1. Structure of an Atom Atoms are composed of 3 subatomic particles: 1.Protons (p + ) – Positively charged, found in the nucleus, number of protons determines the element, has a mass of 1 amu 2.Neutrons (n) – Neutral particle, found in the nucleus, can vary in number, has a mass of 1 amu 3.Electrons (e - ) – Negatively charged, orbit around the nucleus, the farther from the nucleus the higher the energy level (relate to periods on PT), mass is so small relative to protons and neutrons we say it has a mass of 0 amu (it is actually 1/1836 that of a proton)

2. Charges and the Atom A positive and a negative charge cancel each other out….. Protons are positive Electrons are negative When you have an equal number of each you have a neutral atom, if they are different then there is a charge

3. If an atom has more protons (positives)than electrons (negatives) it will be positive…. These are known as CATIONS, usually metals If an atom has more electrons (negatives) than it has protons (positives) it will be negative….. These are known as Anions, usually nonmetals Ca ion = positive ions nion = negative ions If an atom becomes charged we call it an ion

4. Elements and the periodic table How do we read info from the periodic table? 7 N Nitrogen 14.0067 Atomic number (equal to the # of protons) Element symbol Abbreviation of element name Element name Name of the element Average atomic mass Weighted Average mass of 1 mole of all isotopes based on relative abundance

5. When things CHANGE in an atom. WHAT HAPPENS? If the electrons change  the atom has a positive or negative charge. AN ION IS FORMED. If the neutrons change  the mass of the atom changes. Version of the same atom. AN ISOTOPE IS FORMED.

6. Isotopes:

7. Isotope Notations Hyphen NotationNuclear Notation Helium - 3 He 3232 Abbreviation Mass Number Atomic Number Name or Abbreviation Mass Number

8. Mass Number - vs - Average Atomic Mass Mass Number: 1.Protons + Neutrons = Mass # 2.Is a whole number 3.Used when talking about isotopes Average Atomic Mass: 1.Average mass of all the isotopes a.Based on relative abundance 2.Has numbers after a decimal

9. Using Isotope Notations We will use isotope notations to keep track of how many protons and neutrons we have To fill in the table below you will need to use these tricks : Mass # = Protons + Neutrons Neutral atoms have equal numbers of protons and electrons Hyphen Notation Atomic Number Mass Number Number of protons Number of neutrons Number of electrons Nuclear Notation N-14 28 8045

10. Relative Abundance Relative abundance – refers to the abundance of naturally occurring isotopes – An example is Chlorine (Cl), which has 2 isotopes: 1)Chlorine – 35relative abundance = 75.7% 2)Chlorine – 37 relative abundance = 24.3% Calculating the weighted Average atomic mass: – Multiply the mass # by the relative abundance for each isotopes, then add them all together 35 x 0.757 = 26.5 37 x 0.243 = 9.00 Weighted Avg : 26.5 + 9.00 = 35.5 amu Multiply % by mass # Add them all together Slide 12

11. Relative Abundance practice 1.Gallium consists of two naturally occurring isotopes with masses of 68.926 amu (60.1% abundant) and 70.925 amu (39.9% abundant). What is the weighted Average atomic mass? 2.Strontium consists of four isotopes with masses of 84 amu (abundance 0.50%), 86 amu (abundance of 9.9%), 87 amu (abundance of 7.0%), and 88 amu (abundance of 82.6%). Calculate the atomic mass of strontium. 3.Rubidium has two common isotopes, Rb-85 (abundance of 72.5%) and Rb-87 (abundance of 27.8%). What is the average atomic mass? 4.Argon has 3 naturally occurring isotopes: argon-36, argon-38 and argon-40. Based on argon’s reported atomic mass, which isotope do you think is the most abundant in nature? Explain. Slide 13