U Step 4 Review the Knowledge You Need to Score High in C#.NET

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In the hydrogen molecule the electrons are shared equally Each hydrogen nucleus has one proton equally attracting the bonding pair of electrons A bond like this is called a nonpolar covalent bond In cases where the two atoms involved in the covalent bond are not the same, the attraction is not equal, and the bonding electrons are pulled toward the atom with the greater attraction The bond becomes a polar covalent bond, with the atom that has the greater attraction taking on a partial negative charge and the other atom a partial positive charge Consider for example, HF(g) The fluorine has a greater attraction for the bonding pair of electrons (greater electronegativity) and so takes on a partial negative charge Many times, instead of using a single line to indicate the covalent bond, an arrow is used with the arrow head pointing toward the atom that has the greater attraction for the electron pair:
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The electronegativity (EN) is a measure of the attractive force that an atom exerts on a bonding pair of electrons Electronegativity values are tabulated In general, electronegativities increase from left to right on the periodic table, except for the noble gases, and decrease going from top to bottom This means that fluorine has the highest electronegativity of any element If the difference in the electronegativities of the two elements involved in the bond is great (>17), the bond is considered to be mostly ionic in nature If the difference is slight (<04), it is mostly nonpolar covalent Anything in between is polar covalent Many times the Lewis structure will be used to indicate the bonding pattern in a covalent compound In Lewis formulas the valence electrons that are not involved in bonding are shown as dots surrounding the element symbols, while a bonding pair of electrons is represented as a dash There are several ways of deriving the Lewis structure, but here is one that works well for those compounds that obey the octet rule Draw the Lewis structural formula for CH4O First, write a general framework for the molecule In this case the carbon must be bonded to the oxygen, because hydrogen can only form one bond Hydrogen is never central Remember: Carbon forms four bonds H H C O H H To determine where all the electrons are to be placed, apply the N A = S rule where: N = sum of valence electrons needed for each atom The two allowed values are two for hydrogen and eight for all other elements A = sum of all available valence electrons S = # of electrons shared and S/2 = # bonds For CH4O, we would have: 1C 4H N 8 + 4 (2) = 8 A 4 + 4(1) = 4 S = N A = 24 14 = 10 1O + 8 = 24 + 6 = 14 bonds = S/2 = 10/2 = 5
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Place the electron pairs, as dashes, between the adjacent atoms in the framework and then distribute the remaining available electrons so that each atom has its full octet,
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Bonding 151
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eight electrons bonding or nonbonding, shared or not for every atom except hydrogen, which gets two Figure 114 shows the Lewis structural formula of CH4O
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H H C H O H
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Lewis structure of CH4O
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Lewis structures may also be written for polyatomic anions or cations The N A = S rule can be used, but if the ion is an anion, extra electrons equal to the magnitude of the negative charge must be added to the electrons available If the ion is a cation, electrons must be subtracted As we have mentioned previously, there are many exceptions to the octet rule In these cases, the N A = S rule does not apply, as illustrated by the following example Draw the Lewis structure for XeF4 Answer:
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Each of the fluorines will have an additional three pairs of electrons Only the four fluorine atoms have their octets This process will usually result in the correct Lewis structure However, there will be cases when more than one structure may seem to be reasonable One way to eliminate inappropriate structures is by using the formal charge There is a formal charge associated with each atom in a Lewis structure To determine the formal charge for an atom, enter the number of electrons for each atom into the following relationship: Formal Charge = (number of valence electrons) (number of nonbonding electrons + 1/2 number of bonding electrons) A formal charge of zero for each atom in a molecule is a very common result for a favorable Lewis structure In other cases, a favorable Lewis structure will follow these rules: The formal charges are: 1 2 3 4 Small numbers, preferably 0 No like charges are adjacent to each other, but unlike charges are close together The more electronegative element(s), the lower the formal charge(s) will be The total of the formal charges equals the charge on the ion
Now we will apply this formal-charge concept to the cyanate ion OCN We chose this example because many students incorrectly write the formula as CNO , and then try to use this as the atomic arrangement in the Lewis structure Based on the number of electrons needed, the carbon should be the central atom We will work this example using both the incorrect atom arrangement and the correct atom arrangement Notice that in both structures all atoms have a complete octet
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