1. Chemical Bonding

2. Determined by difference in Electronegativity.
Bonding types
Ionic
Covalent
Metallic
Determined by difference in Electronegativity.

4. - Ions + charged atoms (gained or lost e-)
Formed from ve- only (group A #)
Cations- formed when a metal atom loses e- forming a + ion
Anion – formed when a nonmetal atom gains e- to form a - ion + -

5. Ionic compounds Crystalline solids (made of ions)
High melting and boiling points
Conduct electricity
Many soluble in water but not nonpolar liquid

7. Positive and negative IONS Attraction between positive and negative.

8. Positive and negative IONS Becomes brittle when the like charges line up and repel.

9. Ionic

10. Writing Formulas and Naming Ionic Compounds
RULES: Write the symbol and charge of the cation (+ charged metal or hydrogen) Write the symbol and charge of the anion (- charged nonmetal) Check the charges  they must = 0

11. Binary Ionic Compounds (2 atoms: metal & nonmetal)
sodium + iodine
Na I -1
calcium + oxygen
Ca O-2
potassium + sulfur
K S-2
(4) aluminum + oxygen
Al O-2
(5) Chlorine + magnesium
Mg Cl-1
NaI CaO K2S Al2O3 MgCl2

12. Binary Ionic Compounds (2 atoms: metal & nonmetal)
sodium + iodine
Na I -1
calcium + oxygen
Ca O-2
potassium + sulfur
K S-2
(4) aluminum + oxygen
Al O-2
(5) magnesium + chlorine
Mg Cl-1
Sodium iodide NaI calcium oxide CaO potassium sulfide K2S aluminum oxide Al2O3 magnesium chloride MgCl2

13. Binary Ionic Compounds Practice
Write formula and name each
1. chlorine + hydrogen 6. lithium + oxygen
2. calcium + sulfur 7. bromine + sodium
3. selenium + aluminum 8. potassium + sulfur
4. barium + fluorine 9. iodine + lithium
iodine + magnesium 10. strontium + chlorine

14. Covalent Bonds
Sharing of ve- between 2 or more atoms
molecule – smallest chemical unit of a substance capable of a stable independent existence

15. includes: types & #s of atoms
Covalent molecules are represented with a molecular formula – representation of actual composition of a covalent molecule
includes: types & #s of atoms
ex: H20 (water); O2 (oxygen); C12H22O11 (table sugar)

16. Diatomic molecules Element Molecular formula Hydrogen H2 Nitrogen N2
2 atoms of the same element covalently bonded; there are only seven
Element
Molecular formula
Hydrogen H2
Nitrogen N2
Oxygen O2
Fluorine F2
Chlorine
Cl2
Bromine
Br2
Iodine I2

17. Characteristics of Covalent Molecules
Usually liquids or gases at room temperature
Low melting and boiling points
Nonelectrolytes – do not conduct electricity

18. Nonpolar covalent
equal sharing of e-

19. Polar Covalent
unequal sharing of e- 1 atom pulls e- stronger/harder Results in a molecule with partially positive and negative regions

20. Binary covalent Molecules
2 element share e-
Naming Rules:
1st word – use name of element
*use element with lowest electronegativity
2nd word – ALWAYS use prefix to indicate the # of 2nd element (use the ending –ide)

21. drop the “a” or “o” of the prefix if the root word begins with a vowel
EX: penta-  pentoxide
mono-  monoxide

22. Naming Hydrates Prefix Number mono- 1 di- 2 tri - 3 tetra- 4 penta- 5
hexa- 6
hepta- 7
octa- 8
nona- 9
deca - 10
Naming Hydrates

23. Binary Covalent Compounds
SiO2 silicon dioxide S2O4 disulfur tetroxide CF4 carbon tetrafluoride PCl3 phosphorus trichloride

24. Writing formulas rules:
Write the symbol of the 1st element and number for the prefix (use a subscript)
Write the symbol of the 2nd element and number for the prefix (use a subscript)

25. Practice:
Carbon dioxide CO2 Tetranitrogen difluoride N4F2 Diphosphorus trioxide P2O3 Dihydrogen monoxide H2O

26. Metallic Bonding

27. Metallic Bonds
bonds between metals metallic atoms have loosely held electrons

28. electrons are more or less free to move from one atom to another.
metal ions floating in a “sea of mobile electrons” around a positive nucleus

29. Metallic Bonds results in the formation of alloys not compounds
metal atoms do not combine in fixed ratio

30. Properties of Metallic Bonds
Good conductors
Ductile
Malleable
Have luster (shiny)
High melting point
Solid (hard) at room temperature
Alloys: homogeneous mixture of metals (ex. steel, bronze, brass, pewter)

31. Some common alloys

33. Intra vs inter Intramolecular forces. within a molecule
Intra vs inter Intramolecular forces within a molecule “Atom to Atom” Intermolecular forces “Molecule to Molecule”

34. Without these forces there would be no liquids or solids
Without these forces there would be no liquids or solids. Everything would be a gas.

35. Intermolecular Forces 1) London Dispersion forces
Intermolecular Forces 1) London Dispersion forces very weak attraction between electron clusters in Noble gases and nonpolar molecules. There is nothing else holding them together.

36. London Dispersion forces

37. 2) Dipole-Dipole forces Attraction between positive and negative areas of different molecules.

38. Like Dissolves Like sodium chloride (Polar) dissolved in Water (Polar)

39. 3. Hydrogen Bonding. due to polarity. Positive Hydrogen end
3. Hydrogen Bonding due to polarity Positive Hydrogen end attracted to Negative area of another molecule

40. Hydrogen Bonding in DNA

41. 4. Electrostatic bonding
Ionic Bonding!!!!!