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NCSU CH 101 - Molecular bonds and Lewis Structures

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CH 101 1st Edition Lecture 11Outline of Last Lecture I. Types of BondsII. Naming CompoundsIII. Finding the Chemical FormulaIV. Polyatomic IonsA. Oxyanions or (Oxoanions)Outline of Current Lecture I. PolarityII. Molecular BondsA. Formal ChargeIII. Lewis StructuresIV. Drawing out Lewis StructuresCurrent LectureI. Polarity- Polarity is when a more electronegative atom bonds with another atom that is less electronegative. Polarity can be determined by the distance between the elements on the periodic table. For example the bond between K and F is more polar than the bond between Mg and F.Ex) Which bond is more polar? C-O or C-BrC-Br would be more polar because carbon and bromine are further away from eachother on the periodic table than carbon and oxygen.These notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.- There is a certain way to draw the direction of polarity. The polarity goes in the direction that the electrons are being taken, so the arrow points toward the more electronegative element.For example: II. Molecular Bonds- When you talk about “shared pairs” you are really talking about molecular bonds. The more shared pairs there are between two atoms, the stronger the bond. In theory, the stronger bonds have morestrength to pull the two atoms together, making them shorter (double bonds are shorter in length than single bonds). However, when tested it was found that the bonds in an atom are actually the same size because the atoms are being shared throughout the atom but, when it comes to drawing Lewis structures, we still assume that the double and triple bonds are shorter and stronger than thesingle bonds. A. Formal Charge- When two atoms are bonded together they are both receiving the two electrons in the pair. However, when finding the formal charge, you only consider a bond as giving one electron to each atom instead of constantly giving two to each. This number is called the possessed electrons and to find the formal charge you subtract the possessed electrons from the group number of the element. For example, in the molecule:, carbon possesses 4 electrons because for every bond it has it only gets one electron from it, therefore the formal charge of carbon in C2H4 is 4 - 4 = 0.The formal charge for hydrogen in this molecule would also be 0 because 1 – 1 = 0.III. Lewis Structures- The Lewis structure is yet another way to show how electrons move around atoms. Unlike the Bohr model and the Schrodinger model, Lewis structures are used to show how the electrons are distributed throughout a molecule instead of through a single atom. Electrons are represented by dots and electron pairs/bonds are represented as lines, so one line is equivalent to two atoms. The amount of dots/atoms an atom has depends on the amount of valence electrons is has after it has bonded.Octet Rule- This rule describes the atoms that want to have eight electrons in its outer shell. Almost all elements follow this rule except for Hydrogen. Duet Rule – This rule applies to Hydrogen because it only wants two electrons in its outer shell.- Determine the electrons required (ER), valence electrons (VE), shared pairs (SP), and lone pairs (LP) to complete the Lewis structure.ER- Using the octet and duet rule, add up the amount of electrons each atom wants in the moleculetaking into account the subscript. For example the molecule H2O requires 12 electrons because H=2(2) and O=8, the molecule PF3 requires 32 electrons because P=8 and F=3(8).VE-The VE varies depending on what element it is. The number of valence electrons if determined by the group number of the column the element is in. For example the VE for carbon is 4, for hydrogen it is 1, and for oxygen it is 6. Therefore the molecule CH2O has 12 valence electrons to disperse across the molecule.SP- A shared pair means two atoms that are being exchanged back and forth between two atoms. Find the shared pairs by subtracting the valence electrons form the electrons required, then divide by two because you want to know the number of pairs not the number of individual electrons beingshared ER-VE/2=SP. The amount of shared pairs in the compound C6H2 is 7 because the ER=28 the VE=14, 28-14=14, 14/2=7.LP- To find the lone pairs, simply take the number of valence electrons and subtract the total number of electrons sharedVE – SP(2)=LP, so if you find that the SP=7, you would know that there are 14 electrons being shared since there are two electrons in every pair.Ex) Determine the properties of the Lewis structure of a C2H4 molecule.C=2(8), H=4(2), 16 + 8 = 24 ER <- First find the ER, here the octet rule applies to Carbon and the duet rule applies to Hydrogen.C=2(4), H=4(1), 12+4 = 12 VE <- Find the valence electrons by looking at the group number of the elements andmultiplying the group number by the number of atoms of that element there are in the molecule.24 – 12 = 12 SP <- Since the molecule needs each atoms to have a filled outer shell but only has a limited amount of valence electrons, you find the shared electrons through ER-VE=SP to determine how many pairs of electrons are going to be shared between two atoms.12 – 12 = 0 LP <- In this molecule, there are no lone pairs because VE – SP(2) = 0IV. Drawing out Lewis Structures- When you want to take the information you have found about the molecule and actually draw the molecule, start with the SP, starting with the element that has the least amount of atoms in the center and then attaching the rest of the atoms to the center atom/atoms with shared pairs. After you have assembled the “skeleton” of the molecule, calculate the lone pairs and attach those to themolecules that don’t have a full outer shell. Ex) Draw the Lewis structure for the molecule C2H4.You have already found the ER, VE, SP, and LP of this molecule in the previous problem, now just apply these tothe structure of the molecule.C – C <- Notice that Carbon has the least amount of atoms in this molecule so it will become the center of the “skeleton” of the molecule.<- Ignore the number of shared pairs for now and just connect the carbon atoms to the hydrogenatoms, you will want to have the same amount of hydrogen atoms attached to each carbon. <- Now look at the number of shared pairs and see how many more bonds you need to add. Notice that there is a double bond


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