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Lewis Dot Structure For All Elements

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Nasa's 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 Lewis structures of these molecules? Hcn: + = 10 H 3 CCH 3: + = 14 HCCH: + = 10 NH 3: + = 8 Draw skeleton and connect atoms with single bonds. Remember that H is never central atom: HCN: six electrons located on N H 3 CCH 3: no electrons remain HCCH: no terminal atoms capable of accepting electrons. Nh 3: no terminal atoms capable of accepting electrons. Hcn: no electrons remain H 3 CCH 3: no electrons remain HCCH: four electrons placed on carbon NH 3: two electrons placed on nitrogen where needed to rearrange electrons to form multiple bonds in order to obtain octets on each atom: HCN: form two more C - N bonds H 3 CCH 3: all atoms have correct number of electrons HCCH: form triple bond between two carbon atoms NH 3: all atoms have correct number of electrons carbon soot has been know to man since prehistoric times, but it was not until fairly recently that molecular structure of 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. Entire classes of compounds, including spheres and tubes of various shapes, were discovered based on C 60. This type of molecule, called fullerene, consists of a complex network of single - and double - bond carbon atoms arranged in such a way that each carbon atom obtains full octet of electrons. 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. Xenon is a noble gas, but it forms a number of stable compounds. We examined XeF 4 earlier. What are the Lewis structures of XeF 2 and XeF 6? 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. Step 2: Draw skeleton joining atoms by single bonds. Xenon will be the central atom because fluorine cannot be central atom: 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.

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Lewis Structures for Polyatomic Ions

Table

Group NumberIon ChargeExamples
11+Li+, Na+, K+
22+Mg2+,Ca2+, Ba2+
133+Al3+
162-O2-, S2-
171-F-, Cl-, Br-, I-

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 a single bond. Each atom interacts with eight valence electrons: six in lone pairs and two in single bond.


Why Bonds Form

To understand why chemical bonds form, consider a common compound known as water, or H 2 O. It consists of two hydrogen atoms and one oxygen atom. As you can see on the left side of Figure: below, each hydrogen atom has just one electron, which is also its sole valence electron. Oxygen atom has six valence electrons. These are electrons at the outer energy level of the oxygen atom. In water molecule on right in Figure: each hydrogen atom shares a pair of electrons with oxygen atom. By sharing electrons, each atom has electrons available to fill its sole or outer energy level. Hydrogen atoms each have a pair of shared electrons, so their first and only energy level is full. The Oxygen atom has a total of eight valence electrons, so its outer energy level is full. Full outer energy level is the most stable possible arrangement of electrons. It explains why elements form chemical bonds with each other.

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Table2

ChemicalPhysical
Composed of two ions Made of a metal and a non-metal Electric charge of 0 Strong ionic bondsCrystalline structure Can conduct electricity Electrolytes Hard and brittle solids
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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 the third period of the periodic table.

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9.1 Lewis Electron Dot Diagrams

Writing Lewis Structures, NASA's 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 Lewis structures of these molecules? Calculate the number of valence electrons. Hcn: + = 10H 3 CCH 3: + = 14HCCH: + = 10NH 3: + = 8 Draw skeleton and connect atoms with single bonds. Remember that H is never central atom: Where needed to distribute electrons to terminal atoms: HCN: six electrons placed on NH 3 CCH 3: no electrons remainHCCH: no terminal atoms capable of accepting electrons. Nh 3: no terminal atoms capable of accepting electrons Where needed to place remaining electrons on the central atom: HCN: no electrons remainH 3 CCH 3: no electrons remainHCCH: four electrons placed on carbon NH 3: two electrons placed on nitrogen Where needed to rearrange electrons to form multiple bonds in order to to obtain octet on each atom: HCN: form two more C - N bondsH 3 CCH 3: all atoms have correct number of electronsHCCH: form triple bond between two carbon atomsNH 3: all atoms have correct number of electrons check Your Learning Both carbon monoxide, CO, and carbon dioxide, CO 2, are products of combustion of fossil fuels. Both of these gases also cause problems: CO is toxic and CO 2 has been implicated in global climate change. What are Lewis structures of these two molecules?

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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 atoms 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 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 and only needs to form one bond. Transition Elements and inner transition Elements also do not follow the Octet Rule since they have d and F Electrons involved in their Valence shells. Group 15 Elements such as nitrogen have five Valence Electrons in the 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 Octet by forming two Covalent bonds:


Exception 2: Incomplete Octets

The second exception to the Octet Rule is when there are too few valence electrons that results in an incomplete Octet. There are even more occasions where the Octet Rule does not give the most correct depiction of molecule or ion. This is also the case with incomplete events. Species with incomplete octets are pretty rare and generally are only found in some beryllium, aluminum, and boron compounds including boron hydrides. Let's take a look at one such hydride, BH 3. If one were to make a Lewis Structure for BH 3 following basic strategies for drawing Lewis structures, one would probably come up with this Structure Figure: problem with this structure is that boron has an incomplete Octet; it only has six electrons around it. Hydrogen atoms can naturally only have only 2 electrons in their outermost shell, and as such, there are no spare electrons to form double Bond with boron. One might surmise that failure of this structure to form complete octets must mean that this Bond should be ionic instead of covalent. However, boron has electronegativity that is very similar to hydrogen, meaning there is likely very little ionic character in hydrogen to boron bonds, and as such, this Lewis Structure, though it does not fulfill the Octet Rule, is likely the best structure possible for depicting BH 3 with Lewis theory. One of things that may account for BH 3's incomplete Octet is that it is commonly a transitory species, forming temporarily in reactions that involve multiple steps. Let's take a look at another incomplete Octet situation dealing with boron, BF 3. Like with BH 3, initial drawing of the Lewis Structure of BF 3 will form a structure where boron has only six electrons around it. Figure: If you look at Figure: you can see that fluorine atoms possess extra lone pairs that they can use to make additional bonds with boron, and you might think that all you have to do is make one lone pair into bond and the structure will be correct. If we add one double bond between boron and one of the fluorines, we get the following Lewis Structure Figure: each fluorine has eight electrons, and boron atom has eight as well! Each atom has perfect Octet, right? Not so fast. We must examine the formal charges of this structure. Fluorine that shares double bond with boron has six electrons around it. This is one less electron than the number of valence electrons it would have naturally, so it has a formal charge of + 1. Two flourines that share single bonds with boron have seven electrons around them. This is the same amount as the number of valence electrons they would have on their own, so they both have a formal charge of zero. Finally, boron has four electrons around it.

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Double and Triple Bonds

Structural formula is a way of showing the location of atoms or ions relative to one another in molecule, while also showing the number and location of bonds between them. This can tell you many things about the compound. It tells you what kind of atoms are involve, number of them, how they are arranged and bonds between atoms. The steps to writing Lewis dot Structures for compounds are simple. Determine type and number of atoms in molecule. Write Lewis dot structure for each individual atom. Connect atoms by electron pair bonds so that each atom has a full octet. If you have carbon in your molecule, it is always in the middle. Hydrogens are usually on the outside. Double - check your work and make sure every atom has eight electrons and no more.

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Electron-deficient Molecules

The Octet rule is a bonding theory used to predict molecular structure of covalently bonded molecules. According to rule, atoms seek to have eight electrons in their outeror valenceelectron shells. Each atom will share, gain, or lose electrons to fill these outer electron shells with exactly eight electrons. For many elements, this rule works and is a quick and simple way to predict the molecular structure of molecule. But, as saying go, rules are made to be break. And the octet rule has more elements breaking rules than following it. While Lewis Electron dot structures help determine bonding in most compounds, there are three general exceptions: molecules in which atoms have fewer than eight electrons; molecules in which atoms have more than eight electrons; and molecules with an odd number of electrons


Exceptions to the Octet Rule

Although the octet rule has allowed us to draw almost every conceivable Lewis structure, there are certain molecules that do not obey the octet rule. In this section, we will point out the most common exceptions. Boron and aluminum compounds commonly place only six electrons around the metal center. For example, AlH 3 has only six electrons on Al. Compounds with less than an octet of electrons around each atom are called electron deficient. Boron and aluminum compounds are frequently electron deficient while compounds involving most other elements are not. The reason why boron and aluminum can form electron deficient compounds has to do with their low electricity. Because both atoms are not very electronegative, they are not terribly unhappy when they have fewer electrons than they require for full octets. While boron and aluminum can have less than full octet, some atoms like phosphorous and atoms in period three or below on periodic table can exceed their octets. Try to draw a reasonable Lewis structure for either PF 5 or SF 6. You should not find it possible to obey the octet rule on either phosphorous or sulfur. Frequently, textbooks say that atoms like P and S are able to expand their octets by letting extra electrons fill their empty 3d orbitals. Your chemistry course may even require you to memorize this fact. However, this description of bonding in such compounds is completely false. After you have read Molecular Orbital Theory, you should be able to come up with a better reason. Explanation of the expanded octet must wait until then due to its complexity. For now, realize that atoms below period two may expand their octets to accommodate more than eight electrons.


Lewis Structures

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 usually used to indicate share pair of electrons: in the Lewis model, single share pair of electrons is single bond. Each Cl atom interacts with eight valence electrons in total: six in lone pairs and two in single bond.

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Hypervalent Molecules

Another type of exception to Lewis ' approach to bonding is the existence of compounds that possess too few electrons for Lewis ' structure to be write. Such compounds are called electron - deficient compounds. The prime example of an electron - deficient compound is diborane, B 2 H 6. This compound requires at least seven bonds to link its eight atoms together, but it has only 2 3 + 6 1 = 12 valence electrons, which is enough to form only six covalent bonds. Once again, it appears that, as in hypervalent compounds, existence of an electron - deficient compound signifies that a pair of electrons can bond together more than two atoms. Discussion of the quantum mechanical theory of bonding below shows that this is indeed the case. A number of exceptions to Lewis ' theory of bonding have been catalogued here. It has further deficiencies. For example, theory is not quantitative and gives NO clue to how the strength of bonds or their lengths can be assess. In form in which it has been present, it also fails to suggest shapes of molecules. Furthermore, theory offers NO justification for regarding electron pair as central feature of covalent bond. Indeed, there are species that possess bonds that rely on the presence of a single electron. Nevertheless, in spite of these difficulties, Lewis ' approach to bonding has proven exceptionally useful. It predicts when the octet rule is likely to be valid and when hypervalence can be anticipate, and the occurrence of multiple bonds and presence of lone pairs of electrons correlate with chemical properties of a wide variety of species. The Lewiss approach is still widely used as a rule of thumb for assessing structures and properties of covalent species, and modern quantum mechanical theories echo its general content. The following sections discuss how limitations of the Lewis approach can be overcome, first by extending theory to account for molecular shapes and then by developing more thorough quantum mechanical theories of chemical bond.

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Introduction to Lewis Structures

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 connects 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 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 a + 1 charge. Note: octet can expand 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 determined 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 lines 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.


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, {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 occupies 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. 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, {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.

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Key Concepts:

Determine the total number of valence electrons of an element or compound. If a molecule has more than one element, add valence electron of all elements present in the compound. Determine which atom will be the central atom of the Lewis Dot Structure. The central atom is least most electronegative atom in the compound. Remember the trend for electricity on periodic table. Once determine, draw that element by atomic symbol in center and draw single bonds to other atoms. Subtract the full shell of valence electrons of each outer atom from the total number of valence electrons associated with the molecule. Distribute remaining electrons to the central atom as non - bonding pairs form double and triple bonds until the central atom has full octet. Draw nonbonding pairs around outer atoms until they have full octet. Check your work: Ensure that all of your valence electrons and bonds are accounted for.

* 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

How Sciences Interconnect

Table

N5
O (x 3)18
charge1
24

As previously mention, when a pair of atoms share one pair of electrons, we call this a single bond. However, pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet. Double Bond forms when two pairs of electrons are shared between a pair of atoms, as between Carbon and oxygen Atoms in CH 2 O and between two Carbon Atoms in C 2 H 4: triple Bond forms when three electron pairs are shared by a pair of atoms, as in Carbon monoxide and cyanide ion:


The Octet Rule

Lewis was well aware that there are exceptions to the rule of eight, as are exemplify, for example, by PF 5, for which Lewis structure has ten electrons in valence shell of phosphorus, and by BCl 3 for which Lewis structure has only six electrons in valence shell of boron. However, as we have already pointed out, belief in octet rule became so strong over the years, partly due to Langmuir's vigorous promotion of the rule, and partly because exceptions were so few until more recent times, that it came to be regarded as a fundamental law of nature rather than as empirical rule. Thus, molecules that do not obey the octet rule come to be regarded as exceptional and special terms hypervalent and hypovalent were invented to describe them. Reluctance to give up familiar and apparently well - established ideas in the face of contrary experimental facts has often led to the invention of new, and sometimes unnecessary, concepts in an attempt to retain old ideas. Thus, attempts have been made to formulate these exceptions to the octet rule in such a way that they appear to obey the rule, but such concepts become unnecessary once it is recognized that the octet rule as formulated by Lewis is not the law of nature but simply an empirical rule to which there are exceptions. Indeed, octet rule applies strictly only to period 2 elements, C, N, O, and F. For atoms of all other elements, the rule may or may not be obey.

* 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.

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