Chapter 6 Part 1 “ Ionic Bonding ”

1. Chapter 6Part 1“Ionic Bonding”

2. Introduction to Bonding • Chemical bond: an interaction between atoms or ions that results in a reduction of the potential energy of the system, thereby becoming more stable • Three types of bonds: ionic, metallic, and covalent • The bond type can be deduced by the difference in atoms’ electronegativities

4. More • Subtract the electronegativities to determine the nature of the bond • If the difference is greater than 1.7, the bond is ionic • If the difference is from 0 to 0.3, the bond is non-polar covalent • If the difference is from 0.3 to 1.7, the bond is polar covalent

6. Summary • If the atoms have very different electronegativities, then ionic bonding occurs • If they both have high electronegativities, then covalent bonding occurs • If they both have low electronegativities, then metallic bonding occurs

7. Metallic Bonding • Delocalized: refers to valence electrons that are free to move throughout the structure • In metals, valence electrons are delocalized and can move throughout the metal (shared by all atoms) • Metals are therefore composed of cations surrounded by delocalized electrons

8. More About Metallic Bonding • Sometimes referred to as the electron-sea model • Metallic bonding is the electrostatic attraction between the metal ions and the delocalized electrons

9. Physical Properties of Metals • Ductile: can be drawn into a wire • Malleable: can be hammered into thin sheets • Both of these properties are attributed to the metal's ability to have layers pushed so they slide over one another without disrupting the metallic bonding • Some metals are better at this than others

10. Thermal and Electrical Conductivity • When a voltage is applied, the delocalized electrons are free to move • This also aids in conducting heat as the movement of the electrons carries kinetic energy

11. Practice: What kind of bond? • Na and Cl • Sr and O • C and O • Ni and Fe • N and O • H and B • Ti and Cr • ionic • ionic • polar covalent • metallic • polar covalent • non-polar covalent • metallic

12. Comparison of Properties • The physical properties of a substance depend on the forces between the particles of the chemical species that it is composed of • The stronger the bonding and the intermolecular forces, the harder the substances, and the higher the boiling point

13. Comparison of Propeties • Presence of impurities always lowers the melting point as the regular lattice is disrupted • Volatility: a qualitative measure of how readily a liquid or solid is vaporized upon heating or evaporation

14. Volatility

15. Metallic Structure • Variable hardness, malleable rather than brittle • Melting points and boiling points vary, depending on the number of valence electrons, but is generally high • Good electrical and thermal conductivity as solids and liquids • Insoluble, except in other metals to form alloys

16. Ionic Structure • Hard, but brittle • High melting and boiling points • Do not conduct heat and electricity as solids, but do when molten or in solution • More soluble in water than other solvents • Examples: NaCl, CaO • Metal + nonmetal

18. Covalent Structure • Usually soft and malleable unless H-bonded • Low melting and boiling points. Liquids and gases at room temperature are usually molecular covalent • Do not conduct electricity or heat in any state • More soluble in non-aqueous solvents, unless they can H-bond to water or react with it • Examples: CO2, CH3CH2OH, I2 • Nonmetal + nonmetal

19. Simple Molecular Compounds

20. Valence Electrons are…? • The electrons responsible for the chemical properties of atoms, and are those in the outer energy level. • Valence electrons - The s and p electrons in the outer energy level • the highest occupied energy level • Core electrons – are those in the energy levels below.

21. Keeping Track of Electrons • Atoms in the same group have the same outer electronic structure and therefore the same number of valence electrons. • The number of valence electrons is easily determined. It is the group number for groups 1 and 2 • Group 1: H, Li, Na, K, etc. have 1 valence e- • Group 2: Be, Mg, Ca, etc. have 2 valence e-

22. More about Keeping Track • For elements in groups 13-17: • Subtract 10 from the group number • This is the number of valence electrons

24. Electron Dot diagrams are… • A way of showing & keeping track of valence electrons. • How to write them? • Write the symbol - it represents the nucleus and inner (core) electrons • Put one dot for each valence electron (8 maximum) • They don’t pair up until they have to (Hund’s rule) X

25. The Electron Dot diagram for Nitrogen • Nitrogen has 5 valence electrons to show. • First we write the symbol. N • Then add 1 electron at a time to each side. • Now they are forced to pair up. • We have now written the electron dot diagram for nitrogen.

26. The Octet Rule • The noble gases are unreactive in chemical reactions • In 1916, Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules • The Octet Rule: in forming compounds, atoms tend to achieve a noble gas configuration; 8 in the outer level is stable • Each noble gas (except He, which has 2) has 8 electrons in the outer level

27. Formation of Cations • Metals lose electrons to attain a noble gas configuration. • They make positive ions (cations) • If we look at the electron configuration, it makes sense to lose electrons: • Na 1s22s22p63s1 1 valence electron • Na1+1s22s22p6 This is a noble gas configuration with 8 electrons in the outer level.

29. Electron Dots For Cations • Metals will have few valence electrons (usually 3 or less); calcium has only 2 valence electrons Ca

30. Electron Dots For Cations • Metals will have few valence electrons • Metals will lose the valence electrons Ca

31. Electron Dots For Cations • Metals will have few valence electrons • Metals will lose the valence electrons • Forming positive ions Ca2+ This is named the “calcium ion”. No dots are now shown for the cation.

32. Electron Dots For Cations • Let’s do scandium, #21 • The electron configuration is: 1s22s22p63s23p64s23d1 • Thus, it can lose 2e- (making it 2+), or lose 3e- (making 3+) Sc = Sc2+ Sc = Sc3+ scandium (III) ion scandium (II) ion

33. Electron Configurations: Anions • Nonmetals gain electrons to attain noble gas configuration. • They make negative ions (anions) • S = 1s22s22p63s23p4 = 6 valence electrons • S2- = 1s22s22p63s23p6 = noble gas configuration. • Halide ions are ions from chlorine or other halogens that gain electrons

35. Electron Dots For Anions • Nonmetals will have many valence electrons (usually 5 or more) • They will gain electrons to fill outer shell. 3- P (This is called the “phosphide ion”, and should show dots)

36. Stable Electron Configurations • All atoms react to try and achieve a noble gas configuration. • Noble gases have 2 s and 6 p electrons. • 8 valence electrons = already stable! • This is the octet rule (8 in the outer level is particularly stable). Ar

37. Ionic Bonding • Anions and cations are held together by opposite charges (+ and -) • Ionic compounds are called salts. • Simplest ratio of elements in an ionic compound is called the formula unit. • The bond is formed through the transfer of electrons (lose and gain) • Electrons are transferred to achieve noble gas configuration.

39. Ionic Compounds • Also called salts • Made from: a cation with an anion (or literally from a metal combining with a nonmetal)

40. Ionic Bonding Na Cl The metal (sodium) tends to lose its one electron from the outer level. The nonmetal (chlorine) needs to gain one more to fill its outer level, and will accept the one electron that sodium is going to lose.

41. Ionic Bonding Na+ Cl 1- Note: Remember that NO DOTS are now shown for the cation!

42. Ionic Bond • Negative charges are attracted to positive charges. • Negative anions are attracted to positive cations. • The result is an ionic bond. • A three-dimensional crystal lattice of anions and cations is formed.

44. Preserve Electroneutrality • When ions combine, electroneutrality must be preserved. • In the formation of magnesium chloride, • 2 Cl- ions must balance a Mg2+ ion: • Mg2+ + 2 Cl- → MgCl2

45. Ionic Bonding Let’s do an example by combining calcium and phosphorus: • All the electrons must be accounted for, and each atom will have a noble gas configuration (which is stable). Ca P

46. Ionic Bonding Ca P

47. Ionic Bonding Ca2+ P

48. Ionic Bonding Ca2+ P Ca

49. Ionic Bonding Ca2+ P 3- Ca

50. Ionic Bonding Ca2+ P 3- Ca P