What is an ionic compound?

1. What is an ionic compound? • Composed of a metal and a nonmetal • Electrically non-conductive as a solid • Conductive as molten liquids or in solution

2. Ions • Cation: A positive ion • Mg2+, NH4+ • Anion: A negative ion • Cl-, SO42- • Ionic Bonding: Force of attraction between oppositely charged ions.

3. Formation of an ionic compound • Combines a metal and a nonmetal through the transfer of electrons • Example: A compound made from potassium and chlorine

4. Potassium has one valence electron and tends to lose it to become a cation with a charge of +1 K Cl Chlorine has seven valence electrons and tends to gain one to become an anion with a charge of -1

5. K Cl

6. Potassium, being an alkali metal, has an oxidation number of +1 - + K Cl Chlorine, being a halogen, has an oxidation number of -1

7. When these two ions combine to form an ionic compound, they will combine in such a way that the overall charge of the compound is zero. In this case, the smallest ratio that adds up to zero is one potassium ion and one chloride ion. - + K Cl

8. When these two ions combine to form an ionic compound, they will combine in such a way that the overall charge of the compound is zero. In this case, the smallest ratio that adds up to zero is one potassium ion and one chloride ion. KCl Note that the metal ion is ALWAYS written first, and that when there is only one ion of an element, it is not necessary to place a subscript “1” next to that element.

9. Calcium has two valence electron and tends to lose both of them to become a cation with a charge of +2 Ca Cl Chlorine has seven valence electrons and tends to gain one to become an anion with a charge of -1

10. Since calcium has two electrons to give, but chlorine can only accept one, those electrons must go to two separate chlorine atoms. Cl Ca Cl

11. - Calcium, being an alkaline earth metal, has an oxidation number of +2 Cl 2+ Ca - Cl Chlorine, being a halogen, has an oxidation number of -1

12. When these two ions combine to form an ionic compound, they will combine in such a way that the overall charge of the compound is zero. In this case, the smallest ratio that adds up to zero is one calcium ion and two chloride ions. - 2+ Ca Cl

13. When these two ions combine to form an ionic compound, they will combine in such a way that the overall charge of the compound is zero. In this case, the smallest ratio that adds up to zero is one calcium ion and two chloride ions. CaCl2 Note that the metal ion is ALWAYS written first, and that when there is only one ion of an element, it is not necessary to place a subscript “1” next to that element.

14. Hopefully by this point you can see that it is not necessary to draw out the dot structures before writing the formula for an ionic compound. The dot structures provide us with the oxidation number, and the oxidation numbers determine the ratio of ions in the compound. But since we can find the oxidation numbers of the representative elements from the periodic table, we can skip the step of drawing out the dot structures.

15. Beryllium and oxygen • Beryllium is an alkaline earth metal and has an oxidation number of +2 (it loses two electrons) • Oxygen has an oxidation number of -2 (it gains two electrons)

16. Be2+ O2- The smallest ratio that adds up to zero is one beryllium ion to one oxide ion: BeO

17. Lithium and nitrogen • Lithium is an alkali metal and has an oxidation number of +1 (it loses one electron) • Nitrogen has an oxidation number of -3 (it gains three electrons)

18. Li+ N3- The smallest ratio that adds up to zero is three lithium ions to one nitride ion: Li3N

19. Aluminum and sulfur • Aluminum has an oxidation number of +3 (it loses three electrons) • Sulfur has an oxidation number of -2 (it gains two electrons)

20. Al3+ S2- The smallest ratio that adds up to zero is two aluminum ions to two sulfide ions: Al3S2

21. Writing Ionic Compound Formulas Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES! ( ) 2. Check to see if charges are balanced. Ba2+ NO3- 2 Not balanced! 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

22. Writing Ionic Compound Formulas Example: Ammonium sulfate 1. Write the formulas for the cation and anion, including CHARGES! ( ) NH4+ SO42- 2. Check to see if charges are balanced. 2 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!

23. Writing Ionic Compound Formulas Example: Iron(III) chloride 1. Write the formulas for the cation and anion, including CHARGES! Fe3+ Cl- 2. Check to see if charges are balanced. 3 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!

24. Writing Ionic Compound Formulas Example: Aluminum sulfide 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. Al3+ S2- 2 3 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!

25. Writing Ionic Compound Formulas Example: Magnesium carbonate 1. Write the formulas for the cation and anion, including CHARGES! Mg2+ CO32- 2. Check to see if charges are balanced. They are balanced!

26. Writing Ionic Compound Formulas Example: Zinc hydroxide 1. Write the formulas for the cation and anion, including CHARGES! ( ) 2. Check to see if charges are balanced. Zn2+ OH- 2 3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced!

27. Writing Ionic Compound Formulas Example: Aluminum phosphate 1. Write the formulas for the cation and anion, including CHARGES! 2. Check to see if charges are balanced. Al3+ PO43- They ARE balanced!

28. Naming Ionic Compounds • 1. Cation first, then anion • 2. Monatomic cation = name of the element • Ca2+ = calciumion • 3. Monatomic anion = root + -ide • Cl- = chloride • CaCl2= calcium chloride

29. Naming Ionic Compounds(continued) Metals with multiple oxidation states • - some metal forms more than one cation • - use Roman numeralin name • PbCl2 • Pb2+is cation • PbCl2 = lead(II) chloride

31. Naming Ionic Compounds:Examples Na2SO4 sodium sulfate Fe(NO3)2 iron (II) nitrate aluminum chloride AlCl3 Mg3N2 magnesium nitride (NH4)3PO4 ammonium phosphate

32. Formulas Formulas for ionic compounds are ALWAYS empirical (lowest whole number ratio). Examples: NaCl MgCl2 Al2(SO4)3 K2CO3

33. CovalentBonding • Metals can only give electrons, and produce positively charged ions. • Nonmetals can gain or lose electrons to complete their octets. • The chemical bond formed from electron sharing between atoms is called the covalent bond.

34. Covalent Bonding in H2 • Atoms of hydrogen ( 1H : 1) have one valence electron in their first electron shell. • Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. In H2 molecule, which hydrogen atom will donate one electron to which one? • But both hydrogen atoms can attain helium like electron arrangement (2He: 2) by sharing their valence electrons. • Because the hydrogen molecule is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond.

36. When two hydrogen atoms are near each other, the two electrons and the two protons repel each other. But there is an attraction between the proton of one hydrogen atom and the electron of the other. This attraction force brings the atoms closer together. • When the two nuclei are kept a certain distance away from each other, the attractive forces become grater than the repulsive forces, and a stable chemical bond forms.

37. In this situation, the distance between two nuclei is called bond length or bond distance. • The hydrogen atoms in H2 molecule is in a less energetic and therefore more stable state than hydrogen atoms alone. • During the formation of the bond, some amount of energy is released. This energy is called the bond energy. As the magnitude of bond energy increases the strength of the bond increases.

38. Lewis Structures

39. The electrons shared between two identical nuclei such as in H2, Cl2, F2, O2 or N2 are located at equal distances between two nuclei. Neither of the atoms gains or loses electrons, or the positive charge center and the negative charge center in the molecule coincide. • Molecules have no positive or negative poles; is called nonpolar covalent bonding. H2, Cl2, F2, O2 , P4, S8 ... all have nonpolar covalent bonds.

40. Covalent Bonding in HCI • To complete their valence shells, hydrogen and chlorine both need one electron. • Since both have high tendencies to accept electrons, we do not expect one of them to donate one electron to the other. • But they may share a pair of electrons

41. Chlorine has a much greater attraction for the shared electron pair than hydrogen atom. • Such bonds are called polar covalent bonds. • The unequal sharing of electrons in a bond leads to what is referred to as a dipole.

42. The presence of polar covalent bonds does not necessarily mean the molecule will be polar. Some molecules contain two or more dipoles that cancel to give a nonpolar molecule. • For example, in CO2 the two oxygen are attached to carbon by polar covalent bonds. The oxygen atoms having greater tendency to attract electrons pull more strongly the shared electrons that form the double covalent bonds. Thus, two polar double covalent bonds are present in CO2, but because of the symmetry of the dipoles the molecule itself is nonpolar.

43. When determining whether a molecule is polar or nonpolar, it is important to consider the geometry of the molecule. • The tendency of some of the atoms to attract shared electrons is as follows: • F>O>CI>Br>N>S>I>C>H • Whenever we have diatomic molecules consisting of two different elements, the molecule is generally polar. • For example, the polarity order in the molecules HF, HCI, HBr, and HI is HF > HCI > HBr > HI.

44. Covalent Bonding in H2O • Oxygen with the electron configuration of 8O: 2-6 needs two electrons to have eight electrons in its outer most shell. • So two hydrogen atoms and one oxygen atom must combine to form water. • Each bond between O and H atoms is a polar covalent bond. Water is a bent molecule. • In water, the polar covalent bonds lead to dipoles in which the centers of positive and negative charge do not coincide. This makes water a polar molecule.

45. Naming Nonmetal - Nonmetal Compounds • When two nonmetallic elements combine, no ions are formed. However it is generally possible to consider one of the two elements to be the more positive. • Both in writing the name and the formula of a binary compound, the more positive element appears first, followed by the more nonmetallic one.

46. In naming these compounds use the following pattern. • Number of the first nonmetal (in Latin), name of the first nonmetal, number of the second nonmetal (in Latin), ionic name of the second nonmetal.

48. Naming Binary Covalent Compounds:Examples N2S4 dinitrogen tetrasulfide NI3 nitrogen triiodide XeF6 xenon hexafluoride CCl4 carbon tetrachloride P2O5 diphosphorus pentoxide SO3 sulfur trioxide

49. Example

50. SiF4 silicon tetrafluoride Naming Compounds: Practice two nonmetals  covalent  use prefixes Na2CO3 sodium carbonate metal present  ionic  no prefixes Na  group I  no Roman numeral N2O dinitrogen monoxide two nonmetals  covalent  use prefixes K2O potassium oxide metal present  ionic  no prefixes K  group I  no Roman numeral Cu3PO4 copper (I) phosphate metal present  ionic  no prefixes Cu  not group I, II, etc.  add Roman numeral (PO4 is 3-, each Cu must be 1+) CoI3 cobalt (III) iodide metal present  ionic  no prefixes Co  not group I, II, etc.  add Roman numeral (I is 1-, total is 3-, Co must be 3+) PI3 phosphorus trioxide two nonmetals  covalent  use prefixes NH4Cl potassium oxide NH4 polyatomic ion present  ionic  no prefixes