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Chap 8. Unit 3-2 chemical bonds. Chemical Bonds. Three basic types of bonds: Ionic Electrostatic attraction between ions Covalent Sharing of electrons Metallic Metal atoms bonded to each other. Ionic Bonding. Energetics of Ionic Bonding.
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Chap 8 Unit 3-2 chemical bonds
Chemical Bonds • Three basic types of bonds: • Ionic • Electrostatic attraction between ions • Covalent • Sharing of electrons • Metallic • Metal atoms bonded to each other.
Energetics of Ionic Bonding As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.
Energetics of Ionic Bonding We get -349 kJ/mol back by giving electrons to chlorine.
Energetics of Ionic Bonding • But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic! • The formation of ionic bonds is always exothermic!
Energetics of Ionic Bonding • There must be a third piece to the puzzle. • What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.
Q1Q2 d Eel = Lattice Energy • This third piece of the puzzle is the lattice energy: The energy required to completely separate a mole of a solid ionic compound into its gaseous ions. • The energy associated with electrostatic interactions is governed by Coulomb’s law:
Lattice Energy • Lattice energy, then, increases with the charge on the ions. • It also increases with decreasing size of ions.(d between cation and anion decreases.)
Energetics of Ionic Bonding By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.
Energetics of Ionic Bonding • These phenomena also helps explain the “octet rule.” • Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.
Sodium Chloride Crystal Lattice Ionic compounds form solid crystals at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions. All salts are ionic compounds and form crystals.
Covalent Bonding • In these bonds atoms share electrons. • There are several electrostatic interactions in these bonds: • Attractions between electrons and nuclei • Repulsions between electrons • Repulsions between nuclei
Polar Covalent Bonds • Although atoms often form compounds by sharing electrons, the electrons are not always shared equally. • Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. • Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.
Polar-Covalent bonds • Electrons are unequally shared • Electronegativity difference between .3 and 1.7 • Nonpolar-Covalent bonds • Electrons are equally shared • Electronegativity difference of 0 to 0.3 Covalent Bonds
Ionic bond: Page 295 8.7, 8.13, 8.15, 8.17, 8.27 Covalent bond: Page 296 8.37, 8.39, 8.40, 8.45 Homework
The Octet Rule • Combinations of elements tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. • The rule applies to the p block elements only! • For the rest of elements, they try to obtain noble gas configuration Monatomic chlorine Diatomic chlorine
The Octet Rule and Covalent Compounds • Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level. • Covalent compounds involve atoms of nonmetals only. • The term “molecule” is used exclusively for covalent bonding
Comments About the Octet Rule • 2nd row elements C, N, O, F observe the octet rule (HONC rule as well). • 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. • 3rd row and heavier elements CAN exceed the octet rule using empty valence dorbitals. • When writing Lewis structures, satisfy octets first, then place electrons around elements having available dorbitals.
The HONCRule • Hydrogen forms one covalent bond • Oxygen forms two covalent bonds • One double bond, or two single bonds • Nitrogen forms three covalent bonds • One triple bond, or three single bonds, or one double bond and a single bond • Carbon forms four covalent bonds. • Two double bonds, or four single bonds, or a triple and a single, or a double and two singles
Lewis dot Structures Lewis structures are representations of covalent molecules showing all electrons, bonding and nonbonding.
Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge. PCl3 Writing Lewis Structures 5 + 3(7) = 26
Writing Lewis Structures • The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26 6 = 20
Writing Lewis Structures • Fill the octets of the outer atoms. Keep track of the electrons: 26 6 = 20 18 = 2
Writing Lewis Structures • Fill the octet of the central atom. Keep track of the electrons: 26 6 = 20 18 = 2 2 = 0
Writing Lewis Structures • If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.
Writing Lewis Structures • Then assign formal charges. • For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. • Subtract that from the number of valence electrons for that atom: The difference is its formal charge.
Writing Lewis Structures • The best Lewis structure… • …is the one with the fewest charges. • …puts a negative charge on the most electronegative atom.
