CHEM 130 Exp. 8: Molecular Models

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1 CHEM 130 Exp. 8: Molecular Models In this lab, we will learn and practice predicting molecular structures from molecular formulas. The Periodic Table of the Elements IA 1 H IIA IIIA IVA VA VIA VIIA Li 11 Na 19 K 37 Rb 55 Cs 87 Fr Be 1 Mg 0 Ca 38 Sr 56 Ba 88 Ra 1 Sc 39 Y La series Ac series 57 La 89 Ac Ti 0 Zr 7 Hf 58 Ce 90 Th 3 V 1 Nb 73 Ta 59 Pr 91 Pa Cr Mo 7 W 60 Nd 9 U 5 Mn 3 Tc 75 Re 61 Pm 93 Np 6 Fe Ru 76 Os 6 Sm 9 Pu 7 Co 5 Rh 77 Ir 63 Eu 95 Am 8 Ni 6 Pd 78 Pt 6 Gd 96 Cm 9 Cu 7 Ag 79 Au 65 Tb 97 Bk 30 Zn 8 Cd 80 Hg 66 Dy 98 Cf B 13 Al 31 Ga 9 In 81 Tl 67 Ho 99 Es C 1 Si 3 Ge 50 Sn 8 Pb 68 Er 100 Fm N 15 P 33 As 51 Sb 83 Bi 69 Tm 101 Md O 16 S 3 Se 5 Te 8 Po 70 Yb 10 No F 17 Cl 35 Br 53 I 85 At 71 Lu 103 Lr VIIIA He 10 Ne 18 Ar 36 Kr 5 Xe 86 Rn Part I: Valence Electrons To draw a Lewis Structure, we must know how many valence electrons each atom possesses. For main group atoms, we can get this information by looking at the periodic table. Rule 1: A main group atom has a number of valence electrons equal to the position of its main group column from the left of the periodic table (first main group column has one valence electron, second main group column has two valence electrons, etc.). Example: Lithium (Li) is in the first column, so it has one valence electron. The columns of the periodic table are often labeled at the top with a number. This number is called the group number, and is written in roman numerals. Rule : A main group atom will have a number of valence electrons equal to its group number. Determining the number of valence electrons for transition metal, lanthanide, and actinide atoms is much more complicated because of the d and f orbitals. In this course, we will confine ourselves to main group atoms, and we will ignore electrons in d and f orbital when counting valence electrons.

2 Exercise 1: Complete the first column only of Table 1 by writing the number of valence electrons in each of the atoms listed. Part II: Electron Dot Pictures We can write an abbreviated picture of a main group atom with its valence electrons by using the atomic symbol surrounded by dots. An example for neon is shown below. Ne Neon has 8 valence electrons, so we surround the atomic symbol for neon (Ne) with 8 dots. Notice that the dots are written in pairs at four positions around the symbol: right, left, bottom, and top. Rule 3: When writing the electron dot picture of an atom, add one electron dot to all four positions before adding a second electron dot to any positions. An example is shown below. Nitrogen (5 valence electrons) N N N N N 1st electron nd electron 3rd electron th electron 5th electron Atoms in groups I, II, and III will have empty positions; this will influence their bonding and chemistry. Exercise : Complete the second column only of Table 1 by writing the electron dot pictures of the atoms listed. Part III: Bonding Positions and Lone Pairs Examining the electron dot pictures in Table 1, we can see that atoms can have pairs of electrons at some positions and single electrons at other positions. The paired electrons are called lone pairs; in future chemistry courses, especially organic chemistry, lone pairs have a strong effect on the chemistry of an atom and its molecules. The single electrons are called unpaired electrons. Usually, an atom wants to have the same number of bonds attached to it as it has unpaired electrons in its electron dot picture. Exercise 3: Complete columns 3,, and 5 of Table 1 by writing in the number of unpaired electrons, lone pairs, and desired bonds, for each of the listed atoms. Part IV: Lewis Structures Step 1: Count the Electrons

3 1. Count the number of valence electrons in the molecule by adding up the number of valence electrons on all the atoms in the formula, and adjusting for the overall charge: Valence Electrons = [sum of valence electrons on atoms] [overall charge]. Calculate 6N +, where N = number of non-hydrogen atoms. 6N + is the number of valence electrons required to have all single bonds. 3. Compare Valence Electrons to 6N + to determine if there are multiple bonds: If [6N + ] [ Valence Electrons] =, then 1 double bond; If [6N + ] [ Valence Electrons] =, then double bonds or 1 triple bond. For each additional Δ, there will be an additional multiple bond in the structure. (Note: in molecules that contain boron, beryllium, or aluminum, a Δ might indicate that these atoms have less than a full octet.) Step : Draw a Skeleton for the Molecule A skeleton has all the atoms connected together by single bonds, but lacks lone pairs and multiple bonds. It can be difficult to guess the correct skeleton for a molecule, and often there are several possible skeletons, all of which are correct. Try to get clues about the skeleton from the order that the atoms are written in the molecular formula. The following set of guidelines may also be useful, but there are exceptions to all of these guidelines so use them with care. 1. Look for symmetry in the skeleton.. Generally, the less electronegative atoms are closer to the center of the molecule (except hydrogen). 3. When there are several oxygens in the formula along with a different atom, put the different atom in the center and surround it with oxygens.. Avoid oxygen to oxygen bonds if possible. Obviously, this won t be possible for O or O3. 5. Try to give each atom its desired number of bonds. This was determined for each atom in Part III. 6. Add hydrogens last. Each hydrogen should have only one bond. Step 3: Complete the Skeleton by Adding Multiple Bonds and Lone Pairs 7. Add the number of multiple bonds calculated in Part I to the skeleton. Consider the desired number of bonds for each atom when deciding where to place multiple bonds. 8. Determine the remaining number of electrons: Remaining Electrons = [ Valence Electrons] x [number of bonds]

4 9. Complete the octets of the atoms in the structure by distributing the Remaining Electrons as lone pairs, starting with the most electronegative atoms first. Part V: Molecular Structures We will use the Lewis Structures to build models and look at the shapes of atoms in molecules. Most of the molecules that we will deal with in this class follow the octet rule. Because their electrons are generally paired, we will be building molecules with pairs of electrons around each atom. To represent these atoms, we will use balls with equally spaced holes into which we will insert our "electrons". Noteworthy exceptions are hydrogen and helium. Because they desire only valence electrons, and these electrons will be paired, these atoms will be represented by single-holed balls. The pairs of electrons that form bonds between atoms will be represented by a short stick. A single stick represents one bond, which is made of two shared electrons. To form a double bond, it will be necessary to make two connections between two atoms. You will find, as you build your models that you cannot use the sticks to make multiple bonds because they are inflexible. To form a double bond, use two pieces of flexible spring to make the connections. Each spring represents one bond (two electrons), so two springs are two bonds (a double bond, electrons.) A triple bond is formed by three springs, and contains six electrons. A lone pair is a bond that only has an atom on one end. If we wish to represent lone pairs, use bonds but only connect an atom on one end. This will be useful for Part VI below. Part VI: VSEPR (Valence Shell Electron Pair Repulsion) VSEPR theory is used to predict the shapes around atoms in a molecule. VSEPR assumes that pairs of electrons repel each other. Because the bonds around atoms are pairs of electrons, they will repel each other to be as far apart as possible, which will force the atom to have a particular shape. We can see how many pairs of electrons are around a given atom in a molecule by looking at the Lewis Structure. There are two different descriptions of the structure around an atom: the geometry and the shape. The geometry refers to the arrangement of all of the pairs of electrons around an atom, including lone pairs. The shape refers only to the arrangement of bonds around an atom, ignoring lone pairs. Sometimes the two will be the same; often they will be different. When a molecule has multiple bonds in its Lewis Structure, we count the multiple bond as a single pair of electrons for VSEPR. The geometries and shapes around atoms for a variety of different possible arrangements are given in the following table:

5 Total Electron Groups 3 Bonding Groups 3 3 Lewis Diagram Electron Geometry Linear Trigonal Planar Electron Geometry Model Bond Angle 180 o 10 o o o o Molecular Shape Linear Trigonal Planar Trigonal Pyramid Bent Molecular Shape Model Example BeF BF3 CH NH3 HO

6 Date: CHEM 130 NAME: Exp. 8: Molecular Models Prelab Complete the following table: Atom Valence Electrons Electron Dot Picture Unpaired Electrons Lone Pairs Desired Bonds Beryllium (Be) Boron (B) Carbon (C) Nitrogen (N) Oxygen (O) Fluorine (F) Aluminum (Al) Silicon (Si) Phosphorus (P) Sulfur (S) Chlorine (Cl) Selenium (Se)

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