Introduction to Atoms: Development of the Atomic Theory

1. Introduction to Atoms: Development of the Atomic Theory Book K, Ch 4.1 A. Boyle

2. Vocabulary • Atom- the smallest unit of an element that maintains the properties of that element • Electron-subatomic particle with negative charge

3. Vocabulary cont. • Nucleus-in physical science, an atom’s central region, which is made up of protons and neutrons • Electron cloud- a region around the nucleus of an atom where electrons are likely to be found

4. Objectives • 1. Describe some of the experiments that led to the current atomic theory. • 2. Compare the different models of the atom. • 3. Explain how the atomic theory has changed as scientists have discovered new information about the atom.

5. A. The Beginning of Atomic Theory • 440 BCE, Greek philosopher, Democritus, named the atom. “Atomos” means “not able to be divided.”

6. A. The Beginning of Atomic Theory • 1. From Aristotle to Modern Science • Aristotle, another Greek philosopher, disagreed with the early hypothesis. He was incorrect, and Democritus was correct in thinking that the atom is the tiniest particle of matter.

7. B. Dalton’s Atomic Theory Based on Experiments • John Dalton, the British chemist, based his theory on how elements combine, published his atomic theory: • All substances are made of atoms. Atoms are small particles that can’t be created, destroyed, nor divided. • Atoms of the same element are exactly alike, and atoms of different elements are different. • Atoms join with other atoms to make new substances.

8. C. Thompson’s Discovery of Electrons • In 1897, British scientist, J.J. Thompson, using a cathode-ray experiment, discovered there were subatomic particles inside of an atom. He discovered electrons. • He thought the electrons were distributed throughout the atom like plums in plum pudding (our present day chocolate chip ice cream).

9. D. Rutherford’s Atomic “Shooting Gallery” • In 1909, Ernest Rutherford experimented with a thin sheet of gold foil that glowed when hit by positive particles. He proved that atoms are mostly empty space with a dense, positive nucleus.

10. E. Where are the Elements? • 1. Far From the Nucleus • 2. Bohr’s Electron Levels-In 1913, Niels Bohr proposed that electrons are located in levels at certain distances from the nucleus.

11. E. Where are the Elements? 3. Modern Atomic Theory- Electron Cloud Theory Schrodinger & Heisenberg discovered electrons travel in regions of the electron cloud.

12. The Atom Book K, Ch 4.2 A. Boyle

13. Vocabulary • Proton- a subatomic particle that has a positive charge and that is found in the nucleus of an atom • Atomic Mass Unit- a unit of mass that describes the mass of an atom or molecule • Neutron- a subatomic particle that has no charge and that is found in the nucleus of an atom

14. Vocabulary cont. • Atomic number- the number of protons in the nucleus of an atom; the atomic number is the same for all atoms of an element • Isotope- an atom that has the same number of protons (or the same atomic number) as other atoms of the same element do but that has a different number of neutrons (and thus a different atomic mass)

15. Objectives

16. A. How Small is the Atom? • Atoms are extremely small. There are about 50,000 atoms in the thickness of a sheet of aluminum.

17. B. Of What is an Atom Made? • 1. The Nucleus- made up of Protons (+) & Neutrons (+/-) I Proton = 1 atomic mass unit (amu) Neutrons have a very slightly larger mass than Protons, but 1 Neutron is basically equal to 1 amu also. The nucleus is very dense, made up of all the P & N of an atom. • 2. Outside the Nucleus • 3.

18. B. Of What is an Atom Made? • 2. Outside the Nucleus- Electrons (-) are found outside of the nucleus within the electron clouds. Electrons are even smaller in mass than P and N. The Mass of 1 Electron is considered 0 amu. Because the charges of P & E are opposite, their charges cancel out each other. Neutrons do not have a charge, therefore atoms with equal numbers of P & E do not have a charge. The loss of an electron makes the charged atom, ion, positive. Gaining an electron makes the ion negative.

19. C. How Do Atoms of Different Atoms Differ? • Each element has a unique atomic number. This is the number of protons in that element. The number of neutrons may differ.

20. D. Isotopes • 1. Properties of Isotopes • Unstable atoms have a nucleus that will change over time. This means it is radioactive. These atoms will fall apart, giving off energy as this happens. • Stable atoms and isotopes of the same element have the same properties • 2. Telling Isotopes Apart- Isotopes are identified by mass number (P & N) • 3. Naming Isotopes- Element-mass number • 4. Calculating the Mass of an Element

21. D. Isotopes • 3. Naming Isotopes- Element-mass number • Ex: Carbon-12 • 4. Calculating the Mass of an Element

22. D. Isotopes • 4. Calculating the Mass of an Element- • Atomic Mass of an element is the average of naturally occurring atoms & isotopes of that element. • Ex: Chlorine-35 makes up 76% of Chlorine in nature, Chlorine-37 makes up the other 24%. • (35 x 0.76)= 26.60 • (37 x 0.24)= +8.88 • 35.48 amu

23. E. Forces in Atoms • Four basic forces are at work at all times. • Gravitational Force- depends on masses of 2 objects and distance between them • Electromagnetic Force-object with same charges repel, different charges are attracted • Strong Force- works against electromagnetic, keeps like charges together • Weak Force-plays a role in unstable atoms, where a neutron may change into a proton or electron