Chemical Bonding: Lewis Dot Diagrams, Ionic and Covalent Bonds

1. Unit 3: Chemical Formulas and Bonding

2. F F Electron Dot diagrams • Electron dot diagrams show the valence electrons around an atom. In most molecules and compounds a complete octet is achieved for each atom: • Most monatomic ions have an electron configuration of noble gases: Al N 7 valence e-s + e- 8 valence e-s

3. Drawing Lewis Dot Structures • To visualize valence e-, we will use Lewis Dot Structures. • Step 1: The element symbol represents the nucleus and all e- except valence. • Step 2: From the periodic table, determine the number of valence e-. • Step 3: Each “side” of symbol represents an orbital. Draw two dots on one side, then one for each of the remaining three sides. Additional electrons should then be paired.

4. Lewis Dot Structures • Ex: carbon step 1: C step 2: 4 valence e-s step 3: C

5. Lewis Dot Structures • Ex: bromine step 1: Br step 2: 7 valence e-s step 3: Br

6. Chemical BondingWhat holds things together?

7. Conductivity - high Conductivity - low Let’s examine the melting point of compounds across two periods. What is the trend? Nonconductive high low

8. Bonding How can we explain the melting point behavior across a period? Bonding between atoms changes across a period… • Bonding involves the valence electrons or outermost shell (or highest shell) electrons • Atoms form bonds to become more stable –electrons are gained, lost or shared to achieve stability. • The properties of a compound are different from the properties of the atoms that make up the compound. Ex: NaCl

9. Types of Bonds 1. Ionic bond Transfer of e- from a metal to a nonmetal and the resulting electrostatic force that holds them together forms an ionic compound. EX: Na+ + Cl- NaCl (neutral)

10. Ionic Bonding Ionic bonds involve the formation of positive and negative ions that then attract each other. Metals form positive ions by losing electrons Nonmetals form negative ions by gaining electrons Next Slide

11. Na Cl Ionic Bonding Example 1 Sodium has 1 valence electron which it needs to lose. Chlorine has 7 valence electrons and needs to pick up 1 electron. Next Slide

12. Na Cl Ionic Bonding Example 1 The sodium loses its electron to the chlorine. +1 -1 This makes the sodium +1 and the chlorine -1 They attract each other forming the compound NaCl Next Slide

13. Mg O Ionic Bonding Example 2 Magnesium has 2 valence electrons which it needs to lose. Oxygen has 6 valence electrons, It needs to pick up 2 electrons. Next Slide

14. Mg O Ionic Bonding Example 2 Magnesium loses both of its outer electrons to the oxygen. Next Slide

15. -2 +2 Mg O Ionic Bonding Example 2 This gives the magnesium a +2 charge and the oxygen a -2 charge They join together to form the compoundMgO. Next Slide

17. Exchange of Electrons

18. Ionic Bonding When atoms bond, the properties of the new compound are DIFFERENT from the properties of the elements that made them up. Ionic compounds have several characteristics in common due to the presence of the ionic bond. These characteristics include: Crystalline structure (the formula gives the ratio between the ions making up the substance) High melting points, making them solids at room temperature Usually water soluble (can dissolve in water) Electrolytes when in solution (conduct electricity)

19. Cl- Na+ Na + Cl- Cl- Na + Na + Cl- Ionic Bonding Sodium chloride (NaCl) is held together by an ionic bond. The properties of sodium chloride are: Sodium chloride forms a cube shaped crystalline solid. • Melting point = 801˚C • Boiling point = 1413˚C • Highly soluble in water • Strong electrolyte

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21. Types of Bonds 2. Covalent bond Formed from the sharing of e- pairs between two or more nonmetals resulting in a molecule. EX: H2 + O  H2O

22. Covalent Bonding Definition - bond formed due to the sharing of electrons between nonmetals. The high attraction for electrons of nonmetals results in the nonmetals attempting to remove electrons from each other. Since neither nonmetal is able to give up electrons they are forced to share the electrons.

23. Br F Covalent Bonding Example 1 Bromine and Fluorine both have 7 valence electrons and very high attraction for electrons.

24. Br F Covalent Bonding Example 1 Since neither fluorine or bromine are able to lose electrons they get drawn together until their outer orbits overlap and one electron from each atom goes back and forth between the two atoms. This creates the compoundBrF.

25. H O H Covalent Bonding Example 2 Hydrogen and Oxygen can both pick up electrons. (If hydrogen loses its only electron it will end up as a nucleus with no electrons around it.) They will share electrons to form covalent bonds. This results in the formula H2O

26. F F Covalent Bonding Example 3 Seven of the elements have such high attraction for electrons that they will never exist as individual, unattached atoms. Anytime these elements are present in pure form they will bond to other atoms of the same element. For example a fluorine atom will readily bond to a second fluorine atom. Resulting in F2.

27. Covalent Bonding These elements are called diatomic elements, and the molecules they form are called diatomic molecules. The definition of a diatomic molecule is: A molecule made up of 2 atoms of the same element.

28. Covalent Bonding The seven diatomic elements are: Hydrogen, H2 Nitrogen, N2 Oxygen, O2 Fluorine, F2 Chlorine, Cl2 Bromine, Br2 Iodine, I2

29. Covalent Bonding Covalent compounds are made up of small units called molecules. The formula for a covalent compound tells the actual number of atoms, of each element, found in each molecule. For example: The formula for water is H2O. This formula indicates that each water molecule is made up of two Hydrogen atoms and one Oxygen atom.

30. What is a “polar” covalent bond? • Covalent bonds involve sharing electrons. • The electrons may be shard equally (nonpolar covalent) or unequally (polar covalent). • Example: H2 shares electrons equally, but HCl does not. Therefore, H2 contains a nonpolar covalent bond and HCl contains a polar covalent bond.

31. Polarity of water • Take for example H2O. When we draw the structure it looks like: O H H • The oxygen atom pulls electrons away from the hydrogen atoms. This unequal sharing results in polar bonds, which have a “more negative” end and a “more positive” end.

32. Polar Water Molecule EX: O H H H2O is a polar molecule. Bond polarity affects the properties of a material such as melting and boiling points, crystal structure and acidity. more negative more positive

33. Covalent Bonding The characteristics shared by covalent compounds are: Molecular structure –individual units Low melting and boiling points, most covalent compounds are gases or liquids at room temperature (the larger the molecule the higher its melting and boiling point) Soluble in covalent solvents such as alcohol or benzene. Nonelectrolytes gas atoms and molecules

34. Cl H C H C C C C H H C Cl Covalent Bonding Paradichlorobenzene (moth balls) (C6H4Cl2) is a covalent compound. Its properties are: Molecular solid • Melting point = 53.1˚C • Boiling point = 174.55˚C • Soluble in alcohol, ether, acetone and benzene • Nonelectrolyte

35. Comparison of Bonding Types ionic covalent ions molecules The properties of a material depend on the structure -different bond types result in different properties. molten salts conductive non- conductive Both determined by valence electrons transfer of electrons sharing of electrons high mp low mp not usually water soluble water soluble

36. Polyatomic Ions/Radicals Some groups of atoms are covalently bonded together so strongly that the stay together during chemical reactions and act as a single unit. These groups of atoms become charged, with the charge being spread out through out the group. These groups are called polyatomic ions or radicals.

37. Polyatomic Ions/Radicals Definition - groups of atoms bonded together that act as a charged unit. Examples: Ammonium - NH4+1 Sulfate - SO4-2 Acetate - C2H3O2-1 Phosphate - PO4-3 What kind of bond do you think hold the atoms in a polyatomic ion together? What kind would hold two polyatomic ions together?

38. Type of bond? – Ionic, Polar Covalent, or Nonpolar Covalent? TiO2 CH4 NaI CS2 O2 KCl AlCl3 CsF HBr

39. Types of Bonds 3. Metallic bond Metals bonding with other metals do not gain or lose e- or share e- unequally. These bonds are created from the delocalized e- that hold metallic atoms together.

40. Chemical Formulas • A chemical formula is a combination of symbols that represents the composition of a compound. • Chemical symbols are used to indicate types of elements present. • Subscripts are used to indicate the number of atoms for each element present.

41. What are the “parts” of a formula? chemical symbols C8H18 number of atoms of each element • 8 atoms of carbon • 18 atoms of hydrogen

42. Charges of Monatomic Ions • Because atoms want to reach an octet of valence electrons, the oxidation numbers, (positive or negative charges) can be predicted for single atoms (monatomic). • Metals tend to have positive oxidation numbers. (lose e-) • Nonmetals tend to have negative oxidations numbers. (gain e-)

43. Remember... 0 1+ 2+ 3+ varies 3- 2- 1- varies +1 to +7

44. Oxidation Numbers (charges) of Polyatomic Ions • Polyatomic ions are ions that are made up of two or more atoms. • Refer to your table of Polyatomic ions. • Polyatomic ions generally have the following endings: “ate” or “ite” Ex: NO2- nitrite PO43- phosphate SO42- sulfate

45. Polyatomic Ions

46. Oxidation Numbers • The sum of the oxidation numbers in a compound must equal zero. Ex: CaCl2 = Ca+2 + Cl- + Cl- 2 positive charges – 2 negative charges = 0 • The charge on a monatomic ion is its oxidation number. Ex: Ba+2 has an oxidation of +2 Cl- has an oxidation of -1

47. What happens when the predicted charge can vary?? The oxidation number of a transition element is shown using Roman numerals to indicate the charge. The Roman numeral indicating oxidation. Ex: iron (II) is Fe+2 iron (III) is Fe+3

48. Why is aluminum oxide Al2O3?

49. Writing Ionic Formulas Ionic compounds are composed of metals and nonmetals. • Ionic compounds are made from the gaining or losing of electrons and the resulting electrostatic force that holds the ions together. • The sum of the oxidation numbers in a compound must equal zero.

50. Writing Ionic Formulas • When writing formulas, the cation (metal ion) is always written before the anion (nonmetal ion). • When using polyatomic ions, refer to charge given on your table. NOTE: There is only one polyatomic cation (NH4+). The rest are all are polyatomic anions.