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ISOTOPES. Recall from grade 10: # of protons = # of electrons = atomic # # of neutrons = atomic mass – atomic # Example: carbon: atomic = # 6 atomic mass = 12.01 # of protons = 6 # of electrons = 6 # of neutrons = 12- 6 = 6 *always round off. Different variations of atoms of the same
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Recall from grade 10: # of protons = # of electrons = atomic # # of neutrons = atomic mass – atomic # Example: carbon: atomic = # 6 atomic mass = 12.01 # of protons = 6 # of electrons = 6 # of neutrons = 12- 6 = 6 *always round off
Different variations of atoms of the same element occur in nature. These variations are called isotopes. The average mass of the isotopes for each element is a characteristic of that element. Isotopes are atoms of the same element (same # of protons) with different # of neutrons. They have identical atomic #’s but different mass #’s (# of protons and neutrons). A = mass # X = symbol Z = atomic #
Isotopes are usually represented in several ways. sodium – 24 or 24Na The atomic mass unit is defined as 1/12th the mass of a carbon-12 atom. The magnitude of the amu is arbitrary. In fact, 1/24th the mass of a carbon atom or 1/10th the mass of the iron atom could have been selected just as easily.
3 Reasons for using the 1/12th the mass of a C-12 isotope: • Carbon is a very common element • Results in nearly whole-number atomic masses for most elements. • The lightest, H2, has a mass of ~ 1 amu.
Calculating Average Atomic Mass Example: Isotope Relavtive Abundance Mass Cl- 35 75% 35x.75 = 26.25 Cl – 37 25% 37 x .25 = 9.25 total: 26.25 + 9.25 = 35.50
