1. Unit 2: Symbols Say WHAT?

2. The Three Subatomic Particles
After various molecular models of the atoms had been tested, it was determined that three subatomic particles made up the atom
Protons
Neutrons
Electrons

3. Protons Protons are found in the nucleus
Protons have an actual charge of +1.6 x C and a relative charge of +1
The actual mass of a proton is 1.67 x g
The relative mass of a proton is 1 atomic mass unit (amu)
The symbol is p+

4. Neutrons Neutrons are found in the nucleus
Neutrons have an actual charge of 0 C and a relative charge of 0
The actual mass of a neutron is 1.67 x g
The relative mass of a neutron is 1 atomic mass unit (amu)
The symbol is n0

5. Electrons Electrons are found outside the nucleus
Electrons have an actual charge of   -1.6 x C and a relative charge of -1
 The actual mass of an electron is   9.11 x g
The relative mass of a electron is 1/2000 atomic mass unit (amu)
The symbol is e-

6. Atomic Number
Each element has a certain number of protons in its nucleus
The number of protons in the nucleus is called the atomic number
Each element has its own atomic number because each element has its own, unique number of protons
Note: That an element only changes if the atomic number changes. Which means that if the number of protons changes it changes elements

7. On Your Own Which element has a) 87 protons b) 35 protons
c) 50 protons
d) 92 protons
e) 8 protons
f) 19 protons

8. On Your Own:
Determine the number of protons in the following atoms as well as each atom’s identity
a) 6 electrons b) 14 electrons
c) 72 electrons d) 55 electrons

9. Mass Number = Protons + Neutrons
Mass Number: the number of protons and neutrons in an atom added together
Mathematically
Mass Number = Protons + Neutrons

10. Example
What is the mass number of an atom with 16 protons and 16 neutrons?

11. Determine the Mass Number for the following atoms
On Your Own
Determine the Mass Number for the following atoms
a) 17 protons and 18 neutrons
b) 11 protons and 12 neutrons
c) 1 proton and NO neutrons
d) 3 protons and 4 neutrons

12. Nuclide Symbols
Mass Numbers (This is represent by an A) are written in the upper left preceding the chemical symbol
Atomic Number (This is symbolized by a letter Z) is written directly under the mass number

13. Example: Write the correct nuclide symbol for an element with
51 p+ and 71 n0

14. Practice Draw the following Nuclide Symbols for the following: Sr OsZn I Cs K

15. Practice Write the following elements in proper Nuclide Symbols: Pt FrAt K

16. Warm-Up 2/10/2015
-Write the correct Nuclide Symbol for the following elements:
*19 p+, 20 n
*82 e-, 125n
*Mass # 238, neutrons= 146

17. Isotopes
Atoms of the same elements have the same number of protons
HOWEVER there may be different numbers of neutrons and different mass numbers
When an element’s atom has different numbers of neutrons, it is said to have isotopes

18. Hydrogen’s Three Isotopes
Hydrogen has the following isotopes:
Protium-a hydrogen atom with one proton and NO neutrons
Deuterium-a hydrogen atom with one proton and only one neutron
Tritium-a hydrogen atom with one proton and two neutrons

19. The atomic mass unit is based on the relation of standard carbon-12
Average Atomic Mass
Atomic mass is the mass of an atom expressed in atomic mass units or amu
The atomic mass unit is based on the relation of standard carbon-12
Begin 2nd here on 2/13/2012

20. Average Atomic Mass Continued
Carbon-12 has a mass of amu
Example:
If an atoms weighs half as much as carbon-12, its atomic mass will be amu
Begin 4th
begin 5th, begin 7th

21. Example: If an atom weighs four times as much as carbon-12, it will have a mass of 48.000 00 amu

22. What is Atomic Mass?
The atomic mass that is reported in the periodic table is a weighted average based on the relative abundance of each element

23. Relative abundance refers to how common the isotope occurs in nature
Percent Abundance which refers to how many of each isotope are in every hundred

24. To Determine Avg. Atomic Mass
1) First convert relative abundance (%) to decimal equivalent
2) Multiply mass (in amu) by decimal equivalent
3) Add the numbers together
4) The sum (in amu) is the average atomic mass

25. For example, an element has two naturally occurring isotopes
For example, an element has two naturally occurring isotopes. One isotope has a relative abundance of % and a mass of amu. A second isotope has a relative abundance of 80.08% and a mass of amu. Calculate the atomic mass

26. Example
For example, an element has two naturally occurring isotopes. One isotope has a relative abundance of % and a mass of 7.02 amu. A second isotope has a relative abundance of 7.42% and a mass of amu. Calculate the atomic mass

27. Additional Example
Calculate the average atomic masses for the following: Isotope: Rel. Abund. Rel. Mass hydrogen % hydrogen % 2.014
Begin 3rd , 4th
1.007 amu
1.0 amu (s.f.)

28. Practice
Titanium has five common isotopes: If the abundance of Ti- 46 is 8.0%, Ti-47 is 7.8 %, Ti-48 is 73.4 %, Ti-49 is 5.5% and Ti- 50 is 5.3 %. What is the average atomic mass of titanium?

29. Isotope Rel. Abund. Actual Mass O-16 99.762 15.995 O-17 0.038 16.999
Determine Avg. Atomic Mass for oxygen:
Isotope Rel. Abund Actual Mass O O O

30. Review 2/11/2015 What is an Isotope?
What are hydrogen’s three isotopes names?
What is the difference between average atomic mass and mass number?

31. Review Practice
Rubidium is a soft, silvery-white metal that has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium?

32. Practice
Begin 3rd & 4th

33. Practice
Element
Atomic #
Mass # p+ n0 e- Ag
110 19 40
127 53

34. Practice

35. Finding % Abundance
Silver( Atomic Weight ) has two naturally- occurring isotopes with weights of and What is the percentage abundance of the lighter isotope?

36. Practice
Oxygen comes in two stable isotopes, Oxygen- 16 and Oxygen The molar mass of Oxygen- 16 is amu. The molar mass of Oxygen-18 is amu. Determine the percent abundance of each isotope.

37. Finding Percent Abundance
The element indium exists naturally as two isotopes. 113In has a mass of amu, and 115In has a mass of amu. The average atomic mass of indium is amu. Calculate the percent relative abundance of the two isotopes of indium.
Candium Activity Afterwards

38. Review
Antimony has two naturally occurring isotopes. The mass of antimony-121 is amu and the mass of antimony-123 is amu. Using the average mass from the periodic table, find the abundance of each isotope.

39. Practice
Europium has two stable isotopes: Eu-151 with a mass of amu and Eu-153 with a mass of If elemental Europium is found to have an average mass of amu on earth, calculate the percent of each of the two isotopes