Honors Chemistry Unit 6 ( )
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1 Honors Chemistry Unit 6 ( ) Lewis Dot Structures VSEPR Structures 1
2 We are learning to: 1. Represent compounds with Lewis structures. 2. Apply the VSEPR theory to determine the molecular geometry of a compound. We are looking for: 1. Draw Lewis structures for compounds and polyatomics based upon octet, formal charge, and resonance. 2. Determine the molecular geometry of a Lewis structure using VSEPR. 2
3 Name Family Name Electron configuration Valance Shell Electrons Electron Dot Notation Oxidation # Alkali Metals Alkaline Earth Metals Boron Family Carbon Family Nitrogen Family Oxygen Family Halogens Noble Gas 3
4 1. Determine if bond will be ionic or covalent a. Metal + nonmetal Ionic b. Nonmetal + nonmetal Covalent Introduction to Lewis Structures 2. If Ionic: a. Draw Lewis dot notation for each element in the formulas. b. Draw arrows showing the electrons transferring from metal to nonmetal. c. Show the results of the transfer by writing the element symbols again but in brackets and the charge they each have on the outside of the brackets. Also, show the electron dots on the nonmetal. 3. If Covalent: a. Determine the total number of valence electrons available by using all of the elements in given formula. b. Put the single element in the middle. c. Draw a line (single bond; counts as 2 electrons) connecting the other elements to the center element. d. Place the remaining valence electrons (dots) on the surrounding elements to give them an octect (8) of electrons; H will only form a single bond (needs a duet of electrons). e. If you run out of electrons, try moving unshared ones to make double or triple bonds to fulfill the octet rule. f. If there are any leftover electrons (dots) place them on the center element. Examples : (also name each compound) Na2O Ionic or Covalent? CH4 Ionic or Covalent? AlCl3 Ionic or Covalent? NBr3 Ionic or Covalent? CO Ionic or Covalent? 4
5 Resonance Resonance occurs when more than one valid Lewis structure can be drawn for a compound or polyatomic ion. It usually occurs when you have to make a double bond to ensure that the central atom has a full octet. Let s try with sulfur dioxide. You can see after arranging the elements, making the single bonds, and arranging the rest of the electrons, that sulfur only has 6 electrons assigned to it. In order for sulfur to have the full octet, one of the outer oxygen atoms must donate a pair of electrons to make a double bond. Do you think it matters which oxygen atom donates? It doesn t either oxygen atom can form that double bond. This is resonance the double bond could be drawn in either place and you would still have a valid Lewis structure. Let s try to draw all of the valid Lewis structures for the carbon dioxide. Formal Charge Formal charge is a method of evaluating how likely a Lewis structure is. Formal charge is calculated for each atom in a Lewis structure by taking the normal number of valence electrons in an atom and subtracting the number of electrons assigned to it in the drawn Lewis structure. Bonding lines will count only as one for the element when counting for formal charge. Also, C, N, O,F must obey the octet rule even if that means they end up with a nonzero formal charge. 5
6 Let s try this with sulfur dioxide. In a molecule, the most likely Lewis structure is one where each atom has a formal charge of zero. Is that the case for sulfur dioxide? NO! That means another Lewis structure is more likely and that we need to rearrange our electrons. A positive formal charge means that the atom wants to gain another electron. A negative formal charge means that the atoms wants to lose an electron. Since sulfur has a +1 formal charge and oxygen has a -1 formal charge, oxygen can donate a pair of electrons to make another bond and make the formal charges zero. Let s see this. Let s try our formal charges with sulfur trioxide. For ions, the formal charges should add up to the charge of the ion. Negative formal charges are best placed on the outside atoms. Let s try our formal charges with phosphate. The last step when you draw a Lewis structure should be to assign the formal charges to check your work. If the formal charges are not minimized, you need to fix your Lewis structure. 6
7 Lewis Structures If single bonds are used and not all atoms have an octet (Except IA, IIA, and IIIA) Try unshared pairs then double and triple bonds: C2H4 H C C H H H CCl2O Cl C O Cl C2Cl2 Cl C C Cl 1. O 2 4. HClO 2 2. NCCN 5. N 2H 4 3. O 3 7
8 Summary of Writing Covalent Lewis Structures 1. Using whatever information available, write the symbols for the elements in the correct arrangement. 2. Calculate the total number of valence electrons. Remember to adjust amount for polyatomic ions! 3. Place one pair of electrons (single line) between each pair of bonded atoms. 4. Beginning at the outside of the formula, place the remaining electrons in pairs until there are eight electrons around each atom (two for hydrogen) or all electrons have been used. a. If there are extra electrons, place them on the central atom. b. Atoms in the 3 rd period and beyond can have more than eight electrons around them and can form more than 4 bonds, but elements in the 2 nd period cannot. **also of note some Noble gases can form covalent bonds (with the exceptions of He & Ne).** 5. If not enough electrons are available to give all atoms (except H) an octet, move unshared pairs to form double or triple bonds. a. However, Be, B, and other Group IIA elements may have fewer than eight electrons. 6. Check the formal charges of the atoms. It should be at the lowest possible, even if the octet rule is not fulfilled (exception: C, N, O, F must have an octet and H two electrons). Want to try to have a formal charge of zero for the central atom. 7. Examine your Lewis structure to see if resonance structures are needed. (Is more than one position possible for multiple bonds) 8. Check your answer. Does the Lewis structure show: a. the correct number of atoms? b. the correct number of electrons? c. the right number of electrons around each atom? d. the minimum number of formal charges? e. The brackets around and the charge on it if it is a polyatomic ion? 8
9 Lewis Structures Formula (Give the name of the compound under the formula) Ionic Or Covalent? Number Of Valence Electrons Lewis Structure 1. Rb 3N 2. NF 3 3. SCl 2 4. InBr 3 5. PCl 5 6. SF 6 7. SeCl 4 9
10 Formula **(You do NOT have to name these)** Ionic Or Covalent? Number Of Valence Electrons Lewis Structure 8. FeBr 3 9. NH 2Cl 10. CuCl 2 10
11 Lewis Structures for Polyatomic Species **Also, Name Each Ion!** When the charge is (+) Subtract the charge value from the total number of valence electrons When the charge is ( ) Add the charge value to the total number of valence electrons Example: NH 4 + Atom How many x # of e - s Total for each element N 1 5 = 5 H 4 1 = 4 1+ losing 1 = -1 Total valence electrons = 8 Example: PO 4 3- Atom How many x # of e - s Total # of electrons P 1 5 = 5 O 4 6 = gaining 3 = 3 Total valence electrons = 32 11
12 Example: BrO 3 - Atom How many x # of e - s total # of electrons Br 1 7 = 7 O 3 6 = gaining 1 = 1 Total valence electrons = CO 3 Atom How many x # of e - s total # of electrons C 1 4 = 4 O 3 6 = gaining 2 = 2 Total valence electrons = 24 12
13 Show math calculations and then draw the Lewis structure 1. H 3O + 2. NO 3-3. CN - _cyanide 4. SO OH - 13
14 14
15 VSEPR Theory (Valence Shell Electron Pair Repulsion) Predicts the molecular shape of the resulting molecule Electrons on central atom arrange themselves as far apart as possible Unshared pairs on the central atom repel the most Shared pairs on the central atom repel the least To get the shape ONLY LOOK AT WHAT IS CONNECTED TO THE CENTRAL ATOM!!!!! 15
16 A A A A A A A A A A A A A **As mentioned previously, some Noble gases can form covalent bonds (with the exceptions of He & Ne).** 16
17 17
18 1) Linear a. linear= 0 lone pair Draw: VSEPR Practice (molecular geometry) SiS2 CO2 2) Trigonal Planar a. trigonal planar= O lone pair Draw: BF3 H2CO b. bent= 1 lone pair Draw: SO2 O3 18
19 3) Tetrahedral a. Tetrahedral= 0 lone pair Draw: CH4 (SO4) 2- b. trigonal pyramidal= 1 lone pair Draw: NH3 c. bent= 2 lone pairs Draw: H2O <<109 19
20 4) Trigonal bipyramidal a. Trigonal bipyramidal= 0 lone pairs Draw: PCl5 b. See-saw= 1 lone pair Draw: SF4 c. T-shaped= 2 lone pairs Draw: ClF3 20
21 d. linear= 3 lone pairs Draw: XeF2 I3-5) Octahedral a. octahedral= 0 lone pair Draw: SF6 b. square pyramidal= 1 lone pair Draw: IF5 21
22 c. square planar= 2 lone pairs Draw: XeF4 (BrF4) - 22
23 Lewis Structure/VSEPR Problems Formula (write the name under the formula) Lewis Structure Number of Bonded Atoms Number of Lone Electron Pairs Molecular Geometry (name & sketch) 1) SiF4 2) BBr3 3) NF3 4) H2O 5) AsCl5 6) SF6 23
24 Lewis Structures and VSEPR Practice **Only do the VSEPR Shape on Covalent Structures** Formula (write the name under the formula) 1) KCl Ionic or Covalent Lewis Structure VSEPR Shape Drawing Shape Name 2) BH3 3) CCl4 4) CO3 2-5) SF2 6) Na2O 7) AsBr3 8) H3O + 24
25 Name: Lewis Structure/VSEPR Practice WS Formula Compound Name Lewis Structure Resonance? Yes or No VSEPR Structure (include at least 1 angle) VSEPR Name 1) SiI4 2) AlN 3) NO 3-4) BeCl 2 5) PF 3 25
26 Formula Compound Name Lewis Structure Resonance? Yes or No VSEPR Structure (include at least 1 angle) VSEPR Name 6) TeH 6 7) Rb 2O 8) PI 5 9) PO ) XeF 2 26
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