1. Development of Atomic Theory Part I

2. Early Roots Much of science (including chemistry) and mathematics can trace their investigative roots to Arabic countries in the Middle Ages and their philosophical roots to ancient Greece. It was in the Middle and Dark Ages that alchemy as a quasi-science came into practice. As many know, alchemy was the quest to turn common elements (such as lead) into gold-a precious metal. As we hopefully know by now, a transmutation such as this would not occur in a chemical reaction. By the late 1700s most chemists and scientists had accepted a definition of an element as a substance that cannot be further broken down by ordinary chemical means. They also knew that elements combine to form new compounds that have different physical and chemical properties from the elements that formed them. There was still much to know, such as did elements always combine in the same ratio when forming the same compound.

3. First Laws Once scientific methods had developed to the point where quantitative data was being recorded (particularly mass), our understanding of elements and their chemical ratio behavior was furthered. This occurred in the late 1700s and early 1800s. One of the most basic laws that serves as the foundation of chemistry, the Law of Conservation of Mass, came about during this time. The Law of Conservation of Mass states that mass is neither created nor destroyed during ordinary chemical reactions or physical changes.

4. First Laws (cont) Soon thereafter, the Law of Definite Proportions was proposed and supported. This law states that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. In other words, 18.02 g of water (H 2 O) always contains 2.02 g of hydrogen and 16.00 g of oxygen. Continuing the example, 36.04 g of water always contains 4.04 g of hydrogen and 32.00 g of oxygen. These ratios are constant (definite). This is also known as the Law of Constant Composition.

5. First Laws (cont) The next chemical law that became accepted was the Law of Multiple Proportions. It states that if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first is always a ratio of small whole numbers.

6. First Laws (cont) Let’s explain with methane (CH 4 ) and ethane (C 2 H 6 ) If I have 16.05 g of methane then I know from the Law of Definite Proportions that I have 12.01 g of C and 4.04 g of H. When I compare that to a sample that has the same number of ethane particles as the methane did, then it would mass 30.08 g of C 2 H 6. It would contain 24.02 g of C and 6.06 g of H. The Law of Multiple Proportions allows me to compare the C from the methane to the C from the ethane. 12.01 g: 24.02 g or a 1:2 ratio. The H ratio (again methane to ethane) would be 4.04 g: 6.06g or a 2:3 ratio, as they are both reduce when rounded. Essentially the Law of Multiple Proportions says that there is some whole number ratio that relates the same element in two different compounds. We did not innately understand the significance of this fact, but it was one of those odd things that scientists knew would reveal new ideas.

7. Dalton During the turn of the century (1800s-1900s), John Dalton, an English school teacher, proposed the idea of the atom as scientific fact and not esoteric philosophy, and offered these supporting precepts.

8. Dalton Dalton’s Atomic Theory 1. All elements are composed of tiny, indivisible particles called atoms. 2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element. 3. Atoms of different elements can combine with one another in simple whole number ratios to form compounds. 4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element cannot be changed into atoms of another element by a chemical reaction.

9. Dalton If you study Dalton’s Atomic Theory, you should first notice how the Laws of Conservation of Mass, Definite Proportions, and Multiple Proportions all came together and were used to explain the existence and role of atoms of elements. Some of his ideas have had to be further refined and updated.

10. Dalton As science moved forward (and we will get into that later in the lesson), precepts 1 and 2 had to changed. While the atom is still the smallest whole unit of an element, we now know that it is itself comprised of even smaller particles (protons, neutrons, electrons, etc.). Also, we also know that there can be different types of atoms of the same element because they can differ by their neutron count and still be the same atom, because their proton and electron count is still the same. The isotope is the term we have given to atoms of the same element that differ by their neutron count. But Dalton’s Atomic Theory was the first that gave concrete evidential support to the existence of the atom.

11. Thomson Joseph John Thomson, researching in the 1890s, realized that the accepted model of an indivisible atom did not take electrons and protons into account.

12. Thomson His work with the cathode tube ray showed that there were regions of positive and negative charges in the atom. Thomson is credited with discovering the electron (1897).

13. Thomson His atomic model is often called the “Plum Pudding Model.” (Chocolate-chip cookie dough is a good analogy).

14. Thomson The negatively charged electrons (the chocolate chips) are stuck into a lump of positively charged protons (the dough). He was awarded the Nobel Prize in Physics in 1906.

15. Thomson Limitations: His model still does not include neutrons. This model explained some of the electrical properties of the atom, but said nothing about the number of protons or their arrangement in the atom, or how easily atoms are stripped of their electrons to form ions.

16. Rutherford Ernest Rutherford discovered the proton and the nucleus in the early 1900s Proposed the nuclear atom in which electrons surround a dense nucleus comprised of protons and neutrons. He thought of the rest of the atom as empty space. Other work showed that the nuclei of most atoms are composed of protons and neutrons.

17. Rutherford His experiment is known as the “gold foil experiment.” http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf

18. Rutherford This work was performed over 1906-1907. He was awarded the Nobel Prize in Chemistry in 1908.

19. Rutherford This model, though, does not account for the fact that opposite particles attract each other. If this model is correct, the electrons in the empty space around the nucleus in the atom should be drawn into the nucleus due to the attraction to the protons.

20. Chadwick James Chadwick worked for Rutherford in Cambridge, England. In 1932, he finds a subatomic particle that has no charge while studying uranium-235 decay. The particle is eventually named the neutron as it had no charge.

21. Chadwick He won the Nobel Prize in Physics in 1935.

22. Visual Models: Dalton, Thomson, Rutherford