Chemical Bonding

1. Chemical Bonding Unit 3 – Lecture 1 No one wants to be alone…

2. Why atoms bond • Most atoms are not found alone in quantity in the universe, they are bonded to other atoms • All atoms except the ones in Group 18 (the Noble Gases) have an unstable electron configuration • Atoms will gain, lose, or share valence electrons until they become stable

3. Reviewing the “Octet Rule” • Elements will gain, lose, or share electrons to fill their valence shells with 8 e-s • major groups (1,2,13-17) gain, lose or share until 8 or zero valence e-s • zero valence electrons means that layer is stripped and the full inner layer is exposed, making them stable

4. Counting outer shell electrons • An Electron-Dot Diagram allows us to draw atoms and show their valence electron configuration • Write the element’s symbol • Begin placing dots around the symbol equal to the number of outer-shell electrons • 1 to a side, then double up

5. Reviewing Lewis Diagrams • …show valence e- configuration • element symbol in center • surround with # of dots equal to valence e- number • e-s pair if have more than 4

6. Reviewing Oxidation #s • Oxidation number: • charge on the atom after an atom gains, loses, or shares electrons to become stable • the + or – number tells you how many e-s the atom gained or lost • positive = surplus loses e-s • negative = deficit  gains e-s

7. Oxidation/Charge Numbers +1 +2 ----------------------------------------- +3 +/-4 -3 -2 -1 0

8. Practice • Mg • valence electrons? • oxidation number? • how many e-s gained, lost or shared?

9. Practice – cont’d • Ge • valence electrons? • oxidation number? • how many e-s gained, lost or shared?

10. Practice • Cl • valence electrons? • oxidation number? • how many e-s gained, lost or shared?

11. Reviewing Ions • Ion – a charged particle • positively charged particle = cation • negatively charged particle = anion

12. Ionic Bonds • Ionic bond – typically formed btw metal & non-metal • one atom loses 1 or more e-s • Metal • one atom gains 1 or more e-s • Non-metal • attraction between the + and – charge ions cause them to bond with a lot of electromagnetic force

13. Ionic Bonds – cont’d

14. Ionic Bonds – cont’d • Compounds with Ionic Bonds • are created from elements on opposite sidesof the Periodic Table • large difference in electronegativity • example: NaCl • form repeating units

15. Ionic Bonds – cont’d • Compounds with Ionic Bonds • dissolve into aqueous solutions easily • soluble [solid dissolves in liquid] • miscible [liquid dissolves in liquid] • dissociation – particles are pulled apart by water

16. Ionic Bonds – cont’d • Compounds with Ionic Bonds • can have very strong bonds which cause them tohave unusually high melting and boiling points • exist usually as solids at room temperature

17. Covalent Bonds • Covalent Bond – formed between atoms which share valence e-s to complete their octet • once bond is formed, atoms are stuck until they bond to something else [do not dissociate].

18. Covalent Bonds – cont’d

19. Covalent Bonds – contd • Compounds with Covalent Bonds • are created from elements somewhat close to each other on the Periodic Table • can be liquids or solids at room temp • do not dissolve easily into aqueous solutions • immiscible • form distinct molecules

20. Metallic Bonds • Metallic Bond – formed between metals • positive nuclei arrange into crystal latticeelectrons flow freely from atom to atom • sometimes called a “sea of electrons” • this accounts for physical properties of metals, such as strength, malleability, ductility, thermal and electrical conductivity, opacity, and luster.

21. Metallic Bonds – cont’d • Compounds with Metallic Bonds • have strength, malleability, ductility, thermal and electrical conductivity, opacity, and luster because of flow of electrons • are solid at room temp

22. Metallic Bonds – cont’d • Compounds with Metallic Bonds • may lose e-s to substances outside of the bond • causes corrosion or oxidation [rusting] • usually happens with oxygen

23. Reviewing Electronegativity “I’m not stealing, I’m sharing unequally…” • Electronegativity – the ability of an atom to pull electrons to itself • ie, how well an atom holds onto its own electrons or steals from others • increases ↑ the PT • increases  PT

24. Reviewing Reactivity • Reactivity for Metals • increases  and  from the stair step • Cesium/Francium = most reactive metals • Reactivity for Non-Metals • increases  and  from the stair step • Fluorine = most reactive non-metal

25. Polarity &… • Polar-Covalent Bonds – form when atoms in covalent bond do not share e-s equally. • one end of compound = slightly negative • one end of compound = slightly positive

26. Polarity &… – cont’d • Polarity & Solubility • “Like Dissolves Like” • polar & ionic solutes dissolve in polar solvents • non-polar solutes dissolve in non-polar solvents • remember, “surround & separate” • polar cannot dissolve non-polar • non-polar cannot dissolve polar

27. Practice • Most atoms bond because… • lonely too many valence e-s too few valence e-s • What number of valence e-s is suitable for most atoms? • 2 8 0 6 • What model is used to show valence e-s on an atom? • Bohr Lewis Wave Scale • Atoms which have more or less e-s than normal are…? • ions positrons negatrons • How do atoms become stable? • gain e-s lose e-s share e-s all of these

28. Practice – cont’d • Atoms which have gained e-s are called…? • anion cation ion • Atoms which have lose e-s are called …? • anion cation ion • Ionic bonds are held together by a … force. • strong nuclear strong electromagnetic • weak nuclear weak electromagnetic • Atoms far apart on the PT will usually bond …. • ionically covalently metallically not at all

29. Practice – cont’d • Covalent bonds are formed when atoms.. • gain e-s lose e-s share e-s all of these • Ionic bonds are formed when atoms … • gain e-s lose e-s share e-s all of these • Groups 3-12 will use which type of bond? • ionic covalent metallic none of these • Which part of a metallic bond explains why metals are lustrous and good conductor of electricity? • crystal lattice “sea of electrons” • nuclear density metals aren’t good conductors…

30. Practice – cont’d • Atoms with a strong attraction to their electrons are …? • positronicgravometricthaleonic electronegative • What is the most reactive nonmetal? • Cl F O S • A covalent bond where the atoms do not share evenly is called a … bond • polar equatorial longitudinal latitudinal • Most atoms are found … in nature. • bonded unbonded free free-range

31. Practice – cont’d • What is the rule concerning how polar and nonpolar bonds react? • “Opposites Attract” “Like Repel” “Like dissolves like” “Surround and Separate” • How does oxidation affect a metal? • strengthens it weakens it does not affect it • Which of these is the most electronegative? • O Cl F Ar • When all atoms share electrons equally in a covalent bond, it is called • polar non-polar ionic metallic

32. Practice – cont’d • Draw the Lewis diagram for Selenium (Group 16, Period 4, Atomic # 34). • Which group of elements loses electrons most easily? • group 1 group 3 group 15 group 17 • How many dots should be drawn around the element Oxygen in its Lewis diagram?

33. Describe why water will dissolve sugar easily, but sugar will not dissolve in cooking oil.