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Chemistry 111 lecture notes 2.7-3.3 (Dr. Patton)

by: Jamisha Evans

Chemistry 111 lecture notes 2.7-3.3 (Dr. Patton) CHEM 111 003

Marketplace > Eastern Kentucky University > Chemistry > CHEM 111 003 > Chemistry 111 lecture notes 2 7 3 3 Dr Patton
Jamisha Evans
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These notes cover... 2.7-Naming compounds 3.1- Atomic mass 3.2- Avogadros number and molar mass 3.3- Molecular mass and formula mass
General Chemistry 1
Dr. James Patton
Class Notes
Chem, 111
25 ?




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This 4 page Class Notes was uploaded by Jamisha Evans on Wednesday September 14, 2016. The Class Notes belongs to CHEM 111 003 at Eastern Kentucky University taught by Dr. James Patton in Fall 2016. Since its upload, it has received 15 views. For similar materials see General Chemistry 1 in Chemistry at Eastern Kentucky University.


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Date Created: 09/14/16
Chemistry 111 Lecture notes (Dr. Patton) 2.7-Naming Compounds  Acids  Definition: Substance that yields H when dissolved in water  EXAMPLE: HCL is an acid. It contains Hydrogen and yields H when dissolved in water. + - HCL H 2 H + Cl  Acids have several groups  Simple acids: Contains hydrogen and an anion EXAMPLE: HCL (Hydrochloric acid)  Oxoacids: Contains hydrogen, a central atom and oxygen EXAMPLE: HNO (Nitric acid) 3  Oxo-acids often have variable numbers of oxygen EXAMPLE: HNO (Nitr3c acid) HNO (2itrous acid)  Polyprotic acids: Contains more than one ionizable hydrogen EXAMPLE: H SO (Sulfuric acid) 2 4 **The prefix oxo- means oxygen**  Bases -  Definition: Substance that produces OH ions when dissolved in water  EXAMPLE: NaOH + - NaOH H2O Na + OH  Hydrates  Definition: Hydrates are formed from water molecules in the crystal of an ionic compound. They aren’t specifically associated with a cation or anion and they can be removed by heating.  EXAMPLE: FeCL ∙ 6H3O (Ir2n (III) chloride heptahydrate) 3.1- Atomic Mass  Atomic mass  Definition: The mass of an atom of a chemical element. AKA atomic weight. (reported in amu)  Carbon-12 isotope is the reference used for the atomic mass unit. 1 atom of carbon-12 is 12 amu  Mass of atoms are really small. EXAMPLE: There are two isotopes an element has, one having an atomic mass and percent natural abundance of 11.0093 amu (80.22%) and 10.0129amu (19.78%) Add the product of the (amu) (percentage) of each isotope. But make sure you divide the percentages by 100  (11.0093) (.8022) + (10.0129) (.1978) =10.81 amu *This element is Boron because when you look at the periodic table you will see that the atomic mass of boron is 10.81. * 3.2- Avogadro’s number and the molar mass of an element  Avogadro’s number  Definition: A counting number is used to represent very large numbers of atoms, molecules, or ions  The unit used to count large numbers of small particles is the mole  1 mole: 6.022 x 10 (Avogadro’s number; N ) Areported in atoms)  Mol is the SI unit for amount of substance  The mass of a single atom (in amu) is numerically equal to the mass (in grams) of 1 mol of that element. EXAMPLE: 1 atom K has a mass of 39.0983 amu 1 mol of atoms of K has a mass of 39.0983 g  Convert mass to atoms  Mass to moles to atoms  Convert atoms to mass  Atoms to moles to mass  Molar mass  Definition: The mass of 1 mole of atoms of that element (reported in units of grams/mol)  EXAMPLE: Grams of gold in 15.3 moles of gold (Au)? 1. State the given first- 15.3 mol of Au 2. Convert moles of gold to mass of gold 3 15.3 mol Au x 197 g of Au/1 mol of Au = 3.01 x 10 g of Au 3.3- Molecular Mass and formula mass  Molecular mass  Definition: Mass of a molecule (reported in amu)  EXAMPLE: Molecular mass of CCl (Add4mass of C and mass of Cl ) 4 C: 12.01 x 1= 12.01 Cl 4 35.45 x 4= 141.8 153.8 amu  Formula mass  Definition: Sum of atomic weights of the atoms in the empirical formula of the compound. AKA formula weight (reported in grams)  EXAMPLE: Formula mass of NaBr Na: 22.99 x 1=22.99 Br: 79.90 x 1=79.90 102.89g


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