1. Atomic Structure

2. I. Atoms The atom is the basic unit of matter.

3. Smallest particle of an element that retains the characteristics of that element. Smallest particle of an element that retains the characteristics of that element.

4. Three Basic Subatomic Particles: proton neutron electron proton neutron electron Charge: +1 0 -1 Mass Number: 1 amu 1 amu 0 amu Location: nucleus nucleus electron cloud cloud

5. What is an amu? atomic mass unit: used to express very small masses used to express very small masses based on the mass of 1/12 of a carbon-12 atom based on the mass of 1/12 of a carbon-12 atom 1 amu = 1.66054 x 10 -24 grams 1 amu = 1.66054 x 10 -24 grams

6. The atom is divided into two separate areas: Nucleus Nucleus Electron Cloud Electron Cloud

7. 1. Nucleus Found in center of an atom Found in center of an atom Small and dense Small and dense Contains protons and Contains protons and neutrons, so is positively neutrons, so is positively charged charged

8. Contains most of the mass of Contains most of the mass of an atom an atom The number of protons is ALWAYS the same. The number of protons is ALWAYS the same. The number of neutrons can vary. The number of neutrons can vary.

9. 2. Electron Cloud Located outside the nucleus Located outside the nucleus NOT dense NOT dense Contains electrons, so negatively charged Contains electrons, so negatively charged Mass of electrons is negligible Mass of electrons is negligible Electrons determine reactivity Electrons determine reactivity

10. Periodic Table Each box on the periodic table provides three important pieces of information: Each box on the periodic table provides three important pieces of information: 1. Symbol 2. Atomic Number 3. Atomic Mass (Mass #)

11. Atomic number: Atomic number: The number of protons in the nucleus The number of protons identifies the element Examples: Examples: Carbon: atomic # is 6, so 6 p + Carbon: atomic # is 6, so 6 p + Sulfur: atomic # is 16, so 16 p + Sulfur: atomic # is 16, so 16 p +

12. Because atoms are electrically neutral, each atom contains the same number of protons as electrons. Examples: Helium: 2 p +, so 2 e- Helium: 2 p +, so 2 e- Magnesium: 12 p +, so 12 e- Magnesium: 12 p +, so 12 e- Iron: 26 p +, so 26 e- Iron: 26 p +, so 26 e-

13. Atomic mass (mass #): Atomic mass (mass #): The number of protons plus the number of neutrons in an atom Examples: Examples: Carbon-12 = 6 p + + 6 n 0 Carbon-12 = 6 p + + 6 n 0 Mercury-201 = 80 p + + 121 n 0 Mercury-201 = 80 p + + 121 n 0

14. MassNumberAtomicNumber X

15. Example: How many neutrons in C? 12 6 C Mass Number (# p + + # n o ) Atomic Number (# p + ) 12 -6 = 6 neutrons

16. Do Atomic Structure Homework Do Atomic Structure Homework

17. II. Ions An ion is a charged particle An ion is a charged particle # of p + does not equal the # of e - # of p + does not equal the # of e - The ONLY way to form an ion is to change # of ELECTRONS The ONLY way to form an ion is to change # of ELECTRONS

18. Positive ions (cations): # e - < # p + Positive ions (cations): # e - < # p + Negative ions (anions): # e - > # p + Negative ions (anions): # e - > # p +

19. Example 1: Lithium Neutral atom: 3 p + and 3 e - Neutral atom: 3 p + and 3 e - Loses one e - to form Li +1 ion Loses one e - to form Li +1 ion Ion: 3 p + and 2 e - Ion: 3 p + and 2 e -

20. Example 2: Oxygen Neutral atom: 8 p + and 8 e - Neutral atom: 8 p + and 8 e - Gains two e - to form O -2 ion Gains two e - to form O -2 ion Ion: 8 p + and 10 e - Ion: 8 p + and 10 e -

21. Remember: Electrons control behavior in ordinary chemical reactions!

22. III. Isotopes Atoms of the same element that differ from each other by the number of neutrons they contain Atoms of the same element that differ from each other by the number of neutrons they contain

23. Most elements have at least two naturally occurring isotopes. Most elements have at least two naturally occurring isotopes.

24. Isotopes of the same element: 1. Have the same atomic number 2. Have different mass numbers 3. Have similar behavior

25. When you are working with a specific isotope, the mass number on the periodic table will not be exactly correct When you are working with a specific isotope, the mass number on the periodic table will not be exactly correct You need a way to write the symbol so that it shows the atomic number and the correct atomic mass of that particular isotope. You need a way to write the symbol so that it shows the atomic number and the correct atomic mass of that particular isotope.

26. 1.To write an Isotope (Nuclear) Symbol MassNumberAtomicNumber X

27. Examples: How many neutrons in C-12? 12 6 C Mass Number (# p + + # n o ) Atomic Number (# p + ) 12-6 = 6 neutrons

28. How many neutrons in C-13? 13 6 C Mass Number (# p + + # n o ) Atomic Number (# p + ) 13-6 = 7 neutrons

29. 2. To name an isotope using a Hyphen Notation Symbol X – atomic mass of specific isotope C-12 is carbon with a mass of 12 amu C-12 is carbon with a mass of 12 amu (6 protons + 6 neutrons) (6 protons + 6 neutrons) C-13 is carbon with a mass of 13 amu C-13 is carbon with a mass of 13 amu (6 protons + 7 neutrons) (6 protons + 7 neutrons)

30. and do Homework and do Homework

31. IV. Relative Abundance and Average Atomic Mass The naturally occurring isotopes of each element are present in specific amounts known as relative abundance The naturally occurring isotopes of each element are present in specific amounts known as relative abundance

32. The higher the The higher the % Abundance, the more common the isotope.

33. Average Atomic Mass- Mass # on periodic table Mass # on periodic table Weighted average of the mass and abundance of the naturally occurring isotopes of an element Weighted average of the mass and abundance of the naturally occurring isotopes of an element

34. To Calculate Average Atomic Mass: 1. Multiply mass by % abundance for each isotope 2. Add all answers from step #1 to determine average atomic mass

35. Hint: when calculated, average atomic mass should equal the Mass Number from the Periodic Table.

36. Example 1: Calculate the average atomic mass for the element Boron. Boron – 10 19.8% Boron – 10 19.8% Boron – 1180.2% Boron – 1180.2%

37. Step 1: Change % to decimals. Step 1: Change % to decimals. 19.8% =.198 80.2% =.802 Step 2: Multiply mass of each isotope by its abundance. Step 2: Multiply mass of each isotope by its abundance. 10 x.198 = x 11 x.802 = y Step 3: Add. x + y = Step 3: Add. x + y =

38. To make life easier, we will always round atomic masses to the nearest tenth. 10.8 amu Answer: 10 x.198 = 1.98 11 x.802 = 8.82 + 10.802 10.802

39. Example 2: Two isotopes of copper occur in nature. Two isotopes of copper occur in nature. 69.17% of copper atoms have a mass of 62.94 amu 69.17% of copper atoms have a mass of 62.94 amu 30.83% have a mass of 64.93amu. 30.83% have a mass of 64.93amu. What is the average atomic mass? What is the average atomic mass?

40. Answer 62.94 x.6917 = 43.54 64.93 x.3083 = 20.02 + 63.56 amu 63.56 amu 63.6 amu

41. Example 3: Uranium has 3 isotopes with the following masses and % abundance. Calculate its average atomic mass. Uranium has 3 isotopes with the following masses and % abundance. Calculate its average atomic mass. 234 U 0.0058% 234 U 0.0058% 92 92 235 U0.71% 92 92 238 U99.23% 92

42. (234 x.000058) + (235 x.0071) + (238 x.9923) = 237.84 (234 x.000058) + (235 x.0071) + (238 x.9923) = 237.84 237.8 amu

43. Stop and do homework