Electrons in Covalent Molecules The Lewis Model

Transcription

1 BJECTIVES Electrons in Covalent Molecules The Lewis Model To develop a conceptual model of Lewis structures To become proficient at developing Lewis structures To investigate resonance structures INTRDUCTIN Electrons are the glue that hold atoms together in a molecule. The characteristic behavior of electrons has been exploited by physicists and chemists to develop simplistic models of molecular electronic structure. These simple models translate to a systematic set rules that can be used to describe and predict certain chemical and physical properties of molecules. The Lewis model is a widely used and very effective model that describes the arrangement of electrons in molecules. Understanding the Lewis model and learning to generate realistic Lewis models will be one of the most important skills you take from general chemistry to your advanced chemistry courses. As you learn the Lewis model keep in mind that it is a simplistic model and does not explain all of the observable properties of molecules. We use it because it is simple and very effective for many molecules. The Lewis model is based on the idea that atoms tend to achieve an octet of electrons. You have already learned that ionic compounds do this by losing or gaining electrons (metal- nonmetal). When two nonmetal atoms are combined their electronegativity differences are not large enough to result in the transfer of electrons. In order to reach the octet, these atoms will share electrons in bonds. Bonds in which electrons are shared are called covalent as opposed to an ionic bond where electrons are transferred. Remember that the Lewis model can only be used to depict electron arrangements in covalently bonded molecules. INSTRUCTINS Each team will select a role for each person from the list below and record the names and roles of each person. Answer the questions below. Do not skip any questions. Each of you has a copy of the exercise you should complete but you will only turn in the manager s copy along with the recorder s report and the reflector s summary of the group dynamics. Manager Name: Recorder Name: Spokesperson/Ambassador Name: Technician Name: Manages the group. Ensures that members are fulfilling their roles, that the assigned tasks are being accomplished on time, and that all members of the group participate in activities and understand the concepts. Your instructor will respond to questions from the manager only (who must raise his or her hand to be recognized). The instructor responds only to questions from the manager (who must raise his or her hand to be recognized). Records the names and roles of the group members at the beginning of each day. Records the important aspects of group discussions, observations, insights, etc. The recorder's report is a log of the important concepts that the group has learned. The recorder s report will be collected with the manager s copy of the day s exercise. Presents oral reports to the class. These reports should be as concise as possible; the instructor will normally set a time limit. The spokesperson will also act as ambassador to exchange information with other ambassadors when requested in the exercise. Performs all technical operations for the group, including the use of a calculator or computer. Unless otherwise instructed, only the technician in each group may operate equipment such as this. If there are two technicians in a group, the technicians do calculations independently and then compare answers. The technician has sole control of the group s textbook. The technician in this activity is responsible for building molecular models. Page 1 of 9

2 THERY Lewis diagrams are used to represent electrons in covalent molecules. In this model a line drawn between two atoms designates shared electrons and unshared electrons are represented by dots drawn next to the atom. or example consider the Lewis structure of molecular chlorine (igure 1). rom its position in the periodic table we know that each chlorine atom has 7 electrons. In order to obtain an octet each chlorine needs to gain one electron. The molecule is formed by sharing two electrons (one from each chlorine). Both electrons in the bond are counted towards the octet of each atom. The procedure to building Lewis structures is based upon the systematic application of a series of guidelines or rules that explain the nature of electrons in atoms. These rules can only be applied to atoms in the s and p- blocks of the periodic table. The bonding in the transition metals (d- block) and the lanthanides (f- block) is cannot be explained by this simple model. MDEL: GUIDELINES T USE IN DRAWING LEWIS DT STRUCTURES igure 1. Lewis structure of molecular chlorine ( 2 ) H can form only one bond. It is always a terminal atom Halogens or oxygen are usually terminal atoms. When bonds are formed between halogens and oxygen, oxygen is the terminal atom. The atom with the lowest electron affinity is the central atom in a molecule or ion. Count the number of valence electrons in each element. Convert this to valence electron pairs. Place one pair of electrons between bonded atoms (this forms a σ bond) Subtract the number of bonds made from the valence electron pairs. The remaining electrons are used to form lone pairs and π bonds. Place lone pairs about each terminal atom (except H) to satisfy the octet rule If electron pairs are still remaining, assign them to the central atom. If the central atom is from the third period or higher it can accommodate more than 8 electrons. If the central atom does not have 8 electrons around it, convert one or more terminal atom lone pairs to π bond pairs. Remember C, N, and S usually form π bonds. Check the formal charges in your structure. If more than one structure is possible choose the one with the least separation of charge and the overall lowest formal charges. It is clear from these guidelines that not all Lewis structures are as simple as the one for 2. Listed below are some hints that should help you to avoid common mistakes. Some molecules can form double or triple bonds, which represent more than four or six shared electrons between two atoms. Some molecules will have more than one valid Lewis structure and we need to develop a method for determining the most important structure. Many atoms can violate the octet rule. The more electronegative atom will carry the higher formal charge. The Lewis structure does not show the three- dimensional arrangement of the bonds in the molecule. (We will explore Molecular Shapes in the next exercise.) Page 2 of 9

3 MDEL: APPLICATIN THE GUIDELINES T BUILDING STRUCTURES EXAMPLE 1: PHSPHRUS TRICHLRIDE (PCL 3) Add up all of the valence electrons on all atoms. Identify the central atom. (least electronegative) Place the central atom in the center of the molecule and add all other atoms around it. Place one bond (two electrons) between each pair of atoms. 5 + (3 x 7) = 26e - P P P 26 6 = 20e - Complete the octets for all atoms connected to the central atom. P = 2e - P Complete the octet for the central atom. 2 2 = 0e - (1) Why is phosphorus the central atom in this structure? (2) How can you use the periodic table to choose the central atom? Page 3 of 9

4 EXAMPLE 2: BRMINE PENTALRIDE (BR5) Add up all of the valence electrons on all atoms. Identify the central atom. (least electronegative) 7 + (5 x 7) = 42e - Br Place the atom in the center of the molecule and add all other atoms around it. (Bromine is in the fourth period and can violate the octet rule) Br Place one bond (two electrons) between each pair of atoms. Br = 32e - Complete the octets for all atoms connected to the central atom. Br = 2e - Place any leftover electrons on the central atom, even if doing so results in more than an octet. Br 2 2 = 0e - (3) What is the octet rule? (4) Is there an atom in this structure that violates the octet rule? If so, which one? Page 4 of 9

5 EXAMPLE 3: CARBNATE (C 3 2- ) Add up all of the valence electrons on all atoms. - or an anion add electrons equal to the negative charge - or a cation subtract electrons equal to the pos. charge Identify the central atom. (least electronegative) Place the central atom in the center of the molecule and add all other atoms around it. (Bromine is in the fourth period and can violate the octet rule) Place one bond (two electrons) between each pair of atoms. Complete the octets for all atoms connected to the central atom. If there are not enough electrons to give the central atom an octet, try multiple bonds. Use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds. Evaluate formal charges. C (=) 6 6 = 0 C (C) 4-4 = 0 C (- ) 6 7 = (3 x 6) + 2 = 24e - C 24 6 = 18e = 0e = = = = -1 Look for resonance structures. Deduce most stable using formal charges. (These are equivalent) More electronegative atoms carry higher formal charges. The most stable structures will have the least overall formal charge. C C (5) Why isn t the structure in the 5 th step a valid structure? Page 5 of 9

6 KEY QUESTINS (6) You will notice that the rules are applied slightly differently to each of the examples. What rules are applied to every example? What special considerations are made in examples 2 and 3? Why? (7) Why is there a double bond in carbonate? (8) Which bond in the examples will be the shortest bond? Why? (9) How is the formal charge calculated? Why is it important to know the formal charge when considering resonance structures? Page 6 of 9

7 EXERCISES (10) Draw appropriate Lewis diagrams for the following molecules Molecular ormula Lewis Structure Molecular ormula Lewis Structure Molecular ormula Lewis Structure H 2 NH 3 Br 3 B 3 P 4 3- Xe 2 Sn 4 P 5 S 6 CH 4 S 4 I 5 Draw the Lewis diagrams for ammonia, nitrate and nitrite. Calculate the formal charge for every atom. Page 7 of 9

8 PRBLEMS (11) Complete the table below for acetylene (HCCH), ethylene (H 2 C=CH 2 ) and ethane (H 3 C- CH 3 ) Compound: Lewis Structure Calculated HCC bond angle Calculated C- C Bond Length ethane 107.9º 1.53 Å ethylene 117.7º 1.34 Å acetylene 107.9º 1.21 Å (12) Explain the trend you observe in the bond lengths in terms of the Lewis structures. (13) What observations can you make about the bond angles in these three molecules? Page 8 of 9

9 (14) Benzene, C 6 H 6, is in an aromatic ring. The carbon atoms are arranged in a hexagonal ring. Each carbon is bound to two other carbons in the ring and a hydrogen atom. The bond lengths between each carbon are listed in the table below. C1- C2 140 pm C2- C3 140 pm C3- C4 140 pm C4- C5 140 pm C5- C6 140 pm C6- C1 140 pm (15) Draw two resonance structures for benzene and calculate the formal charges on each carbon. (16) Use your resonance structures to explain the observed C- C bond lengths with respect to the bond lengths you reported in problem (11). Page 9 of 9