2Chemical Bonds Three basic types of bonds Ionic Covalent Metallic Electrostatic attraction between ionsCovalentSharing of electronsMetallicMetal atoms bonded to several other atoms
3Why do atoms bond? Why DON’T some atoms bond? The noble gases – why?Why do other atoms bond, then?They are more chemically stable when bonded
4How do atoms bond? The octet rule The octet rule, or rule of eight, says that an atom will strive for a full s and p subshellAtoms will either lose or gain electrons to get 8 in the outer shellNOTE: when an atom loses or gains electrons, it’s nucleus remains the same – only the outer electron shell has changed!!!
5Bonding and energy changes Energy is the ability to do workStability is a measure of inability to do workSo, the lower the energy, the more stable something is!When atoms bond, the process favors stability (lower energy). Things will never go from a stable to an unstable state on their own!
6Electrons, bonding, and IONS When they do this, they get a CHARGE, because protons (+) and electrons (-) are no longer equal. They are now IONSPositive and negative IONS come together and balance each other out in IONIC BONDS.
7Cations and AnionsRemember:+ ++An “antion”A “plussy cat”
8Ionic BondingSodium wants to GIVE an electron, Chlorine wants to GET an electron.
9Ionic BondingThe low ionization energy of sodium and the high electronegativity of chlorine is one reason this works so well.
10Energetics of Ionic Bonding As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium.
11Energetics 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!
12Energetics 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.
13Lattice 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.
14Lattice EnergyLattice energy, then, increases with the charge on the ions.It also increases with decreasing size of ions.
15Energetics of Ionic Bonding By accounting for all three energies (ionization energy, eletronegativity, and lattice energy), we can get a good idea of the energetics involved in such a process.
16Energetics 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.
17Naming ions Monatomic ions = Polyatomic ions = Naming monatomic ions One atom ionsPolyatomic ions =Many atom ionsNaming monatomic ionsTo name positive ions, just add the word “ion”To name negative ions, drop the last part of the word, and add “-ide ion”
18Naming monatomic ions Rubidium loses an electron to become Rb+ Rubidium ionCalcium loses two electrons to become Ca2+Calcium ionChlorine gains an electron to become Cl-Chloride ionOxygen gains two electrons to become O2-Oxide ionNitrogen loses three electrons to become N3-Nitride ion
19Compounds made of two monatomic ions These are called BINARY COMPOUNDSYou always put the positive part first and the negative part last: Na+ + Cl- NaClNames = name of the positive ion + name of the negative ion: Sodium Chloride
22Back to ions: Writing Ionic Formulas The nomenclature (naming system):Write the symbols for the ions side by side. Write the cation first.Al3+ O2-Find the lowest common multiple that will make the charges on each ion cancel outAl O2-Check the subscripts for the lowest whole number ratio of ions. Then write the formula. Al2O3
23d-block naming Write the electron configuration for Iron. Predict the oxidation number
24d-block The d-block (yo) has its own rules Metals in the d-block have variable chargesAll d-block metals must get a ROMAN NUMERAL to indicate the chargeEXAMPLE: copper (I) chloride is made of Cu1+ and Cl-EXAMPLE: copper (II) chloride is made of Cu2+ and Cl-Don’t use roman numerals if you don’t have to
25ExamplesWrite the formulas forTin(II) iodideCobalt(III) chloride
26Working backwardIf you are given the formula you need to calculate the charge of the d-block metal.Assume the anion did not change its charge (they are very consistent)Example: FeO, to write the name we need the charge of iron.
28More d-block (old school) Roman numerals are the ‘new’ way.The ‘old’ way is based on Latin namesTwo endingsicousic is for the highest chargeous is for the lower chargeExampleFerric = Iron(III)Ferrous = Iron(II)
29Old School The only three I expect you to know are: Tin (Sn) Stannic = Tin(IV)Stannous = Tin(II)Copper (Cu)Cupric = Copper(II)Cuprous = Copper(I)
30Polyatomic ionsWhen two or more ions are clumped together it is a polyatomic ions.They usually end with –ates or -ites
31Nomenclature of Oxyanions They are not standard!Example Sulfate vs PhosphateNomenclature examplesPerchlorateChlorate NitrateChlorite Nitritehypochlorite
32Writing formulas for compounds with polyatomic ions Polyatomic ions should ALWAYS be treated like BOY BAND. Don’t ever break it up!If you need more than one polyatomic ion to balance a charge, use PARENTHESES ( )
33Polyatomic ions Naming compounds that contain polyatomic ions: The steps are the same:the name of the first ion + the name of the second:NH4+ = ammonium ion (polyatomic)Cl- = chloride ion (monatomic)NH4Cl = ammonium chloride
34Write the formulas for: ExampleWrite the formulas for:potassium perchloratetin(IV) sulfateIron(II) chromateammonium sulfate
35Ionic vs. Metallic bonds In an IONIC BOND, the electrons of one atom (that wants to lose electrons) are donated to the electrons of another atom (that wants to gain electrons). The charges on each ion balance each other out and equal ZERO.In a METALLIC BOND, all the atoms are the same (all copper, for example) and the electrons don’t belong to any one atom. They move around a lot – that’s why electricity is conducted.
36Metallic Bonds A “sea” of mobile outer electrons. Low ionization energies means the atoms don’t hold electrons well.