrangement step by step. Valence electrons are the outer-shell electrons that help show how atoms bond. For ions, a positive charge means one electron is removed for each plus sign, and a negative charge means one electron is added for each minus sign.

The next step is to sketch the skeleton structure. The least electronegative atom usually goes in the center, while hydrogen is almost never central. Fluorine also cannot be a central atom because it is the most electronegative element. Ions are written in brackets, with the charge shown outside the brackets.

After the skeleton is drawn, connect the atoms to the central atom with single bonds. Each single bond represents one electron pair. Then place the remaining electrons as lone pairs on the terminal atoms, except hydrogen, until those atoms have octets when possible. Any electrons still left go on the central atom.

If the central atom still does not have an octet, move lone pairs to make double or triple bonds. This can create resonance structures, which are different valid arrangements of electrons. For the formate ion, CHO2−, one oxygen can donate a lone pair to make a double bond with carbon. That gives two possible resonance forms.

The examples in the model show how the procedure works in real cases. SiH4 has 8 valence electrons total, so its structure does not need more changes once the single bonds are drawn. CHO2− has 18 valence electrons. NO+ has 10 valence electrons, and its atoms must be rearranged into a triple bond to satisfy octets as much as possible. OF2 has 20 valence electrons, and after the lone pairs are placed, its structure stays unchanged because each atom already has an octet.

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Learn the step-by-step method for drawing Lewis structures, counting valence electrons, placing lone pairs, and handling ions, resonance, and multiple bonds.

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Drawing Lewis Structures with Valence Electron Counts

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