This preview shows pages 1–3. Sign up to view the full content.
This preview has intentionally blurred sections. Sign up to view the full version.View Full Document
Unformatted text preview: 1 Atomic Mass c Atoms are so small, it is difficult to discuss how much they weigh in grams c Use atomic mass units. c an atomic mass unit (amu) is one twelfth the mass of a carbon-12 atom c This gives us a basis for comparison c The decimal numbers on the table are atomic masses in amu They are not whole numbers c Because they are based on averages of atoms and of isotopes. c can figure out the average atomic mass from the mass of the isotopes and their relative abundance. c add up the percent as decimals times the masses of the isotopes. Examples c There are two isotopes of carbon 12 C with a mass of 12.00000 amu(98.892%), and 13 C with a mass of 13.00335 amu (1.108%) c There are two isotopes of nitrogen , one with an atomic mass of 14.0031 amu and one with a mass of 15.0001 amu. What is the percent abundance of each? The Mole c The mole is a number c a very large number, but still, just a number c 6.022 x 10 23 of anything is a mole c a large dozen c The number of atoms in exactly 12 grams of carbon-12 The Mole c Makes the numbers on the table the mass of the average atom c Average atomic mass c Just atomic mass 2 Molar mass c mass of 1 mole of a substance c often called molecular weight. c To determine the molar mass of an element, look on the table. c To determine the molar mass of a compound, add up the molar masses of the elements that make it up. Find the molar mass of c CH 4 c Mg 3 P 2 c Ca(NO 3 ) 2 c Al 2 (Cr 2 O 7 ) 3 c CaSO 4 2H 2 O Percent Composition c Percent of each element a compound is composed of....
View Full Document
This note was uploaded on 09/28/2009 for the course CHEM 102 taught by Professor Freeman during the Spring '08 term at South Carolina.
- Spring '08
- Atomic Mass