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Please read the following article:


and write a response in 1-2 pages (minimum of 300 words, maximum of 600 words, not including the exam question) using the 3-2-1 format described below:

3: Find 3 concepts from within the article and relate them to 3 concepts within CHEM 210 we have discussed in class and cite 3 textbook references using the chapter and page number.

2: Find 2 concepts from within the article that you want to know more about (i.e. muddy points, have questions about, did not quite understand).

1: Write an exam question with the answer about 1 concept discussed from within the article. The exam question must be well thought out and appropriate to the subject matter.

I upload the lecture notes to use for Find 3 concepts from within the article and relate them to 3 concepts within CHEM 210 we have discussed in class and cite 3 textbook references using the chapter and page number.

please make it clear and follow the format described.

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Chapter 8. Covalent Bonding and Molecular Structure 8.1 An Introduction to Covalent Bonding 8.2 Lewis Structures 8.3 Bond Properties 8.4 Electron Distribution in Molecules 8.5 Valence-Shell Electron-Pair Repulsion Theory and Molecular Shape 8.6 Molecular Polarity Examine chemical bonding in detail by applying what you have learned in Chapters 6 and 7 (atomic structure, electron configurations, and periodic trends) to the chemical bonds formed between atoms and ions and the shapes of molecules and ions that contain covalent bonds. Covalent Bonding and Molecular Structure 8.1 An Introduction to Covalent Bonding 2 8.1 Forms of Bonding Atoms • Ionic bonding: Complete transfer of 1 or more electrons from one atom to another – Contains strong attractive forces among cations and anions use electrostatic forces • Covalent bonding: Valence electrons shared between two adjacent atoms – Attractive forces between electrons and the nuclei of adjacent atoms within a molecule • Metallic bond: Attractive forces that exist between electrons and the nuclei – Hold pure metals together – Cation exist in a “sea” of electrons • Transfer/sharing of e- results in each atom/ion attaining an octet or noble gas electron configuration 3 Relationship Between Potential Energy and Interatomic Distance qA  qB   Force  r2 Changing the distance between H atoms affects potential energy 4 Covalent Bonding and Molecular Structure 8.2 Lewis Structures 5 8.2 Covalent Bonding and Lewis Structures Lewis Structures (Electron-Dot Structures): Lewis symbol: Simplest Lewis structure for an element • Element symbol represents nucleus and electrons are arranged around its four sides • Dots represent valence electrons • Can be drawn to reflect electron configuration • To form bonds, elements gain, lose, or share e- to achieve 8 valence e- 8.2 Lewis Dot Symbols Main group elements, the number of dots in the Lewis dot symbol is the same as the group number 7 8.2 Covalent Bonding Lewis summarized much of his theory of chemical bonding with the octet rule. According to the octet rule, atoms will lose, gain, or share electrons to achieve a noble gas electron configuration. Bonding pair 8 8.2 Covalent Bonding H-atom: H-atom: H H H2 molecule: HH H–H Two H atoms move close enough to each other to share the e-pair. Arrangement allows each H atom to “count” both electrons as its own and to “feel” as though it has the noble gas e- configuration of He. Number of unpaired valence electrons gives general indication of the number of bonds an atom will likely form: • Hydrogen has only 1 electron and can only make 1 covalent bond • Group 7A has only 1 unpaired electron, generally forms 1 covalent bond • Group 6A had 2 unpaired electrons, generally forms of 2 covalent bonds 9 8.2 Covalent Bonding with Multiple Bonds 10 8.2 Electron-Dot Structures Multiple bonds are both shorter and stronger than their corresponding single-bond counterparts because there are more shared electrons holding the atoms together. 8.2 Guidelines for Writing Lewis Structures Summary 1. Count the valence e- for each atom in the molecule. Electron configuration, Nobel Gas notation 2. Draw a skeleton structure 1st listed atom goes in the middle except for H and halogens. Join atoms with single lines (pairs of e-). 3. Add e- pairs to form octets (except H). Start with terminal atoms. 4. Extra e- Place around the central atom. 5. Too few e- Convert lone pairs into multiple bonds. 6. Self-Check, all atoms have an octet? Are all valence e- used? 8.2 Electron-Dot Structures of Polyatomic Molecules Draw an electron-dot structure for CH2O. C H O Step 1: 4 + 2(1) + 6 = 12 valence electrons O O Step 2: H C H Step 5: H O Step 3: H C C H O H H C H 8.2 Electron-Dot Structures of Polyatomic Molecules Draw an electron-dot structure for H3O1+. Step 1: 3(1) + 6 - 1 = 8 valence electrons H Step 2: H O 1+ H H Step 4: H O H 8.2 Exceptions to the Octet Rule • H and He form e- deficient compounds, only need 2 e• Be and B form e- deficient compounds, very reactive molecules: F H Be H F B 2 + 2(1) = 4 valence e- F 3 + 3(7) = 24 valence e- 8.2 Odd Number of Valence Electrons Some stable molecules have an odd number of eNO 5 + 6 = 11 valence e- NO2 5 + 2(6) = 17 valence e- N O O N O Free radical atom or molecule with unpaired e-. Very reactive. Most stable molecules have paired e- 8.2 More Than Eight Valence Electrons “Expanded octets” are relatively common. ONLY 3p to 6p can have more than 8! Resulting from the d-orbitals accepting extra e- 8.2 Resonance A resonance structure Molecules that have more than one valid Lewis structures that differ in the arrangement of e• Atom arrangement remains the same • Different location and/or types of BONDING N : : O – O: :O: : :O: O – : : N : : :O : :O: O: : : N : : : : :O – 18 Covalent Bonding and Molecular Structure 8.3 Bond Properties 19 8.3 Bond Order • Number of e- pairs between the bonding atoms # of AB bonding pairs Bond Order = # of AB bond locations 3 N - O bonding pairs BO = = 1.5 2 N - O bonding locations 20 8.3 Interactive Table 8.3.1 - Average Bond Lengths (pm) 21 8.3 Bond Length Trends • Bond length increases with increasing atomic size • As the bond order increases, the bond length decreases – e- density between the two nuclei increase with each added pair of e– Attractive force between e- and the nuclei increases – Distance between the bonding nuclei decreases C  O 143 pm > C = O 122 pm > C  O 113 pm 22 8.3 Bond Enthalpy Bond energy increases with increasing bond order and decreasing bond length Greater the bond order, the higher the bond strength and the shorter the bond 23 8.3 Bond Enthalpy aka Bond Energy bond enthalpy, energy required to break a chemical bond 1 mole of gaseous molecules. Always Endothermic! ∆H° = [S # bonds*mol* HReactant bonds] - [S# bonds*mol* HProduct bonds] 24 8.3 Bond Dissociation Energies H2(g) + Cl2(g) 2HCl(g) H-H (g) + Cl-Cl (g)  2 H-Cl ∆H° = [S # bonds*mol* HReactant bonds] - [S# bonds*mol* HProduct bonds] ∆H° = (#H-H*molH2*HH-H + (#Cl-Cl*molCl2*HCl-Cl) - ((#H-Cl*molHCl*HH-Cl) ∆H° = [(1bond)(1 mol)(436 kJ/mol) + (1bond)(1mol)(243 kJ/mol)] - [(1bond)(2mol)(432 kJ/mol)] = -185 kJ 8.3 Bond Enthalpy 26 8.3 Problem: Exercise 2: Using Bond Energies Estimate the DH for the following reaction using average bond dissociation enthalpies: CH 4  g  + 2O 2  g   CO 2  g  + 2H 2 O(l) Lewis Structures Count Bonds of each type: Formula: Substitution: Bond Values from tables Answer 4  C  H  + 2  O=O   2  C=O  + 4  O  H  DH = S (bonds broken)  S (bonds formed) DH = (4(C-H) + 2(O=O))  (2(C=O) + 4(O-H)) DH = [(4(413) + 2(498))  (2(732) + 4(463)] = -668 kJ 27 8.3 Ionic Bonding, Lattice Energy Increase lattice energy, the more stable the compound and the higher the melting point, need more thermal energy, heat (q). 28 8.3 Comparison of Ionic and Covalent Bonding 29 Covalent Bonding and Molecular Structure 8.4 Electron Distribution in Molecules 30 8.4 Lewis Structures and Formal Charge Formal charge can be used to determine the most plausible Lewis structures when more than one possibility exists for a compound. • All the atom’s nonbonding e- are associated with the atom. • Half of the atom’s bonding e- are associated with the atom. 31 8.4 Lewis Structures and Formal Charge Lewis Structures and Formal Charge Note: sum of formal charges = molecular charge 32 8.4 Formal Charges If there is choice between Lewis structures: • Lewis structure in which all formal charges are ZERO is preferred • SMALLER formal charges are favored. • Negative formal charges should be on the MOST electronegative atoms • Like charges should NOT be on adjacent atoms Which N2O structure is preferred? O N N Formal -1 charges: +1 0 Preferred. ENO > O N N 0 +1 ENN -1 8.4 Example - Formic Acid There are two possible Lewis structures for this molecule Each has the same number of bonds. Which structure is better? Determine the formal charge on each atom in the 2 structures 34 8.4 Electronegativity and Polarity Electronegativity (): is the ability of an atom in a compound to draw electrons to itself. Electronegativity is related to electron affinity (makes anions) and ionization energy (makes cation). 35 8.4 Electronegativity and Polarity Ionic and covalent bonds are simply the extremes in bonding. Bonds that fall between these two extremes are polar, meaning that electrons are shared but are not shared equally. Such bonds are referred to as polar covalent bonds. Δ = 0.0 < 0.4 Δ = 0.4 < 1.9 Δ > 2.0 increasingly ionic increasingly covalent 36 8.4 Exercise: Bond Polarity • Which of the following bonds are nonpolar? C–Cl, H–H, H–Cl, P–H, S–O, B–F, and F–F Bond Difference in χ Polar/Nonpolar C—Cl 3.0 – 2.5 = 0.5 Polar covalent H—H 2.2 – 2.2 = 0 Nonpolar H—Cl 3.0 – 2.2 = 0.8 Polar covalent P—H 2.2 – 2.1 = 0.1 Nonpolar S—O 3.5 – 2.5 = 1.0 Polar B—F 4.0 – 2.0 = 2.0 Polar F—F 4.0 – 4.0 = 0 Nonpolar 37 Covalent Bonding and Molecular Structure 8.5 Valence-Shell Electron-Pair Repulsion Theory and Molecular Shape 38 8.5 Molecular Shapes: The VSEPR Model VSEPR: Valence-Shell Electron-Pair Repulsion model Electrons in bonds and in lone pairs can be thought of as “charge clouds” (areas of e- density) that repel one another and stay as far apart as possible, this causing molecules to assume specific shapes. Working from a the Lewis electron-dot structure: 1. count the number of “charge clouds,” • domains = bonding or lone e- pair 2. then determine the molecular shape. 8.5 VSEPR Theory • Electron-pair geometry: which is the arrangement of electron domains (bonds and e- lone pairs) around the central atom, only 5 choices • Molecular Geometry (Shape): defined by the positions of the atoms in the molecule, Lone pair electrons alter molecular shape – Arrangement of bonded atoms, NO lone e- pairs are shown. • Bond angle: Formed by the nuclei of two atoms with a central atom at the vertex 40 8.5 Linear 180o Electron Pair Geometry Trigonal planar 120o Tetrahedral 109.5o Trigonal bipyramidal 120o 180o Octahedral 90o VSEPR model predicts the electron domains repel one another, arrange themselves to be as far apart as possible, thus minimizing the repulsive interactions between them. 41 8.5 Valence Shell Electron Pair Repulsion model predicts shapes. 1. e- pairs stay as far apart as possible to minimize repulsions. 2. Shape of a molecule is governed by the number of bonds and lone e- pairs present. 3. Treat a multiple bond like a single bond when determining a shape. • Multiple bonds is 1 area of e- density. 4. Lone e- pairs occupy more volume than bonds due to electrostatic repulsion interactions. 8.5 Molecular Geometry Electron-Pair Geometry and Molecular Geometry Steps to determine the electron-pair and molecular geometries are as follows: 1. Draw the Lewis structure of the molecule or polyatomic ion. (e- configuration is needed) 2. Count the number of electron domains on the central atom. 3. Determine the electron-pair geometry by applying the VSEPR model on central atom 4. Determine the molecular geometry by considering the positions of the atoms only and number of lone pairs on the central atom. 43 8.5 Molecular Geometry Deviation from Ideal Bond Angles A lone pair takes up more space than the bonding pairs. • They contain more electron density • Multiple bonds repel more strongly than single bonds. 44 8.5 Electron-Pair Geometry & Molecular Geometry If there are NO e- lone pairs, the Electron-Domain Geometry and Molecular Geometry are the SAME! 117o 107.5o 104.5o 45 8.5 Electron-Pair Geometry & Molecular Geometry 117o 114o 87o 84o Axial Two positions that are directly across from each other, like the axis of the earth Equatorial Three positions in a plane, midway between the axial positions, are in the region that is like the equator 46 8.5 Electron-Pair Geometry & Molecular Geometry 47 Summary Number of groups Electron-Pair geometry Composition of groups Molecular Geometry 2 Linear 2 atoms Linear 3 Trigonal Planar 3 atoms Trigonal Planar 2 atoms, 1 LP Bent 4 atoms Tetrahedral 3 atoms, 1 LP Trigonal pyramidal 2 atoms, 2 LP Bent 5 atoms Trigonal Bipyramidal 4 atoms, 1 LP See-Saw 3 atoms, 1 LP T-Shaped 2 atoms, 1 LP Linear 6 atoms Octahedral 5 atoms, 1 LP Square Pyramidal 4 atoms, 2 LP Square Planar 4 5 6 Tetrahedral Trigonal Bipyramidal Octahedral 48 Table 9-1, p. 384 9.2 Which if any of the bond angles would you expect to be smaller than the ideal values? 51 Covalent Bonding and Molecular Structure 8.6 Molecular Polarity 52 8.6 Molecular Geometry and Polarity • Covalent bonds are polar when there is an uneven attraction for e- between the bonded atoms • Polar bonds in a molecule can result in a polar molecule – Affects the physical properties of a compound – Polar molecules are often very soluble in water, whereas nonpolar molecules are not Polarity depends of the individual bonds and its molecular geometry. 53 8.6 Molecular Polarity To Determine the molecular polarity, ask these question(s): Q1: Is the e- domain geometry and the molecular geometry the same? NO : POLAR YES: Ask Q2 Q2: Are all the terminal atoms (X atoms) bonded to the central atom the same? NO: POLAR YES: NONPOLAR Exception- higher level symmetry broken down into simpler symmetry Trigonal bipyramidal of linear and trigonal planar Octahedral broken down into simpler symmetry of linear 8.6 Molecular Geometry and Polarity e- Pair Geo: Trigonal Planar Mol Geo: Trigonal Planar Polarity: Nonpolar e- Pair Geo: Tetrahedral Mol Geo: Bent Polarity: Polar 55 Chapter 9. Theories of Chemical Bonding 9.1 9.2 9.3 9.4 Valence Bond Theory Hybrid Orbitals Pi Bonding Molecular Orbital Theory In this chapter, we learn more about why molecules have predictable shapes. This deeper understanding involves a model of chemical bonding called valence bond theory and will allow us to predict not only expected structures, but also expected exceptions to the usual rules. In the second part of this chapter, we explore a second theory of chemical bonding, called molecular orbital theory. This theory can be used to explain structures of molecules as well as the energetics of chemical processes. © 2017 Cengage Learning. All Rights Reserved. Theories of Chemical Bonding 9.1 Valence Bond Theory © 2017 Cengage Learning. All Rights Reserved. 2 9.1 Valence Bond Theory Valence Bond Theory: A quantum mechanical model which shows how electron pairs in atomic orbitals (s, p, d, f) are shared in a covalent bond, greater the overlap, the stronger the bond. 9.1 Valence Bond Theory Energetics and Directionality of Bonding In summary, the important features of valence bond theory are as follows: • A bond forms when singly occupied atomic orbitals on two atoms overlap. • The two electrons shared in the region of orbital overlap must be of opposite spin. • Formation of a bond results in a lower potential energy for the system. 4 9.1 Sigma Bond Formation sigma (s) bonds – head to head overlap of atomic orbitals, edensity along the axis of the bond, ALL 1st bonds are s bonds Two s orbitals overlap Overlap of s and p orbitals Two p orbitals overlap 5 Theories of Chemical Bonding 9.2 Hybrid Orbitals © 2017 Cengage Learning. All Rights Reserved. 6 9.2 Hybridization of Atomic Orbitals: s, p *Promoted, exited state Each sp hybrid (180° apart) holds one e-. Two equivalent covalent bonds form. Linear Electron Pair Geometry 7 9.2 Hybridization of Atomic Orbitals: s, p, p *Promoted, exited state Each sp2 hybrid (120° apart) holds one e-. Three equivalent covalent bonds form. Trigonal Planar Electron Pair Geometry 8 9.2 Hybridization of Atomic Orbitals: s, p, p, p *Promoted, exited state Each sp3 hybrid (109.5° apart) holds one e-. Four equivalent covalent bonds form. Tetrahedral Electron Pair Geometry 9 9.2 Hybridization of Atomic Orbitals: s, p, p, p, d *Promoted, exited state Each sp3d hybrid (120° or 90o apart) holds one e-. 5 equivalent covalent bonds form. Trigonal bipyramidal Electron Pair Geometry 10 9.2 Hybridization of Atomic Orbitals: s, p, p, p, d, d *Promoted, exited state Each sp3d2 hybrid (90° apart) holds one e-. 6 equivalent covalent bonds form. Octahedral Electron Pair Geometry 11 9.2 Hybridization of Atomic Orbitals Mixed s+p s+p+p s+p+p+p Hybrids (#) sp (2) sp2 (3) sp3 (4) Remaining p,p p Geometry Linear Trigonal planar Tetrahedral s+p+p+p+d sp3d (5) d,d,d,d Trigonal bipyramidal s+p+p+p+d+d sp3d2 (6) d,d,d Octahedral 12 9.2 Predicting Hybridization • Start by drawing the Lewis structure (check formal charges) • Use VSEPR theory to check the electron group geometry around the central atom • Relate the central atom electron group geometry to the corresponding hybridization • Identify and label the orbital overlap in each bond • Label the bonds with s and  bonds 13 9.2 Exercise: Hybridization • Indicate the central atom hybridization for the following: – XeF4 – CH2O – BrF5 – SF6 – Br3 14 Theories of Chemical Bonding 9.3 Pi Bonding © 2017 Cengage Learning. All Rights Reserved. 1 9.3 Hybridization in Molecules Containing Multiple Bonds Sigma bond (s) – bonds line head-to-head orbital overlap so electron density of bind lies along the bonding axis, as in diatomics Pi bond () – formed sideways (edgewise) overlap parallel to atomic porbitals, parallel p orbitals (py and py) overlap above and below bond axis 16 9.3 Pi () Bonding • Two orbitals overlap to form a bond where the bonding region is above and below the internuclear axis – Electrons occupy the space above and below the nuclei • Pi bonds form when two unhybridized p orbitals on adjacent atoms overlap 17 SAMPLE PROBLEM 9.5 Determine the number of carbon-carbon sigma bonds and the total number of pi bonds in thalidomide. Solution Thalidomide contains 12 carbon-carbon sigma bonds and a total of seven pi bonds (three in carbon-carbon double bonds and four in carbon-oxygen double bonds). 18 Pi Bonding Summary • Pi bonds are formed by the overlap of: – Two or more p orbitals – p and d orbitals – Two d orbitals Relationship between hybridization and number of possible p–p Pi Bonds Number of Hybridizatio Unhybridized Structural Pairs n p Orbitals on the Central Atom 2 sp Two p Number of Possible p–p Pi Bonds 2 3 sp2 One p 1 4 sp3 None 0 5 sp3d None 0 6 sp3d2 None 0 19 Theories of Chemical Bonding 9.4 Molecular Orbital Theory © 2017 Cengage Learning. All Rights Reserved. 20 9.4 Molecular Orbital Theory Bonding and Antibonding Molecular Orbitals Species in which all the electrons are paired are diamagnetic, whereas species that contain one or more unpaired electrons are paramagnetic. O2 exhibits paramagnetism, it must contain unpaired e- 21 9.4 Molecular Orbital Theory Atomic Orbital (AO): A wave function whose square (Y2) gives the probability of finding an electron within a given region of space in an atom. Molecular Orbital (MO): A wave function whose square (Y2) gives the probability of finding an electron within a given region of space in a molecule. 9.4 Molecular Orbital Theory Molecular Orbital Theory, the atomic orbitals (AOs) involved in bonding actually combine to form new orbitals that are the “property” of the entire molecule, rather than of the atoms forming the bonds. Bonding elect ...
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