Stoichiometry
AP Chemistry Unit 2 Notes
Stoichiometry You should understand all that is presented in chapter 3 of your text (Zumdahl: Chemistry, 8th edition). Some of the highlights are presented below.
Atomic Masses (Section 3.1)
Nearly every element is made up of atoms of more than one isotope for that element. A few, like Be, only have one isotope. Others can have a large number of isotopes. Tin (Sn) has ten isotopes. (No pun intended.) Isotopic abundance is determined by the use of a mass spectrometer. Using a magnetic field, this device can sort the isotopes of a pure element into sub-groups based on mass. This way, we can see what percentage of an element is composed of each of the isotopes. The atomic mass values given on the periodic table are average masses, reflecting the weighted average of the isotopes for each element.
average atomic mass = % of each isotope atomic mass of each isotope
For example, Gallium has two isotopes. Ga-69 is 60.108% abundant and has a mass of 68.92558 amu. Ga-71 is 39.892% abundant and has a mass of 70.9247005 amu.
average atomic mass = .60108 x 68.92558amu .39892 x 70.9247005 average atomic mass = 69.723amu
The Mole (Section 3.2)
In chemistry, we use the counting unit of a mole (mol). This is defined as the amount of particles that are in exactly 12.00 g of the isotope 12C. This number is 6.022 x 1023. Using this same relationship, it can be reasoned that 6.022 x 10 23 amu = 1.000 g. Therefore, the average mass of one mole of any element in grams is equal to the average atomic weight, also called the average molar mass.
Percent Composition of Compounds and Empirical Formulas (Sections 3.5 and 3.6)
The chemical formula of a compound shows the exact molar ratio of the different elements in the compound. The numbers of each element are recorded using a subscript after each chemical symbol. The subscript of one is understood and never written. The percent
Stoichiometry You should understand all that is presented in chapter 3 of your text (Zumdahl: Chemistry, 8th edition). Some of the highlights are presented below.
Atomic Masses (Section 3.1)
Nearly every element is made up of atoms of more than one isotope for that element. A few, like Be, only have one isotope. Others can have a large number of isotopes. Tin (Sn) has ten isotopes. (No pun intended.) Isotopic abundance is determined by the use of a mass spectrometer. Using a magnetic field, this device can sort the isotopes of a pure element into sub-groups based on mass. This way, we can see what percentage of an element is composed of each of the isotopes. The atomic mass values given on the periodic table are average masses, reflecting the weighted average of the isotopes for each element.
average atomic mass = % of each isotope atomic mass of each isotope
For example, Gallium has two isotopes. Ga-69 is 60.108% abundant and has a mass of 68.92558 amu. Ga-71 is 39.892% abundant and has a mass of 70.9247005 amu.
average atomic mass = .60108 x 68.92558amu .39892 x 70.9247005 average atomic mass = 69.723amu
The Mole (Section 3.2)
In chemistry, we use the counting unit of a mole (mol). This is defined as the amount of particles that are in exactly 12.00 g of the isotope 12C. This number is 6.022 x 1023. Using this same relationship, it can be reasoned that 6.022 x 10 23 amu = 1.000 g. Therefore, the average mass of one mole of any element in grams is equal to the average atomic weight, also called the average molar mass.
Percent Composition of Compounds and Empirical Formulas (Sections 3.5 and 3.6)
The chemical formula of a compound shows the exact molar ratio of the different elements in the compound. The numbers of each element are recorded using a subscript after each chemical symbol. The subscript of one is understood and never written. The percent