1. Chapter 5 “Atomic Structure”
Chemistry
Olympic High School
Mr. Daniel
Credits: Stephen L. Cotton
Charles Page High School

2. Section 5.1 Defining the Atom
OBJECTIVES:
Describe the history of the development of ideas about atoms.
Explain Dalton’s atomic theory.
Describe the size of the atom.

3. Section 5.1 Defining the Atom
The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos” – meaning “unable to be cut”)

4. Antoine Laurent Lavoisier ( ) The Father of Modern Chemistry: discovered oxygen & hydrogen, developed modern thermodynamics, invented the first periodic table

5. Dalton’s Atomic Theory (experiment based)
All elements are composed of tiny indivisible particles called atoms
Atoms of the same element are identical. Atoms of any one element are different from those of any other element.
John Dalton
(1766 – 1844)
Atoms of different elements combine in simple whole- number ratios to form chemical compounds
In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.

6. Section 5.2 Structure of the Nuclear Atom
OBJECTIVES:
Identify three types of subatomic particles and their properties.
Describe the structure of atoms, according to the Rutherford Model.

7. Section 5.2 Structure of the Nuclear Atom
One change to Dalton’s atomic theory is that atoms are divisible into subatomic particles:
Electrons, protons, and neutrons are examples of these fundamental particles
There are many other types of particles, but we will study these three

8. The electron’s charge-to-mass ratio = 1.76  108 C/g
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron
JJ Thomsonl
The electron’s charge-to-mass ratio = 1.76  108 C/g

9. A Standard Television Tube (First common in the 1950’s)
Cathode Ray Tubes
A Standard Television Tube (First common in the 1950’s)
A CRT from a Radar Scope (WW II)
Standard Computer Monitor (often called a “CRT”)
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

10. Mass of the Electron
simulation
Robert Millikan
Mass of the electron is
9.11 x g
The oil drop apparatus
1916 – Robert Millikan determined the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge

11. Conclusions from the Study of the Electron:
Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.
Atoms are neutral, so there must also be positive particles in the atom to balance the negative charge of the electrons
Electrons have very little mass, therefore atoms must contain other particles that account for most of the mass

12. Thomson’s Atomic Model
J. J. Thomson
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

13. Other Particle Discoveries:
Eugen Goldstein first observed evidence of what is now called the “proton” in particles with a positive charge, and a mass of 1840 times that of an electron. It’s existance was later confirmed by Ernest Rutherford in 1919.
1932 – James Chadwick confirmed the existence of the “neutron” – a particle with no charge, but a mass nearly equal to a proton

14. The problem:
In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?
Target #1
The Answer:

15. The Answer
Target #2

16. Ernest Rutherford’s Gold Foil Experiment - 1911
Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil
Particle that hit on the detecting screen (film) are recorded

17. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The nucleus is small
The nucleus is dense
The nucleus is positively charged

18. The Rutherford Atomic Model
Based on his experimental evidence, Rutherford’s Nuclear Model was developed and stated:
The atom is mostly empty space
The electrons are distributed around the nucleus, and occupy most of the volume
Because of the exceptionally high mass of the nucleus, it must contain particles in addition to protons (neutrons were discovered later)
All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “nucleus”
Ernest Rutherford video

19. Subatomic Particles Particle Charge Mass (g) Location
Crash Course Chemistry - History of Chemical Concepts
Particle
Charge
Mass (g)
Location
Electron x Electron
(e-) cloud
Proton x Nucleus
(p+) (1 atomic mass unit)
Neutron x Nucleus
(no) (1 atomic mass unit)

20. Sizing up the Atom
Elements can be subdivided into smaller and smaller pieces until there are only single atoms; The smallest particles of an element that still have the properties of that element.
100,000,000 copper atoms in a single file, would be approximately 1 cm long
Despite their incredibly small size, individual atoms have recently become observable with scanning tunneling microscopes
The Beginning: Xe on Ni
Atomic Kanji “original child”
Forming a Image: Fe atoms on a Cu surface
Image: Fe atoms on a Cu surface
Image: Fe atoms on a Cu surface
An Artistic View of Ni
Carbon Monoxide Man

21. Section 5.3 Distinguishing Among Atoms
OBJECTIVES:
Explain what makes isotopes different from each other.
Calculate the number of protons, neutrons and electrons in an atom using atomic number and mass
Calculate the average atomic mass and atomic number.

22. Atomic Number Element # of protons # of electrons Atomic # (Z) Carbon
Atoms are composed of protons, neutrons, and electrons
How then are atoms of one element different from another element?
Elements are different because they contain different numbers of PROTONS
Atomic number (Z) : the number of protons in the nucleus of each atom of that element.
An atom is always neutral, therefore:
# protons in an atom = # electrons
Element
# of protons
# of electrons
Atomic # (Z)
Carbon 6
Phosphorus 15
Iodine 53
Gold 79

23. Mass Number Nuclide p+ n0 e- Mass #
Mass number is the number of protons and neutrons in the nucleus of an atom:
Mass # = p+ + n0
Nuclide p+ n0 e-
Mass #
Oxygen - 10 - 33 42
- 31 15 18 8 8 18
Arsenic 75 33 75
Phosphorus 16 15 31

24. X Complete Symbols Mass U
Complete Symbols contain the symbol of the element, the mass number and the atomic number.
Examples: C U
Mass
number
Superscript → 12 X 6 14 6
235
Atomic
number
Subscript → 92

25. Symbols continued…
We can also put the mass number after the name of the element:
carbon-12
carbon-14
uranium-235

26. Br 80 35 Symbols Find each of these: number of protons
number of neutrons
number of electrons
Atomic number
Mass Number
35 b) 45 c) 35 d) 35 e) 80

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 b) 44 c) 34 d) Se 78 34

28. Symbols If an element has 91 protons and 140 neutrons what is the
Atomic number
Mass number
number of electrons
complete symbol
91 b) 231 c) 91 d) Pa
231 91

29. Symbols If an element has 78 electrons and 117 neutrons what is the
Atomic number
Mass number
number of protons
complete symbol
78 b) 195 c) 78 d) Pt
195 78

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, they have different mass numbers.
These are called isotopes.

31. Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons.
Protons
Electrons
Neutrons
Nucleus
Hydrogen–1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2

32. Isotopes
Elements occur in nature as mixtures of isotopes.

33. Atomic Mass How heavy is an atom of oxygen?
It depends, because there are different kinds of oxygen atoms.
We generally refer to the average atomic mass.
Average atomic mass is based on the abundance (percentage) of each isotope of an element as it is found in nature.
It is the number (red) that we find on the periodic table

34. Measuring Atomic Mass
Instead of grams, the mass unit we use for atoms is the Atomic Mass Unit (amu or µ)
It is defined as one-twelfth the mass of a carbon-12 atom.
Carbon-12 is chosen because of its isotope purity.
Each isotope has its own atomic mass, thus we determine the average mass from the percent abundance of each isotope.

35. 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 = µ

36. To calculate the average atomic mass:
Multiply the mass number of each isotope by its abundance (expressed as a decimal), then add the results.
The mass of the isotope is expressed in atomic mass units (amu or µ)
Isotope
Symbol
% in nature
Carbon-12
12C
98.89%
Carbon C 1.11%
Carbon C <0.01%
12.01 µ
(12.00 µ) (.9889) +
(13.00 µ) (.0111) =

37. Atomic Mass Calculation
Element X has two natural isotopes. The isotope with a mass of amu has a relative abundance of 19.91%. The isotope with a mass of amu has a relative abundance of 80.09%. Calculate the atomic mass of this element.
( µ) (.1991) + ( µ) (.8009) =
10.81 µ
Boron

38. Section 5.4 Organizing the Elements
OBJECTIVES:
Explain how elements are organized in a periodic table.
Compare early and modern periodic tables.
Identify three broad classes of elements.
Distinguish different areas of the periodic table.

39. Section 5.4 Organizing the Elements
A few elements, such as gold and copper, have been known for thousands of years.
Yet, only about 13 (of 90) had been identified by the year 1700.
As more were discovered, chemists realized they needed a way to organize the elements.

40. Section 5.4 Organizing the Elements
Chemists used the properties of elements to sort them into groups.
In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties
    ( ) ÷ 2 = 88
Ca 40 Sr 88 Ba 137
Li Na 23 K 39
Cl Br I

41. John Newlands ( )
Law of Octaves: noted that after interval of eight elements, similar physical/chemical properties reappeared. 
Newlands was the first to formulate the concept of periodicity (repeating patterns) in the properties of the chemical elements.

42. Mendeleev’s Periodic Table
Lothar Meyer
Mendeleev’s Periodic Table
By the mid-1800s, about 70 elements were known to exist
Dmitri Mendeleev – A Russian chemist arranged the elements in order of increasing atomic mass
It was the beginning of the modern “Periodic Table”

43. Co and Ni; Ar and K; Te and I
His Completed Work…
His Original Table…
Co and Ni; Ar and K; Te and I
Mendeleev left blanks for elements he predicted that existed such as Germanium
But, there were problems: some elements did not fit with their groups
When discovered, the elements generally matched his predictions

44. A better arrangement…
In 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number
His basic arrangement is still used today
The symbol, atomic number & mass are the basic items included in the periodic table

45. Squares in the Periodic Table
The periodic table displays the symbols and names of the elements, along with atomic number and average atomic mass
Black symbol = solid
Red symbol = Blue symbol = liquid
25º C

47. There are other possible arrangements. How about a:
Spiral Periodic Table

48. The Periodic Table:
Your “best friend” is an arrangement of elements in which they are separated into groups based on a set of repeating properties.
The periodic table allows you to easily compare the properties of one element to another

49. The Periodic Law Vertical column = group (or family)
Similar physical & chemical properties
Identified by number & letter
Horizontal rows = periods
Repeated properties in each row
There are 7 periods
When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

50. Areas of the periodic table
Three classes of elements are:
1) Metals 2) Nonmetals 3) Metalloids

51. Metals
Metals: electrical conductors, have luster, are ductile and malleable (in pink below)

52. Nonmetals: generally brittle and dull, are poor conductors of heat and electricity (in blue below)
Some nonmetals…
are gases (O, N, Cl);
are brittle solids (S, C);
one is dark red liquid (Br)

53. Notice the heavy, stair-step line…
Metalloids: border the line
Their properties are intermediate between metals and nonmetals (in green below)

54. Elements in the 1A-7A groups are called the Representative Elements
Have very predictable properties and patterns of behavior 1A 8A 2A 3A 4A 5A 6A 7A

55. Groups of elements Group IA – Alkali Metals Alkali Metals
Are the most reactive metals
Form a “base” when reacting with water
Alkali Metals

56. Groups of Elements Alkaline Earth Metals
Group 2A – Alkaline Earth Metals
Are the second most reactive metals
Also form bases with water, but do not dissolve well; hence “earth metals”
Alkaline Earth Metals

57. Groups of elements The Halogen Group Group 7A – Halogens
The most reactive non-metal group
Means “salt-forming”
The Halogen Group

58. Groups of elements Noble Gases Group 8A- Noble Gases
Previously called “inert gases” because they rarely take part in a reaction
Noble gases have a completely full electron arrangement
Noble Gases

59. The “B” groups are called the Transition ElementsLa Ac
Lanthanide Series
Actinide Series
The “Inner Transition Metals” actually belong here
Crash Course Chemistry - The Periodic Table

60. End of Chapter 5