Chemistry 1B Fall 2012 Lecture 9
bonding in molecules
• Chapter 13 (pp 593-612)– Overview of bonding and Chemistry 1B ionic bonding (lect 9)
• Chapter 13 (pp 612-650)- “Classical” picture of Fall 2012 bonding and molecular geometry (lect 10-12) Lecture 9 • Chapter 19 (pp 933-937;- Bonding in transition metal 946-948; complexes (lect 13-14) (chapter 13; pp 593-612) 958-966)
[596-614]7th • Chapter 14- Quantum mechanical description of bonding
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types of bonding general properties of the 3-types of bonding (Silberberg fig.9.2)
• Ionic
• Covalent
• Metallic
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general properties of the 3-types of bonding (Silberberg fig.9.2) hello Lewis electron-dot diagrams
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1 Chemistry 1B Fall 2012 Lecture 9
G. N. Lewis- UC Berkeley G.N. Lewis 1916
The Atom and the Molecule by Gilbert N. Lewis Journal of the American Chemical Society Volume 38, 1916, pages 762-786 Received January 26, 1916
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G.N. Lewis 1916 G.N. Lewis 1916
The Cubical Atom. The Cubical Atom. A number of years ago, to account for the striking fact which has become known as Abegg's law of valence and countervalence, and according to which the total difference between the maximum negative and positive …….. valences or polar numbers of an element is frequently eight and is in no case more than eight, I designed what may be called the theory of the cubical atom. This theory, while it has become familiar to a number of my The pictures of atomic structure which are reproduced colleagues, has never been published, partly because it was in many in Fig. 2,1 and in which the circles represent the respects incomplete. Although many of these elements of incompleteness electrons in the outer shell of the remain, and although the theory lacks to-day much of the novelty which it originally possessed, it seems to me more probable intrinsically than some of the other theories of atomic structure which have been proposed, and I cannot discuss more fully the nature of the differences between polar and nonpolar compounds without a brief discussion of this theory.
The pictures of atomic structure which are reproduced in Fig. 2,1 and in which the circles represent the electrons in the outer shell of the 10 11
Lewis electron-dot symbols for atom Lewis hypothesis (first look) (kernels or Lewis Valence Electron Diagrams LVEDs )
To form compounds, atoms will gain, lose, or share electrons to attain “complete outer shells”.
For hydrogen, a “complete shell” corresponds to 2 electrons (1s2).
For atoms in period ‘n’, a “complete shell” often corresponds to 8 electrons (ns2 np6) octet structure.
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2 Chemistry 1B Fall 2012 Lecture 9
order of material: remember metals vs nonmetals and the periodic table
• ionic bonding (pp. 603-612)
non-metals • covalent bonding (pp. 612-650; 597-603)
• metallic bonding (extra fun, but no extra tuition charge $$$’s)
metals
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size of ions (Zumdahl, figure 13.8) Electronegativity (more: pp. 587-590, later)
• Electronegativity- the tendency of an atom to attract electrons and to ‘hold on to’ its own electrons n2 r 52.9 pm • Mulliken: (EN) = (IE EA)/2 (arbitrary units) Zeff MUL (see ch13 prob 13.18)
n and Zeff of outermost electron • e.g. for Na (EN) MUL = [(496) - (-52.9 )]/2=274 (kJ/mol) for Cl (EN) MUL = [(1256) - (-349 )]/2=802 (kJ/mol)
• High electronegativity- wants to accept electrons Low electronegativity- will donate electrons atoms with high electronegativity are electronegative atoms with low electronegativity are electropositive
• non-metals are electronegative metals are electropositive
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ionic bonding WebAssign HW#3 prob 29 (Zumdahl 13.22)
• metallic atoms lose electrons to attain an only one state for: ‘octet’ structure s2, p6 d10 s1, p1, d1 p5, d9 • nonmetallic atoms gain electrons to attain an ‘octet’ structure all others several states ambiguous
Na + Cl Na+ + Cl [Ne]3s [Ne]3s23p5 → [Ne] [Ne]3s23p6
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3 Chemistry 1B Fall 2012 Lecture 9
factoids relating to the heat of a reaction (ch 9): Why should (does) Na+ Cl - exist?
Table 12.7 • H, the change in enthalpy for a reaction, is the IE=+495 kJ HEAT given off or absorbed by the reaction energy in (for our purposes H º energy change) 147146 kJ/mol + Na(g) + Cl(g) → Na+ (g) + Cl (g)
• if heat is given off by the reaction [surroundings heat up], energy out the reaction is EXOTHERMIC and H < 0 EA=-349 kJ/mol [products MORE STABLE than reactants] Table 12.8
• if heat is absorbed by the reaction [surroundings cool], the reaction is ENDOTHERMIC and H > 0 NET ENERGY CHANGE = + 146 kJ/mol (+495kJ-349kJ) [reactants MORE STABLE than products; ionization is endothermic, IE > 0] ENDOTHERMIC • H for a complex process can be calculated by gas phase ions unstable relative to atoms adding H for the individual step of the process
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does Li+ F - exist? but Li+F- is a salt (solid) figure 13.10
Table 12.7 + + IE=520 kJ 192kJ +Li(g) + F(g) → Li (g) + F (g) → Li F (s) + energy [lattice energy] 147192 kJ/mol + Li(g) + F(g) → Li+ (g) + F (g)
energy out EA=-328 kJ/mol Table 12.8
NET ENERGY CHANGE = + 192 kJ/mol (520 – 328)
ENDOTHERMIC gas phase ions unstable relative to atoms
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Coulombs Law, electrostatic energy (p 594, 609) lattice energy (LE): ions in gas ô ions in solid (crystal lattice)
+ A B Li+ (g) + F (g) → Li+F (s) + 1047 kJ/mole
RAB
QQ lattice energy of LiF(s)= -1047 kJ/mole E AB exothermic 4R 0AB QQ stabilizes ionic solids 2.31 1019 Jnm AB RAB QQ k AB RAB
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4 Chemistry 1B Fall 2012 Lecture 9
Born cycle: measuring lattice energies neutral atoms ô ionic solid stabilization by lattice energy (are responsible for this concept)
- 855 kJ/mol Li(g) + F(g) Li+F (s) + 855 kJ/mole heat heat given off by Li+ (g) + F- (g) ö Li+F- (s)
+ 192 kJ/mol -1047 kJ/mol difficult to measure directly (IE+EA) lattice energy endothermic exothermic Born cycle: find an alternative set of reactions where heat of Li+ (g) + F (g) reaction CAN be measured for each step and the combinations of these reactions leads to H= +192 kJ/mol – 1047 kJ/mol=-855 kJ/mole Li+ (g) + F- (g) ö Li+F- (s)
exothermic Use this cycle to compute Lattice Energy (LE) Li+F-(s) STABLE relative to Li(g) + F(g) 26 27
Born cycle and lattice energies (13.9) trends in lattice energy (you ARE responsible for this) (HW#3 prob 30 (not responsible for calculation ; ch 9 in Chem 1C this year) #13.26)
Li+ (g) + F− (g) → LiF (s) + 1047 kJ/mol
ΔHLE= − 1047 kJ/mol (p. 600) LiF MgO NaF not directly measurable QQ but ΔHf for Li (s) + ½ F2 (g) → LiF (s) is and Lattice Energy=k AB RAB Li (s) → Li (g) ΔHvap = + 161 kJ/mol (p. 600)
½ F2 (g) → F (g) ½ ΔHBE of F2 = + 77 kJ/mol (p. 600) + Li (g) → Li (g) ΔHIE1 of Li = + 520 kJ/mol (p. 600) − • ionic charge Q Q (p. 609) (usually more F (g) → F (g) ΔHEA1 of F = - 328 kJ/mol (p. 600) A B important) + − Li (s) + ½ F2 (g) → Li (g) + F (g) ΔH1= +430 kJ/mol • ionic size (R ) figure Silb. 9.7 Li (s) + ½ F2 (g) → LiF (s) ΔHf =− 617 kJ/mol (p. 600) AB
+ − Li (g) + F (g) → LiF (s) lowest LE greatest LE ΔH = − ΔH + ΔH = − 1047 kJ/mol (least negative, << (most negative LE 1 f least exothermic) most exothermic) -923 kJ -1047 kJ -3916 kJ 28 lattice energy (molar) 29
WebAssign HW#3 prob 30 (Zumdahl 13.26) skip for later (pp 616-620)
• Bond energies, bond lengths, bond order (after Lewis structures chapter 13.10-13.12)
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5 Chemistry 1B Fall 2012 Lecture 9
factoids about ionic compounds covalent bonding (sharing of electrons): MUCH MORE LATER
•‘strong’ metals and ‘strong’ nonmetals are likely to form ionic compounds • sharing of electrons leads to lower energy than two isolated atoms (figure 13.01) • lattice energy stabilizes solids
• hydration of ions in aqueous solvents can contribute to • lone or non-bonding pairs solubility (Olmstead figure 2.2) • more than one-pair of electrons may be shared to form stable ‘octet’ • ionic compounds ‘crack’ (fig Silb 9.8) (single, double, triple bonds with bond orders 1, 2, 3 respectively)
• ionic compounds have high boiling and melting points • covalent bonding CANNOT be satisfactorily explained by classical (table Silb 9.1) (fig. Silb 9.10) electrostatics, but we need quantum mechanics chapter 14 • ionic compounds conduct electricity in molten (liquid) phase or in solution (Silb fig. 9.9)
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factoids about covalent bonding in molecules fully ionic vs fully covalent the TRUTH lies in between (sec 13.6)
• usually bonding between atoms of similar electronegativity • electron transfer in ionic compounds may be incomplete (metallic bonding will be special case) • two atoms may not equally share electrons in a covalent • many covalently bound molecules have strong intramolecular forces (the covalent bonds) but weak intermolecular forces; bond thus relatively low melting and boiling points (figure 9.14) • the greater the ΔEN the more ionic (figure S9.22) • poor conductors • polar covalent bonds and covalency in ionic bonds • some atoms form extended networks of covalent bonds with high melting/ boiling points and hardness (figure 9.15) (figure 13.11) (figure S9.21), (figure 13.12), graphene
• continuum across period (more covalent as ΔEN decreases) (figure S9.25), (figure S9.24)
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more about electronegativity bonding in metals
• the degree of ‘attraction’ of a given atom for electrons electron sea model (its own and from other atoms)
• Mulliken scale: (EN)MUL = (IE EA)/2 (arbitrary units)
• Pauling electronegativities (section 13.2) trends (figure Zumdahl 13.3, figure Silb 9.19) HW#3
• oxidation number and electronegativity (common valences) (table 13.5, figure Silb 9.3) HW#3
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6 Chemistry 1B Fall 2012 Lecture 9
properties of metals related to electron sea model big picture: structure and properties
• electrical and thermal conductivity • Ionic • moderate melting point; high boiling point • Covalent (table S9.7) (figure S9.26) • Metallic • malleability (figure S9.27)
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Silberberg figure 9.6; Zumdahl figure 13.9
END OF LECTURE 9 ? ?
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Lattice Energy Born Cycle: MgO vs NaF (Zumdahl, figure 13-11) Silberberg figure 9.7
LE (MgO)= -3916 kJ/mol
LE (NaF)= -923 kJ/mol
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7 Chemistry 1B Fall 2012 Lecture 9
Olmstead figure 2.2: hydration of ions Silberberg figure 9.8
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Silberberg table 9.1 Silberberg figure 9.10
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Silberberg figure 9.9 Zumdahl figure 13.1
lonely (not sharing e-’ s)
lower energy=happy (sharing 2 e- ’s)
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8 Chemistry 1B Fall 2012 Lecture 9
Silberberg figure 9.14 Silberberg figure 9.15
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Zumdahl, Table 13.1; Silberberg figure 9.22 Pauling electronegativity (figure 13.3)
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Silberberg figure 9.19 HW#3 prob 28
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9 Chemistry 1B Fall 2012 Lecture 9
Zumdahl Table 2.5 and 13.5, Silberberg figure 9.3 HW#3 prob 31
KNOW: + - 2- - NH4 , NO3 , SO4 , HSO4 , - - 3- 2- CN , OH , PO4 , CO3 , - - - HCO3 , C2H3O2 , MnO4
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Bond polarity (figure 13.12 Zumdahl) Silberberg figure 9.21 (covalent electron density)
pure covalent
polar covalent
pure ionic
fig 13.11, 5th edition red= highest electron density blue= lowest electron density 58 59
percent ionic character (Zumdahl figure 13.12; Silberberg figure 9.23) Silberberg figure 9.25
LiF is highly ionic see the ‘rest’ of 3rd row p. 357 text more covalent as one goes across row
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10 Chemistry 1B Fall 2012 Lecture 9
Silberberg figure 9.24 Silberberg figure table 9.7
ionic covalent
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Silberberg figure 9.26 Silberberg figure 9.27
2nd period metals harder to melt
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graphene 2010 Nobel Prize in physics
Graphene is a one-atom-thick planar sheet of sp2-bonded carbon atoms that are densely packed in a honeycomb crystal lattice.
Potential applications Nobel Prize in physics won by Russian duo working in Manchester Graphene transistors Integrated circuits "for groundbreaking experiments regarding the Anti-bacterial two-dimensional material graphene", Single molecule gas detection 66 67
11 Chemistry 1B Fall 2012 Lecture 9
Born cycle and lattice energies (figure 13.10)
Mg2+ (g) + O2− (g) → MgO (s) + 3916 kJ/mol
ΔHLE= − 3916 kJ/mol (p. 601) not directly measurable
but ΔHf (MgO) is and
Mg (s) → Mg (g) ΔHvap = + 150 kJ/mol (p. 601)
½ O2 (g) → O (g) ½ ΔHBE of O2 = + 247 kJ/mol (p. 601) 2+ Mg (g) → Mg (g) ΔHIE1+IE2 of Mg= +2180 kJ/mol (p. 601) 2− O (g) → O (g) ΔHEA1+EA2of O= + 737 kJ/mol (p.601) ▬▬▬▬▬▬▬▬▬▬▬▬ 2+ 2− Mg (s) + ½ O2 (g) → Mg (g) + O (g) ΔH1= +3322 kJ/mol
Mg (s) + ½ O2 (g) → MgO (s) ΔHf =− 602 kJ/mol (p. A22)
Mg2+ (g) + O2− (g) → MgO (s)
ΔHLE= − ΔH1 + ΔHf = − 3916 kJ/mol
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