Masses of Atoms

1. Masses of Atoms Chapter 19-2 Pages 584-587

2. Atomic Mass - Nucleus of atom contains most of the mass  Protons and neutrons are far more massive than electrons

3. Atomic Mass • Atomic mass unit (amu) – unit of measurement used for atomic particles  Mass of a proton = 1 amu = mass of 1 neutron  1 amu ≈ 1/12 mass of a carbon atom containing 6p+ and 6 n

4. Numbers • Atomic number – number of protons in an atom  Each element has a unique number of protons (atomic number)

5. Numbers - Mass number – the sum of the number of protons and number of neutrons in the nucleus - Mass # = Atomic # + # of neutrons

6. Isotopes – atoms of the same element with different number of neutrons  Isotopes of an element have slightly different properties - Examples: Carbon-12 (C-12)  most common and stable Carbon-14 (C-14)  unstable and radioactive - To identify isotopes, use the element name followed by the mass number  Examples: Boron-10, Boron-11

7. - Average atomic mass – the weighted-average mass of the mixture of all isotopes of an element

8. Boron-11  80% (4/5) of all boron atoms are the isotope Boron-11 Boron-10  20% (1/5) of all boron atoms are the isotope Boron-10 4/5 x (11 amu) + 1/5 x (10 amu) = 10.8 amu this is the average atomic mass ↑

9. ** The average atomic mass is always very close to the mass number of the most common (most abundant) isotope, and can be rounded to the nearest whole number when using mass numbers.