Chapter 6 Student Notes PHW

Chapter 6 Student Notes PHW - Chapter 6 Objectives 6.1....

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Unformatted text preview: Chapter 6 Objectives 6.1. Explain the basic concepts surrounding bonding including the Lennard – Jones potential, electronegativity, ionization, and electron affinity. electronegativity, 6.2. Illustrate the spectrum of bonding based upon percent ionic character. character. 6.3. Discuss Lewis Structures. 6.4. Describe resonance and formal charges. 6.5. Outline the characteristics of polar bonds and polar molecules. molecules. 6.6. Discuss VSEPR theory. 8/12/2009 Zumdahl Chapter 13 1 “What we represent when we draw a molecule is what we want to represent: We abstract a piece of reality to show it to another person.” Roald Hoffmann Chemistry Nobel Laureate Cornell University 8/12/2009 Zumdahl Chapter 13 2 1 Preliminaries • Our knowledge of atomic structure, electron configurations, & periodic properties is a foundation for understanding bonding. • Electrons can be divided into – – (e− in a filled shell) (e− in an unfilled shell; outermost electrons) • Valence electrons participate in bonding through – Sharing of e− by atoms: – Transfer of e− from one atom to another: 8/12/2009 Zumdahl Chapter 13 3 Ionic Bonds Ionic substances are formed when an atom that loses electrons relatively easily reacts with an atom that has a high affinity for electrons Na · e− + Cl Na+ 8/12/2009 Loss of a valence electron Na+ + e− .. → : Cl : Gain of a valence electron .. .. Combination to form the → NaCl : Cl : + ionic compound NaCl .. → Zumdahl Chapter 13 4 2 The Coulomb Potential The energy of interaction between a pair of ions can be calculated by using Coulomb's law: If Q1 and Q2 have opposite signs, V is negative ( interaction). If Q1 and Q2 have the same sign, V is positive ( interaction). 8/12/2009 Zumdahl Chapter 13 5 The Lennard − Jones Potential The potential energy of interaction between neutral atoms and/or molecules is conveniently described using the Lennard − Jones potential: potential: ε ≡ Well Depth re 8/12/2009 r0 ≡ Separation at which V = 0 ≡ Separation at the well minimum = Zumdahl Chapter 13 6 3 • Atoms or molecules approach at large distance (V → 0 as r → ∞). • V goes negative as intermolecular forces come into play • Minimum energy at r with . maximum (r) • V goes positive as forces become dominant at small r. 8/12/2009 Zumdahl Chapter 13 7 Covalent Bonds Covalent Bonding: Whenever possible, the valence electrons in a compound are distributed in such a way that each main-group mainelement in a molecule (except hydrogen) is surrounded by electrons (an should have .. H· + ·Cl: .. 8/12/2009 → of electrons). Hydrogen electrons). electrons in such a structure. structure. .. H:Cl: .. Zumdahl Chapter 13 or .. .. H―Cl: 8 4 The Dipole Moment Polar Covalent Bonds • Bonded atoms share electrons unequally, whenever the atoms differ in differ • Example − HF: The F atom carries a slightly electric charge and the H HF: charge atom a slightly charge of equal magnitude. Aligns in an electric field. magnitude. • Polar molecules posses a 8/12/2009 , µ. Zumdahl Chapter 13 9 Electronegativity (EN) vs. Electron Affinity (EA) • Electron affinity is a measure of the energy required to detach an electron from an atom or molecule: A− → A + e− , ∆E ≡ EA (a large EA means a strong attraction of electrons) • Electronegativity is a measure of the ability of an atom in a molecule to attract shared electrons to itself. 8/12/2009 Zumdahl Chapter 13 10 5 The Person Behind the Science Linus Pauling (1901-1994 ) Moments in a Life • • • • 1954 Nobel Prize in chemistry – Chemical Bonding, molecular biology • 1962 Nobel Peace Prize – Health (Vitamin C advocate) and Nuclear testing (banning atmospheric tests ) Electronegativity Scale: concept of partially ionic bonds. Scale: – fluorine with x = 4 is the most electronegative element, francium with x = 0.7 the least. Percent Ionic Character – x values can also be used to estimate the dipole moment and ionic character of bonds. Pauling: Electronegativity is the power of an atom in a molecule to attract electrons to itself. – x differs from the electron affinity of the free atom although the two run roughly parallel. 8/12/2009 Zumdahl Chapter 13 11 The Pauling Electronegativity Values as Updated by A.L. Allred in 1961 Note: 8/12/2009 Zumdahl Chapter 13 12 6 Expected H−X bond energy = [(H−H bond energy) (X−X bond energy)]1/2 [(H energy) (X energy)] (Geometric Mean of H−H and X−X Bond Energies) H− X− Bond energy = expected bond energy if ∆ = EN(H) −EN(X) = 0. The larger ∆ is, the stronger the H−X bond is. H− Ionic vs. Covalent Bonding 8/12/2009 Zumdahl Chapter 13 13 Ionic Character = 100% if µ = Q r where r = bond length (µ & r det’d experimentally) det’ Q = electron charge 8/12/2009 Zumdahl Chapter 13 14 7 Non-Polar Covalent Bonding 8/12/2009 Zumdahl Chapter 13 Ca: [Ar]4s2 O: Lose 2 electrons 15 Ca2+: [Ar] 2− 2 6 [He] 2s22p4 Gain 2 electrons O : [He] 2s 2p or [Ne] 8/12/2009 Zumdahl Chapter 13 16 8 Sizes of ions related to positions of elements in the periodic table. Atomic Radii In picometers Cations: Anions: Isoelectronic series: O 2FNa+ Mg2+ 8/12/2009 Zumdahl Chapter 13 Al3+ 17 Formation of Binary Ionic Compounds Lattice Energy (5) can be calculated using a modified version of Coulomb’s Law Coulomb’ (see text, p. 229) 8/12/2009 Zumdahl Chapter 13 18 9 Partial Ionic Character of Covalent Bonds Percent Ionic Character Covalent e.g., H2, Cl2 N2 Polar Covalent e.g., HF, H2O Ionic Bond The relationship between the ionic character of a covalent bond and the electronegativity difference between the bonded atoms e.g., LiF, NaCl 8/12/2009 Zumdahl Chapter 13 19 Covalent Bond Energy and Enthalpy Bond Enthalpy (∆H) is the enthalpy change in a reaction in which a chemical bond is broken in the gas phase. Bond Energy (∆E) is the energy needed to break a chemical bond. – Energy is released when bonds are formed ( – Energy must be supplied when bonds are broken ( ) ) Gas Phase: ∆H = Condensed Phase: ∆H = 8/12/2009 Zumdahl Chapter 13 20 10 The breaking of chemical bonds in stable substances often generates highly reactive products (or intermediates, i.e., radicals) CH4 → ·CH3 + ·H (∆H° = + 439 kJ mol-1) Bond Enthalpy 8/12/2009 Zumdahl Chapter 13 21 Average Bond Enthalpies C2H6 → ·C2H5 + ·H ∆H° = + 410 kJ mol-1 CHF3 → ·CF3 + ·H ∆H° = + 429 kJ mol-1 CHCl3 → ·CCl3 + ·H ∆H° = + 380 kJ mol-1 CHBr3 → ·CBr3 + ·H ∆H° = + 377 kJ mol-1 average ∆H°(C−H) ~ + 413 kJ mol-1 8/12/2009 Zumdahl Chapter 13 22 11 8/12/2009 Zumdahl Chapter 13 23 Estimate the Standard Enthalpy of Reaction for the gas-phase reaction that forms methanol from methane and water CH4(g) + H2O(g) → CH3OH(g) + H2(g) Approach: (1) (2) 8/12/2009 Zumdahl Chapter 13 24 12 CH4(g) + H2O(g) → CH3OH(g) + H2(g) H H H H + C H O H O H H C H H + H H Broken Formed 4 C−H = 4 x 413 C− 2 O−H = 2 x 467 O− 3 C−H or 3 x (−413) C− (− 1 O−H or 1 x (−467) O− (− 1 C−O or 1 x (−358) C− (− 1 H−H or 1 x (−432) H− (− −2496 kJ mol-1 Exothermic +2586 kJ mol-1 Endothermic ∆H = = 8/12/2009 Zumdahl Chapter 13 25 CCl2F2 + 2H2 → CH2Cl2 + 2HF ∆Hr = ? = ∆Hbond breaking – ∆Hbond making = –114 kJ mol-1 8/12/2009 Zumdahl Chapter 13 26 13 8/12/2009 Zumdahl Chapter 13 27 Lewis Structures to Represent Covalent Bonding Lewis structures (also called Lewis dot diagrams) are diagrams diagrams) that show the bonding of a molecule and the lone pairs of electrons that may exist in the molecule. A Lewis structure can be drawn for any covalently-bonded molecule. covalentlyThe Lewis structure is named after Gilbert N. Lewis, who Lewis, introduced it in his 1916 article entitled The Atom and the Molecule [Journal of the American Chemical Society 1916, Vol. 38, p. 762]. 1916, 38, 762]. 8/12/2009 Zumdahl Chapter 13 28 14 Drawing Lewis Structures 1. Count up the total number of valence electrons available (A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it. ion, 2. Calculate the total number of electrons needed (N) for each atom to have its own noble-gas set of electrons around it (two for hydrogen, eight for the nobleelements from carbon on in the periodic table). 3. Subtract the number in step 1 from the number in step 2. This is the number is of shared (or bonding) electrons present (S = ). 4. Assign bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion. 8/12/2009 Zumdahl Chapter 13 29 Drawing Lewis Structures (con’t) 5. If any of the electrons earmarked for sharing remain, assign them in pairs them by making some of the bonds double or triple bonds. In some cases, there bonds. may be more than one way to do this. Typically, double or triple bonds form only between atoms of carbon, nitrogen, oxygen, and sulfur. carbon, nitrogen, oxygen, sulfur. 6. Assign the remaining electrons ( octets to all atoms except hydrogen. ) as lone pairs to the atoms, giving 7. Determine the (C) on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion. C ≡ group no. − no. of lone pair electrons − ½ (no. of electrons in bonding pairs) When choosing between more than one possible Lewis structure, the one with the is usually preferred. preferred. 8/12/2009 Zumdahl Chapter 13 30 15 Lewis Structure of Methane (CH4) 1. Count up the total number of valence electrons available (A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it. H C H H H H C H H H 8/12/2009 Zumdahl Chapter 13 31 2. Calculate the total number of electrons needed (N) for each atom to have its own noble-gas set of electrons around it (two for noblehydrogen, eight for the elements from carbon on in the periodic table). H H C H H H H C H H 8/12/2009 Zumdahl Chapter 13 32 16 3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S). (S H C H H N−A = H H C H H H 8/12/2009 Zumdahl Chapter 13 33 4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion. H H C H H 2 2 2 2 All bonding electrons assigned 8/12/2009 Zumdahl Chapter 13 34 17 Lewis Structure of Ethylene (C2H4) 1. Count up the total number of valence electrons available (A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it. H H C C H H H H C C H H 8/12/2009 Zumdahl Chapter 13 35 2. Calculate the total number of electrons needed (N) for each atom to have its own noble-gas set of electrons around it (two for noblehydrogen, eight for the elements from carbon on in the periodic table). H H C C H H H H C C H H 8/12/2009 Zumdahl Chapter 13 36 18 3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S). (S H H C C H H H H C C H N−A = H 8/12/2009 Zumdahl Chapter 13 37 4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion. H H C C H H H 2 2 H C 2 C 2 2 H 8/12/2009 H Zumdahl Chapter 13 38 19 5. If any of the electrons earmarked for sharing remain, assign them in pairs by making some of the bonds double or triple bonds. In some cases, there may be more than one way to do bonds. this. Typically, double or triple bonds form only between atoms of carbon, nitrogen, oxygen, and sulfur. sulfur. H 2 2 H 8/12/2009 C 2 H 2 C 2 H Zumdahl Chapter 13 39 Lewis Structure of Carbon Monoxide (CO) 1. Count up the total number of valence electrons available (A) by first adding the group numbers of all the atoms present. If the species is a negative ion, add the absolute value of the total charge; if it is a positive ion, subtract it. · ·C· · · ·· ·O: 8/12/2009 Zumdahl Chapter 13 40 20 2. Calculate the total number of electrons needed (N) for each atom to have its own noble-gas set of electrons around it (two for noblehydrogen, eight for the elements from carbon on in the periodic table). 8/12/2009 Zumdahl Chapter 13 41 3. Subtract the number in step 1 from the number in step 2. This is the number of shared (or bonding) electrons present (S). 8/12/2009 Zumdahl Chapter 13 42 21 4. Assign two bonding electrons (as one shared pair) to each connection between two atoms in the molecule or ion. 2 C──O 8/12/2009 Zumdahl Chapter 13 43 5. If any of the electrons earmarked for sharing remain, assign them in pairs by making some of the bonds double or triple bonds. In some cases, there may be more than one way to do bonds. this. Typically, double or triple bonds form only between atoms of carbon, nitrogen, oxygen, and sulfur. carbon, nitrogen, oxygen, sulfur. C 8/12/2009 2 O Zumdahl Chapter 13 44 22 6. Assign the remaining electrons (A − S) as lone pairs to the (A atoms, giving octets to all atoms except hydrogen. (2 lone pairs) 8/12/2009 O 2 O C O C 2 C C O Zumdahl Chapter 13 45 7. Determine the formal charge (C) on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion. What is the formal charge on the carbon atom in methane (CH4)? C ≡ group no. − no. of lone pair electrons − ½ (no. of electrons in bonding pairs) 8/12/2009 Zumdahl Chapter 13 46 23 What is the formal charge on each hydrogen atom in methane (CH4)? C ≡ group no. − no. of lone pair electrons − ½ (no. of electrons in bonding pairs) As should be the case for the neutral molecule CH4, 8/12/2009 Zumdahl Chapter 13 47 7. Determine the formal charge (C) on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion. What is the formal charge on each carbon atom in ethylene (C2H4)? C ≡ group no. − no. of lone pair electrons − ½ (no. of electrons in bonding pairs) 8/12/2009 Zumdahl Chapter 13 48 24 What is the formal charge on each hydrogen atom in ethylene (C2H4)? C ≡ group no. − no. of lone pair electrons − ½ (no. of electrons in bonding pairs) As should be the case for the neutral molecule C2H4, 8/12/2009 Zumdahl Chapter 13 49 7. Determine the formal charge (C) on each atom, and write it next to that atom. Check that the formal charges add to give a correct total charge on the molecule or molecular ion. What are the formal charges on the atoms in carbon monoxide (CO)? C O C ≡ group no. − no. of lone pair electrons − ½ (no. of electrons in bonding pairs) As should be the case for the neutral molecule CO 8/12/2009 Zumdahl Chapter 13 50 25 Resonance Structures There are cases where more than one Lewis structure can be .. .. drawn for a molecule: O O O O O O Equivalent Structures Typical O−O single bond length: 1.49 Å O− Typical O=O double bond length: 1.21 Å In O3 the experimental bond lengths are the same: 1.28 Å The actual structure can be described as a combination of the two Lewis structures, i.e., a .. .. O O O 8/12/2009 O O O Zumdahl Chapter 13 51 More Examples of Resonance Structures Carbonate anion, CO32− -2 2− O -2 2− O C C O O -2 2− O C O O O O Thiocyanate anion, NCS− N C S -1 − N C S -1 − N2O (nitrous oxide, laughing gas) N 8/12/2009 N O N N O Zumdahl Chapter 13 N N O 52 26 Favored Resonance Structures When comparing resonance structures for the same molecule, usually those with the contribute more to the overall resonance hybrid. When formal charges are necessary, resonance structures that have and positive charges on the less electronegative elements typically typically contribute more to the overall resonance hybrid. Examples: N2O and NCS 8/12/2009 Zumdahl Chapter 13 53 Breakdown of the Octet Rule Case 1: (free radicals) such as NO: radicals) NO: # of valence e− = 5 (N) + 6 (O) = 11 N O Lewis structure does not satisfy the octet rule but the molecule is stable, although somewhat reactive. Case 2: (mainly Be, B, and Al compounds): Be, (A) F B F F F -1 B F F (B) +1 Even though structure (B) obeys the octet rule, structure (A) is favored (Formal charges are lower & experimental evidence shows that there is no there double bond in BF3 ). 8/12/2009 Zumdahl Chapter 13 54 27 Breakdown of the Octet Rule (con’t.) Case 3: Third and higher period elements can exhibit bonding where an octet on the central atom is exceeded -- can expand up 12 e−! F F F S F F SF6 F SF4 34 total valence e− 8 e− in bonds 24 e− in lone pairs 2 e− left over F F S F F Rule for Lewis structures: 8/12/2009 If electrons remain after satisfying the octet rule, Zumdahl central add lone pairs to the Chapter 13 atom! 55 The Shapes of Molecules: The VSEPR Theory VSEPR ≡ Valence Shell Electron-Pair Repulsion Theory 8/12/2009 Zumdahl Chapter 13 56 28 VSEPR Theory: The Basic Idea Electron pairs in the valence shell of an atom repel each other on a spherical surface formed by the underlying core of the atom. The geometry which applies to a particular arrangement is determined by the determined steric number (SN) of the central atom. The favored geometry is the one that SN) minimizes electron − electron repulsions. repulsions. “Steric” means “having to do with space.” The steric number of an atom in a Steric” space. molecule can be determined by drawing the Lewis structure of the molecule and adding the number of atoms that are bonded to it and the number of lone pairs that it has. has. 8/12/2009 Zumdahl Chapter 13 57 Geometry and Steric Number • • • • SN = 2 SN = 3 SN = 4 SN = 5 Linear, 180° 180° Trigonal planar, 120° 120° Tetrahedral, 109.5° 109.5° Trigonal bipyramidal 90° (equatorial − axial) 90° 120° (equatorial – equatorial) 120° Lone pairs occupy equatorial positions in preference to axial positions. • SN = 6 Octahedral, 90° 90° • SN for Double and Triple bonds count the same as single bonded atoms • When lone pairs are present, the situation is more complicated due due to repulsive forces Lone pair vs. lone pair > Lone pair vs. bonding pair > Bonding pair vs. bonding pair 8/12/2009 Zumdahl Chapter 13 58 29 Examples with no Lone Pairs on the Central Atom 8/12/2009 Zumdahl Chapter 13 59 Geometry is 8/12/2009 Zumdahl Chapter 13 60 30 Geometry is 8/12/2009 Zumdahl Chapter 13 61 Geometry is 8/12/2009 Zumdahl Chapter 13 62 31 Geometry is 8/12/2009 Zumdahl Chapter 13 63 Geometry is 8/12/2009 Zumdahl Chapter 13 64 32 Example: NH3 Geometry is 8/12/2009 Zumdahl Chapter 13 65 Example: H2O Geometry is 8/12/2009 Zumdahl Chapter 13 66 33 8/12/2009 Zumdahl Chapter 13 67 Chapter 6 Bonding; General Concepts 6.1 Types of Chemical Bonds 6.2 Electronegativity 6.3 Bond Polarity and Dipole Moments 6.4 Ions: Electron Configurations and Sizes 6.5 Formation of Binary Ionic Compounds 6.6 Partial Ionic Character of Covalent Bonds 6.7 The Covalent Chemical Bond: A Model 6.8 Covalent Bond Energies and Chemical Reactions 6.9 The Localized Electron Bonding Model 6.10 Lewis Structure 6.11 Resonance 6.12 Exceptions to the Octet Rule 6.13 Molecular Structure: The VSEPR Model 8/12/2009 Zumdahl Chapter 13 68 34 ...
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This note was uploaded on 10/06/2009 for the course CHEM 1101 taught by Professor Bottomley during the Fall '08 term at Georgia Tech.

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