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Chemical Bonds

Chemical Bonds. Mutual attraction that binds atoms together to form compounds. Types of Chemical Bonds. Ionic bond = bond that results from attraction between oppositely charged ions (transfer of electrons) Covalent bond = bond resulting from the sharing of electrons. Metallic Bonds.

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Chemical Bonds

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  1. Chemical Bonds Mutual attraction that binds atoms together to form compounds

  2. Types of Chemical Bonds • Ionic bond = bond that results from attraction between oppositely charged ions (transfer of electrons) • Covalent bond = bond resulting from the sharing of electrons

  3. Metallic Bonds = metal atoms in a delocalized cloud of electrons • Remember: metals like to give up electrons, so no atom “wants” the free e-

  4. Bonding happens on a spectrum • Bonds are rarely purely ionic or purely covalent • Remember Electronegativity (the ability of atoms to attract electrons) • Comparing the electronegativities of atoms involved in bond, can determine whether bond is ionic or covalent see page 161 PT for electronegativity values

  5. Electronegativity difference above 1.7 = ionic • Ex. Cs and F • Cs = 0.7 • F = 3.3

  6. Electronegativity difference of 1.7 or less = covalent • Polar-covalent (0.3-1.7), • ex. H2O: H-2.1, O-3.5, diff = 1.4 • Non-polar covalent (under 0.3) • Ex. Bonds between atoms of the same element are always non-polar (purely) covalent, ex. H2, O2, N2

  7. Covalent Bonding Forms molecular compounds that consist of molecules

  8. Molecular Compounds • Molecule = a neutral group of atoms that are held together by covalent bonds • Chemical formula (molecular formula) = indicates the number of atoms of each element in a compound • H2O • O2 • CO2 • HCl • C6H12O6 Diatomic molecule = made up of 2 atoms of the same element

  9. Diatomic Molecules (Hydrogen and the magic 7) • Some elements always exist as diatomic molecules

  10. Empirical Formula • The “reduced” form of a compound • Has the lowest ratio • Molecular formula = C6H12O6 • Empirical formula = CH2O

  11. Potential Energy and Bonding • Nature favors covalent bonding • Atoms have lower PE when they are bonded • As atoms near each other, their charged particles interact • Nucleus is attracted to e- + - PE↓ • e- repel each other - - PE↑ • 2 nuclei repel each other + + PE↑

  12. Characteristics of Covalent Bonds • Bond length = avg. distance btw 2 bonded atoms • Bond energy = E required to break the bond How are they related? Covalent Bond = overlap of orbitals where shared electrons live

  13. Octet Rule • Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet (8) electrons in its highest occupied energy level (valence) • F: ↑↓↑↓↑↓↑↓↑__ 1s 2s 2p • F: ↑↓↑↓↑↓↑↓__↓ Bonding electron pair in overlapping orbitals

  14. Example Bonding electron pair in overlapping orbitals 1s H: ↑__ 1s 2s 2p 3s 3p Cl: ↑↓↑↓↑↓↑↓↑↓↑↓↑↓↑↓__↓

  15. Exceptions to the Octet Rule • Hydrogen (1 valence) and Helium (2 valence) experience stability with only 2 electrons in the 1s-orbital (it is their highest and ONLY energy level) • Boron (3 valence) tends to form stable compounds with 6 valence e- • Ex. BF3 • Other elements can be surrounded by more than 8 e- when they bond with highly electronegative elements (like halogens) • This bonding will involve d orbitals In addition to s and p • Ex. PF5 and SF6

  16. Electron Dot Notation # valence e- electron-dot notation example

  17. Lewis Structures = formulas in which atomic symbols represent nuclei and inner shell electrons; only valence electrons are drawn • Use electron-dot notation to represent molecules • Examples: Structural Formulas indicate the kind, number, arrangement, and bonds, but not the unshared pairs of electrons

  18. Types of Covalent Bonds • Single bonds = result from the sharing of 1 electron pair • Double bonds = result from the sharing of 2 electron pairs • Triple bonds = result from the sharing of 3 electron pairs

  19. Bond Length and Bond Energy • REMEMBER: bond energy is energy required to break the bond • Shorter bonds have higher energies and are harder to break Shorter length single bonds double bonds triple bonds Higher energy

  20. Molecular Geometry = the 3-dimensional arrangement of a molecule’s atoms in space Geometry and bond polarity will determine molecular polarity and intermolecular forces. Forces of attraction between molecules

  21. Intramolecular Forces • Forces between atoms or ions that keep compounds together • Ionic attraction • Sharing electrons

  22. Intermolecular Forces • Dipole-Dipole forces = attractive forces between polar molecules • Hydrogen-bonds (H-bonds) = type of D-D force btw H and certain other elements • H-N, H-O, or H-F

  23. Intermolecular Forces (cont) 2. Ion-Dipole forces = attraction btw an ion and a polar molecule

  24. Intermolecular Forces (cont) 3. London Dispersion forces = attractive forces resulting from temporary dipoles induced by ions

  25. Naming Covalent Compounds • Use the prefixes based on the number of each type of atom and suffix –ide for 2nd element 1 = mono- CO = carbon monoxide 2 = di- CO2 = carbon dioxide 3 =tri- BF3 = boron trifluoride 4 = tetra- CCl4 = carbon tetrachloride 5 = penta- N2O5 = dinitrogenpentoxide 6 = hexa- B3F6= triboronhexafluoride 7 = hepta- 8 = octa- 9 = nona- 10 = deca Only used in naming organic compounds

  26. Naming Covalent Compounds • Exceptions: covalent compounds containing hydrogen, e.g. • H2S (hydrogen sulfide) • CH4 (methane) • NH3 (ammonia)

  27. Organic Compounds • Compounds containing carbon • Hydrocarbons = carbon chain + hydrogen • Carbohydrates = carbon + hydrogen + oxygen Propane Glucose

  28. Naming Hydrocarbons Consist of chains of carbon atoms Name = prefix + -ane

  29. Drawing Hydrocarbons

  30. Ionic Compounds = Compounds consisting of a positive ion and a negative ion held together by the attraction of opposite electrical charges The overall charge of an ionic compound = 0 The + and - ions cancel each other

  31. Formula Units • Molecules are the basic unit of covalent compounds (a molecule of water, H2O) • The smallest unit of an ionic compound is the formula unit (a formula unit of sodium chloride, NaCl)

  32. Crystal Lattice = an orderly arrangement of ions Attractive forces btw opposite ions btw nuclei and e- of adjacent ions Repulsive forces btw like ions btw e- and e- btw nucleus and nucleus

  33. Lattice Energy and Bond Strength To compare bond strength in ionic compounds, we compare lattice energy Related to the energy in the bond = how tightly the ions in a crystal are held together Remember: higher energy bonds are harder to break

  34. Molecular (covalent) vs. Ionic • Covalent and ionic bonds are strong attractions that hold the atoms in a compound together Molecules in Covalent Compounds Held together by intermolecular forces Formula Units in Ionic Compounds Held together by attractive forces in a crystalline lattice STRONGER

  35. Molecular v. Ionic Properties Molecular • Lower melting points (liquids and gases) Ionic • Higher melting points (solids) • Hard and brittle • Conduct electricity as liquids or when dissolved in water • Many are soluble in water • Some aren’t. WHY?

  36. Types of Ions • Monatomic = ions that consist of one element Cl- = chloride Br- = bromide Na+ Mg2+ Al3+ • Polyatomic = ions that consist of more than one type of atom (CO3)2- = carbonate (PO4)3- = phosphate (C2H3O2)1- = acetate (ClO3)1- chlorate (NO3)1- = nitrate

  37. Naming Binary (2 elements) Ionic Compounds Cation (metal) comes 1st • Element name Anion (nonmetal) comes 2nd • Element root + -ide

  38. Naming Ionic Compounds that Contain Polyatomic Ions • Cation (metal) • Element name • Polyatomic ion • Ion name

  39. Stock System for Transition Metals • Some transition metals (Group 3-12) can form more than one type of cation • Ex. Copper can form +1, +2, or +3 • When writing the name, use a Roman numeral to indicate the charge of the transition metal ion • Ex. Copper (I) chloride, copper (II) chloride, etc.

  40. Writing Formulas for Ionic Compounds 1. Symbol for cation + symbol/formula for cation 2. Write the charges as super scripts 3. Cross the numbers down to subscripts

  41. Naming Acids • Binary Acids = consist of 2 elements Hydrogen + a halogen • HF = hydrofluoric acid • HCl = hydrochloric acid • HBr = hydrobromic acid • HI = hydroiodic acid

  42. Naming Acids • Oxyacids = consist of H, O, and another nonmetal H2SO4 = sulfuric acid H2SO3 = sulfurous acid H2CO3 = carbonic acid HNO3 = nitric acid HNO2 = nitrous acid HC2H3O2 = acetic acid HClO4 = perchlorous acid HClO3 = chloric acid HClO2 = chlorous acid HClO = hypochlorous acid

  43. Assigning Oxidation Numbers • The oxidation # for any uncombined element or diatomic molecule is ZERO • The oxidation # for any monatomic ion is its CHARGE • H is usually +1, unless it’s combined with a metal to form a metal hydride then it’s -1 • O is usually -2, unless it’s a peroxide, then it’s -1 Ex. Zn, Cu, O2, N2, Cl2 Ex. Ba+2, K+1 HCl (+1), NaH (-1) MgO (-2), H2O2 (-1)

  44. Assigning Oxidation Numbers • In binary covalent compounds (nonmetal + nonmetal) the positive one is first and the negative one is second • The sum of the oxidation numbers for all atoms in a neutral compound is ZERO • The sum of the oxidation numbers in a polyatomic ion is equal to the CHARGE of the polyatomic ion

  45. Practice • H2 • CaCl2 • KClO4

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