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Chemical Bonding

Chemical Bonding. What is a Bond?. Force that holds atoms together Results from the simultaneous attraction of electrons (-) to the nucleus (+). Breaking/Forming Bonds. When a bond is broken energy is absorbed Endothermic When a bond is formed energy is released Exothermic

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Chemical Bonding

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  1. Chemical Bonding

  2. What is a Bond? • Force that holds atoms together • Results from the simultaneous attraction of electrons (-) to the nucleus (+)

  3. Breaking/Forming Bonds • When a bond is broken energy is absorbed • Endothermic • When a bond is formed energy is released • Exothermic • The greater the energy released during the formation of the bond, the greater its stability • Stable bonds require a great deal of energy to break

  4. Lewis Dot Diagrams • Use dots to represent the number of valence electrons • How to write: • Write the symbol. • Put one dot for each valence electron • Electrons go on the 4 sides, no more than 2 per side

  5. Dot Diagram Examples: • Draw dot-diagrams for the following • Mg • C • Ne

  6. Dot Diagrams - Ions • For ions, use brackets and place the charge outside the brackets • Examples: • Na+ • O2- • H+

  7. Octet Rule • Atoms will gain or lose electrons in order to have a full valence shell – like the nobles gases • “Take the shortest route” • Metals lose electrons to form positive ions (Cations) • Nonmetals gain electrons to form negative ions (Anions)

  8. Exceptions • 1st principle energy level only holds 2 electrons • Transition elements can lose valence (s) and inner (d) electrons – this is why they have multiple oxidation states • Some atoms may be stable with less than an octet – many compounds with B • Some atoms may be stable with more than an octet – elements beyond period 2, especially P and S, the additional electrons are added to the d sublevel • Molecules with an odd number of electrons – they will be unstable

  9. Types of Bonds • Ionic - Electrons are transferred from a metal to a nonmetal • Covalent - Electrons are shared between 2 nonmetals • Polar Covalent – electrons are shared unequally • Nonpolar Covalent – electrons are shared equally • Metallic - Electrons are mobile within a metal, “Sea of Electrons”

  10. Identifying Bond Type • Ionic – metal and a nonmetal • Covalent – 2 nonmetals • Nonpolar Covalent (equal) • Same atoms (diatomics, triatomics) • Polar Covalent (unequal) • Unequal electronegativity • Metallic – metals

  11. Identifying Bond Types • Indicate the type of bond present in each: • HCl • CCl4 • MgCl2 • O2 • Hg • H2O

  12. Ionic Bonds • Transfer of 1 or more electrons from a metal to a nonmetal • The greater the electronegativity difference between atoms, the greater the ionic character • Example: Sodium Chloride (NaCl) Na electron transferred to Cl Na Cl X

  13. Monatomic Ions • One atom in an ion • Look at the valence electrons to determine the charges • Examples: K+, O2-

  14. Polyatomic Ions • More than one atom in the ion • Reference Table E • Charge belongs to the entire ion, not an individual atom • Within the polyatomic ion the atoms are held together by covalent bonds • When writing it, place ( ) around the entire ion, with the charge outside • Examples: (NH4)+, (H3O)+, (CO3)2-

  15. Writing Ionic Formulas • You need an equal amount of positive and negative charges, so that the compound is neutral • Ionic Formulas are always written as empirical formulas (reduced)

  16. Examples • Na1+ + Cl1- • Mg2+ + Cl1- • Ca2+ + CO32- • Al3+ + O2-

  17. Criss Cross Method • Write the symbol for the cation and anion • Write each ion’s charge as a superscript • Criss-cross the charges to become subscripts of the other ion • Do not put (+) or (-) charges in the final formula • Reduce to least common multiple (empirical formula)

  18. Ionic Formulas • Write the formula for the compound formed from the following ions: • Mg2+ + Cl- • Ca2+ + CO32- • Al3+ + O2- • Calcium ion + hydroxide ion

  19. Naming Ionic Compounds • Name the cation first, the anion second • Cation keeps its name, anion changes its ending to –ide (Chlorine → Chloride) • Do not change the ending of polyatomic ions • Examples: • NaCl • CaCO3 • MgF2

  20. Stock System – only used for positive ions • Some cations have more than one positive oxidation states • A roman numeral is used to indicate the charge of the positive ion

  21. Stock System Examples • Iron (II) Chloride • Iron (III) Oxide • Copper (II) Oxide • a. What charge does copper have in copper II sulfate? b. What is the formula for copper II sulfate?

  22. Ionic Salts • Salts are ionic compounds made up of cations and anions • The ratio of cations to anions is always such that an ionic compound has no overall charge • Many of the ions are bonded together to form a crystal

  23. Properties of Ionic Salts • Ionic Bonds are very strong • Very high melting and boiling points • All solids at STP • Hard • Brittle

  24. Melting and Boiling Points of Compounds

  25. Properties of Salts (cont’d) • Do not conduct electricity as solids • Do conduct electricity when the salt melts or is dissolved in water (liquid phase or aqueous) • In order to conduct electricity a substance must have free moving charged particles • In the solid phase the ions are not free to move

  26. Covalent Bonds • Sharing of electrons between 2 nonmetals

  27. Non-Polar Covalent • Electrons are shared equally • Equal distribution of electrons • All diatomic molecules have non-polar covalent bonds

  28. Nonpolar Covalent Examples • Flourine (F2) • Hydrogen (H2)

  29. Polar Covalent • Unequal Sharing of electrons • Unequal distribution of electrons • Partial positive and partial negative charges • The side with the higher electronegativity will have a greater share of the electron(s) resulting in a partial negative charge • The greater the electronegativity difference, the more polar the bond is

  30. Polar Covalent Examples • HCl • H2O

  31. Dipoles • Form when the charge in a bond is asymmetrical • Present in polar bonds • Partial positive and partial negative charges

  32. Polar Bonds / Dipoles • Isn’t a whole charge just a partial charge • means a partially positive • means a partially negative Example: H - Cl +---→ • The Cl pulls harder on the electrons (more eneg) • The electrons spend more time near the Cl d+ d- d+ d-

  33. Dipole Examples • Which molecule contains more polar bonds? a. CCl4 b. CH4 2. Which has a stronger dipole? • HCl • HBr

  34. Properties of Molecular Substances (Covalent Compounds) • Soft • Low melting points and boiling points • Many exist as gases or liquids at STP • Poor conductors of heat and electricity (in all phases) Examples: H2O, CCl4, NH3, C6H12O6, O2

  35. Molecular Formulas (Covalent Compounds) • Contain covalent bonds • Tells you how many atoms are present in a single molecule • Named similarly to ionic compounds, except use prefixes to indicate the number of atoms per molecule

  36. Prefixes • Mono- is only used for the second element • Example: CO = carbon monoxide

  37. Examples • CCl4 • H2O • NO • N2O5 • BBr3

  38. Structural Formulas • Specifies how atoms are bonded together • Dashes represent bonds • 2 atoms can share up to 3 pairs of electrons

  39. Single Bonds • 2 atoms share 1 pair of electrons (2 electrons) Examples: • Ammonia (NH3) • Chlorine (Cl2) • Hydrochloric Acid (HCl)

  40. Double Covalent Bonds • 2 atoms share 2 pairs of electrons (4 electrons) • 2 bonds between 2 atoms Examples: • Carbon Dioxide (CO2) • Oxygen (O2)

  41. Triple Covalent Bond • 2 atoms share 3 pairs of electrons (6 electrons) • 3 bonds between 2 atoms Examples: • Nitrogen (N2) • Ethyne (C2H2)

  42. Bond Length/Strength • Length: • Single > Double > Triple • The more electrons in a bond, the greater the attraction, therefore shorter • As you move down a group bond length increases • Due to increasing molecular size • Strength: • Triple is the strongest, most stable, requires the most energy to break

  43. Network Solids • Covalently bonded atoms are linked into a giant network (macromolecules) • Examples: Diamond (C), Graphite (C), Silicon Carbide (SiC), and Silicon Dioxide (SiO2)

  44. Network Solids • Properties: • Hard • High melting and boiling points • Do not conduct heat and electricity

  45. + + + + + + + + + + + + Metallic Bonding • Sea of Electrons • Electrons are free to move through the solid.

  46. Properties of Metallic Solids • Very Strong • Good conductors of heat and electricity because electrons are free to move about • Luster • High melting point (except Hg) • All solids at STP (except Hg) • Malleable, Ductile

  47. VSEPR Theory • In a small molecule, the electron pairs are as far away from each other as possible • VSEPR = Valence Shell Electron Pair Repulsion

  48. Linear • Drawn on a straight line • All molecules of only 2 atoms are linear • Many 3 atom molecules are linear, if there are no unshared electron pairs on the central atom • If both ends are the same, the molecule is nonpolar (Symmetrical = Nonpolar) • If the ends are different, the molecule will be polar (Asymmetrical = Polar) • Bond Angle = 180o See Molecules • Examples: H2, CO2, HCl

  49. Tetrahedral

  50. Tetrahedral • A central atom bonded to 4 other atoms • 3-D shape allows the electron pairs to get as far away from each other as possible H 109.5º C H H H

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