1. Atomic Structure Atoms: The Building Blocks of Matter

2. Early Atomic Theory 400 BC to the early 1800’s

3. The Ancient Greeks Democritus – the particle theory of matter, nature’s basic particle is the atom from the Greek word atomos which means indivisible Aristotle – succeeded Democritus and did not believe in atoms. He thought all matter was continuous and made of a substance called hyle.

4. Antoine Lavoisier – the Father of Chemistry Lavoisier used the balance and measured everything which lead to the Law of conservation of Mass He named oxygen which means “acid former” and helped determine that air is a mixture not an element

5. John Dalton’s Atomic Theory 1808 All matter is composed of small particles called atoms and empty space All atoms of an element are identical Atoms of different elements are different Atoms combine in simple whole number ratios to form compounds

6. Modern Atomic Theory The late 1800’s to Present

7. J. J. Thompson in 1898 discovered the electron The atom is a mass of positive charge which contains regions of negative charge (electrons) like the plums in plum pudding Plum Pudding Model

9. Rutherford, Geiger, and Marsden – discovered the nucleus The center of an atom is dense and postively charged The positive “nucleus” of the atom is very small compared to the total volume of the atom Rutherford suggested that the electrons surround the nucleus like planets around the sun

10. The Atomic Nucleus Rutherford’s Gold-Foil Experiment 4.2

11. The Atomic Nucleus Alpha particles scatter from the gold foil. 4.2

12. The Atomic Nucleus The Rutherford Atomic Model Rutherford concluded that the atom is mostly empty space. All the positive charge and almost all of the mass are concentrated in a small region called the nucleus. The nucleus is the tiny central core of an atom and is composed of protons and neutrons. 4.2

13. Rutherford-Bohr Model Electron Proton Neutron

14. The Periodic Law In the modern periodic table, elements are arranged in order of increasing atomic number. 6.1

15. Atomic Number (Z) Is the number of protons in the nucleus of an atom Is the whole number in the periodic table Is used the calculate the number of neutrons in an atom by its subtraction from the Mass Number

16. Squares in the Periodic Table 6.2

17. Mass Number Au is the chemical symbol for gold. 4.3 The total number of protons and neutrons in a single atom of an element 79 Protons 79 Electrons 118 Neutrons

18. for Sample Problem 4.1

19. Atomic Mass It is useful to compare the relative masses of atoms to a standard reference isotope. Carbon-12 is the standard reference isotope. Cabon- 12 has a mass of exactly 12 atomic mass units. An atomic mass unit (amu) is defined as one twelfth of the mass of a carbon-12 atom. 4.3

20. Distinguishing Among Atoms Just as apples come in different varieties, a chemical element can come in different “varieties” called isotopes. 4.3

21. The Nucleus of the Atom All elements are identified by their number of protons Atoms with the same number of protons but different numbers of neutrons are called isotopes

22. ISOTOPES

23. Average Atomic Mass Number The average of the total number of protons and neutrons in the nuclei of all the isotopes of an element The decimal number in the periodic table

24. Relative Abundance Tells how common a particular isotope of an element is in nature It is used to calculate the average atomic mass of an element

25. Average Atomic Mass Some Elements and Their Isotopes 4.3

26. Average Atomic Mass To calculate the average atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. 4.3

27. Average Atomic Mass For example, carbon has two stable isotopes: − Carbon-12, which has a natural abundance of 98.89%, and − Carbon-13, which has a natural abundance of 1.11%. 4.3

29. Conceptual Problem for Conceptual Problem 4.3