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Objectives: 1. To draw Lewis structures for simple molecules and ions. 2. To count electron regions on the central atom of the compound. (These will later be used to predict molecular shape. Discussion: Chemistry deals with the study of the properties, composition, and structure of matter. Knowledge of the chemical and physical properties of samples of matter provides information about their composition. This information helps us describe the structure of the particles that make up the sample. You will do this experiment in three parts. After learning about Lewis electron dot structures you will use this experiment to practice drawing Lewis structures for a variety of the molecules and ions we will be considering in this course. After learning about the shapes of molecules you will use your knowledge of the Lewis structure to predict molecular shape. Finally, after learning about electronegativity and polar bonds you will combine this information with your knowledge about the shape of the molecule to predict its polarity. The first step in predicting a compound's structure is to determine whether the bonding in that species is likely to be ionic or covalent. Any compound containing one or more metal atoms and one or more nonmetal atoms is most likely an ionic compound. lonic bonds, as the name implies, result from the attraction of charged particles (ions) for one another. lons are formed when electrons are gained or lost by a single atom or by a group of atoms (polyatomic ions). In contrast, substances consisting of only nonmetal or metalloid atoms are generally held together by covalent bonds, which result from atoms sharing electrons. When deciding whether a formula corresponds to an ionic or a covalent compound, remember that many common polyatomic ions exist. Examples are ammonium (NH4*), sulfate (SO4²"), carbonate (CO3²), nitrite (NO2"), and hydroxide (OH`). When one or more of these common groups of atoms appears in the formula of a compound, the compound is probably ionic. Thus, the compounds NH4CI, NaNO2, Fe2(SO4)3, KOH, and CU2CO3 are all ionic compounds. However, these polyatomic ions are formed between two nonmetals so the bonding within the ion is covalent. Once you have determined a substance uses covalent bonding, you can write a Lewis electron dot structure for the species. (The covalent species might be the molecules of a covalent compound or the polyatomic ions of an ionic compound.) This procedure was described in an activity sheet and only a brief description will be given below. Steps for Writing Lewis Structures: 1. Count the number of valence electrons available for bonding. 2. Determine the most likely bonding arrangement for the atoms based on the number of bonds they most commonly form. We can determine the number of bonds normally formed by subtracting the number of valence electrons from 8. For example, carbon has 4 valence electrons; it would like to have an octet or 8 electrons and thus would most likely form 4 covalent bonds. When considering bonding arrangements try to draw the structure that is most symmetric. 3. Connect each bonding pair of atoms with a line to represent a covalent bond. Remember that each bond represents a shared pair of electrons. 4. Place three pairs of electrons on each outer atom, except hydrogen, to complete the octets of the atoms involved. (Remember, however, that this is a tentative structure and that there are exceptions to the octet rule.) 5. Assign any remaining electrons to form lone pairs on the inner atoms. 6. Create multiple bonds if necessary to give each non-hydrogen atom a share in eight electrons (an octet of electrons). Each hydrogen atom can only form a single covalent bond. Multiple bonds can be formed only between atoms of C, N, S, and O. Form the multiple bonds by moving a lone pair of electrons from an exterior atom to form an additional bond to the central atom. You may be able to draw more than one Lewis structure for a given molecular formula. If two (or more) structures show approximately the same arrangement of bonds and electrons but have the atoms in different positions, they are isomers of one another. Isomers represent distinctly different chemical compounds. For example, consider a compound with the molecular formula C2H2Cl2. There are 24 valence electrons available for bonding (1 from each H, 4 from each C, and 7 from each CI) for this molecule. Carbon usually forms four bonds, and hydrogen and chlorine usually form only one. This means that the C is most likely to be near the center of the molecule (to form more bonds), and the H and CI will be at the outside edges. The arrangement below is chosen because it is symmetric. CI H C C H CI First place one pair of electrons (represented by a line to indicate the covalent bond) between each pair of adjacent atoms and place three lone pairs on each CI atom. The resulting structure is CI H C H CI: This arrangement uses up 22 of the 24 electrons available. Because there are two central atoms we can place the remaining pair of electrons as an additional bonding pair between the carbon atoms (to give them both their octets and their normal number of bonds). The resulting Lewis structure of C2H2Cl2 is: CI H C C H CI: If you look at the arrangement of atoms it should be reasonably apparent that two additional Lewis structures are possible that do not violate any of our guidelines: CI CI : CI H C==C C C H H CI : H Since we can draw three different structures that all seem to follow the guidelines, what does this mean? Which one is correct? Actually, there is no "correct one" in this case. There are actually three different compounds, with molecular formula C2H2Cl2, each with different physical properties such as boiling points. This is an example of three isomers (different structural formulas) corresponding to one molecular formula. A more complete discussion of isomers will be given in the Introductory Chemistry Il when we study organic compounds. If two (or more) structures show the atoms in the same relative positions but have the electrons shifted (different arrangements of electron pairs), they are resonance structures. The thiocyanate anion, SCN, provides an example of a pair of resonance structures. The two possible structures are shown below: S-CE Note that the arrangement of atoms is the same but the arrangement of bonding pairs and lone pairs is different. Go to the Data Sheet for this experiment (located at the end of this handout) and draw the Lewis structure for the indicated covalent compounds and polyatomic ions. Once you have completed this, in order for me to determine whether you have produced the correct Lewis structure, please fill in the next two columns. Please note that you are counting only the electron groups on the central atom. A group is either a single pair of electrons or multiple pairs of electrons (a double bond or a triple bond). If a group is a single pair of electron it is either a bonding pair connecting the two atoms or a lone pair belonging only to the central atom. When counting the total groups on the central atom each of the above possibilities represents a group, thus in the SCN structures above the central carbon atom has two groups (in the left hand structure it is two double bonds and in the right hand structure it is a single bond and a triple bond but in either case there are two groups). Notice that we are not counting the lone pairs of electrons on the exterior S and N atoms, only the electron groups belonging to the central atom are counted. When counting the bonding groups, any lone pairs of electrons on the central atoms are not counted, that is we count only those electron groups that are used to form bonds (single, double or triple). Part I Part II Part III No. of No. of bonding Compound groups groups Electronic Molecular Polar or (formula and Lewis Structure(s) on the on the Geometry shape nonpolar name) central central atom atom CH4 methane H2O water H3O+ hydronium ion NH3 ammonia NH4* ammonium ion CCI3F trichloro-fluoro methane CO2 carbon dioxide carbonate ion Part I Part II Part III No. of No. of Compound bonding groups Molecular (formula and Lewis Structure(s) on the groups Electronic Polar or on the Geometry Shape name) central nonpolar central atom atom O3 ozone SO2 sulfur dioxide SO3 sulfur trioxide SO2- sulfite ion N2O dinitrogen monoxide PCI3 phosphorus trichloride SF2 sulfur difluroide CH2O formaldehyde

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