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Chem 1151: Ch. 4

Chem 1151: Ch. 4

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Chem 1151: Ch. 4

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  1. Chem 1151: Ch. 4 Forces between particles

  2. Noble Gas Configuration • The first widely-accepted theory for chemical bonding was based on noble gas configurations. • Noble gases have a filled valence shell (2e- for He and 8 e- for everything else). • Because noble gases are very unreactive, and chemical reactivity depends on electronic structure, two scientists (Lewis and Kossel) concluded that this represented stable (i.e., low energy) configuration. • Formed the basis for the “octet rule” (1916): 8 electrons in valence shell results in greater stability. Noble gases have filled s and p subshells. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  3. Valence Shell e- • Valence shell is the outermost shell where electrons reside. • From the electronic configuration (or other means) you can determine that the valence shell will have the highest number n. • This means that for transition metals, even though the last electrons may be added to the d subshell for n, the valence shell e- will be found in the s subshell at n+1. • Ex. Yttrium (Y) 1s22s22p63s23p64s23d104p65s24d1 2p 1s 1s 2s 1s 2s 3s

  4. Valence Shell e- in Lewis Structures • Lewis Structure (electron dot formula): Simplified way to represent valence shell e-. • Element symbol represents nucleus, with valence shell e- represented by dots surrounding it. • The number of valence e- can be found by looking at the group number, except for the transition metals, which all have filled s orbitals at n+1. Atom # Val Shell e - Lewis Structure C 4 C 2 Sr Sr 5 In In O 6 O Po 6 Po

  5. Ions • Ion: Atom or molecule that has either lost or gained electrons from valence shell resulting in a net charge (positive or negative) compared to the number of protons. • Ionization Energy: Energy required to remove an e- from an atom. • Common atomic ions you should know: • H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+, • N3-, P3-, O2-, S2-, F-, Cl-, Br- • For Group A elements, the number of e- gained or lost results in an electronic configuration like that of the noble gases (valence shell octet). • If Na lose e-, has electronic configuration of Ne. • If Cl gains e-, has electronic configuration of Ar. • If Ca loses 2 e-, has electronic configuration of Ar.

  6. Ionic Compounds • Ionic compounds are formed when valence electrons lost by metal are gained by non-metal with which it is reacting. • Electron(s) cannot be lost from one atom unless there is another atom available to accept the electron(s). • Common atomic ions you should know: • H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+, • N3-, P3-, O2-, S2-, F-, Cl-, Br- • Common molecular/polyatomic ions you should memorize: • OH- (hydroxide), NH4+ (ammonium), SO42- (sulfate), SO32-(sulfite), PO43- (phosphate), NO2- (nitrite), NO3-(nitrate), CO32- (carbonate) • Ionic compounds formed by these ions will have neutral charges. NH3 HBr NaOH Pb(OH)2 MgCl2 HCl (NH4+)2SO4 H3PO4 AgNO3

  7. Chemical Bonding • Ionic bond: Attractive force that holds ions of opposite charge together. • Involves transfer of e- from one component to the other. • Occurs between positively-charged metal (loses 1 or more e-) and non-metal atom or molecule (gains 1 or more e-). • Usually satisfies octet rule • Common to inorganic chemistry • Covalent bond: Formed by sharing of electrons. • Occurs between: • Two non-metals • Nonmetal and metalloid • Two metalloids • Usually satisfies octet rule • Common to organic chemistry

  8. Chemical Bonding • Ionic bond: Attractive force that holds ions of opposite charge together. • Covalent bond: Formed by sharing of electrons.

  9. Electron-half-equations (for redox rxns) Na  Na+ + 1e- oxidation Each of these represents ½ of the rxn Cl + 1e- Cl- reduction

  10. Naming Ionic Binary Compounds [Name of Metal] + [nonmetal stem + ide] = • potassium (K+) + chlorine (Cl-)  potassium chloride (KCl) • strontium (Sr2+) + oxygen (O2-)  strontium oxide (SrO) • 3 calcium (Ca2+) + 2 nitrogen (N3-)  calcium nitride (Ca3N2) • Some metals may form more than 1 type of charged ion. • Exs: Cu+ and Cu2+; Fe2+ and Fe3+ • Compounds with these ions would be named by adding a roman numeral equivalent to charge in parentheses after metal name: [Name of Metal] + [nonmetal stem + ide] = • copper (Cu+) + chlorine (Cl-)  copper(I) chloride (CuCl) • iron (Fe2+) + 2 chlorine (Cl-)  iron(II) chloride (FeCl2) • iron (Fe3+) + 3 chlorine (Cl-)  iron(III) chloride (FeCl2)

  11. Units of Ionic Compounds • Stable form of ionic compound is not a molecule, but a crystal lattice where ions occupy lattice sites. • Molecular compounds have molecular weight • Ionic compounds have formula weight • Although ionic compounds form crystal lattices, we still represent them the same as molecular compounds when discussing formula weight Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  12. Lewis Structures for Covalent Compounds • Use molecular formula to determine how many atoms of each type. • Draw a structure with the elements relative to each other. • Determine total number of valence e- for all atoms. • Put one pair of e- between each bonded pair of atoms, subtract this number of e- from total. • Use remaining e- to form octets (except for H and He) for all atoms. • If all octets cannot be satisfied with available e-, move nonbonding e- pairs between bonded atoms to complete octets (will form double or triple bonds). Cl + Cl  Cl2 7 + 7 = 14 e- Lone pairs or nonbonding e- Cl Cl Cl Cl Now fill in remaining 12 e- Use 2 e- to form bond

  13. More Lewis Structures for Covalent Compounds N + N  N2 + N N N N H N + 3H  NH3 + 3 N H N H H O + 2H  H2O H + + H H O H O

  14. Still More Lewis Structures for Covalent Compounds 6 + 6 + 6 + 6 = 24 e- S + 3O  SO3 O O O S O S O S O O O O 24 - 6 = 18 e- Confirm all octets satisfied. Confirm all e- accounted for.

  15. …And One More Lewis Structures for Covalent Compounds 4 + 4 + 1 + 1 = 10 e- C2H2 H C C H H C C H 10 - 6 = 4 e- H C C H H C C H Confirm all octets satisfied. Confirm all e- accounted for.

  16. Lewis Structures for Polyatomic Ions Follow the same instructions indicated for covalent compounds, but add or subtract the number of e- indicated by the ion charge. Use molecular formula to determine how many atoms of each type. Draw a structure with the elements relative to each other. Determine total number of valence e- for all atoms. Put one pair of e- between each bonded pair of atoms, subtract this number of e- from total. Use remaining e- to form octets (except for H and He) for all atoms. If all octets cannot be satisfied with available e-, move nonbonding e- pairs between bonded atoms to complete octets (will form double or triple bonds).

  17. Lewis Structures for Polyatomic Ions 6 + 6 + 6 + 6 +6 + 2 = 32 e- SO42- 2- O O S O S O O O O O 32 - 8 = 24 e- Confirm all octets satisfied. Confirm all e- accounted for.

  18. Lewis Structures for Polyatomic Ions 5 + 6 + 6 + 6 +6 + 3 = 32 e- PO43- 3- O O P O P O O O O O 32 - 8 = 24 e- Confirm all octets satisfied. Confirm all e- accounted for.

  19. Shapes of molecules and Polyatomic Ions Molecules have 3-D shapes. • VSEPR (valence-shell electron-pair repulsion) Theory: Electron pairs in valence shell of atoms are repelled by other pairs and try to get as far away from each as possible. • Shapes around central atom (any atom bonded to 2 or more other atoms) can be predicted by VSEPR. 2 Rules for determining shape: • All valence shell electron pairs around central atom are counted equally (both bonding and non-bonding). • Double or triple bonds are treated like a single bond when predicting shapes.

  20. Electron Pair Arrangements • According to the VSEPR theory, the arrangement of electron pairs around the central atom (represented by E) depends on the number of electron pairs. • Two pairs locate opposite each other. • Three pairs arrange themselves in a flat triangle around the central atom. • Four pairs become located at the four corners of a pyramid-like shape called a tetrahedron. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  21. Shapes of molecules and Polyatomic Ions Linear Triangular Tetrahedral

  22. One Molecule, Different Geometries

  23. Polarity of Bond • Nonpolar Covalent Bond: Electrons forming bond between 2 atoms spend nearly equal time between both atoms. • Polar Covalent Bond: Covalent bond where electrons polarized and remain closer to atom with higher EN (bond polarization). • Electronegativity (EN): Ability of an atom to attract shared e- of covalent bond. • Ionic Bond: Electrons are transferred (highly polar). Polarity indicated by arrow pointed towards destination of e- Cl Cl H Cl Na Cl nonpolar polar ionic

  24. Polarity of Molecules • Nonpolar Molecules: Charge distribution from bond polarizations is symmetric. • Polar Molecules (dipoles): Charge distribution from bond polarizations is not symmetric. • Need to consider all charges to determine direction of polarity. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  25. Binary Covalent Compounds • Hazards of dihydrogen monoxide • http://www.dhmo.org/facts.html Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  26. Naming Binary Covalent Compounds • Give the name of the less EN element first. • Give the stem of the name of the more EN element next, and add suffix –ide. • Indicate the number of each type of atom in molecule using numeric prefix. CO carbon monoxide N2O5 dinitrogen pentoxide CO2 carbon dioxide CCl4 carbon tetrachloride H2O dihydrogen monoxide S2O7 disulfur heptoxide Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

  27. Naming ionic compounds containing polyatomic ions • Give the name of the metal first. • Make sure that charges add up to zero. • Put parentheses around polyatomic ions if more than 1 used. Write formulas for the following and name them: K and PO43- K3PO4 potassium phosphate Na and NO3- NaNO3 sodium nitrate Mg and PO43- Mg3(PO43-)2 magnesium phosphate NO3- and NH4+ NH4NO3 Ammonium nitrate

  28. Crystal Lattices of non-ionic compounds • Most pure substances (elements or compounds) form crystal lattices in solid state. • These lattice sites may be occupied by neutral atoms or molecules instead of ions. • Components of lattice sites held together by covalent bonds. • Network Solids: Lattice formed by atoms bound in covalent bonds (ex. Si and O). • Metallic Bond: Lattice of metal ions • In this structure, valence e- can move about more freely, which is how metals can conduct heat and electricity. Copper Graphite: Each Carbon is covalently bonded to 3 other carbons in ring Diamond: Each carbon is bonded to 4 other carbons http://www.eduys.com/Copper-Molecular-Structure-Model-303.html

  29. Interparticle Forces δ- δ+ • Dipolar forces: Attraction between positive end of one polar molecule and negative end of another (usually weak, ~0.5 -2.0 kcal) • Hydrogen Bonding: Attractive interaction of a hydrogen atom with an electronegative atom (e.g. N, O, F) from another molecule or chemical group. • The H must be covalently bonded to another electronegative atom. • Stronger than dipolar and dispersion (12-16 kcal) • Dispersion Forces: Momentary nonsymmetric electron distributions in molecules (very weak) δ- Hydrogen bond δ+ δ- δ+ Intermolecular dipolar δ- δ+

  30. Relative Strengths of Interparticle Forces Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011