1. Section 4.3

2. Standard: 1a, 11c, 3b Mastering Concept: 112(47-52) Terms: 98 Practice Problems: 99(11-13), 101(14), 104(15-17) homework: Cornell Notes: 4.3 Section Assessment: 104(18-20) Mastering Problems: 113 (59-68)

3. Explain why atomic masses are not whole numbers.
Section 4-3
Section 4.3 How Atoms Differ
Explain the role of atomic number in determining the identity of an atom.
Define an isotope.
Explain why atomic masses are not whole numbers.
Calculate the number of electrons, protons, and neutrons in an atom given its mass number and atomic number.

4. The number of protons and the mass number define the type of atom.
Section 4-3
Section 4.3 How Atoms Differ (cont.)
periodic table: a chart that organizes all known elements into a grid of horizontal rows (periods) and vertical columns (groups or families) arranged by increasing atomic number
atomic number
isotopes
mass number
atomic mass unit (amu)
atomic mass
The number of protons and the mass number define the type of atom.

5. Each element contains a unique positive charge in their nucleus.
Section 4-3
Atomic Number
Each element contains a unique positive charge in their nucleus.
The number of protons in the nucleus of an atom identifies the element and is known as the element’s atomic number.

7. Isotopes and Mass Number
Section 4-3
All atoms of a particular element have the same number of protons and electrons but the number of neutrons in the nucleus can differ.

8. Atoms with the same number of protons but different numbers of neutrons are called isotopes.

9. The relative abundance of each isotope is usually constant.
Isotopes and Mass Number (cont.)
Section 4-3
The relative abundance of each isotope is usually constant.
Isotopes containing more neutrons have a greater mass.
Isotopes have the same chemical behavior.
The mass number is the sum of the protons and neutrons in the nucleus.

10. Section 4-3
Isotopes and Mass Number (cont.)

11. Section 4-3
Mass of Atoms
One atomic mass unit (amu) is defined as 1/12th the mass of a carbon-12 atom.
One amu is nearly, but not exactly, equal to one proton and one neutron.

13. Mass of Atoms (cont.)
Section 4-3
The _____________ of an element is the weighted average mass of the isotopes of that element.
atomic mass

15. Mastering Concept: 112(47-52)
47. Does the existence of isotopes contradict part of Dalton’s original atomic theory? Explain. (4.3) Yes; not all atoms of an element are identical in mass.

16. Mastering Concept: 112(47-52)
48. How do isotopes of a given element differ? How are they similar? (4.3) Different number of neutrons, masses; same: chemical properties, number of protons and electrons

17. Mastering Concept: 112(47-52)
49. How is the mass number related to the number of protons and neutrons an atom has? (4.3) Mass number = number of p+ + number of n0

18. Mastering Concept: 112(47-52)
50. What do the superscript and subscript in the notation 40 K represent? (4.3) 19 The superscript represents the mass number (40) and the subscript represents the atomic number (19).

19. Mastering Concept: 112(47-52)
51. Explain how to determine the number of neutrons an atom contains if you know its mass number and its atomic number. (4.3) Number of n0 = mass number – atomic number

20. Mastering Concept: 112(47-52)
52. Define the atomic mass unit. What were the benefits of developing the atomic mass unit as a standard unit of mass? (4.3) amu = 1/12 of the mass of a C-12 atom; Scientists defined the atomic mass unit as a relative standard that was closer in size to atomic and subatomic masses.

21. Calculating Average Atomic Massx
Sum of the products = average atomic mass
Example:
The isotopes of element “Bob” are found below:
Bob-18, 25%
Bob-19, 60%
Bob-20, 15%
What is the average atomic mass of naturally occurring Bob?  x  x 
1amu = 1.66x10-27kg

22. Isotopes, Ions, and Allotropes (Oh my)
A review…
Isotopes, Ions, and Allotropes (Oh my)
Isotopes
atoms of the same element with different number of neutrons.
Ions
atoms of the same element with different number of electrons.
Ions are easy to create; adding or removing electrons can be done with electric current.
Allotropes
forms of the same element, bonded in different structures.
Diamond and pencil graphite are allotropes. They are both pure carbon, but in different structures.

23. Practice Problems: 99(11-13)
11. How many protons and electrons are in each of the following atoms?
Boron c. Platinum
Radon d. Magnesium
12. An atom of an element contains 66 electrons. What element is it?
13. An atom of an element contains 14 protons. What element is it?

24. Practice Problems: 99(11-13)
11. How many protons and electrons are in each of the following atoms?
Boron
a. boron, 5

25. Practice Problems: 99(11-13)
11. How many protons and electrons are in each of the following atoms? b. Radon Radon, 86

26. Practice Problems: 99(11-13)
11. How many protons and electrons are in each of the following atoms? c. Platinum platinum, 78 78 Pt
Platinum

27. Practice Problems: 99(11-13)
11. How many protons and electrons are in each of the following atoms? d. Magnesium magnesium, 12

28. Practice Problems: 99(11-13)
12. An atom of an element contains 66 electrons. What element is it? 12. dysprosium
word=dysprosium

29. Practice Problems: 99(11-13)
13. An atom of an element contains 14 protons. What element is it? 13. silicon

30. Isotope Composition Data
Practice Problems: 101 (14)
14. Determine the number of protons, electrons, and neutrons, Name each isotope, and write its symbol.
Isotope Composition Data
Element
Atomic Number
Mass Number
P/ e n
isotope
symbol
b. Calcium 20 46 26
Calcium-46
c. Oxygen 8 17 9
Oxygen-17
d. Iron 57 31
Iron-57
e. Zinc 30 64 34
Zinc-64
f. Mercury 80
204
124
Mercury-204

31. Practice Problems: 104 (15-17)
15. Boron has two naturally occurring isotopes: boron-10 (abundance = 19.8%, mass = amu), boron-11 (abundance = 80.2%, mass = amu). Calculate the atomic mass of boron.
For isotope boron-10 :
Mass: amu
abundance 19.8% = 0.198
Mass contribution = (mass)(percent abundance)
= amu (0.198) =

32. Practice Problems: 104 (15-17)
15. Boron has two naturally occurring isotopes: boron-10 (abundance = 19.8%, mass = amu), boron-11 (abundance = 80.2%, mass = amu). Calculate the atomic mass of boron.
For isotope boron-11 :
Mass: amu
abundance 80.2% = 0.802
Mass contribution = (mass)(percent abundance)
= amu (0.802) = amu
Sum the mass contributions to find the atomic mass.
amu amu = amu

33. Alternative Isotope % Abundance Mass Mass contribution boron-10 19.8%
0.198
10.013
boron-11
80.2%
0.802
11.009
8.829

34. Practice Problems: 104 (15-17)
16. Helium has two naturally occurring isotopes, helium-3 and helium-4. the atomic mass of helium is amu. Which isotope is more abundant in nature? Explain. 16. Helium-4 is more abundant in nature because the atomic mass of naturally occurring helium is closer to the mass of helium-4 (~4 amu) than to the mass of helium-3 (~3 amu).

35. Practice Problems: 104 (15-17)
Calculate the atomic mass of magnesium. The three magnesium isotopes have atomic masses and relative abundances of amu (78.99%), amu (10.00%), and amu (11.01%)
amu (0.7899) =
amu (0.1000) =
amu (0.1101) =
24.31 amu