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Chemical Bonds

Chemical Bonds. What are chemical bonds? A Chemical Bond is a link between atoms that results in a structure with Lower Energy that the energy sum of the individual atoms Valence electrons are the ‘currency’ of chemical bonds Keeping this in mind, let’s look at the types of bonds….

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Chemical Bonds

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  1. Chemical Bonds • What are chemical bonds? • A Chemical Bond is a link between atoms that results in a structure with Lower Energy that the energy sum of the individual atoms • Valence electrons are the ‘currency’ of chemical bonds • Keeping this in mind, let’s look at the types of bonds…

  2. Types of Bonds • Ionic Bonds • Lower energy can be achieved by transferring one or more electrons between atoms • The resulting atoms are held together by Coulombic attraction • Covalent Bond • Two or more atoms share electrons between them • Metallic Bond • Cationic atoms held together by a ‘sea’ of electrons

  3. Na n=3 1s22s22p63s1 Na+ n=3 1s22s22p6 Ions: Cations (s-block) • Cations of the s-block elements form when the atom loses its valence electrons in the s-shell and is stripped down to its noble gas core, or octet • Let’s look at sodium

  4. Ions: Cations (p-block) • Metals in the p-block lose any p electrons in addition to the s-shell electrons from their valence shell • Galium (n=4) does not lose its 3d electrons • Fe (n=4) loses its 2 4s shell electrons forming Fe2+ and can lose one of its 3d electrons to give five singly occuped d-orbitals • Less repulsion with a single electron in an orbital 

  5. Ions: Anions • Nonmetals in the p-block do not lose electrons due to their ionization energies • Instead, they gain electrons to fill their valence shell • Example: Nitrogen N = [He] 2s22p3 --> N3- = [He] 2s22p6

  6. Ions: Summary • Cations are formed by removing electrons in the following order: np --> ns --> (n-1)d • Once enough electrons have been removed to reach the noble gas core: Stop! • d-shell electrons are the exceptio. You lose enough of these to reach a set of singly occupied orbitals or none at all. • Anions are formed from the non-metallic elements by adding electrons untils the valence shell is ull • An octet is reached (Full s-shell and full p-shell)

  7. Lewis Symbols • A convenient way of working with atoms in bonding is to draw a representation of the valence electrons of individual atoms • Lewis structures or Lewis symbols are pictorial ways of keeping track of valence electrons • Examples:

  8. Using Lewis Symbols • We can use Lewis symbols in chemical reactions to represent the electronic configuration of the compounds formed • eg: Ca + Cl2 --> CaCl2 C + 2H2 --> CH4 • When working with metals, assume they have fully ionized and then draw the Lewis structure

  9. Understanding Ionic Bonds • Why does a crystal of sodium chloride have lower energy than a lump of sodium metal and a cloud of chlorine gas? • We can look at the ionization energy of sodium and the electron affinity of chlorine

  10. Understanding Ionic Bonds Na (g) --> Na+(g) + e-energy req. = 494 kJ/mole • It takes 494 kJ to strip 1 mole of sodium atoms of their 3s electrons (ionizaiton energy) Cl (g) + e- --> Cl-(g) energy released = 349 kJ/mole • When 1 mole of chlorine atoms absorbs 1 mole of electrons, 349 kJ of energy is released (electron affinity)

  11. Understanding Ionic Bonds Na+(g) + Cl-(g) --> NaCl (s)energy released =787 kJ/mole • To form one mole of NaCl from one mole of sodium ions and one mole of chlorine ions, 787kJ of energy is released • The crystallization process releases energy!

  12. Formation of Sodium chloride Na (g) --> Na+(g) + e-494 kJ/mole Cl (g) + e- --> Cl-(g) -349 kJ/mole Na+(g) + Cl-(g) --> NaCl (s)-787 kJ/mole Total Energy = -642 kJ/mole • 642 kJ of energy is released when we form sodium chloride form its constituent elements • The attraction of the cationic sodium ions and anionic chloride ions holds the solid together

  13. Interaction between ions • In an ionic solid, the formula unit is what we write to describe the simplest interaction group of the solid • HOWEVER: All of the ions in the solid are interacting (either atracting or repelling) with the ions immediately around them • We use the term Lattice Energy to describe the total interaction energy of the solid relative to the ions existing by themselves in a gas phase

  14. Ionic Solids: Interactions • We know that we can crush salt crystals in our fingers and if we look a the pieces, they form smooth breaks along specific planes • When we apply force to a face of a crystal…

  15. Lattice Energy • We can use a variant of the Coulombic potential energy equation to calculate the lattice energies of ionic solids Where: A=Madelung Constant (how atoms are packed) 0 = Vacuum permitivity constant d = rcation + ranion • Gives us the energy per mole of ionic solid

  16. Lattice Energy: How can we use it? • For crystals with different ions, but similar packing, we can use the value of d to determine the energy • Smaller values of d, higher energy • We can also use the equation for similarly packed crystals with ions of different charge by taking the charges z1 and z2 into account

  17. Lattice Energy: Take Home Message • Ionic solids have higher metling temperatures and are brittle • The Coulombic interactions between ions is highest when the ions are small and highly charged

  18. Covalent Bonds • Most bonds relevant to life are Covalent bonds • Arise form the sharing of an electron pair between 2 atoms • Lewis proposed this in 1916 • What was happening in the world at this time?

  19. The Octet Rule • In covalent bond formation, atoms go as far as possible to complete their octets by sharing electrons • Completing the octet? • Nitrogen needs 3 electrons • Chlorine needs 1 electron • Hydrogen needs 1 electron • General Rule: The # of valence electrons an atom needs is equal to the # of bonds it can form 

  20. The Octet Rule and Valence • Nitrogen needs 3 electrons to fill its valence shell • Compounds of Nitrogen: NH3 (ammonia) N2H4 (hydrazine) • Each nitrogen makes 3 covalent bonds, thereby filling its valence shell

  21. Lewis Structures • We symbolize a bond in a Lewis Structure with a line • Each Fluorine has 3 Lone Pairs of Electrons and a single electron in its valence shell • A Lone pair is a non-bonding electron pair held around an atom  

  22. Lewis Structures and Polyatomic Species • Most compounds in the universe consist of more just 2 atoms • Many have multiple bonds between 2 atoms • We can use Lewis Structures to depict these molecules

  23. Lewis Structures and Polyatomic Species • We can arrange the dots around the central atom • How do we know which atom is central?

  24. Drawing Lewis Structures How do we know which atom is central? • Hydrogen is ALWAYS TERMINAL • The atom with the lowest ionization energy is usually central • If there are more than one of a particular atom and just one of a non-hydrogen atom, that single atom is usually central • Arrange the atoms symmetrically 

  25. Resonance • Some compounds have multiple ways they could be drawn as a Lewis Structure • This implies that we can’t say exactly where an electron pair may be, but rather, that it is delocalized around the molecule • It is shared between multiple atoms • Structures that are equivalent are resonant

  26. Resonance: An Example All three structures are correct and the actual molecule is a blend of the three resonance structures. We use the double-headed arrow between structures to show that they are resonant

  27. Resonance: Benzene • Prototype cyclical resonance structure • The delocalized electron shuttles around the ring • It also gives the compound unique properties with respect to reactivity, bond lenghts and structure

  28. Resonance: Summary Resonance is a blending of structures with the same arrangement of atoms, but different arrangements of electrons. The delocalized electron pair spreads multiple bond character over a molecule giving it a lower energy

  29. Formal Charge • Formal charge is helpful in determining which structure in a resonance group is actually the one with lowest energy • Formal charge takes into account the electrons an atom actually “owns” as well as the electrons shared by covalent bonding • Remember that when we draw Lewis Structures, we just throw the valence electrons into a pile and then distribute them around the molecule to satisfy the octet rule. • Formal Charge sets the record straight, so to speak.

  30. Calculating Formal Charge Formal Charge = # of valence electrons - (# of lone pair electrons + 0.5(# of bonding electrons) FC = V - (L + 1/2(B)) • A Lewis Structure with the formal charges of all the atoms closest to zero represents that lowest energy arrangement of atoms and electrons 

  31. Drawing Lewis Structures 1. Count Valence Shell Electrons 2. Arrange atoms appropriately 3. Insert single bond pairs 4. Distribute lone pairs 5. Calculate formal charges 6. Modify if necessary 7. Resonance

  32. Exceptions to the Octet Rule • Some atoms just don’t follow the octet rule and can form compounds that give odd numbers of electrons or more than 8 electrons in the valence shell • Radicals • Phosphorous, sulfur, chlorine and Group 13/III (Boron and Aluminum)

  33. Radicals • Radicals are VERY highly reactive compounds in which a single unpaired electron exists • Biradicals are molecules with two unpaired electrons • Radicals are responsible for the effects of aging in living cells • The compounds may form from metabolic reactions (Redox reactions in the electron transfer chain, for example) or from chemical exposure • Antioxidants are compounds that can quickly and safely react with radicals before they can damage tissues

  34. Expanded Valence Shells / Expanded Octets • Central atoms that have empty d-orbitals sometimes form compounds that give them 10 or more valence electrons • Only non-metals in periods 3 or later may expand their valence shells

  35. Expanded Valence Shells / Expanded Octets • Phosphorous frequently forms PCl5 • Sulfur sometimes forms 4 covalent bonds and STILL has a lone pair

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