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Week Three (2/4) Notes

by: Kelly Johnson

Week Three (2/4) Notes Chem 107

Kelly Johnson
GPA 3.63

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These notes cover what should be on quiz two and part of exam 2
General Chemistry for Health Science
Jacqueline Butler
Class Notes
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This 8 page Class Notes was uploaded by Kelly Johnson on Sunday January 31, 2016. The Class Notes belongs to Chem 107 at West Chester University of Pennsylvania taught by Jacqueline Butler in Winter 2016. Since its upload, it has received 117 views. For similar materials see General Chemistry for Health Science in Chemistry at West Chester University of Pennsylvania.


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Date Created: 01/31/16
Chemistry Unit 2- Chapter 3 Notes up to Drawing Lewis Structures and VSEPR Theory 1. Chemical Bonding a. Molecular and Ionic bonding i. All molecules can be classified by the bonds that hold them together 1. Molecular a. These are covalent bonds, meaning they share the electrons instead of donating b. Combination of a Nonmetal + nonmetal c. Can exist as one molecule 2. Ionic a. In an ionic bond, electrons from one molecule transfers to another electron. b. This transfer will give the atoms a charge i. Positive charge- cation ii. Negative charge- anion c. Combination of a metal and a nonmetal d. Only exist in a lattice structure made of multiple of the same ions ii. Examples 1. On slide 5, this shows how NaCl is created a. Sodium loses an electron becoming isoelectric with neon, thus gaining a positive charge and becomes the cation b. Chlorine accepts this electron, and becomes isoelectric with argon, this becomes the anion c. These molecules will combine in a lattice structure to stabilize b. Chemical Formulas i. About 1. Chemical formulas indicate the number and types of atoms contained in a molecule 2. This is just to show the composition and does not indicate structure a. C 3 6 – this shows that the element has 3 Carbon, 6 Hydrogen, and 1 Oxygen. It also shows that this is a combination of nonmetals, making it a molecular compound. 3. Chemical formulas are known as the simplest structural unit ii. Examples 1. Na O2 sodium oxide, a metal and a nonmetal making it ionic 2. Al (SO ) - Aluminum sulfate, a metal and nonmetal 2 4 3 compound, ionic c. Formulas and Names of Compounds i. As stated before, a combination of nonmetals forms a molecular compound, and a metal and a nonmetal form an ionic compound ii. Transition metals may have multiple charge states, so the charge is represented by a roman numeral in the name of the compound iii. Zinc and Silver are charge exceptions, Zinc is always +2 and Silver is always +1 iv. Writing Ionic formulas 1. Rules a. The charges of the two molecules must be equal to zero b. Cation symbol will be written first, followed by the anion c. You can cross check the charges and number of molecules to see if they are equal 2. Examples a. Magnesium Chloride i. Magnesium always has a +2 charge, and Chlorine always has a -1 charge. For these sum of these charges to equal zero, there must be two chlorine molecules for each magnesium. ii. MgCl2 b. Cobalt (III) oxide i. Oxygen always has a -2 charge, and since cobalt is a transition metal, its charge can vary. However, the III stands for its +3 charge. In this case, you need 2 cobalt for every 3 oxygen. ii. Co2O3 v. Naming Ionic Compounds 1. Rules a. Make sure you have a metal and a nonmetal b. The cation will always be named first c. If the element is a transition metal, remember to add its charge as a roman numeral d. Name the anion second, and if it is binary end in -ide 2. Examples a. NaCl i. Ionic ii. Na is sodium (a cation) and Cl is chlorine (an anion) iii. It is binary, so chlorine will end in –ide iv. Sodium Chloride b. K 2 i. Ionic ii. K is potassium (a cation) and S is sulfur (an anion) iii. It is binary, so sulfur will end in –ide iv. Potassium Sulfide c. Al2O3 i. Ionic ii. Al is aluminum (a cation) and O is oxygen (an anion) iii. It is binary, so oxygen will end in –ide iv. Aluminum Oxide d. Cu 2 i. Ionic ii. Cu is copper (a transition metal, cation) and O is oxygen (an anion) 1. Copper has two molecules that must equal oxygens -2 charge, making it +1 each iii. It is binary, so oxygen will end in -ide iv. Copper (I) oxide vi. Naming molecular compounds 1. Rules a. Make sure all elements are nonmetals b. Atoms listed first will only receive prefixed if there is more than one c. The second element will always have a prefix d. The second element will also end in –ide if it is binary 2. Prefixes you must know a. Mono- 1 b. Di- 2 c. Tri- 3 d. Tetra- 4 e. Penta- 5 f. Hexa- 6 g. Hepta- 7 h. Octa- 8 i. Nona- 9 j. Deca- 10 3. Examples a. CO 2 i. Both Nonmetals, binary ii. C-Carbon O- Oxygen iii. Carbon Dioxide b. CO i. Both nonmetals, binary ii. C-Carbon O- Oxygen iii. Carbon Monoxide c. P4S10 i. Both nonmetals, binary ii. P- Phosphorus S-Sulfur iii. Tetraphosphorus Decasulfide d. H 2 i. Both nonmetals, binary ii. H- hydrogen O- oxygen iii. Dihydrogen Monoxide vii. Writing Formulas of Molecular Compounds 1. Rules a. The least electronegative element is written first i. Left side of the periodic table ii. The trifecta (Oxygen, Chlorine, and Fluorine**) along with nitrogen are the most electronegative iii. Do not like to share electrons b. If carbon is listed, it will go first (THIS IS AN EXCEPTION) c. If there is only one of the first atom, no prefix will be listed, and no subscript will be needed. If there is a prefix listed, a subscript equal to that prefix will be needed in your formula. d. Determine the correct number of atoms of the second element via the prefix of the second listed element. Write the correct subscript 2. Examples a. Diarsenic Pentabromide i. There are 2 arsenic molecules (Ar) ii. There are 5 bromine molecules (Br) iii. Ar2Br5 b. Carbon Tetrachloride i. There is 1 carbon molecule (C) ii. There are 4 chlorine molecules (Cl) iii. CCl4 c. Sulfur Hexafluoride i. There is 1 sulfur molecule (S) ii. There are 6 fluoride molecules (F) iii. SF6 d. Iodine Trioxide i. There is 1 iodine molecule (I) ii. There are 3 oxygen molecules (O) iii. IO3 viii. Polyatomic Ions 1. Definition a. Charged groups of covalently bonded atoms 2. Ones we must+know a. NH4 - Ammonium b. OH - Hydroxide c. CN - Cyanide 2- d. SO4 - Sulfate e. ClO4 - Perchlorate f. NO3 - Nitrate 2- g. O2 - Peroxide h. PO4 - Phosphate i. CO3 - Carbonate - j. HCO3 - Bicarbonate or Hydrogen Carbonate ix. Writing Polyatomic Compound Formulas 1. Rules a. The sum of the charges in the formula must equal 0 b. The cation will be written first, followed by the anion c. If there is more than 1 polyatomic needed to make the formula be equivalent to 0, use parenthesis around the whole ion 2. Examples a. Lithium Hydroxide i. Lithium (Li) has a charge of +1 and Hydroxide (OH) has a -1 charge. These equal 0 when added. ii. LiOH b. Calcium Nitrate i. Calcium (Ca) has a charge of +2 and Nitrate (NO ) has a -1 charge. To equal 0, you need 3 2 nitrate ions ii. Ca(NO )3 2 c. Aluminum Phosphate i. Aluminum (Al) has a charge of +3 and phosphate (PO ) 4as a -3 charge. These are equal. ii. AlPO 4 x. Writing Polyatomic Compound Names 1. Rules a. Name the cation first b. Name the anion second 2. Examples a. Na 2O 4 i. This contains sodium (a cation) and Sulfate (an anion) ii. Sodium Sulfate b. Mg(OH) 2 i. This contains Magnesium (cation) and hydroxide (anion) ii. Magnesium Hydroxide c. (NH 4 3O 4 i. This contains ammonium (cation) and phosphate (anion) ii. Ammonium Phosphate d. NiCO 3 i. This contains Nickle (cation) and carbonate (anion) ii. Nickle Carbonate 3. CHART TO HELP NAMING IT IN HER SLIDESHOW ON SLIDE 26 xi. Formulas of Acids 1. Rules/Characteristics a. Acids begin with an H (hydrogen) b. Charges add up to 0 2. Naming Acids a. If it is binary, it is a hydro____ic acid b. If it is not binary, it is a _____ic acid i. _____ate=_____ic ii. _____ite=_____ous 3. Examples a. HCl i. Binary ii. Hydrochloric Acid b. HI i. Binary ii. Hydroiodic Acid c. H2S i. Binary ii. Hydrosulfuric Acid d. H3PO 4 i. Not Binary ii. Phosphoric Acid e. H2SO 4 i. Sulfuric Acid xii. Acids 1. Definition a. Molecular compounds that form H+ when dissolved it water i. (aq) is often in the formula to indicate this b. Contain H+ cation and an anion i. Binary Acids have the H+ cation and a nonmetal anion ii. Oxyacids have the H+ cation and a polyatomic anion 2. Characteristics a. Sour taste b. Many metals dissolve when in contact with an acid c. Generally have a formula starting with H xiii. Naming Oxyacids 1. Rules a. If the polyatomic ion ends in –ate, it should be changed to –ic b. If the polyatomic ion ends in –ite, it should be changed to –ous c. Write acid after all 2. Examples a. H 2O 4 i. Anion is sulfate, which ends it –ate so it will end in –ic ii. Sulfuric Acid b. H 2O 3 i. Anion is sulfite, which ends in –ite, so it will end it –ous ii. Sulfurous Acid c. H S 2 i. Anion is Sulfur, not a polyatomic ion ii. Hydrosulfuric acid d. HClO 3 i. Anion is Chlorate ii. Chloric Acid e. HNO 2 i. Anion is Nitrite ii. Nitrous acid xiv. Writing formulas for acids 1. Rules a. If the name ends in acid, the formula starts with H b. Write the formula as if ionic, even though it is molecular c. If binary, it will have a hydro prefix. d. Ends in –ic, it is a polyatomic ending in –ate e. Ends in –ous, it is a polyatomic ending in –ite 2. Examples a. Hydrosulfuric acid i. Hydro indicates binary ii. Acid indicates it starts with H iii. Sulfuric means sulfur iv. H+1, S-2 v. H 2 b. Carbonic Acid i. Not binary ii. Acid indicates H iii. –ic means carbonate iv. H+1, CO -3 v. H 2O 3 c. Sulfurous Acid i. Not binary ii. Acid indicates H iii. –ous means sulfite iv. H+1, SO 32 v. H 2O 3 d. Chlorous acid i. Not binary ii. Acid indicates H iii. –ous means chlorite iv. H+1, ClO -2 2 v. H 2lO 2 e. Phosphoric acid i. Not binary ii. Acid indicates H iii. –ic means phosphate iv. H+1, PO 43 v. H 3O 4 f. Hydrobromic acid i. Binary ii. Acid indicates H iii. Bromic indicates bromine iv. H+1, Br-1 v. HBr d.


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