Unit 7. Bonds and Naming

Transcription

1 Unit 7 Bonds and Naming

2 I. Ionic Bonds Positive ion is attracted to a negative ion; usually a metal & a nonmetal Ionic compound: a substance that has ionic bonds Cation: positive ion Anion: negative ion Ionic compounds are neutral so cation charges must equal anion charges

3 Octet Rule Atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons

4 A. Lewis dot diagrams Diagrams that show the bonding between atoms of a molecule, and the lone pairs of valence electrons that may exist in the molecule Magnesium [Ne] 3s 2 Mg

5

6 Lewis dot diagrams

7 Review 1. What is an anion? a cation? 2. How do ionic bonds form? 3. What are valence electrons? 4. What is the octet rule? 5. What is a Lewis dot diagram? 6. How would you draw a Lewis dot diagram for Carbon?

8 B. Types of Ions Monatomic ions: one atom ions Polyatomic ions: many atom ions Monatomic cations (+): name after element; for transition elements include a roman numeral for charge Na + Mg 2+ Fe 2+ Fe 3+ sodium ion magnesium ion iron (II) ion iron (III) ion

9 Table 4.2

10 B. Types of Ions Monatomic anions: name after element but change ending to ide Cl - O 2- N 3- chloride oxide nitride

11 Types of Ions Polyatomic ions: atoms that are covalently bonded, but as a unit they form an ion. Use special names

12 Table 4.1

13 Table 4.4

14 C. Binary Ionic Compounds Contain ions of only two elements Name: cation name & anion name CaF 2 calcium flouride MgO magnesium oxide Name does not tell ratio of ions Empirical formula: the empirical formula of a chemical compound is the simplest whole number ratio of atoms of each element present in a compound Al 2 O 3 aluminum oxide two Al 3+ three O 2- Li 2 O lithium oxide Two Li + one O 2- no subscript means one Ionic compounds are neutral charges must equal

15 D. Crisscross method Calcium fluoride Ca 2+ F - \ / CaF 2 Magnesium Oxide Mg 2+ O 2- \ / MgO

16 Crisscross Method Calcium nitrate Ca 2+ NO 3 - \ / Ca(NO 3 ) 2 Aluminum nitrate Al 3+ NO 3 - \ / Al(NO 3 ) 3

17 Review 1. What is a monoatomic cation? 2. What is a polyatomic anion? 3. What is a binary ionic compound? 4. What is an empirical formula? 5. What does an empirical formula of H 2 SO 4 mean? 6. Explain the crisscross method. 7. What would the empirical formula be if ammonium (NH 4+ ) and phosphate (PO 4 3- ) bond together?

18 II. Covalent Bonding Electrons are shared between atoms A. Molecules and their Formulas Molecules: covalently bonded atoms Molecular formulas: tells how many atoms of each type are in a molecule oxygen O 2 2 O atoms sucrose C 12 H 22 O C atoms 22 H atoms 11 O atoms

19 More formulas Glucose Molecular formula C 6 H 12 O 6 Empirical formula CH 2 O Lactic Acid Molecular formula C 3 H 6 O 3 Empirical formula CH 2 O

20 Structural Formula Shows where atoms are bonded Lewis structures:

21 More Lewis Structures

22 Describing covalent bonds Use octet rule

23 C. Multiple Bonds Single bonds share one pair of electrons between atoms Double bonds share two pairs of electrons between atoms Triple bonds share three pairs of electrons between atoms

24 Double bonds

25 Triple bonds

26 Dots or Dashes? Dash represents a shared pair of electrons

27 More Lewis Structures

28 Exceptions to Octet Rule 1. Less than an octet

29 Exceptions to Octet Rule 2. More than Octet:

30 Exceptions to Octet Rule 3. Odd number of electrons:.... N.. + Ȯ.. :.... N=Ȯ :

31 Review 1. What is a molecule? 2. What is a covalent bond? 3. What does the molecular formula C 12 H 22 O 11 tell you? 4. What are the molecular and empirical formulas of glucose (6 sugars, 12 hydrogens, 6 oxygens)? 5. What do dashes mean in Lewis structures?

32 E. Properties of Covalent Bonds Polar covalent bond: results from an unequal sharing of electrons; depends on electronegativity (elements with bigger electronegativites hog electrons)

33 The electronegativity of Oxygen is 3.5 and of Hydrogen is 2.1 Figure 12.5: Charge distribution in the water molecule.

34 Electronegativities

35 E. Properties of Covalent Bonds Nonpolar covalent bonds: results from equal sharing of electrons; occurs between same element or elements with the same electronegativity Examples: F 2, O 2, N 2

36 E. Properties of Covalent Bonds Electronegativity Difference Bond Type 0.4 nonpolar covalent Between 0.5 and 1.9 polar covalent 2.0 ionic

37 Review 1. What is a polar covalent bond? 2. How much of a difference in electronegativity between elements will cause a polar covalent bond? 3. Which atoms hog the electrons in polar covalent bonds? 4. How much of a difference in electronegativity between elements will cause an ionic bond? 5. How much of a difference in electronegativity between elements will cause a nonpolar covalent bond?

38 Table 12.1

39 III. Naming Ionic Compounds Type I the metal present only forms one kind of cation (H, Li, etc.) 1. cation named first, anion second 2. a simple cation takes its name from the name of the element (Na is just sodium) 3. A simple anion is named by taking the first part of the element and adding ide. (chlorine becomes chloride)

40 Naming Ionic Compounds Type II Ionic Compounds the metal present can have different charges Cation first, anion second The name of the cation is followed by a roman numeral indicating its charge (Iron II is 2 + or Iron III is 3 + ) The name of the anion is the same as with Type I

41 III. Extra Names Involving Water Hydrates: compounds that absorb water into their solid structure Anhydrous: no water Use a dot to separate water molecules CuSO 4 5H 2 O 1. name like ionic compound 2. Use prefix and hydrate Copper (II) sulfate pentahydrate

42

43 C. Naming Molecular Compounds 1. The first element in the formula is named first, and the full element name is used 2. The second element is named as though it were an anion (-ide suffix) 3. Use prefixes (don t use mono on first element) to denote the numbers of atoms present 4. common names exist O 2 H 2 O NH 3 oxygen water ammonia

44 Diatomic molecules

45 D. Naming Acids A molecular substance that dissolves in water to produce H + ions. Named after anion HCl H + + Cl - End in ide change to Hydro ic acid End in ate change to ic acid End in ite change to ous acid

46

47 Review 1. Give the names for these Type I compounds: 1. CsF, AlCl 3, MgI 2 2. Name the following Type II compounds: 1. CuCl, Fe 2 O 3, HgO 3. Name the following molecules: 1. BF 3, NO, N 2 O 5