1. Chapter 4 “Atomic Structure”

2. Section 4.1 Defining the Atom
OBJECTIVES:
Describe Democritus’s ideas about atoms.
Explain Dalton’s atomic theory.
Identify what instrument is used to observe individual atoms.

3. Section 4.1 Defining the Atom
Democritus
First to suggest the existence of atoms (from the Greek word “atomos”)
He believed that atoms were indivisible and indestructible.
-Greek philosopher
-not based on the scientific method – but just philosophy

4. Dalton’s Atomic Theory
All elements are composed of tiny indivisible particles called atoms.
John Dalton
(1766 – 1844)
2) Atoms of the same element are identical.
--Atoms of any one element are different from those of any other element.

5. Dalton’s Atomic Theory
Atoms of different elements combine in whole-number ratios to form compounds.
4) In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.

6. Sizing up the Atom 100,000,000 atoms = 1 cm
1,000,000 atoms = width of hair
Can be observed with scanning tunneling (electron) microscopes

7. Section 4.2 Structure of the Nuclear Atom
OBJECTIVES:
Identify three types of subatomic particles.
Describe the structure of atoms, according to the Rutherford atomic model.

8. Section 4.2 Structure of the Nuclear Atom
Atoms are divisible into three subatomic particles:
Electrons
Protons
Neutrons

9. Discovery of the Electron
J.J. Thomson used a cathode ray tube to discover the negatively charged electron
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

10. Mass of the Electron
Mass of the electron is
9.11 x g
The oil drop apparatus
Robert Millikan determined the mass of the electron: 1/1840 the mass of a hydrogen atom

11. Conclusions from the Study of the Electron:
Atoms have no charge, so there must be positive particles to balance the negative charge of the electrons
Electrons have so little mass that other particles must account for most of the mass

12. Conclusions from the Study of the Electron:
Eugen Goldstein observed positive proton
Mass of 1 (or 1840 times that of an electron)
James Chadwick confirmed the neutral neutron
Mass nearly equal to a proton
Protons have a positive charge. Neutrons have no charge. Mass is same.

13. Subatomic Particles Particle Charge Mass (g) Location Electron (e-) -1
9.11 x 10-28
Electron cloud
Proton (p+) +1
1.67 x 10-24
Nucleus
Neutron
(no)
Don’t have to know numbers, just charge and how masses compare, and location.

14. Thomson’s Atomic Model
J. J. Thomson
Thomson - plum pudding model.
Electrons were like plums embedded in a positively charged pudding.

15. Ernest Rutherford’s Gold Foil Experiment - 1911
Alpha particles (helium nuclei) fired at a thin gold foil.
Particles that hit on the detecting screen are recorded

16. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected.
Conclusions:
The nucleus is small, dense, and, positively charged

17. The Rutherford Atomic Model
Based on his experimental evidence:
Atom is mostly empty space.
All the positive charge, and almost all the mass is in the center at the nucleus.

18. The Rutherford Atomic Model
Nucleus is made of protons and neutrons
Electrons surround the nucleus.
Called the “nuclear model”

19. Section 4.3 Distinguishing Among Atoms
OBJECTIVES:
Explain what makes elements and isotopes different from each other.
Calculate the number of neutrons in an atom.
Calculate the atomic mass of an element.
Explain why chemists use the periodic table.

20. Atomic Number
Atoms are composed of identical protons, neutrons, and electrons
How then are atoms of one element different from another element?

21. Atomic Number
Elements are different because they contain different numbers of PROTONS
Atomic number - number of protons in the nucleus (smaller #)
# protons = # electrons
Atomic Number:
# p+ : # e- : 35 35 35 53 53 53

22. Atomic Number
Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.
Element
# of protons
Atomic # (Z)
Carbon (C)
Phosphorus (P)
Gold (Au) 6 6 15 15 79 79

23. Mass Number Mass # = p+ + n0
Mass number is the number of protons and neutrons in the nucleus of an isotope:
Mass # = p+ + n0
Atomic Number: Mass Number:
# p+ : # e- : #n0 :
# p+ : # e- : #n0 : 35
79.9 35 35 45 53
127 53 53 74

24. Mass Number Practice Atom p+ n0 e- Mass # 16 8 8 8 33 41 33 74 15 16
Oxygen 16 8 8 8
Arsenic 33 41 33 74
Phosphorus 15 16 15 31

25. Complete Symbols
Contain the symbol of the element, the mass number and the atomic number.
Atomic
number X
Superscript →
Subscript →
Mass
number *

26. Na Symbols 11 11 12 23 11 11 23 Find each of these: number of protons
number of neutrons
number of electrons
Atomic number
Mass Number 11 11 12 Na 23 11 11 23 *

27. Symbols
If an element has an atomic number of 34 and a mass number of 78, what is the:
number of protons =
number of neutrons =
number of electrons =
complete symbol 34 43 34 34 X 78 *

28. If an element has 91 protons and 140 neutrons what is the
Symbols
If an element has 91 protons and 140 neutrons what is the
Atomic number =
Mass number =
number of electrons =
complete symbol 91
131 91 *

29. If an element has 78 electrons and 117 neutrons what is the
Symbols
If an element has 78 electrons and 117 neutrons what is the
Atomic number
Mass number
number of protons
complete symbol *

30. Isotopes Dalton was wrong about all elements of
the same type being identical
Atoms of the same element can have different numbers of neutrons.
Thus, different mass numbers.
These are called isotopes.
Atomic #:
Mass #:
# p+:
#n0:
Atomic #:
Mass #:
# p+:
#n0:
Atomic #:
Mass #:
# p+:
#n0: *

31. Isotopes
Isotopes are atoms of the same element with different masses, due to varying numbers of neutrons.

32. We can also put the mass number after the name of the element:
Naming Isotopes
We can also put the mass number after the name of the element:
carbon Mass:
carbon-14 Mass:
uranium-235 Mass: 12 14
235 *

33. What’s the only thing that changes? # of neutrons
Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Nucleus
Hydrogen–1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2
What’s the only thing that changes? # of neutrons

34. Atomic Mass How heavy is an atom of oxygen?
Depends - there are different masses of oxygen atoms.
We want the average atomic mass.
Based on abundance (percentage) of each variety of that element in nature. *

35. Measuring Atomic Mass
Measure atomic mass with the Atomic Mass Unit (amu)
Defined as one-twelfth the mass of a carbon-12 atom.
Each isotope has its own atomic mass, thus we determine the average from percent abundance.
We don’t use grams for this mass because the numbers would be too small.
Carbon-12 chosen because of its isotope purity. – 6 p+, 6 n0 almost all mass and high % C12. *

36. To calculate the average:
Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results.
Expressed as amu.
C-12 = 12 amu. *

37. Composition of the nucleus
Atomic Masses
Atomic mass is the average of all the naturally occurring isotopes of that element.
Isotope
Symbol
Composition of the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
7 neutrons
1.11%
Carbon-14
14C
8 neutrons
<0.01%
Carbon = *

39. Atomic Mass Example (11.0)(.802) 10.8 amu (10.0)(.198) B-10 = 19.8%
At. Mass = =
(10.0)(.198)
(11.0)(.802)
10.8 amu

40. The Periodic Table: A Preview
Periodic table - arrangement of elements in which the elements are separated into groups based on a set of repeating properties.
Allows easy comparison of the properties of different elements

41. The Periodic Table: A Preview
Period - horizontal row (there are 7 of them)
Group - vertical column
Also called a family
Elements in a group have similar chemical and physical properties
Identified with number and “A” or “B”

42. Draw an arrow and label a period and a group.