Chapter 4: Atomic Structure

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

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

3. Dalton’s Atomic Theory • 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: • Nucleus: small region located at the center of an atom that contains protons (positive particles) and neutrons (neutral particles). • 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: • Electron’s have a negative charge. • Atoms are electrically neutral, thus they must contain positive particles too. • The other particles in an atom account for most of the mass. JJ Thomson Cathode Ray Tube 1 Cathode Ray Tube 2 Millikan’s Experiment Robert Millikan

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

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

11. 3 H 1 Tritium (1 proton, 2 neutrons, 1 electron) 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: • The abundance of each of the element’s isotopes. • 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: Atomic mass (amu) Percent natural abundance Isotope (.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 16 O 8

15. Pennium Lab Calculations #1. Determine the number of isotopes of Pennium. <Hint> 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) + …