2. Ionic Bonds Chapter 4

3. Ionic Compounds What are Chemical Bonds Force that holds 2 atoms together Attraction between + nucleus and – electron Attraction between + ion and – ion Valence electrons make bonds

4. Ionic compounds What are chemical bonds? (cont) Elements react to form a stable octet (noble gas configuration). The + and – charges act like opposite poles of a magnet. Opposites attract strongly Likes repel strongly Strength diminishes with distance

5. Typical Ions Oxidation number = oxidation state = number of electrons transferred from an atom to make a compound Na+ oxidation number = +1 O 2- oxidation number = 2- Used to determine compound formulas

6. Ionic Compounds How are positive ions formed? Atom loses one or more VALENCE electrons Called a CATION Ion becomes more stable by losing electrons (octet rule) Not a change in atom, Just an ion Loses all electrons in outer shell Reactivity depends on ease of losing electrons Transition metals usually form 2+ or 3+ ions shown with a (II) or (III)

7. Ionic Bonds How do negative ions form? Atoms gain negative electrons Nonmetals have a great attraction for electrons Adding electron fills up the shell = stable Called an ANION Naming: change name to end in –ide Gaining enough electrons to fill outer shell (octet rule) 7A gains 1 6A gains 2 5A gains 3

8. Ionic Compounds What are oxidation states Oxidation state is the charge of the ‘normal’ ion formed Group 1 loses 1 valence electron (+1) Group 2 loses 2 valence electrons (+2) Group 13 loses 3 valence electrons (+3) Group 14 does not generally make ionic compounds Group 17 gains 1 valence electron (-1) Group 16 gains 2 valence electrons (-2) Group 15 gains 3 valence electrons (-3)

9. Ionic Compounds Ionic compounds are Metal+ and Nonmetal- Metals make Cations Groups 1A (1) – 3A (13) and all Group D elements Form + ions Nonmetals make Anions Group 5A (15) - 7A (17) Nobel Gasses (Group 8A/18) do not form compounds. Why?

10. Properties of Ionic Compounds Crystal shape Alternate positive and negative ions in patterns Crystals are all the same shape for each compound High melting points Table Salt Melting point is 801 o C Conduct electricity Dissolve in water Ions become more loosely associated Pass electrical charges along

11. Counting Atoms in a Compound CH 4 Subscript Element Symbol

12. Formula Names Remember: total number of electrons lost by cations must equal total number of electrons gained by anions! Metal name is stated first. Number of cations in ratio is subcripted Nonmetal name is stated second Suffix –ide is used Number of anions in ration is subscripted

13. Formula Names Examples Sodium (Na + ) and Chlorine (Cl - ) NaCl = ratio 1:1 Sodium Chloride Calcium (Ca2+) and Fluorine (Fl-) CaFl 2 = ratio 1:2 Calcium Fluoride Aluminum (3+) and Sulfur (2-) Find lowest common dominator (6) 2(Al 3+) + 3 (S 2+) = both transfer 6 electrons Al 2 S 3 Aluminum Sulfide

14. Determining Formulas from Oxidation States 1. Find oxidation numbers for elements Ca, O 2. Put elements together with oxidation numbers as superscripts Ca +2 O -2 3. Criss-Cross the oxidation numbers removing the signs Ca +2 O -2  Ca 2 O 2 4. Take lowest common denominator CaO

15. Formulas with Polyatomic Ions Polyatomic ions stay together as a single group Example Calcium and Phosphate Calcium (2+) Phosphate PO 4 (3-) 3 (Ca) + 2 ( PO 4 ) Ca 3 (PO 4 ) 2 Calcium Phosphate

16. Covalent Bonds Chapter 9

17. Electron Sharing Covalent bonds are nonmetal/nonmetal Covalent (Co = together, Valence Electrons) Covalent Bonds – Formed when two atoms SHARE a PAIR of electrons Both atoms attract the shared electron(s) at the same time

18. CoValent Bonds Every atom MUST have a full valence electron shell: Carbon = 8 Hydrogen = 2 Shared electrons count for BOTH atoms

19. Covalent Bonds Some compounds share more than one pair of electrons

20. Molecular Compounds Molecular Compounds have covalent bonds Much lower boiling points than ionic Poor conductors of electricity Pure water does not conduct electricity Water with sugar does not conduct electricity Water with SALT DOES conduct electricity

21. Naming Covalent Molecules Chapter 9.2

22. Basic Rules First element in formula is always named first, using the entire element name Second element in formula is named second, using the root of the element name and adding the suffix –ide Prefixes are added to each name to indicate number of atoms of each type.

23. Prefixes 1 = mono- 2 = di- 3 = tri- 4 = tetra- 5 = penta- 6 = hexa- 7 = hepta- 8 = octa- 9 = nona- 10 = deca

24. Example H 2 O (water) Hydrogen – Oxygen Oxide DIhydrogen oxide NH 3 (ammonia) Nitrogen – Hydrogen 1 nitrogen, 3 hydrogens Nitrogen, 3 HydrIDE Nitrogen TRI hydride

25. Metallic Bonds and Metals Chapter 4

26. Metallic Bonds In solid state,m Metals do not bond ionically but for form lattices Similar to ionic crystal lattices Metals have at least one valence electron but: do not share these electrons Do not lose electrons

27. Metallic Bonds In the solid crystal lattice of metals Electrons are crowded Outer energy levels overlap Like "an array of positive ions in a sea of electrons". Electrons are delocalized Can move from one atom to another easily

28. Metallic Bonds Metallic Bond Attraction of a metallic cation for a delocalized electron

29. Properties of Metals Melting points vary greatly Metals are malleable – can be hammered into sheets Ductile – can be drawn into wires Generally durable – but with some variation Good conductors – due to delocalized electrons Mobile electrons consist of: ‘d’-level electrons 2 outer ‘s’ electrons

30. Metals Malleability and ductility Metals are described as malleable (can be beaten into sheets) and ductile (can be pulled out into wires). This is because of the ability of the atoms to roll over each other into new positions without breaking the metallic bond. If a small stress is put onto the metal, the layers of atoms will start to roll over each other. If the stress is released again, they will fall back to their original positions. Under these circumstances, the metal is said to be elastic.

31. Metals If a larger stress is put on, the atoms roll over each other into a new position, and the metal is permanently changed.

32. Metal Alloys Easy to introduce other metals into the metallic crystal Mixture called an alloy Properties of alloys differ some from either base metal Steel is iron mixed with another element Alloys form when: Elements involved are similar in size Or one is much smaller than the other

33. Metal Alloys 2 types of alloys Substitutional Interstitial Sustitutional alloys Original metallic atoms are replaced by other metal atoms of similar size Sterling silver (copper and silver), brass, pewter, 10-carat gold

34. Metal Alloys Interstitial alloys Small holes in metal crystal are filled with other, smaller atoms Like pouring sand in a bucket of gravel Example: carbon steel (iron crystal is filled with carbon) Physical properties of steel are changed. Iron is relatively soft and brittle. Adding Carbon makes the solid harder, stronger less ductile, giving it different uses.

35. Making Ionic Compounds Write the electron configuration of Magnesium. Based on this configuration, will Mg gain or lose electrons? Write the electron configuration for oxygen Based on this configuration, will oxygen gain or lose electrons? When we burn magnesium, What compound results? Magnesium + oxygen Which atom donated electrons? How Many? What is the formula for the new compound? Do you think this new compound will conduct electricity?