1. Chapter 4 Chemistry 1K Cypress Creek High School
Unit 2: Atomic Theory
Chapter 4
Chemistry 1K
Cypress Creek High School

2. Table of Contents Chapter 4: The Structure of the Atom
4.1: Early Theories of Matter
4.2: Subatomic Particles & the Nuclear Atom
4.3: How Atoms Differ
4.4: Unstable Nuclei & Radioactive Decay

3. 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

4. Development of the Modern Atomic Theory
4.1
Early Theories of Matter
Development of the Modern Atomic Theory
In 1782, a French chemist, Antoine Lavoisier ( ), made measurements of chemical change in a sealed container.
He observed that the mass of reactants in the container before a chemical reaction was equal to the mass of the products after the reaction.
Lavoisier concluded that when a chemical reaction occurs, mass is neither created nor destroyed but only changed.
Lavoisier’s conclusion became known as the law of conservation of mass.

5. Development of the Modern Atomic Theory
4.1
Early Theories of Matter
Development of the Modern Atomic Theory
In 1799, another French chemist, Joseph Proust, observed that the composition of water is always 11% hydrogen and 89% oxygen by mass.
Regardless of the source of the water, it always contains these same percentages of hydrogen and oxygen.
Proust studied many other compounds and observed that the elements that composed the compounds were always in a certain proportion by mass. This principle is now referred to as the law of definite proportions.

6. Dalton’s Atomic Theory
4.1
Early Theories of Matter
Dalton’s Atomic Theory
John Dalton ( ), an English schoolteacher and chemist, studied the results of experiments by Lavoisier, Proust, and many other scientists.
Dalton 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.

7. Dalton’s Billiard Ball Model
4.1
Early Theories of Matter
Dalton’s Billiard Ball Model
Called the father of Atomic Theory
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.

8. 4.1
Early Theories of Matter
J. J. Thomson
Discovered the electron (1st subatomic particle) through experiments with cathode ray tube
Plum Pudding model (or Chocolate Chip Cookie model)
“Pudding” or “Cookie” is the positive charge and most of the mass of the atom
“Plums” or “Chocolate Chips” are the scattered electrons
POSITIVE CHARGE
ELECTRONS

9. 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 somewhere around the nucleus

10. 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

11. 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

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. TAKS Review
The graph shows the results of a study testing chemical pesticides on a pest species common to cotton plants. Different chemical pesticides were used in five different areas. According to these results, which of the following is the most effective chemical for controlling this pest species?
a. R
b. S
c. T
d. V

14. 4.2
Subatomic Particles & the Nuclear Atom
The Electron
Because of Dalton’s atomic theory, most scientists in the 1800s believed that the atom was like a tiny solid ball that could not be broken up into parts.
In 1897, a British physicist, J.J. Thomson, discovered that this solid-ball model was not accurate.
Thomson’s experiments used a vacuum tube.
At each end of the tube is a metal piece called an electrode.

15. 4.2
Subatomic Particles & the Nuclear Atom
Cathode-Ray Tube
When the electrodes are charged, rays travel in the tube from the negative electrode, which is the cathode, to the positive electrode, the anode.
Thomson found that the rays bent toward a positively charged plate and away from a negatively charged plate.
He knew that objects with like charges repel each other, and objects with unlike charges attract each other.
Click box to view movie clip.

16. 4.2
Subatomic Particles & the Nuclear Atom
Cathode-Ray Tube
Thomson concluded that cathode rays are made up of invisible, negatively charged particles referred to as electrons.
These electrons had to come from the matter (atoms) of the negative electrode.
Matter is not negatively charged, so atoms can’t be negatively charged either.
If atoms contained extremely light, negatively charged particles, then they must also contain positively charged particles—probably with a much greater mass than electrons.

17. 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.
These rays travel in a direction opposite to that of cathode rays.
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.

18. 4.2
Subatomic Particles & the Nuclear Atom
Protons
At this point, it seemed that atoms were made up of equal numbers of electrons and protons.
However, in 1910, Thomson discovered that neon consisted of atoms of two different masses.
Atoms of an element that are chemically alike but differ in mass are called isotopes of the element.

19. 4.2
Subatomic Particles & the Nuclear Atom
Neutrons
Because of the discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained these differences in mass.
Calculations showed that such a particle should have a mass equal to that of a proton but no electrical charge.
The existence of this neutral particle, called a neutron, was confirmed in the early 1930s.

20. 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

21. Rutherford’s Gold Foil Experiment
4.2
Subatomic Particles & the Nuclear Atom
Rutherford’s Gold Foil Experiment
In 1909, a team of scientists led by Ernest Rutherford in England carried out the first of several important experiments that revealed an arrangement far different from the cookie-dough model of the atom.
The experimenters set up a lead-shielded box containing radioactive polonium, which emitted a beam of positively charged subatomic particles through a small hole.

22. Rutherford’s Gold Foil Experiment
4.2
Subatomic Particles & the Nuclear Atom
Rutherford’s Gold Foil Experiment
Today, we know that the particles of the beam consisted of clusters containing two protons and two neutrons and are called alpha particles.
The sheet of gold foil was surrounded by a screen coated with zinc sulfide, which glows when struck by the positively charged particles of the beam.

23. The Nuclear Model of the Atom
4.2
Subatomic Particles & the Nuclear Atom
The Nuclear Model of the Atom
To explain the results of the experiment, Rutherford’s team proposed a new model of the atom.
Because most of the particles passed through the foil, they concluded that the atom is nearly all empty space.
Because so few particles were deflected, they proposed that the atom has a small, dense, positively charged central core, called a nucleus.
The new model of the atom as pictured by Rutherford’s group in 1911 is shown below.

24. TAKS Review
Trade winds blow from east to west across the Pacific Ocean. The winds move surface waters westward across the ocean. This causes deeper, colder water to rise to the surface along the coast. This upwelling of deep ocean waters brings with it nutrients that would otherwise lie near the bottom of the ocean. Which of the following conclusions is supported by the information above?
a. Trade winds help maintain some food chains.
b. Trade winds can reverse parts of the water cycle.
c. Trade winds produce useful minerals in some oceans.
d. Trade winds may be able to reduce greenhouse gases

25. 4.3
How Atoms Differ
Atomic Numbers
The atomic number of an element is the number of protons in the nucleus of an atom of that element.
It is the number of protons that determines the identity of an element, as well as many of its chemical and physical properties.
Because atoms have no overall electrical charge, an atom must have as many electrons as there are protons in its nucleus.
Therefore, the atomic number of an element also tells the number of electrons in a neutral atom of that element.

26. 4.3
How Atoms Differ
Masses
The mass of a neutron is almost the same as the mass of a proton.
The sum of the protons and neutrons in the nucleus is the mass number of that particular atom.
Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number.

27. 4.3
How Atoms Differ
Atomic Mass
In order to have a simpler way of comparing the masses of individual atoms, chemists have devised a different unit of mass called an atomic mass unit, which is given the symbol m.
An atom of the carbon-12 isotope contains six protons and six neutrons and has a mass number of 12.
Chemists have defined the carbon-12 atom as having a mass of 12 atomic mass units.
Therefore, 1 m = 1/12 the mass of a carbon-12 atom.
1 m is approximately the mass of a single proton or neutron.

28. Information in the Periodic Table
4.3
How Atoms Differ
Information in the Periodic Table
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.

29. X Isotope Notation Element- Mass
4.3
How Atoms Differ
Isotope Notation
Element Symbol with mass number and atomic number
Can also be the element name dash mass number
Mass Number X
Element- Mass or
Atomic Number

30. Practice 9 10 19 28 59 150 94 30 65 Symbol # Protons # neutrons
4.3
How Atoms Differ
Practice
Symbol
# Protons
# neutrons
# electrons
Atomic Number
Mass Number 9 10 19 28 59
150 94 30 65

31. Isotopes Atoms of the same element with different numbers of neutrons
4.3
How Atoms Differ
Isotopes
Atoms of the same element with different numbers of neutrons
Think of it as different sized shirts!
6 neutrons Carbon-12
7 neutrons Carbon-13
8 neutrons Carbon-14

32. Calculating Atomic Mass
4.3
How Atoms Differ
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.

33. Calculating Atomic Mass
4.3
How Atoms Differ
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.

34. Calculating Atomic Mass
4.3
How Atoms Differ
Calculating Atomic Mass
The average atomic mass of the element is the sum of the mass contributions of each isotope.

35. 4.3
How Atoms Differ
Percent Abundance
If you are given information about an elements isotope you can estimate the most abundant isotope!
Example: Carbon-12, Carbon-13, Carbon-14. Look at the atomic mass on the periodic table. Which isotope is the mass closest to?

36. Question 1 Calculate the atomic mass of germanium. Answer 72.59 amu
4.3
How Atoms Differ
Question 1
Calculate the atomic mass of germanium.
Answer
72.59 amu

37. 4.4
Unstable Nuclei & Radioactive Decay
Radioactivity
Reactions which involve a change in an atom’s nucleus are called nuclear reactions.
In the late 1890’s scientists noticed that some substances spontaneously emitted radiation in a process they called radioactivity.
The rays and particles emitted by the radioactive material were called radiation.
Unstable nuclei lose energy by emitting radiation in a spontaneous process called radioactive decay.

38. 4.4
Unstable Nuclei & Radioactive Decay
Types of Radiation
Scientists named the radiation that was deflected toward the negatively charged plates alpha radiation.
This radiation is made up of alpha particles contains 2 protons & 2 neutrons.
Scientists named the radiation that was deflected toward the positively charged plate beta radiation.
This radiation consists of fast moving electrons called beta particles containing an electron.
The third type of radiation is called gamma radiation or gamma rays, high energy radiation that has no mass.

39. TAKS Review
A medical researcher hypothesizes that a newly developed medication can reduce high blood pressure. Which of these would most likely be the dependent variable in a study involving this medication?
a. The number of participants in the study
b. The ages of people treated for high blood pressure with other medications
c. The blood pressure of the participants in the study
d. The number of people treated for high blood pressure with other medications

40. Be Prepared for Unit 2 Test.
End of Unit 2
Be Prepared for Unit 2 Test.