Counting Atoms

1. CountingAtoms

2. Atomic Number • Atoms of different elements have different numbers of protons. Atoms of the same element all have the same number of protons. • The atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. • An element’s atomic number is indicated above its symbol on the periodic table.

3. Elements are placed in order of increasing atomic number. • The atomic number identifies an element.

4. Isotopes • Isotopes are atoms of the same element that have different masses. • The isotopes of a particular element all have the same number of protons and electrons, but different numbers of neutrons. Isotopes of Hydrogen

5. Mass Number • The mass number is the total number of protons and neutrons in the nucleus of an isotope. • This isotope of oxygen has 8 protons and 8 neutrons in the nucleus. Its mass number is 16.

6. Designating Isotopes • There are two methods for specifying isotopes. • In the first method, the massnumber is written with a hyphen after the name of the element (uranium-235). • The second method shows the composition of a nucleus as the isotope’s nuclear symbol ( ). Two different isotopes of carbon.

7. The superscript indicates the massnumber and the subscript indicates the atomic number.

8. The number of neutrons is found by subtracting the atomic number from the mass number • (mass # - atomic # = # of neutrons). • Example: 235 (protons + neutrons) – 92 (protons) = 143 neutrons. • For a neutral atom, the number of protons equals the number of electrons. Uranium

9. Example: How many protons, electrons, and neutrons are there in an atom of chlorine-37? Given: name and mass number of chlorine-37 Remember: atomic number = number of protons = number of electrons From the periodic table, the atomic number of chlorine is 17. Mass # - atomic # = # of neutrons 37 – 17 = 20 An atom of chlorine-37 is made up of 17 protons, 17 electrons, and 20 neutrons.

10. The charge on an ion (an atom with a positive or negative charge) indicates an imbalance between protons and electrons. • Too many electrons produces a negative charge, too few, a positive charge.

11. Nuclide is a general term for any isotope of any element. • Below are three hydrogen nuclides.

12. Relative Atomic Mass • Masses of atoms expressed in grams are very small. For example, an atom of oxygen-16 has a mass of 2.657 x 10-23 g. • For simplicity, scientists use the atomic mass unit (amu) to express the mass of atoms. 1 amu is 1/12 the mass of a carbon-12 atom. • The atomic mass of any nuclide is determined by comparing it with the mass of the carbon-12 atom.

13. Why use carbon? • Remember scientists use standards of measurement that are constant and are the same everywhere. • In order to set up a relative scale of atomic mass, one atom was arbitrarily chosen as the standard and assigned a value. • The carbon-12 atom was arbitrarily assigned a mass of exactly 12 atomic mass units, or 12 amu. • Remember one atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom.

14. Average Atomic Masses of the Elements • Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. • Calculating the average atomic mass depends on both the mass and the relative abundance of each of the element’s isotopes.