1. “Structure of Matter” Covalent Bonds Ch. 6

2. Matter  Matter is anything that has mass and occupies space. Matter is made of atoms which are the smallest particles that have the properties of an element.

3. Matter  Pure substances are any matter that has a fixed composition and definite properties. Cannot be broken down by physical changes. There are about 100 million pure substances that have been identified Out of these pure substances, only 118 of them are elements, the rest are compounds

4. Matter

5.  Elements are substances that cannot be broken down into simpler substances.  Compounds are substances made of atoms of more than one element bound together. Every compound is made up of a chemical formula

7. Chemical Formulas  A chemical formula tells us: the type of atoms present the number of atoms present the type of compound

8. Chemical Formulas  Example: table salt: Sodium Chloride  Chemical formula: NaCl  Count the atoms present: 1 Na atom 1 Cl atom

9. Chemical Formulas  Sometimes there are subscripts present. A subscript is a small number that is in a chemical formula. Example - water: H 2 O  2 H atoms  1 O atom Subscript

10. Chemical Formulas  Sometimes there are parentheses with a subscript. The subscript only applies to the atoms within the parentheses.  Example - calcium hydroxide (kidney stones): Ca(OH) 2. 1 Ca atom 2 O atoms 2 H atoms

11. Chemical Formulas  Sometimes there are subscripts in the parentheses. Multiply the subscript outside the parentheses by the subscript of each element within the parentheses. If no subscript is present assume that it is 1.  Example - calcium nitrate: Ca(NO 3 ) 2 1 Ca atom 2 N atoms 6 O atoms (3 oxygens x 2 = 6)

12. Structure of Matter  Nuances in molecular structure can affect its properties. Chemical formulas can be visually represented using chemical structures which can show bond length, bond angles and atomic sizes.

13. Structure of Matter  The structure of a compound affects its properties. Example: strong bonds = high melting points.

14. Types of Molecular Structures  Network Structures: Structure: strong, rigid structure Bond Strength: strong Boiling and Melting Points: high

15. Types of Molecular Structures  Ionic network structures: Structure: regularly shaped crystals Bond Strength: strong Boiling and Melting Points: high

16. Types of Molecular Structures  Molecular structures: Structure: molecules weakly bonded together. Bond Strength: weak Boiling and Melting Points: low

17. Types of Molecular Structures  Molecular structures typically experience two types of attractive force: The attraction between molecules is called intermolecular force. It is rarely as strong as intramolecular force which is inside the molecule.

18. Atomic Bonds  Atoms form atomic bonds to become more stable. Atoms become more stable by filling their valence shell or at least meeting the octet rule by getting 8 valence electrons.

19. Atomic Bonds  There are three main types of chemical bonds used by atoms to fill their valence shell: Covalent Metallic Ionic “Bond, Chemical Bond”

20. Covalent Bonds  In covalent bonds, nonmetal atoms meet the octet rule by sharing one or more pairs of electrons. The shared electron pair is called a bonding pair and represented by a line on a Lewis structure.

21. Covalent Bonds Chlorine forms a covalent bond with itself.

22. Covalent Bonds Each chlorine atom wants to gain one electron to achieve an octet.

23. Covalent Bonds Each chlorine atom wants to gain one electron to achieve an octet.

24. Covalent Bonds Each chlorine atom wants to gain one electron to achieve an octet.

25. Covalent Bonds The octet is achieved by each atom sharing the electron pair in the middle.

26. Covalent Bonds This is the bonding pair.

27. Covalent Bonds It is a single bonding pair so it is called a single bond.

28. Covalent Bonds Single bonds are abbreviated with a dash

29. Covalent Bonds This is now a chlorine molecule.

30. Covalent Bonds Oxygen is also a diatomic molecule (a molecule with 2 of the same element bonded together).

31. Covalent Bonds How will oxygen bond?

32. Covalent Bonds How will oxygen bond?

33. Covalent Bonds How will oxygen bond?

34. Covalent Bonds Since each oxygen has 6 valence, they would each need to gain 2 more electrons to be stable.

35. Covalent Bonds Both pairs of electrons are shared.

36. Covalent Bonds 6 valence electrons + 2 shared electrons = full octet

37. Covalent Bonds Two bonding pairs, making a double bond.

38. Covalent Bonds The double bond can be shown as two dashes.

39. Covalent Bonds This is now an oxygen molecule.

40. Covalent Bonds  Elements can share up to three pairs (6 electrons). Single Bond (2e) Double Bond (4e) Triple Bond (6e)

41. Covalent Bonds  Equal sharing of electrons creates nonpolar covalent bonds. Ex. Ethane, C 2 H 6  Unequal sharing of electrons is called polar covalent bonds and can lead to molecules having a positively and negatively charged side. Ex. Water, H 2 0

42. Covalent Bonds  The slight charges on a polar molecule can cause a loose atomic bond called polar or hydrogen bond.

43. Covalent Bonds Nomenclature  Naming binary covalent compounds: Two nonmetals Name each element End the last element in –ide Add prefixes to show more than 1 atom or 1 atom on the second element. # of AtomsPrefix 1mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- 9nona- 10deca-

44. Covalent Bonds Nomenclature  CO carbon monoxide  CO 2 carbon dioxide  PCl 3 phosphorus trichloride  CCl 4 carbon tetrachloride N2ON2O dinitrogen monoxide # of AtomsPrefix 1mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- 9nona- 10deca-

45. Covalent Bonds Nomenclature  dihydrogen monoxide H 2 O  nitrogen dioxide NO 2  carbon tetrahydride CH 4 # of AtomsPrefix 1mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- 9nona- 10deca-

46. Metallic Bonds  Metallic bonds are metal to metal bonds formed by the attraction between positively charged metal ions and the electrons around them. Atoms are packed tightly together to the point where outermost energy levels overlap.  This allows electrons to move freely from one atom to the next making them great conductors of electricity.

47. Ionic Bonds  An ion is a charged atom or molecule. It is charged because the number of electrons do not equal the number of protons in the atom or molecule. Atoms with ADDED electrons are negative (anions). Atoms with LESS electrons are positive (cations).

48. Ionic Bonds  The normal charge of an ion can be quickly determined using the oxidation number of an element. The oxidation number of an atom is the charge that atom would have if the compound was composed of ions.

49. Ionic Bonds  To find oxidation number: All elements with a valence number less than four will lose all of their electrons to achieve a full valence or the octet rule.  Example: Beryllium has 2 e- Loses the 2 e- Gains a charge of +2

50. Ionic Bonds  To find oxidation number: All elements with a valence number greater than four will gain electrons until they have achieved a full valence or the octet rule.  Example: Nitrogen has 5 e- Gains 3 e- Gains a charge of -3

51. Ionic Bonds  Examples: Oxygen – Group 16  -2 Calcium – Group 2  +2 Aluminum – Group 13  +3 Chlorine – Group 17  -1

52. Ionic Bonds  Ionic bonds are bonds formed by the attraction between oppositely charged ions. Electrons are transferred from one element to another.

53. Potassium (metal – cation) needs to lose 1 valence electron to drop down to a full valence shell. Fluorine (nonmetal – anion) only needs 1 electron to complete its valence shell.

58. Once the transfer is complete, the potassium will have a +1 charge (K + ) and the fluorine will have a -1 charge (F - ).

60. The ionic bond is formed because of the electrostatic forces between the positive and negatively charged ions and the new overall charge is 0.

61. Magnesium (metal – cation) needs to lose 2 valence electron to drop down to a full valence shell. Iodine (nonmetal – anion) only needs 1 electron to complete its valence shell, but Mg can give to two different atoms.

65. Once the transfer is complete, the magnesium will have a +2 charge (Mg 2+) and each iodine will have a -1 charge (I - ).

68. Ionic Bonds  Ionic bonds form strong network structures with high melting and boiling points.  When melted or dissolved in water ionic compounds conduct electricity because ions are free to move.

69. Ionic Bonds Nomenclature. Name the cation (metal). If the first ion is a transition element other than zinc, cadmium, or silver, you must use a Roman Numeral with the name – we’ll discuss this later. Name the anion (nonmetal) by changing the suffix to -ide.

70. Examples NaCl Name the metal ion Sodium Name the nonmetal ion, changing the suffix to –ide. Chloride CaO Calcium Oxide Al 2 S 3 AluminumSulfide MgI 2 MagnesiumIodide BaNa 2 You should recognize a problem with this one This is two metals – not a binary ionic compound Banana The name of this is Banana (JOKE – haha) What is the name of this compound: HIJKLMNO? WATER – “H” to “O” You have to admit – that was funny!

71. Ionic Bonds Nomenclature.  To go backwards from the name to the formula you can use the “Swap and Drop” method.: 1. Write the symbols for each ion. 2. Determine the oxidation number of each ion. 3. Swap and Drop 4. Reduce (if necessary). 5. Rewrite

72. Ionic Bonds Nomenclature.  To go backwards from the name to the formula you can use the “Swap and Drop” method.: 1. Write the symbols for each ion. 2. Determine the oxidation number of each ion. 3. Swap and Drop 4. Reduce (if necessary). 5. Rewrite

73. Polyatomic Ions  A polyatomic ion is a group of covalently bonded atoms that have lost or gained an electron. (Example: Nitrate NO 3 - and Ammonium NH 4 + ). Oppositely charge polyatomic ions can form compounds. (Example: Ammonium nitrate NH 4 NO 3 ).

74. Polyatomic Ions  Naming of these compounds follows the same rules as binary ionic compounds. The most important part is recognizing there is a polyatomic ion present.

75. Polyatomic bonds  To go from the formula to the name: 1. Name the cation. 2. Name the anion.

76. Polyatomic bonds  To go from the formula to the name: 1. Name the cation. 2. Name the anion.

77. Polyatomic bonds  To go from name to formula: 1. Write the symbols for each ion. 2. Determine the oxidation number of each ion. 3. Swap and Drop 4. Reduce (if necessary). 5. If a subscript greater than one is added to the polyatomic ion use parentheses. 6. Rewrite

78. Polyatomic bonds  To go from name to formula: 1. Write the symbols for each ion. 2. Determine the oxidation number of each ion. 3. Swap and Drop 4. Reduce (if necessary). 5. If a subscript greater than one is added to the polyatomic ion use parentheses. 6. Rewrite

79.  Transition metals are cations that have variable charges that makes them hard to name. We use Roman numerals to indicate the charge of a transition metal.  Example: copper (II) oxide – charge of copper is +2 titanium ( IV) sulfide – charge of titanium is +4 Transition Metal Ionic Compounds

80.  To go from formula to name you need to determine the Roman numeral for your transition metal: 1. If there are subscripts present use the reverse “Swap and Drop.” 2. Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY. Transition Metal Ionic Compounds

81.  To go from formula to name you need to determine the Roman numeral for your transition metal. 1. If there are no subscripts, simply give the transition metal the equal and opposite charge to the nonmetal. 2. Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY. Transition Metal Ionic Compounds