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Chapter 6

Chapter 6. Manipulating Polyatomic Ions and Chemical Bonding. Basic Polyatomics. https://www.youtube.com/watch?v=mlRhLicNo8Q. Ways to expand your polyatomics. Polyatomic ions vary in their charges, number of oxygen atoms, and number of hydrogen atoms. 1. To change the number of oxygens:

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Chapter 6

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  1. Chapter 6 Manipulating Polyatomic Ions and Chemical Bonding

  2. Basic Polyatomics https://www.youtube.com/watch?v=mlRhLicNo8Q

  3. Ways to expand your polyatomics • Polyatomic ions vary in their charges, number of oxygen atoms, and number of hydrogen atoms. • 1. To change the number of oxygens: • One more oxygen ClO4-1perchlorate • Memorized ClO3-1chlorate • One less oxygen ClO2-1chlorite • Two less oxygens ClO-1hypochlorite

  4. Ways to expand your polyatomics • 2. Family Members • Whatever is true for chlorine, is also true for fluorine, bromine, and iodine. • Memorized ClO3-1 chlorate • F substitution FO3-1 fluorate • Br substitution BrO3-1 bromate • I substitution IO3-1 iodate

  5. Ways to expand your polyatomics • 3. If you add a hydrogen, you have to make the ion more positive (one less negative) and call the ion “bi________” • Memorized CO3-2 - carbonate HCO3-1 – bicarbonate • Memorized SO4-2 - sulfate HSO4-1 - bisulfate

  6. Ways to expand your polyatomics • 4. Combinations of #1, #2, and #3 are possible: • HSO3-1 is called bisulfite • FO2-1 is called fluorite

  7. Rules for expanding your list of polyatomic ions Rule #1 To Change the number of oxygens: • Remove one oxygen = change ending of name to –ite • Remove two oxygens = change ending of name to –ite and beginning of name to Hypo- • Add one oxygen = change beginning of name to Per-

  8. Rules for expanding your list of polyatomic ions Examples:

  9. Rules for expanding your list of polyatomic ions Rule #2 Other Family Members • Elements near each other in the same column tend to form similar polyatomic ions.

  10. Rules for expanding your list of polyatomic ions Rule #3 Add a hydrogen • Add only one H = change the beginning of the name to bi- and make the charge one less negative (due to hydrogen’s positive one charge)

  11. Rules for expanding your list of polyatomic ions Rule #4 Combinations of 1, 2, & 3 • Combinations of #1, #2, and #3 are possible: • HSO3-1 • Memorized SO4-2 Sulfate • Lose an “O” SO3-2 Sulfite • Add an “H” HSO3-1 Bisulfite • HFO2 • Memorized ClO3-1 Chlorate • Substitute an FFO3-1 Fluorate • Lose an “O” FO2-1 Fluorite • Add an “H” HFO2Bifluorite

  12. Rules for expanding your list of polyatomic ions • Combos Cont’d • Ex1: What is the formula for hypoiodite? • Find I on the P-table (near Cl). Chlorine forms chlorate (ClO3-1). Thus, Iodine forms iodate (IO3-1). The –ite and hypo- in hypoiodite mean that iodate lost two oxygens. • Hypoiodite = IO-1

  13. Rules for expanding your list of polyatomic ions • Combos Cont’d • Ex2: What is the formula for Biperselenate? • Find Se on the periodic table. It is near S. Sulfur forms sulfate (SO4-2). Therefore, selenium forms selenate (SeO4-2). The per- in biperselenate means that selenate has gained one oxygen. Also, the bi- means that it has gained a hydrogen (don’t forget to change the charge!). • Biperselenate = HSeO5-1

  14. Monatomic Ions • For nonmetals, almost all single names that end with –ide indicates a single charged atom. • Simply write the symbol and the charge. The periodic table column indirectly indicates the element’s charge. Remember, elements want to have 8 electrons in their outer shell (Octet Rule).

  15. Monatomic Ions • Column #1 elements have a +1 charge • Column #2 elements have a +2 charge • Column #3 = +3 • Column #15 = -3 • Column #16 = -2 • Column #17 = -1

  16. Monatomic Ions • Ex1: What is the formula for chloride? Cl-1 • Ex2: What is the formula for an aluminum ion? Al+3 • Ex3: What is the name of the S-2 anion? Sulfide • Ex4: What is the name of the Mg+2 cation? Magnesium Ion

  17. 6.1 Introduction to chemical bonding • Most elements are not found alone in nature. They are “stuck” to other atoms. • Chemical Bond - Link between atoms that results from the mutual attraction of their nuclei for their electrons. • Types of chemical bonds: • Ionic - transfer of electrons (metal + nonmetal) • Covalent - sharing of electrons (2 nonmetals) • Metallic - happens in metals when there is only one type of element https://www.youtube.com/watch?v=QXT4OVM4vXI

  18. Introduction to Chemical Bonding • Covalent bonds may be polar or nonpolar • Polar - unequal sharing of electrons (HCl) • Nonpolar - equal sharing of electrons (H2) • There are two ways to predict polar vs. nonpolar ( and covalent vs. ionic)

  19. Introduction to Chemical Bonding • #1 Use electronegativity difference • 0 = nonpolar covalent • 0.4 - 1.7 = polar covalent • greater than 1.7 = ionic • Examples • NaCl Cl = 3.16 HCl Cl = 3.16 Na= - 0.93 H = - 2.20 2.23 .96 Ionic Polar Covalent • Cl2 Cl = 3.16 Cl = - 3.16 0 Nonpolar Covalent

  20. Introduction to chemical bonding • #2 - There is an easier way to predict • Ionic = metal + nonmetal or metal + p ion • Polar Covalent = 2 different nonmetals • Nonpolar Covalent = 2 of the same nonmetals

  21. Ionic Bonds • Ionic compound - a substance composed of positive and neg. ions so that the charges are equal. It involves a transfer of electrons. • Ca+2 with Cl–1 will form the compound CaCl2. • It takes two chlorine ions to cancel out the the +2 charge on the calcium ion. • Formula unit - lowest whole # ratio of ions • Ionic Bond = a METAL + a NONMETAL • Metals - lose e- - why? low IE • NM - gain electrons - why? high electronegativity

  22. Ionic Bonds • Metals lose electrons until they become like a noble gas . (8 valence e-) • Nonmetals gain e- until they do the same. • Both go to s2p6 - 8 valence e- - called a stable octet • The tendency to arrange e- so each atom has 8 is called the octet rule or rule of 8

  23. Ionic Bonds • The formation of an ionic bond: • Na to Cl = [Na]+1[Cl]-1 • Na 1s 2s 2p 3s • Cl 1s 2s 2p 3s 3p

  24. Ionic Bonds • Ionic bonding picture: • Ex1: Na to Cl= [Na]+1[ Cl ]-1 NaCl Na Cl • Ex2: Ba to Cl = [Ba]+2 2[ Cl ]-1 BaCl2 Ba Cl Cl

  25. Ionic Bonds • Ionic bonding picture: • Ex3: Al to N • Ex4: Al to S

  26. Ionic Bonds • The easy way: • Find the charge of each atom • “criss cross” the charges – charge cancels out and you are left with a neutral compound Formula Name • EX1: Al N • EX2: Na S • EX3: Al S

  27. Ionic Bonds • A few more examples Formula Name • Li and NO3-1 • Ca and C2H3O2-1 • Magnesium and Phosphite • Aluminum and hyponitrite • Calcium bromide • Aluminum sulfide

  28. Ionic Bonds • Energy is involved in all chemical reactions. • Na + Cl yields NaCl + 769 kJ • Lattice energy - energy released when an ionic compound forms. • NaCl = - 769 kJ/mole NaF = - 922 kJ/mole KCl = -718 kJ/mole • smaller ions have higher lattice energies

  29. Ionic Bonds • Properties of ionic compounds: • Hard • Shatter • Conduct electricity • High melting point • Odorless

  30. 6.4 Metallic Bonding - “Sea of electrons theory” • The nuclei are arranged in a systematic lattice. • The bond strength relies on the nuclear charge and the number of valence e- • Ex. Mg is stronger than Na • The valence electrons form a sea of free moving electrons that are attracted to multiple positive nuclei.

  31. Metallic Bonding • Conducts Electricity as a result of free electrons. • Malleability and ductility results from the nuclei's ability to move passed each other

  32. Metallic Bonding • Remember: • in ionic bonds some atoms want e- and some don’t • in covalent bonds, all atoms share – • in metals, no one atom wants the e-

  33. 6.2 Covalent Bonding • In covalent bonding atoms share electrons. In the H2 molecule, each H atom says, "I only need one more e- to be like a noble gas (helium)." Since each hydrogen has only one electron, when two hydrogens bond they can share their electrons. https://www.youtube.com/watch?v=a8LF7JEb0IA

  34. Covalent Bonding • Molecule - smallest quantity of matter that exists by itself and retains the properties of that substance. Describes a covalently bonded substance. • monatomic molecules - He, Ne, Ar, (noble gases are always monatomic) • diatomic molecules – H2 O2 N2 Cl2 Br2 I2 F2 (you must memorize these!!) • polyatomic molecules - P4, S8, C6H12O6

  35. Covalent Bonding • The formation of a covalent bond • Bond Length vs. Bond Energy •  Bond length =  Bond Energy =

  36. Covalent Bonding • Diatomic Molecules and Orbital Notation (Orbital overlap or notation diagrams): • H2 O2 H 1s O 1s 2s 2p H 1s O 1s 2s 2p N2 N 1s 2s 2p N 1s 2s 2p

  37. Covalent Bonding • Why are these atoms forming bonds? • Octet Rule- Atoms lose, gain, or share electrons to have 8 electrons in their outer shell. • HF – orbital notation H 1s F 1s 2s 2p

  38. Lewis Dot Diagrams of molecules (covalent compounds) and polyatomic ions • Basic rules • Each atom wants 8 electrons (except H wants 2). • Each atom goes for close to the right # of bonds. • The least electronegative atoms goes in the middle OR • The atom that makes the most bonds goes in the middle. (H always on the outside.) OR • The “single guy” (the atom that does not have a subscript after it) goes in the middle. • Symmetry is key!!! • Place the atoms in order (left, right, bottom, and top) around a central atom

  39. Lewis Dot Diagrams • To determine the numberof bonds: • S = N – AS = shared electrons (bonds) 2 N = needed e- (all elements need 8 except for H which needs 2) A = how many e- an atom actually has (# of valence e-) • If using S=N-A, add the charge into the Actual amount of electrons. • Put [ ] around the molecule and include the charge

  40. Lewis Dot Diagrams • Examples: Draw the following Lewis structures • EX1: CH4 • Ex2: H2O • Ex3: PCl3

  41. Lewis Dot Diagrams • Ex4: SiH2F2 • Ex5: CS2 • Ex6: C2H6

  42. Lewis Dot Diagrams • Ex7: C2H4 • Ex8: C2H2 • Ex9: CH2O

  43. Lewis Dot Diagrams • Ex10: HCN • Ex11: FON

  44. Drawing polyatomic ions • count electrons: if the charge is - 3, add 3 electrons to A • EX: PO4-3 • less bonds than atoms want = negative charge • more bonds than atoms want = positive charge • P wants 3 bonds, has 4: + 1 charge • Each O wants 2, has 1 so each O = -1 • Total = - 3 • Ex11: PO4-3

  45. Coordinate covalent bond • Coordinate covalent bond- 2 shared electrons in a bond are donated by 1 atom • Examples: • NH4+ • OH-1 • sulfate • nitrate • nitrite • carbonate • bicarbonate • H2SO4 • H3PO4

  46. 6.5 The Properties of Molecular Compounds • Valence shell electron pair repulsion theory (VESPER) – e- pairs get as far away from each other as possible • Because of this we can predict the shape of molecules based on how many bonds and lone pairs are on the central atom

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