1. Chemistry Chapter 4 Notes #2 (sec. 3)

2. Periodic Table  All atoms are composed of smaller particles including equal numbers of protons (+) and electrons (-)  Atomic Number = # of protons Distinguishes one element from another Distinguishes one element from another  Elements are in sequence by atomic number  The table also reflects the arrangement of electrons....

3. Periodic Table  Periods 7 Horizontal Rows (across) 7 Horizontal Rows (across) By increasing atomic number By increasing atomic number  Groups/Families Columns (up and down) Columns (up and down) Have elements with similar properties Have elements with similar properties

4. Periodic Table  Chemical Name  Chemical symbol – letters to rep. element  Atomic Number – # of protons in nucleus  Atomic Mass – The mass of an atom measured in amu’s (atomic mass units)

5. How atoms differ….  For now, electrons and protons will be equal  # of neutrons can be different Atoms with the same number of protons and different numbers of neutrons are called isotopes Atoms with the same number of protons and different numbers of neutrons are called isotopes  Mass Number = #protons + #neutrons Isotopes #p + #n 0 #e - Hydrogen-1101 Hydrogen-2111 Hydrogen-3121

6. Isotopes  Abundance of isotopes is constant in nature. Potassium (K)….K-39, K-40, and K-41 Potassium (K)….K-39, K-40, and K-41 93.25% have 20 neutrons, Mass # = 39 93.25% have 20 neutrons, Mass # = 39 6.7502% have 22 neutrons, Mass # = 41 6.7502% have 22 neutrons, Mass # = 41.0117% have 21 neutrons, Mass # = 40.0117% have 21 neutrons, Mass # = 40 # neutrons = mass number – atomic # # neutrons = mass number – atomic #

7. Mass of Atoms  Mass of an atom is in the nucleus  Protons and Neutrons = 1.67 x 10 -24 g  Electrons = about 1/1840 of a p + or n 0  Chemists method for measuring mass of atoms – easier numbers Atomic standard = carbon 12 atom (6 p and 6n) = 12 atomic mass units Atomic standard = carbon 12 atom (6 p and 6n) = 12 atomic mass units One amu (atomic mass unit) = 1/12 carbon 12 atom One amu (atomic mass unit) = 1/12 carbon 12 atom

8. Mass of Atoms  Mass is close to the mass of protons, but not exact. Ex. Silicon 30 = 29.974 amu Ex. Silicon 30 = 29.974 amu  Masses of subatomic particles in amu’s Electron = 0.000549 Electron = 0.000549 Neutron = 1.00865 Neutron = 1.00865 Proton = 1.004276 Proton = 1.004276  Atomic Mass of an element is the weighted average mass of the isotopes of that element

9. Mass of Atoms #15 page 104  (0.198 x 10.013) + (.802 x 11.009)  = 10.81 amu  Notice it is closer to 11 amu than 10. Why?  THE END! IsotopeAbundance Mass (amu) Boron 10 19.8 % 10.013 Boron 11 80.2%11.009