CHM1020 Week 4 | General Chemistry in Chemistry - University of Florida

Week 4 Assignmenta, b, and c Atomic Structure, Chemical Bonding, Lewis Structure, and 3D Molecular Shape Objectives In this lab, you will apply valence bond theory to draw appropriate Lewis structures, use electronegativity differences to classify bonds as ionic, polar covalent, or nonpolar covalent, andapply valence shell electron pair repulsion theory (VSEPR) to predict molecular geometry. We will also use molecular kits to create ball and stick models of molecules of interest to you. Theory Atoms of certain species tend to bond together. An atom is more stable if its valence shell (electrons in its outermost energy level) is similar to that of a Noble Gas (typically eight electrons), and this number is most often achieved when atoms combine. Notice that atoms of elements in the same group on the periodic table tend to have the same number of valence electrons; for example, halogens, Group 7A elements (Figure 1), have seven valence electrons. Figure 1: Dot Diagrams For most elements, a full outer energy level has eight electrons, an octet. The elements in group 8A have a full outer energy level. Helium (He), in period 1, is an exception, requiring only two electrons. Because group 8A elements’ atoms already have a full outer energy level, those elements tend to be nonreactive—they rarely combine with other atoms to form compounds. Atoms can fill their outer energy level by transferring or sharing electrons to form either ionic or covalent compounds. Figure 2: Two Bonding Types Whether two given atoms tend to bond ionically or covalently is determined by the difference in their electronegativity. Electronegativity is a dimensionless number that is a measure of an atom’s attraction for bonding valence electrons. Electronegativities show periodic trends on a periodic table. Figure 3: Table of Electronegativities Excluding the noble gases, the most electronegative element is fluorine, which is assigned a value of 4.0. The other elements’ values are calculated on the basis of that of fluorine. Across each period, electronegativities tend to increase. The nonmetal families of nitrogen, oxygen, and fluorine have the highest values. Due to their atoms’ small radii, the positive nuclei exert a greater attraction for bonding electrons. The alkali metals and alkaline earth metals (the groups on the left side of the periodic table) have the lowest electronegativities because their atoms have the largest radii. Cesium (and francium – not shown in Figure 3), with the largest radii, have the lowest electronegativity, at 0.7. Lewis Dot Structures A Lewis dot structure is one way to represent the arrangement of valence electrons in a molecule (Figure 1). In Lewis structures, an element symbol is surrounded by a specific number of dots representing valence electrons. Most atoms obey the octet rule; they need eight valence electrons to fill their outer shell. (Exceptions include hydrogen and helium, which need only two. Also, boron and beryllium may form compounds with fewer than eight, and elements in periods 3–6 may use more than eight.) To fill their outer shells, elements can form covalent bonds by sharing electrons. To show those bonds, Lewis structures are often used. Covalent compounds may have single, double, or triple bonds between atoms. These bonds are represented in a Lewis structure with dashes between the chemical symbols of the bonded elements. The number of dashes corresponds to the number of bonds. Lewis structures also indicate the lone pairs of electrons on different atoms. Use the guide below if it helps in drawing Lewis Structures. Parta: Lewis Structures Procedure