1. The Structure of an Atom
Chapter 4

2. 4.1: Defining the Atom
Essential Question: How did the concept of the atom change from the time of Democritus to the time of John Dalton?

3. 4.1: Defining the Atom
All matter is composed of particles, which are called atoms.
Atoms are the smallest particle of an element that retains its identity in a chemical reaction.

4. 4.1: Defining the Atom
The Greek philosopher Democritus (460 B.C.-370 B.C) was among the first to suggest the existence of atoms.
Democritus reasoned that atoms were indivisible and indestructible.

5. 4.1: Defining the Atom
Democritus’s ideas agreed with later scientific theory.
His ideas did NOT explain chemical behavior.
His ideas lacked experimental support because his approach was not based on the scientific method. (It was based in philosophy and logic)

6. 4.1: Defining the Atom
The modern process of discovery regarding atoms began with John Dalton ( ), an English chemist. By using experimental methods, Dalton transferred Democritus’s ideas on atoms into scientific theory.

7. 4.1: Defining the Atom John Dalton’s theory stated that:
1. All elements are made up of tiny indivisible particles called atoms.
2. Atoms of the same element are identical.
3. Atoms of different elements can physically mix together OR can chemically combine to form compounds in simple whole # ratios.
4. Chemical reactions occur when atoms are separated and rearrange.

8. How small is an atom? Atoms are very small.
The radii of most atoms fall within a range of 5 x 10-11m to 2 x 10-10m
Observable with scanning electron microscopes.
Electron microscopes are capable of higher magnification than light microscopes.

9. Scanning Electron Microscope

10. 4.2: Structure of the Nuclear Atom
Subatomic Particles:
Most of Dalton’s theory is acceptable today. However, one important change is that atoms are now known to be divisible. They can be broken down into even smaller, most fundamental particles, called subatomic particles. There are 3 types:
1. Electrons
2. Protons
3. Neutrons

11. 4.2: Structure of the Nuclear Atom
J. J. Thompson ( )
1897: Discovered the electron.
Electrons are negatively charged subatomic particles.

12. 4.2: Structure of the Nuclear Atom
Thomson performed experiments that involved passing electric current through gases with low pressure. He sealed the gases in the glass tubes fitted at both ends with electrodes. One electrode, the anode, became positively charged. The other electrode, the cathode became negatively charged. The result was a glowing beam, or cathode ray that traveled from the cathode to the anode.
Show video clip here

13. 4.2: Structure of the Nuclear Atom
Protons and Neutrons:
In 1886, Eugene Goldstein observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays.
He called those rays, canal rays and concluded that they were composed on positive particles called protons.
Protons are positively charged subatomic particles.

14. 4.2: Structure of the Nuclear Atom
Protons and Neutrons:
In 1932, James Chadwick confirmed the existence of the neutron. Neutrons are subatomic particles with no charge.

15. 4.2: Structure of the Nuclear Atom
The Atomic Nucleus:
Rutherford’s Gold Foil Experiment:
Physicist Ernest Rutherford established the nuclear theory of the atom with his gold-foil experiment.
When he shot a beam of alpha particles as a sheet of gold foil, a few of the particles were deflected.
He concluded that a tiny, dense nucleus was causing the deflections.

16. Ernest Rutherford
Discovered the nucleus of an atom

17. Matter is mostly empty space

18. 4.2: Structure of the Nuclear Atom
The Atomic Nucleus:
The Rutherford Atomic Model:
Rutherford suggested a new theory of the atom. He proposed that the atom is mostly empty space, thus explaining the lack of deflection of most of the alpha particles.
He concluded that all of the positive charge and almost all of the mas are concentrated in a small region that has enough positive charge to account for the great deflection of some of the alpha particles.

19. 4.2: Structure of the Nuclear Atom
The Atomic Nucleus:
The Rutherford Atomic Model:
The nucleus is the tiny central core of an atom and is composed of protons and neutrons.

20. 4.2: Structure of the Nuclear Atom
The Atomic Nucleus:
The Rutherford Atomic Model:
The Rutherford atomic model is known as the nuclear atom. In the nuclear atom, the protons and neutrons are located in the positively charged nucleus.
The electrons are distributed around the nucleus and occupy almost all of the volume of the atom.
Do Rutherford Scattering Activity Here

21. End Lesson #1

22. 4.3: Distinguishing Among Atoms
Atomic Number and Mass Number:
Atoms are composed of protons, neutrons and electrons.
Protons and neutrons make up the nucleus and electrons surround the nucleus.
Elements are different because they contain different numbers of protons.
An element’s atomic number is the number of protons in the nucleus.

23. } A. Basic Structure + + proton (+) neutron (Ø) Nucleus electron (-)
electron cloud
neutron (Ø) } + +
Nucleus
Nucleus: smallest yet heaviest part of the atom

24. 4.3: Distinguishing Among Atoms
Mass Number:
The total number of protons and neutrons in an atom is called the mass number.
Number of neutrons = Mass Number – Atomic number

25. 4.3: Distinguishing Among Atoms
Key Point: Elements are different because they contain different numbers of protons.
An element’s atomic number is the number of protons in the nucleus of an atom of that element.
Mass number or atomic mass: The total number of protons + neutrons in an atom.

26. Hydrogen 1 H 1.008 Element Name Atomic number Element Symbol
Avg. Atomic Mass

27. End Lesson #2

28. 4.3: Distinguishing Among Atoms
EQ: How do isotopes of an element differ?
1. Isotopes:
Isotopes are: atoms of the same element that have different numbers of neutrons. (which means they also have different mass numbers)
Isotopes are chemically alike because they have the same number of protons.

29. 4.3: Distinguishing Among Atoms
2. Analogy: Isotopes are like Skittles – Skittles have different “flavors” just like atoms of the same element have different “flavors” or isotopes.

30. 4.3: Distinguishing Among Atoms

31. 4.3: Distinguishing Among Atoms
3. Isotopic Notation
Mass Notation
Charge
Lithium - 7
Mass # → 7 + Li
mass #
element name 3
Atomic # →

32. 4.3: Distinguishing Among Atoms
4. Average Atomic Mass is:
Weighted average of all isotopes
The isotope that is the most abundant has the greatest effect on the avg. atomic mass found on the periodic table.
Example: The average atomic mass of copper is amu. Which of copper’s two isotopes is more abundant: Copper-63 or Copper-65?
Why?

33. 4.3: Distinguishing Among Atoms
3. Avg. Atomic mass = ∑ (mass x abundance)

34. 4.3: Distinguishing Among Atoms
Avg. Atomic mass = ∑ (mass x abundance) Example: Boron has 2 naturally occurring isotopes. Boron-10 has a mass of amu and a relative abundance of 19.91%. Boron-11 has a mass of amu and a relative abundance of 80.09%. What is the average atomic mass of Boron?

35. 4.3: Distinguishing Among Atoms
Example: Boron has 2 naturally occurring isotopes. Boron-10 has a mass of amu and a relative abundance of 19.91%. Boron-11 has a mass of amu and a relative abundance of 80.09%. What is the average atomic mass of Boron?
Mass x Abundance = Atomic Mass
B-10 = amu x =
B-11 = amu x = +
Avg. Atomic mass

36. End Lesson #3

37. Robert Millikan (1916)
Discovered the quantity of charge of an electron (1.60 x coulomb)
An electron has one unit of negative charge.
An electron’s mass is 1/1840th mass of one hydrogen atom (actual mass = 9.11 x 10-28)

38. Hydrogen 1 H 1.008 Element Name Element Symbol Avg. Atomic Mass
Atomic number
# of protons
Element Symbol
Avg. Atomic Mass

39. EQ: How are ions and atoms different?
Ion Notes
EQ: How are ions and atoms different?

40. A. Definitions
Ion: When an atom gains or loses an electron, it has an imbalance of charge
Cation (+) = loses e- to become positive
Anion (-) = gains e- to become negative
In an ion, the number of protons do not equal the number of electrons

41. B. Examples
Element
Atomic #
Mass #
Protons
Neutrons
Electrons
O2- F-
Ca2+ 16 8 8 16 8 8 10 8 19 9 19 9 10 10 9 41 20 41 20 21 18 20

42. Unstable Nuclei

43. A. Normal Reactions
atoms rearrange, the elements do not change

44. B. Nuclear Reactions 14 14  C N + 6 7 -1
Radioactive decay- atom breaks apart spontaneously 14 14 C N  + 6 7 -1
*Note: Please fit all reactions in one line

46. B. Nuclear Reactions 9 4 12 1 Be n He C + + 4 6 2
Radioactive bombardment: Particle hits atom & it splits 9 4 12 1 Be n He C + + 4 6 2

48. C. Types of Radioactive Particles
Symbol
Composition
Penetrating Power
Alpha,  He
2 P & 2 N
Low
Beta,  
electron
100 x alpha
gamma, 0
EM waves
Very great 4 2 -1

49. Penetrating Power Alpha Beta Gamma

50. Penetrating Power

51. D. Miscellaneous Notation
1. Positron
2. Neutron e +1 1 n

52. E. Transmutation
1. Fission : a very heavy-mass nucleus splits to form two medium-mass (size) nuclei.

53. E. Transmutation
2. Fusion : two very light-mass nuclei combine to form heavier, more stable nuclei

54. F. Balancing Nuclear Equations
1. mass # & atomic #’s must add up the same on both sides of the equation 31 4 27 1 30 Al + He + H Si
____ 2 13 1 14 15

55. Example #1 14 14 C N ? + 6 7
14 + ? = 14
7 + ? = 6 -1 e -1

56. Example #2
230 4 Th ? He  + + 2 90
? = 230
226 88
? = 90
226 Ra 88

57. Example #3 4 27 30 Al + He Si + ? 2 13 14
= 30 + ? 1 1
= 14 + ? 1 1 p or H 1 1