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  1. 1. BondingHow 92 Naturally Occurring Elements Combine to Create Earth’s Tremendous Variety
  3. 3. Definition• Valence Electrons: are electrons in the outermost shell (energy level). – The Electrons in the S-Block and P-Block – DOES NOT INCLUDE electrons in the D-Block and F-Block • They are only filled after a new valence shell has been started.• They are the electrons available for bonding.• The number of valence determines what type of bonds will usually be formed
  4. 4. Group 1 (alkali metals) have 1valence electron
  5. 5. Group 2 (alkaline earth metals)have 2 valence electrons
  6. 6. Group 13 elements have 3valence electrons
  7. 7. Group 14 elements have 4valence electrons
  8. 8. Group 15 elements have 5valence electrons
  9. 9. Group 16 elements have 6valence electrons
  10. 10. Group 17 (halogens) have 7valence electrons
  11. 11. Group 18 (Noble gases) have 8valence electrons, excepthelium, which has only 2
  12. 12. Transition metals (“d” block)have 1 or 2 valence electrons
  13. 13. Lanthanides and actinides(“f” block) have 1 or 2 valenceelectrons
  14. 14. Lewis Dot DiagramsLewis Dot Diagrams are a tool to help you visually represent the valence shell.Note: Lewis Dot structures follow Hund’s Rule and the Pauli Exclusion Principle 1 valence e- 2 valence e- 3 valence e- 4 valence e- X X X X 5 valence e- 6 valence e- 7 valence e- 8 valence e- X X X X
  15. 15. Dot Notations: Period 2Lewis dot notations for the valence electrons ofthe elements of Period 2. lithium beryllium boron carbon Li Be B C nitrogen oxygen fluorine neon N O F Ne
  16. 16. The Octet Rule• Octet Rule: Atoms tend to gain, lose, or share electrons in order to gain a full set of 8 valence electrons. – Want to be like Noble Gas – What about Period 1?• When atoms gain/loose electrons they become ions. – Anion: Gains Electrons (Net – Charge) – Cation: Looses Electron (Net + Charge)• Oxidation Number: Represents the charge of a typical ion of the element. – Tells you how many electrons are typically gained or lost.
  17. 17. Bonding• Chemical Bond: Any force that holds two atoms together. – Created when two (or more) atoms interact with one another in a way that fulfills the Octet Rule.• Two or more different Elements bond together to form Compounds. – Compounds can have very different chemical and physical properties than their component elements.
  18. 18. Sodium (Na)
  19. 19. Chlorine (Cl)
  20. 20. Horses Need Gas Masks Too…
  21. 21. Sodium Chloride (NaCl)
  22. 22. Metal to Non-Metal BondingIONIC BONDING
  23. 23. Ionic Bonding
  24. 24. Ionic Bonding• Ionic Bonding: Occurs when one atom donates an electron and another receives the electron.• The exchange (gain/loss) of an electron creates an Anion (- Charge) and Cation (+ Charge). – Remember opposites attract! • The electrostatic attraction between the Anion and Cation causes them to stick together forming an ionic bond. – Also known as electrostatic bonds.• 1 or more electrons are actually exchanged between the two atoms.• Occur when there are large differences in ionization energy and electronegativity. – Between metals and non-metals.
  25. 25. Examples of Ionic compoundsMg2+Cl-2 Magnesium chloride: Magnesium loses two electrons and each chlorine gains one electronNa+2O2- Sodium oxide: Each sodium loses one electron and the oxygen gains two electronsAl3+2S2-3 Aluminum sulfide: Each aluminum loses two electrons (six total) and each sulfur gains two electrons (six total)
  26. 26. Metal Monatomic Ion name CationsLithium Li+ LithiumSodium Na+ SodiumPotassium K+ PotassiumMagnesium Mg2+ MagnesiumCalcium Ca2+ CalciumBarium Ba2+ BariumAluminum Al3+ Aluminum
  27. 27. Nonmetal Monatomic Ion Name AnionsFluorine F- FluorideChlorine Cl- ChlorideBromine Br- BromideIodine I- IodideOxygen O2- OxideSulfur S2- SulfideNitrogen N3- NitridePhosphorus P3- Phosphide
  28. 28. Sodium Chloride Crystal LatticeIonic compounds form solidcrystals at ordinarytemperatures.Ionic compounds organizein a characteristic crystallattice of alternatingpositive and negative ions. These ionic bonded crystals are known as salts.
  29. 29. Properties of Ionic CompoundsStructure: Crystalline solidsMelting point: Generally high (strong bonds)Boiling Point: Generally high (strong bonds)Electrical No conductivity in solid Conductivity: form. Excellent conductors when molten and aqueous (dissolved in liquid).Solubility: Water: Soluble Alcohol: Insoluble
  30. 30. Non-Metal to Non-Metal BondingCOVALENT BONDING
  31. 31. Covalent Bonding• What happens when both atoms want to gain electrons? – Covalent Bonding: Bond formed when two atoms share one or more valence electrons to achieve full octets. – Occurs when both have have high ionization energies and electronegativity. • Both try to take electron from each other… tug-of-war. • Non-Metal to Non-Metal• Typically less strong than ionic bonds.• The term “molecule” is used exclusively for covalent bonding F F
  32. 32. The Octet Rule: The Diatomic Fluorine MoleculeF Each has seven 1s 2s 2p valence electrons Each Wants 1 MoreF Electron 1s 2s 2p F F Single Bond: 1 Electron is shared by each atom (total of 2 shared electrons).
  33. 33. The Octet Rule: The Diatomic Oxygen MoleculeO Each has Six 1s 2s 2p valence electrons Each Wants 2 MoreO Electrons 1s 2s 2p O ODouble Bond: 2 Electrons are shared by each atom (total of 4 shared electrons).
  34. 34. The Octet Rule: The Diatomic Nitrogen MoleculeN Each has Five 1s 2s 2p valence electrons Each Wants 3 MoreN Electrons 1s 2s 2p N NTriple Bond: 3 Electrons are shared by each atom (total of 6 shared electrons).
  35. 35. Representing Covalent Bonds Covalent Bonds (shared electrons pairs) can berepresented by two dots (:) or by a single line ( - ) O= O O O
  36. 36. Properties of Covalent MoleculesStructure: Molecular ChainsMelting point: Generally low (weak bonds)Boiling Point: Generally low (weak bonds)Electrical Poor conductors (no free Conductivity: ions or electrons)Solubility: Water: Insoluble Alcohol: Soluble
  37. 37. Metal to Metal BondingMETALLIC BONDING
  38. 38. Metallic Bonding• Metals tend to be electron donors.• When metals bond with other metals their valence electrons become delocalized (free to move around) and form what is known as an electron sea. o Valence electrons do not belong to any one atom.• A metallic bond results from the attraction between the metal cations and the surrounding sea of delocalized electrons
  39. 39. Packing in MetalsLattice Model: Packinguniform, hard spheres tobest use available space.This is called closestpacking. Each atom has12 nearest neighbors.• Vacant p and d orbitals in metals outer energy levels overlap, and allow outer electrons to move freely throughout the metal
  40. 40. Metal AlloysSubstitutional Alloy: some metal atoms replaced by others of similar size.Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms.
  41. 41. Properties of Metals Metals are good conductorsof heat and electricity (lots offree electrons). Metals are malleable Metals are ductile Metals have high tensilestrength Metals have luster