1. Chapter 4: Atomic Structure Democritus believed that matter was made up of particles. he called nature’s basic particle an “atom”. The …… Aristotle’s idea was accepted for nearly 2000 years! (poor Democritus  ) Aristotle believed that everything was made up of 4 substances: “People’s Choice” FireWater Air Earth

2. 3. Law of multiple proportions: 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 element is always a ratio of small whole numbers. Foundation of Atomic Theory Basic laws of chemistry: 1.Law of conservation of mass: states that mass is neither created nor destroyed during ordinary chemical or physical reactions. 2.Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions by mass regardless of the sample size.

3. Dalton’s Atomic Theory 1.All matter is composed of extremely small particles called atoms. (atom: the smallest particle of an element that retains the chemical and physical properties of that element.) 2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. 3. Atoms cannot be subdivided, created, or destroyed. 4. Atoms of different elements combine in simple whole-number ratios to form compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged.

4. Note: Not all aspects of Dalton’s atomic theory are correct. atoms ARE divisible into even smaller particles. atoms of a given element CAN have different masses. The Structure of the Atom all atoms consist of two regions: 1.Nucleus: small region located at the center of an atom that contains protons (positive particles) and neutrons (neutral particles). 2.Outside nucleus: very large region surrounding the nucleus that contains electrons (negative particles). Protons, neutrons, and electrons are all considered to be subatomic particles.

5. Discovery of the Electron JJ Thomson conducted experiments using Cathode Ray Tubes (CRT). Results: All cathode rays are composed of identical negatively charged particles. Charge and Mass of the Electron Robert Millikan experimentally determined the mass of an electron to be 9.109 x 10 -31 kg. also confirmed the following: 1.Electron’s have a negative charge. 2.Atoms are electrically neutral, thus they must contain positive particles too. 3.The other particles in an atom account for most of the mass. Cathode Ray Tube 1 Cathode Ray Tube 2 Millikan’s Experiment JJ Thomson Robert Millikan

6. JJ Thomson JJ’s “Plum – Pudding” Model of the Atom

7. Discovery of the Atomic Nucleus Ernest Rutherford’s Gold Foil Experiment bombarded a thin, gold foil with alpha particles. he expected to see the heavier alpha particles pass through relatively undisturbed. most particles did, however a few bounced straight back. Result: there must be a very densely packed bundle of matter with a positive charge  nucleus

8. ParticleSymbolChargeRelative Mass Location Electron e-e- 0 amu Outside (e - cloud) Proton p+11 amunucleus Neutron n01 amunucleus If the nucleus contains positive protons, what keeps the nucleus together? neutrons provide the “glue” of the nucleus and are responsible for nuclear forces (prevent the nucleus from breaking apart)

9. Isotopes the identity of the atom is determined by the number of protons. ex: 1 proton atom  Hydrogen however, like many other elements, hydrogen atoms can contain different numbers of neutrons  ISOTOPES ex: 3 isotopes of H protium  1 proton, 1 electron deuterium  1 proton, 1 neutron, 1 electron tritium  1 proton, 2 neutrons, 1 electron Atomic Number: indicates the number of protons in an atom. Mass Number: indicates the total number of protons and neutrons present in an atom.

10. Let’s look back at the isotopes of hydrogen. Protium (1 proton, 1 electron) 1 H 1 Mass Number Atomic Number Hyphen notation Hydrogen - 1 Deuterium (1 proton, 1 neutron, 1 electron) 2 H 1 Hyphen notation Hydrogen - 2 Nuclear Symbol

11. Tritium (1 proton, 2 neutrons, 1 electron) 3 H 1 Hyphen notation Hydrogen - 3 My definition of an isotope: Isotopes are atoms of the same element that have different masses. Example: “Tin” has 10 stable isotopes……the most of any element!

12. Relative Atomic Mass because atoms have such a very small mass, chemists use atomic mass units (amu) to describe the mass of atoms. 1 amu = 1/12 the mass of a carbon-12 atom Average Atomic Mass since most elements occur naturally as mixtures of isotopes, scientists refer to the average atomic mass of an element. the average atomic mass of an element depends on two factors: 1.The abundance of each of the element’s isotopes. 2.The mass of the element’s isotopes.

13. Average atomic mass = (% abundance #1)(amu of #1) + (% abundance #2)(amu of #2) + … Calculate the average atomic mass for copper: Copper – 6362.929 59969.17 Copper – 6564.927 79330.83 Isotope Atomic mass (amu) Percent natural abundance (.6917 x 62.929 599) + (.3083 x 64.927 793) Average atomic mass for copper (Cu) = 63.54 amu

14. Do now “Isotope Problems”  How many protons, neutrons and electrons are in an atom of carbon – 13? 6 protons 7 neutrons 6 electrons  Write the nuclear symbol for oxygen – 16  Write the hyphen notation for the element whose atoms contain 7 electrons and 9 neutrons. Nitrogen - 16 O 16 8

15. Pennium Lab Calculations #1. Determine the number of isotopes of Pennium. Group your data so that there is no more than 5 isotopes. #2. % abundance of each isotope. % abundance = (total mass of the isotope) x 100 mass of 20 pennies #3. Average atomic mass of each isotope. Isotope mass = (total mass of each isotope) # of pennies of that isotope

16. #4. Atomic mass of Pennium Atomic mass of Pennium = (% abundance #1)(amu of #1) + (% abundance #2) )(amu of #2) + …