1. Unit 6: Atomic Structure and nuclear chemistry

2. Guiding Questions How did the concept of the atom develop?
Which chemist do you think contributed the most in discovering the structure of the atom?

3. Predict the Timeline of Atomic Theory
4.1
Early Theories of Matter
Predict the Timeline of Atomic Theory

4. Democritus Lived around 400 B.C. Came up with the concept of the atom
4.1
Early Theories of Matter
Democritus
Lived around 400 B.C.
Came up with the concept of the atom
Believed atoms could not be divided (broken down)

5. Dalton’s Atomic Theory
4.1
Early Theories of Matter
Dalton’s Atomic Theory
John Dalton ( ), an English schoolteacher and chemist, proposed his atomic theory of matter in 1803.
Although his theory has been modified slightly to accommodate new discoveries, Dalton’s theory was so insightful that it has remained essentially intact up to the present time.

6. Dalton’s Billiard Ball Model
4.1
Early Theories of Matter
Dalton’s Billiard Ball Model
The following statements are the main points of Dalton’s atomic theory.
All matter is made up of atoms.
Atoms are indestructible and cannot be divided into smaller particles. (Atoms are indivisible.)
All atoms of one element are exactly alike, but are different from atoms of other elements.

7. 4.1
Early Theories of Matter
J. J. Thomson
Plum Pudding model (or Chocolate Chip Cookie model)
Discovered the electron (negative charge)
POSITIVE CHARGE
ELECTRONS

8. Ernest Rutherford - 1911 Nuclear Model (atom contains a nucleus)
4.1
Early Theories of Matter
Ernest Rutherford
Nuclear Model (atom contains a nucleus)
Gold foil Experiment
Atoms have:
A nucleus
Protons (positive charge) in nucleus
Mostly open space
Electrons found around the nucleus

9. Niels Bohr - 1913 Planetary Model
4.1
Early Theories of Matter
Niels Bohr
Planetary Model
Electrons (e-) have definite path around the nucleus (orbit)
e- arranged around the nucleus according to energy level
e- with lowest energy level are closest to nucleus

10. Quantum Mechanical Model - 1923
4.1
Early Theories of Matter
Quantum Mechanical Model
Electron Cloud (modern theory)
Calculates the probability of finding the electron within a given space
Electrons, instead of traveling in defined orbits, travel in diffuse clouds around the nucleus

11. Predict the Timeline of Atomic Theory
4.1
Early Theories of Matter
Predict the Timeline of Atomic Theory

12. Stepwise Timeline of Atomic Theory
4.1
Early Theories of Matter
Stepwise Timeline of Atomic Theory
Dalton
1803
Rutherford
1909
Modern Theory
Thomson
1897
Bohr
1913

13. 4.2
Subatomic Particles & the Nuclear Atom
The Electron
J.J. Thomson, discovered that cathode rays are made up of invisible, negatively charged particles referred to as electrons.
Matter is not negatively charged.
If atoms contained negatively charged particles, then they must also contain positively charged particles.

14. 4.2
Subatomic Particles & the Nuclear Atom
Protons
In 1886, scientists discovered that a cathode-ray tube emitted rays not only from the cathode but also from the positively charged anode.
Thomson was able to show that these rays had a positive electrical charge.
Years later, scientists determined that the rays were composed of positively charged subatomic particles called protons.

15. 4.2
Subatomic Particles & the Nuclear Atom
Neutrons
In 1910, Thomson discovered that neon consisted of atoms of two different masses.
Atoms of an element that differ in mass are called isotopes.
Because of isotopes, scientists hypothesized that atoms contained a third type of particle.
Such a particle should have a mass equal to that of a proton but no charge.
The existence of this neutral particle, called a neutron, was confirmed in the early 1930s.

16. Chemistry Jargon
element - a basic substance that can't be simplified (hydrogen, oxygen, gold, etc...)
atom - the smallest amount of an element
molecule - two or more atoms that are chemically joined together (H2, O2, H2O, etc...)
compound - a molecule that contains more than one element (H2O, C6H12O6, etc...)

17. Subatomic Particles Name Symbol Relative Mass Charge Position Proton
4.2
Subatomic Particles & the Nuclear Atom
Subatomic Particles
Name
Symbol
Relative Mass
Charge
Position
Proton
1H or p+
1 amu 1
Nucleus
Electron e-
0 amu -1
Outside
Neutron 1n
amu – atomic mass unit; based on carbon-12
1 amu = 1/12 mass of C-12 = mass H
Impractical to use actual mass of subatomic particles

18. Atomic Numbers
The atomic number of an element determines the identity of the atom
Equals the # of protons
Equals the # of electrons (for a neutral atom)
Lithium has 3 protons and 3 neutrons.

19. Average Atomic Mass
The number at the bottom of each box is the average atomic mass of that element.
This number is the weighted average mass of all the naturally occurring isotopes of that element, in units of amu.

20. Isotopes Atoms of the same element with different numbers of neutrons
Protons and electrons are the same.
Oxygen-15
Oxygen-16
Oxygen-17
Protons 8
Neutrons 7 9
Electrons

21. Masses Mass number = # Protons + # Neutrons
# Neutrons = Mass # - Atomic #
Isotopes of an element have different mass numbers because they have different numbers of neutrons.
Li-7 has a mass # of 7, because it has 4 neutrons.
Li-6 has a mass # of 6, because it has 3 neutrons.

22. X Isotope Notation Symbol – Mass #
Element Symbol with mass number and atomic number
Can also be the element name dash mass number
Mass Number X
Atomic Number
Symbol – Mass #

23. Isotope Abundance
Recall, the atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element.
It reflects both the mass and the relative abundance of the isotopes as they occur in nature.
Carbon has isotopes of Carbon-12, Carbon-13, Carbon-14. Which isotope is most abundant? Refer to your Periodic Table.

24. Isotope Abundance
Boron has two isotopes: boron-10 and boron-11. Which is more abundant, given that the atomic mass of boron is amu?
There are three isotopes of silicon; they have mass numbers of 28, 29, and 30. The atomic mass of silicon is amu. Comment on the relative abundance of these three isotopes.
Given relative abundance of isotopes, we can calculate the average atomic mass.

25. Calculating Atomic Mass
Copper exists as a mixture of two isotopes.
The lighter isotope (Cu-63), with 29 protons and 34 neutrons, makes up 69.17% of copper atoms.
The heavier isotope (Cu-65), with 29 protons and 36 neutrons, constitutes the remaining 30.83% of copper atoms.

26. Calculating Atomic Mass
The atomic mass of Cu-63 is amu, and the atomic mass of Cu-65 is amu.
Use the data above to compute the atomic mass of copper.
First, calculate the contribution of each isotope to the average atomic mass, being sure to convert each percent to a fractional abundance.

27. Calculating Atomic Mass
Second, add the mass contributions of each isotope together for the average atomic mass.

28. Calculating Atomic Mass
Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of amu (50.69%) and amu (49.31%).

29. Calculating Atomic Mass
Isotope
Protons
Neutrons
% Abundance 1 82
122
1.37 2
124
26.25 3
125
20.82 4
126
51.55
Above is the data for four isotopes of an unknown atom. Calculate the approximate atomic mass. What is the identity of the atom? Write the isotope symbols for each isotope.

30. Calculating Atomic Mass
Element X has two natural isotopes. The isotope with a mass of amu (10X) has a relative abundance of 19.91%. The isotope with a mass of amu (11X) has a relative abundance of 80.09%. Calculate the atomic mass of this element. What is the identity of element X?

31. Calculating Atomic Mass
The element copper has naturally occurring isotopes with mass numbers of 63 and 65. The relative abundance and atomic masses are 69.2% for mass = amu, and 30.8% for mass = amu. Calculate the average atomic mass of copper.

32. Nuclear Radiation
Nuclear Radiation
Nuclear chemistry is the study of the structure of atomic nuclei and the changes they undergo.
In 1895, Wilhelm Roentgen (1845–1923) found that invisible rays were emitted when electrons bombarded certain materials.
The emitted rays were discovered because they caused photographic plates to darken. Roentgen named these invisible high-energy emissions X rays.

33. The Discovery of Radioactivity
Nuclear Radiation
The Discovery of Radioactivity
Marie Curie (1867–1934) and her husband Pierre (1859–1906) took mineral samples (called pitchblende) and isolated the components emitting the rays.
Marie Curie named the process by which materials give off such rays radioactivity; the rays and particles emitted by a radioactive source are called radiation.

34. Nuclear Radiation
Types of Radiation
Unstable nuclei emit radiation to attain more stable configurations in a process called radioactive decay.

35. Nuclear Radiation
Penetrating Ability

36. Types of Radiation

37. Nuclear Radiation
Types of Radiation
Because of their mass and charge, alpha particles are relatively slow-moving compared with other types of radiation.
A beta particle is a very-fast moving electron that has been emitted from a neutron of an unstable nucleus.
Gamma rays are high-energy (short wavelength) electromagnetic radiation.
Gamma rays almost always accompany alpha and beta radiation.
Because gamma rays have no effect on mass number or atomic number, it is customary to omit them from nuclear equations.

38. Radioactive Decay Rates
Transmutation
Radioactive Decay Rates
Radioactive decay rates are measured in half-lives.
A half-life is the time required for one-half of a radioisotope’s nuclei to decay into its products.

39. Radioactive Decay
Nuclear Stability
Nuclear force is a force that acts on subatomic particles overcoming the repulsion between protons.
For atoms with low atomic numbers (< 20), stable nuclei have neutron-to-proton ratios of 1:1.
As atomic number increases, more and more neutrons are needed to balance the electrostatic repulsion forces.

40. Radioactive Decay
Nuclear Stability
All nuclei with more than 83 protons are radioactive and decay spontaneously.
These very heavy nuclei often decay by emitting alpha particles.

41. Fission & Fusion of Atomic Nuclei
Nuclear Fission
The splitting of a nucleus into fragments is called nuclear fission.
Nuclear fission releases a large amount of energy.
One fission reaction can lead to more fission reactions, a process called a chain reaction.
The mass needed to sustain a chain reaction is called critical mass.

42. Fission & Fusion of Atomic Nuclei
Nuclear Reactors
Nuclear power plants use the process of nuclear fission to produce heat in nuclear reactors.
The heat generates steam, which drives turbines that produce electricity.
Fissionable uranium (IV) oxide (UO2) is used as fuel.
Cadmium and boron are used to keep the fission process under control.

43. Fission & Fusion of Atomic Nuclei
Nuclear Fusion
The combining of atomic nuclei is called nuclear fusion.
For example, nuclear fusion occurs within the Sun, where hydrogen atoms fuse to form helium atoms.
Fusion reactions can release very large amounts of energy but require extremely high temperatures. For this reason, they are also called thermonuclear reactions.

44. Fusion Attempts Fusion Excessive heat can not be contained.
Fission & Fusion of Atomic Nuclei
Fusion Attempts
Fusion
Excessive heat can not be contained.
Attempts at “cold” fusion have FAILED.
“Hot” fusion is difficult to contain.