1. Objectives:  1. Name and describe the three subatomic particles in an atom.  2. Determine the number of protons, neutrons, and electrons in an atom or ion.  3. Define isotope and atomic mass. Key Terms:  Proton, neutron, atomic mass unit, atomic number, ion, mass number, isotope, atomic mass Atomic Theory

2. Atoms are made of even smaller subatomic particles called protons, neutrons, and electrons.  protons - have a positive charge, found in the nucleus and have an atomic weight of 1 amu  electrons - have a negative charge, move in the space around the nucleus, the number is equal to protons in a neutral atom and they have no appreciable mass  neutrons - have no net charge, found in the nucleus, and have an atomic weight of 1 amu Subatomic Particles

3. A. Atomic number – number of protons in the nucleus.  The number of protons determines the positive charge of the atom  The number of protons determines the atom’s identity  The electrons and neutrons of an atom can change but it cannot change the amount of protons and still be of the same element  In neutral atoms the # of protons = # of electrons B. Mass Number = protons + neutrons  32amu – 16(+) = 16(+) C. & D. – Chemical symbol and name Atomic Number

4. Whenever an atom gains or loses electrons it becomes an ion. You can find the net charge of the atom by subtracting the number of electrons from the atomic number.  Loses an e-… the result is positive (atomic # > electrons) called a cation - overall positive charge  example: Ca… 20(+) - 18(-) = 2+ or Ca 2+  Gains an e-… the result is negative (atomic # < electrons) called an anion - overall negative charge  example: O… 8(+) - 10(-) = 2- or O 2- Ions

5. Most elements in the first two rows of the periodic table have at least two known isotopes.  Isotopes have the same number of protons and electrons but differ in the number of neutrons. (mass - protons = neutrons)  Atoms in an element normally contain mixtures of isotopes in specific ratios  Isotopes are named after their masses (neutrons + protons)  Example: hydrogen-1 (H-1), hydrogen-2 (H-2), hydrogen (H-3) Isotopes

6. Determining Average Atomic Mass from the Isotopes Found in Nature  Look at the three common isotopes of silicon. Multiply the masses of the isotopes by their fractional abundances (percent found in nature) and add the products together. Element Symbol Mass amu Fractional Abundance Contributing Mass Si-28 27.977 92.21% 25.80 Si-29 28.976 4.70% 1.36 Si-30 29.974 3.09% 0.93 Average Atomic Mass 28.09amu  Since most atoms exist in isotopes with known ratios, the mass number expressed on the periodic table will usually not agree exactly with the average amu number.

7. The relative atomic mass of an atom is expressed in atomic mass units (amu).  This unit is derived from the carbon atom and is measured to be 1 / 12 of the mass of the carbon-12 atom.  The amu expresses the most common isotope of the atom  C-12 is more common than C-13 or C-14 Relative Atomic Mass

8.  Chemical notation  Displays atomic mass, atomic number and charge  If a charge is not listed then the number of electrons is equal to the number of protons  In the example to the left  11(+), 23-11 = (+), 10 (-)  1+ charge tells you that there is 1 more proton than electrons Chemical Notation