Bell Work Compare and contrast the Bohr model and the Quantum Mechanical Model.
Sticky Note Matching Game • Each one of you has been given an index card with a name, facts, or diagram. Now you have to find the name, facts and diagrams that match (not every name has a diagram). • The Catch: you can only use ‘yes’ or ‘no’ questions to try to find the matches. • When you find all the matches, read over the cards and prepare a 1-3 minute explanation.
Distinguishing Among Atoms • Structure • Atomic/ Mass Number • Isotopes • Atomic mass
The atomic number is the number of protons in an atom. This number is unique for all elements and is used to identify each element. Atomsare electrically neutral therefore: # PROTONS = # ELECTRONS.
Average Mass number = # protons + # neutrons Atomic number = # protons = # electrons Mass number – Atomic number = # neutrons
10 protons 10 neutrons 10 electrons 10 protons 11 neutrons 10 electrons 10 protons 12 neutrons 10 electrons
Isotopesare atoms of the same element that differ in the number of neutrons.Isotopesof the same element have the same chemical properties, because they have the same number of protons and electrons. Isotopes are identified by mass number. Neutrons affect mass, so, isotopes with more neutrons are heavier
Sample Problem Three isotopes of oxygen are oxygen-16, oxygen-17, and oxygen-18. Write the symbol for each, including the atomic number and mass number.
Atomic Mass • The actual masses of individual atoms are very very very very small (arsenic = 1.244 x 10 -23 )…so to measure the relative mass of atoms, a reference isotope is used. • The reference isotope is carbon-12. It has a mass of exactly 12 atomic mass units. • Atomic mass units (amu) is the unit used.
Atomic Masses Atomic mass is the weighted average of all the naturally occurring isotopes of that element. Carbon = 12.011amu
To calculate relative atomic mass, multiply the massnumber of each isotope by its percent abundance changed to a decimal and total. (Mass #)(isotope’s relative abundance) + (Mass #)(another isotopes Rel. abundance) = Relative atomic mass of the element
Example 1: Neon has 2 isotopes, Ne-20 with an abundance of 90% and Ne-22 with an abundance of 10%. Calculate the average atomic mass of neon. 20 x 0.9= 18 22 x 0.1= 2.2 18 + 2.2 =20.2 amu Why is it closer to 20 than 22? 20 is more abundant
Example 2: Carbon occurs in nature as a mixture of atoms of which 98.89% have a mass of 12.00 u and 1.11% have a mass of 13.00335 u. Calculate the atomic mass of carbon. 12.00 x 0.9889= 11.87 13.00335 x 0.0111= 0.144 11.87+ 0.144 = 12.014 Why is it closer to 12? 12.00 is more abundant
Sample Problem Carbon has two stable isotopes: carbon-12 which has a natural abundance of 98.89 % and carbon-13 which has a natural abundance of 1.11 %. The mass of carbon-12 is 12.000 amu and the mass of carbon-13 is 13.003 amu. What is the atomic mass?
Sample Problem The element antimony (Sb) has naturally occurring isotopes with mass numbers of 121 and 123. The relative abundance and atomic masses are 57.12 % for mass = 120.90 amu, and 47.29% for mass = 122.90 amu. Calculate the atomic mass of antimony.
Sample Problems In book, page 119 numbers 24 and 25
Meet 115, the newest element on the Periodic Table, Ununpentium • Man-made element first by Russian scientist 10 years ago. It was replicated by chemists in Sweden at the Lund University. Therefore making it officially a new element. • No official name has been given yet, so scientists are calling it Ununpentium, based on the Latin and Greek words for its atomic number 115. • http://www.rsc.org/periodic-table/element/115/ununpentium