1. Atoms – The Basics

2. Inside the Atom  Atoms are made up of smaller particles  These particles are found in different regions of the atom

3. Atomic Number & Mass Number  Atomic number = # p + For atoms atomic number also = number of e-  Mass number = # p + + # n 0 Mass number is a whole number Number on the periodic table is average atomic mass (not a whole number)  How can you solve for numbers of neutrons?

4. How Atoms Differ: Ions Different elements have different numbers of p + Ions have charges  the number of e- ≠ number of p+  charge = # p + - # e - Ions are the SAME element, but DIFFER because they have either a positive or negative charge. If an atom GAINS electrons, its overall charge becomes more negative. If it LOSES electrons, its charge becomes more positive

5. Isotopes Atoms with the same number of p + but different number of n 0 are isotopes.  Mass number of isotopes of the same element changes  Isotopes are the SAME element, but DIFFER because they have different masses.

6. EX:

7. Symbols of Ions and Isotopes  For ions chemists use the following notation: Or just:

8.  For isotopes chemists use the following notation:  Name of element – mass number  Examples: carbon-12 carbon-14 uranium-236

9. Determine the Numbers of p +, n 0, and e - for the following symbols given: calcium - 46 nickel - 60

10. Complete the following table ProtonsNeutronsElectrons Na + Bromine- 84 O 2- with an atomic mass of 13amu

11. Average Atomic Mass  The masses of p + and n o are the same and they are actually tiny (1.67 x 10-24 g). We round it to 1 amu.  The mass of an e- is even smaller 1/1840 of a proton or neutron so we say it is approximately 0 amu.

12.  We can do this because scientists measure the mass of an atom relative to a standard mass and that is carbon-12.  Scientists agreed that carbon-12 has a mass of EXACTLY 12 atomic mass units (amu).

13. So, why do the elements on the Periodic Table have masses with decimals??? They are Average Atomic Masses

14. Average Atomic Masses  This is the weighted average mass of all the isotopes of an element.  Examples of weighted averages: Semester and yearly grades Taxes Budget

15. Calculating the Weighted Average Atomic Mass 1. Multiply the mass of each isotope by the % of the isotope 2. Add the products

16. Example  Calculate the atomic mass of magnesium. The three magnesium isotopes have atomic masses and relative abundances of 23.985 amu (78.99%), 24.986 amu (10.00%) and 25.982 amu (11.01%).  24.31 amu

17.  You can also estimate which isotope is the most abundant.  Ex: Fluorine has an atomic mass of 18.99840 amu. It has several naturally occuring isotopes Fluorine-14, Fluorine- 16, Fluorine-19, Fluorine-21, Fluorine-22.  Guess which one is the most abundant?  Helium has two naturally occurring isotopes: helium-3 and helium-4. The atomic mass of helium is 4.003amu. Which isotope is more abundant in nature and why?