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There are three types of bonding in chemistry….

There are three types of bonding in chemistry…. Ion. ic. Covalent. (metal + nonmetal). (nonmetal + nonmetal). Metallic. (metal + metal). When TWO NONMETALS combine together, they SHARE electrons to form COVALENT bonds.

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There are three types of bonding in chemistry….

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  1. There are three types of bonding in chemistry…. Ion ic Covalent (metal + nonmetal) (nonmetal + nonmetal) Metallic (metal + metal)

  2. When TWO NONMETALS combine together, they SHARE electrons to form COVALENT bonds. Properties: Because the electrons are localized between the nuclei, molecules do not stick together as tightly as if they had permanent positive and negative charges. Molecular compounds tend to have low boiling points, and be soft solidsor liquids. Covalent network solids: solids with extensive covalent bonding e.g. SiO2, diamond When NONMETALS and METALS combine together, they TRANSFER electrons to form IONIC bonds. Properties: Ionic solids are held together by strong attractive forces between positive and negative ions. This causes them to have high melting points, be brittle, and to make good electrolytes when dissolved in water.

  3. When two metals combine together, they form METALLIC bonds. The low ionization energy of a metal atom causes the electrons to only be loosely held by the nuclei. The valence electrons are delocalized and are distributed in an ‘electron sea’ around the positively charged metal centers. + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + + • Properties: It is because the electrons are not localized around a particular nucleus that metals are shiny, malleable, ductile and good conductors of heat and electricity.

  4. Now you try some… Write the formulas for the following combinations • Sodium and phosphorus • Magnesium and sulfur • Carbon and silver • Oxygen and zinc • Phosphorus and iron(II) Na3P MgS Ag4C • Remember that the cation always appears first in the formula. • Silver (almost) always has a charge of +1 when in a compound. ZnO • Zinc always has a charge of +2 when in a compound. Fe3P2

  5. The Name Game: Ionic Compounds • The cation comes first in the name (and in the formula) and it retains the name of the element. • If the compound contains a transition metal cation, the oxidation number of the cation is indicated in parentheses in roman numerals. • Ex: Au3+ gold(III) • 3. The name of the anion is changed so that it ends in -ide • Ex: chlorine  chloride, oxygen  oxide AuCl3 Na3N Sodium nitride Gold(III) chloride PbO2 Lead(IV) oxide PbBr2 Lead(II) bromide Metals in the p-block can often have more than one oxidation number

  6. Naming ionic compounds containing polyatomic ions • The cation comes first in the name (and in the formula) and it retains the name of the element or the ion if ammonium. • The polyatomic anion retains its name. Ex: CaSO4 would be calcium sulfate NH4Cl would be ammonium chloride would be ammonium perchlorate NH4ClO4

  7. The Name Game: Covalent Molecules • The first element in the formula is also first in the name and retains the name of the element. • The ending of the second element is changed to –ide. • Prefixes are used to indicate the number of each element in the molecule.

  8. What is the formula and name of the molecule formed from… 1. Carbon and chlorine Carbon tetrachloride CCl4 2. Nitrogen and bromine Nitrogen tribromide NBr3 3. Nitrogen and oxygen Dinitrogen trioxide N2O3 4. Phosphorus(V) and oxygen P2O5 Diphosphorus pentoxide

  9. Suffixes: -ate -ic -ite -ous Naming Acids Binary acids: contain hydrogen and one other element • First part of name is hydro- • Second part is the root of the second element with the suffix –ic • Ex: HCl is hydrochloric acid Oxyacids: acid form of a polyatomic anion that contains oxygen Format: Root of anion + suffix acid HNO3 HNO2 Nitric acid Nitrous acid

  10. A few examples HBr Hydrobromic acid Perchloric acid HClO4 H3PO4 Phosphoric acid Phosphorous acid H3PO3 HClO2 Chlorous acid HClO Hypochlorous acid Hydrosulfuric acid H2S HCl Hydrochloric acid PO4-3 phosphate ClO4-1 perchlorate ClO2-1 chlorite PO3-3 phosphite ClO3-1 chlorate ClO-1 hypochlorite

  11. Warming up with a few ionic compounds: 1. Ca2C Calcium carbide 2. (NH4)3PO4 Ammonium phosphate 3. Mg(OH)2 Magnesium hydroxide Potassium oxide 4. K2O 5. NH4Cl Ammonium chloride Tin(II) nitrate 6. Sn(NO3)2 Iron(II) hydrogen phosphate 7. FeHPO4 Copper(I) sulfide 8. Cu2S Nickel(II) acetate 9. Ni(C2H3O2)2 Mercury(I) perchlorate 10. HgClO4

  12. Working out the kinks with a few covalent compounds: Carbon dioxide 1. CO2 Dinitrogen tetroxide 2. N2O4 Sulfur hexafluoride 3. SF6 Phosphorus pentachloride 4. PCl5 Sulfur trioxide 5. SO3 6. H2S Dihydrogen sulfide 7. BCl3 Boron trichloride 8. NO Nitrogen monoxide 9. CF4 Carbon tetrafluoride 10. XeF4 Xenon tetrafluoride

  13. Feel the burn! Ionic and covalent together: 1. K2SO4 Potassium sulfate 2. PCl3 Phosphorus trichloride 3. BF3 Boron trifluoride 4. (NH4)3PO4 Ammonium phosphate 5. ZnCl2 Zinc chloride 6. Au2O3 Gold(III) oxide 7. SeF6 Selenium hexafluoride 8. NCl5 Nitrogen pentachloride 9. OCl2 Oxygen dichloride 10. Ca(OH)2 Calcium hydroxide

  14. Cooling down with some acid naming: 1. HNO3 Nitric acid 2. HBr Hydrobromic acid 3. H2SO4 (battery acid) Sulfuric acid 4. HF (dissolves glass) Hydrofluoric acid 5. HClO4 Perchloric acid (ignites organic cmpds) 6. H2SO3 Sulfurous acid 7. HBrO3 Bromic acid 8. HCl (stomach acid) Hydrochloric acid 9. H3PO4 Phosphoric acid (acid in colas) 10. HClO Hypochlorous acid

  15. •• •• • • • •CC• •Si• •O• •B• •N• •• •• • • • • • Si H H H H •• •• •• •• •• •• •• • CC • • • • • B H● ●H ●Cl● O N H● ●H •• •• • ●F● • •• •• •• •• •• •• •• •• •• •• •• • • • • ●Cl● ●Cl● • • • • • • • • • • • • ●F● ●F● ●F● ●F● ●F● ●F● •• •• •• •• •• •• •• •• Simple Lewis Dot Structures 1. Determine the central atom. This is usually the atom that you have the least of in the formula. Ex: H2O BF3 NCl3 C2H6 SiF4 2. Draw the electron dot structure of the central atom. 3. Wherever there are unpaired electrons on the central atom, pair them up with unpaired electrons on the attached atoms. 8 24 26 14 32 H2O: 2+6 = 8 NCl3: 5 + 3(7) = 26 SiF4: 4 + 4(7) = 32 BF3: 3 + 3(7) = 24 C2H6: 2(4) + 6(1) = 14 4. Finally, check your electron count.

  16. The ‘Keep ‘em Happy’ approach to Lewis structures: 1. Add up the number of valence electrons present 2. Draw the stick-skeleton of the molecule 3. Satisfy the octet rule for all atoms in the molecule Exceptions: H only needs 2e- and B only needs 6e- 4. Count up the number of electrons present in the Lewis structure - If there aren’t enough e- (the molecule is ‘unhappy’), add the missing electrons to the central atom - If there are too many e- (the molecule is ‘too happy’), take the excess away from the central atom and then form double bonds with the terminal atoms to satisfy the octet.

  17. O S O Examples: Fc = 7-7=0 Fc = 6 – 6 = 0 This is the only possible structure because Fluorine never forms double bonds SF4 F F S #VE = 34 F F 32 34 Formal charge: #VE – (#e in LP + 1e/bond) Note that sulfur’s ox# is +4 and that there are 4 bonds to S. • However if we take into account the fact that sulfur’s ox# is +4, the structure with 2 double bonds would seem to be better 0 +1 -1 SO2 O S O #VE = 18 20 18 57% of the time, SO2 has this structure 0 0 0

  18. Molecular Geometry The shape of molecules The shape of a molecule plays a very important role in determining its properties. Properties such as smell, taste, and proper targeting of drugs are all the result of molecular shape.

  19. Lewis structures tell us nothing about how atoms in a molecule are arranged in 3-dimensional space. Could you have predicted the molecular geometry of carbon tetrachloride from its Lewis structure?

  20. A useful model for predicting the shape of molecules is the… VSEPR Theory Valence Shell Electron Pair Repulsion Theory • Molecules will adopt a shape that is lowest in energy • A low energy shape is one that minimizes the valence shell electron pair repulsion (VSEPR) between adjacent atoms (electrons in bonds and in lone pairs repel each other).

  21. 109.5° 90° Methane (CH4) You might think this is the farthest that the hydrogens can get away from each other But if you think in 3 dimensions, this shape actually causes less repulsion between the bonding pairs of electrons.

  22. The 5 Main Shapes Linear 180° Trigonal planar 120° Tetrahedral 109.5° Octahedral 90°, 180° Trigonal bipyramidal 120°, 180° Molecules adopt a geometry that minimizes electron-electron repulsions  this occurs when e- pairs are as far apart as possible.

  23. Steps to determining molecular geometry: • Draw a Lewis structure • Count the # of bonds and # of lone pairs around the central domain (electronic domains) • -Single, double and triple bonds count as ONE domain • -Each lone pair counts as ONE domain • Use AXE chart to determine shape • (the name of the molecular geometry is based on position of the atoms, not on the domains)

  24. VSEPR Notation: • Also known as “AXE” notation • A = central atom • X = # atoms bonded to the central atom • E = # of lone pairs on central atom Examples: CH4 is AX4 NH3 is AX3E AX2E2 H2O is

  25. Let’s look at a few examples…

  26. O H H Trigonal planar bent bent Trigonal pyramidal

  27. Note: Lone pairs take up more space than bonding pairs and thus decrease the predicted bond angle CH4: Tetrahedral, 109.5° NH3: Trigonal pyramidal, 107° …the bond angle has been reduced by the one lone pair H2O: Bent, 104.5° …the bond angle has been reduced by the two lone pairs

  28. Hybrid Orbitals: atomic orbitals formed by blending different orbitals together Just like with other hybrids, the characteristics of the hybrid orbitals will depend upon the traits of the ‘parent’ orbitals.

  29. C 2s2 1s2 2p2 Consider carbon… How many valence electrons does carbon have? 4 According to its electron dot structure, how many unpaired electrons does carbon have? What does carbon’s orbital filling diagram look like? Atomic carbon only has 2 upaired electrons! In order to form chemical bonds, carbon’s atomic orbitals must hybridize to form molecular orbitals.

  30. one p orbital one s orbital one s orbital two p orbitals one s orbital three p orbitals + two sp hybrid orbitals (AX2) + three sp2 hybrid orbitals (AX3) + four sp3 hybrid orbitals (AX4)

  31. H H H H H H O O H H Sigma bonding When two atomic orbitals (hybridized or not) overlap end-on, they form a single sigma bond. + s bond + 2 sp3 AX2E2

  32. Pi bonding When two atomic orbitals overlap in a side-to-side fashion, they form a pi bond. Orbitals that form pi bonds are usually NOT hybridized. p bond • One unhybridized p orbital results in one pi bond. • A sigma and a pi bond form a double bond. O O s bond

  33. Multiple pi bonds p bond 2 s N N s bond

  34. 2- 2- O O 2- 2- 2- 2- C C O O O O O O O O 2- O O S S O O O O S S O O O O O O Resonance Structures : equivalent Lewis dot structures Draw the stick structure for CO32- AX3 so sp2 hybrid. O C O O AX4 so sp3 hybrid. Draw the resonance structures for SO42-

  35. Effect of resonance structures: C6H6 Unhybridized p-orbitals Localized p bonds Delocalized p bonds

  36. Cl Cl Ca2+ 2 Cl(g) Ca Cl- Cl- Ca(g) Lattice energy Ca Ca Ca Ca Ca Ca Cl Cl Ca(s) Cl2(g) CaCl2(s) Calculating Lattice Energy: The Born-Haber Cycle IE1+IE2 Enthalpy 2 EA 0 + DHf

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