2. Na Mass # Atomic # Electric charge # of atoms

3.  Also referred to as a “salt”  Formation involves a transfer of electrons  Usually made up of a metal and a non-metal  Are good conductors when they can be melted or dissolved  Typically have extremely high melting points

4. Electron acceptor (Cl) meets electron donor (Na) Ions attract to form a neutral pair e- jumps from Na to Cl

5.  Smallest building blocks are ions, NOT MOLECULES  Large numbers of ions can attract to form clusters and eventually crystals Ion pair Ion cluster Crystal lattice

6.  Cations – positively charged ions ◦ Na + Ca 2+ Al 3+  Anions – negatively charged ions ◦ Cl - O 2-  Polyatomic ions – ions made up of more than one type of atom ◦ NO 3 - SO 4 -2 PO 4 -3

7.  The number of e- gained, lost or shared ub compound formations ◦ Alkali metals +1 ◦ Alkaline earth metals +2 ◦ Oxygen group -2 ◦ Halogens -1

8.  K + and N 3- ◦K3N◦K3N  Ca 2+ and N 3- ◦ Ca 3 N 2  Ba 2+ and NO 3 - ◦ Ba(NO 3 ) 2  Criss-cross rule

9.  Binary – made of 2 ions  Write cation first  Change anion ending to –ide  Na + and Cl - ◦ Sodium chloride  H + and F - ◦ Hydrogen fluoride  CaBr 2 ◦ Calcium bromide

10.  Name the cation  Polyatomic ion name is unchanged  NaNO 3 ◦ Sodium nitrate  Zinc carbonate ◦ ZnCO 3

11.  Also called covalent compounds  A molecule is a neutral group of atoms that are held together by covalent bonds  The valence e- are shared by the atoms  Covalent bonding usually occurs between 2 non-metals ◦ H 2 O, CO 2, O 2, NO

12.  Use prefixes 1mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- 9nona- 10deca-

13.  P 4 O 10 N2O3N2O3  As 2 O 5  OF 2 Tetraphosphorous decoxide Dinitrogen trioxide Diarsenic pentoxide Oxygen difluoride

14. H2H2 O2O2 N2N2  Cl 2  Br 2 I2I2 F2F2  7 diatomic molecules  No noble gases  Halogens and N, O, H  They are all gases (not noble gases) except for Br and I  “Honcl brif”

15.  H 2 SO 4  HF  H 3 PO 4  H 2 SO 3  H 2 CO 3  HNO 3 Sulfuric Acid Hydrofluoric Acid Phosphoric Acid Sulfurous Acid Carbonic Acid Nitric Acid

16.  Calcium bromide  Chromium (III) acetate  Barium sulfate  Copper (I) sulfide  Sulfur hexafluoride CaBr 2 Cr(C 2 H 3 O 2 ) 3 BaSO 4 Cu 2 S SF 6

17.  Cr 2 (C 2 O 4 ) 3  Hg(CN) 2  Cu(ClO 4 ) 2  ZnC 4 H 4 O 6 Chromium (III) oxalate Mercury (II) cyanide Copper (II) perchlorate Zinc tartrate

18.  The mass of a compound  In order to calculate molar mass (also called molecular weight) you add up the masses of each element in the compound ◦ Be aware of subscript numbers that designate the amount of atoms per element  You get the masses from the periodic table  **be careful when rounding the mass

19.  NaCl ◦ Na = 23 g/mol ◦ Cl = 35.5 g/mol H2OH2O ◦ H = 1 g/mol (but there are 2) = 2 g/mol ◦ O = 16 g/mol  HNO 3 ◦ H = 1 g/mol ◦ N = 14 g/mol ◦ O = 16 g/mol (but there are 3) = 48 g/mol  Ba(NO 3 ) 2 ◦ Ba = 137.3 g/mol ◦ N = 14 g/mol (but there are 2) = 28 g/mol ◦ O = 16 g/mol (but there are 6) = 96 g/mol 58.5 g/mol 18 g/mol 63 g/mol 261.3 g/mol

20.  All metal atoms in a metallic solid contribute their valence e- to form a “sea” of e- ◦ These e- move easily and freely because they are not tied to a specific atom  Delocalized electrons ◦ Metallic cation is formed All empty space is evenly distributed v.e-

21.  The attraction of a metallic cation for delocalized electrons  This accounts for a lot of theproperties of metals ◦ Range of melting points ◦ Malleability ◦ Ductile ◦ Durable  Hard to remove metallic cation because of the strong e- attraction ◦ Mobile e-  Explains why they are good conductors

22.  Find the difference in electronegativities of the two elements 0.51.7 Pure Covalent -share e- evenly -2 non metals and/or metalloids Non-polar Polar Covalent -Share e- but not evenly -One element holds e- more Polar Ionic -Metal and non-metal

24.  Count total valence electrons available  Place electrons around atoms  Ensure each atom has an octet (8) ◦ Or a pair for H (2)

25.  Draw the Lewis Structure for the molecule  Count the total number of... ◦ Bonded regions around the central atom  DOUBLE and TRIPLE bonds count as ONE REGION ◦ Unshared e- pair  Count as ONE REGION

26. Molecular Lewis Dot electron pairs around central atom Structure structure total shared unshared H CH 4 H-C-H 4 4 0 “tetrahedral” H NH 3 H-N-H 4 3 1 “trigonal H pyramidal” H 2 O H-O-H 4 2 2 “bent”

27. Molecule Total no. of electron pairs No. of shared pairs No. of unshared pairs Molecular shape

28.  A molecule is polar if ◦ There is a polar bond ◦ It is ASSYMETRICAL (not symmetric) O HH (-) (+) H H H H C Polar Non-Polar

29.  Symmetric (non-polar) ◦ Linear ◦ Tetrahedral ◦ Trigonal planar  If all elements around the center atom are the same  Asymmetric (polar) ◦ Bent ◦ Trigonal pyramidal

30.  Van der Waals forces (London Dispersion forces) ◦ Weak forces between non-polar molecules ◦ These forces determine volatility  Doesn’t take much nrg to break apart (liquid  gas)  Most likely to be a gas  Like playing red rover and only holding pinkies together

31.  Dipole-Dipole ◦ Attraction between polar molecules  Most likely to be a liquid  Play red rover and hold hands

32.  Hydrogen Bonding (H-Bonds) ◦ Between hydrogen (H) and a highly electronegative element  F, O, N ◦ Extreme case of dipole-dipole ◦ Strongest of the intermolecular forces  Play red rover and link elbows  Needs A LOT of nrg to break bonds

33.  Carbon has a mass of 12 g  Oxygen has a mass of 16 g  H 2 O molecules has a mass of 18 g  How do these #’s relate to the atom or compound? ◦ Atomic mass

34.  Amedeo Avogadro (1776-1856)  1 mole = 6.0221415 x 10 23 ◦ Particles ◦ Molecules ◦ Atoms ◦ Ions ◦ Formula units ◦ Etc, etc

35.  Determine the mass percentage of each element in the compound.

36.  Gives the lowest whole # ratio of elements in a compound.  The empirical formula for C 6 H 12 O 6 is  The empirical formula for C 2 H 6 is  * most basic ratio of elements in the compound CH 2 O CH 3