1. Chapter 3 Atoms: The Building Blocks of Matter

2. Objectives: Students should be able to: Summarize the essential points of Dalton’s atomic theory. Describe the particle theory of matter. Use the Bohr model to differentiate among the three basic particles in an atom Compare the Bohr atomic model to the electron cloud.

3. Atomic Models The Bohr Model The nucleus is orbited by electrons. The electrons are in different energy levels.

4.  Early theories of matter Philosophers (ideas) 1.Democritus (460-370 B.C.) a. First to propose the idea that matter was not divided into smaller parts b. Atomos: Atoms: tiny particles could not be divided

5. 2. Aristotle (384-322 B.C.) a. Rejected Democritus b. believed in 4 elements air, earth, water and wind (everything came from these 4 elements)

6. 3. Alchemist a. Changed metal into gold b. Tried to prolong life c. Through their mistakes they learned chemistry d. First to do actual experiments

7. Dalton’s Model - Performed a number of experiments that eventually led the idea of atoms - Theory: 1. All elements are composed of atoms. Atoms are indivisible and indestructible particles. 2. Atoms of the same element are exactly alike. 3. Atoms of different elements are different.

8. 4. Compounds are formed by the joining of atoms of two or more elements. This theory became the foundation of modern chemistry.

9.  J.J. Thomson - Discovered the electron - Plum-pudding model of the atom Stated atoms were made from a positively charged substance with negatively charged electrons scattered about

10. His Experiment 1. Passed an electric current through a gas. 2. As the current passed through the gas, it gave off rays of negatively charged particles 3. Passed a magnet over the gas. The beam bent away or moved towards the magnet.

11. Conclusions: 1. The negative charges came from within the atom. 2. A particle of matter smaller than the atom had to exist. 3. The atom was divisible. 4. Called the negatively particles “corpuscles” (now called electrons) 5. Since the gas was known to be neutral, there had to be positive charged particles in the gas.

12. His experiment:

13.  Rutherford - Discovered the nucleus and it’s charge, positive - His experiment: Gold-Foil experiment

14. Results: 1. Most of the positively charged particles passed straight through 2. A few of the alpha particles were slightly deflected 3. A very small number of alpha particles hit and bounced back off the foil Conclusions: 1. Most of matter is empty space

15. 2. Scattered throughout- not everywhere- are small areas of positive charge 3. Scattered throughout are areas of very dense matter 4. He called the center the nucleus; it is positive and tiny compared to the rest of the atom

16. Rutherford’ model  Chadwick: discovered the neutron

17. Completing the atom 1.Atomic number (Z) - tells the number of protons - in a neutral atom: protons = electrons 2.Mass number (A) protons + neutrons 3. AzAz x ex. 40 20 Ca nucleotide

18. 4. Protons and neutrons make up 99.7% of the atoms mass 5. Chemical name Hydrogen atomic number 1 chemical symbol H average atomic mass 1.008

19. Examples:  How many protons, neutrons and electrons in the following: 1. 2. 3. 4.

20. Properties of subatomic particles

21.  Isotopes 1. atoms with the same number of protons but different number of neutrons 2. Identified by the mass 3. Example: Chlorine-35 Chlorine-37 mass number

22. Examples of neutral atoms and isotopes Atomic # Mass # p n e 27 Al 3+ 13 27 13 14 10 13 70 Zn 30 70304030 30

23. Relative Atomic Masses 1.Atomic mass unit (u) : Is exactly 1/12 the mass of a carbon-12 atom 2. Atomic mass: expression of atomic mass unit 3. Atomic mass unit– weighted average mass of the isotopes of that element

24. Average atomic mass formula (mass)(% of element) + (mass)(% of element) = (total mass) Examples: 1. A carbon isotope has an abundance of 98.90% with a amu of 12.0 g and the other isotope has an abundance of 1.10% with a amu 13.003 g. What is the average atomic mass of the two isotopes?

25. 2. An element, E, has an atomic mass of 18.40 and consists of two isotopes: E-17 with a mass of 16.95 and E-20 with a mass of 19.35. How much E-20 does this element contain? (amu)(%) + (amu)(%) = total mass (100)

26.  Unstable Nuclei and Radioactive Decay Radioactivity 1. Substances spontaneously emitting radiation 2. radiation – ray and particle emitted by radioactive materials

27. 3. Elements (atoms) can change into atoms of another element Radioactive elements 1. Nuclei are unstable 2. Gain stability by losing electrons

28. 3. Radioactive decay – process of becoming stable by losing energy Type of radiation 1. alpha radiation (α) a. Deflected towards (-) charged plate

29. b. alpha particles have 2p + and 2n 0 c. α particles equivalent to He-4 d. Nuclear equation (Shows how decay occurs) 226 4 + 222 Ra  α Rn 88 2 86

30. 2. Beta radiation a. Deflected towards the positive plate b. Beta particles ( ) c. 14 0 14 C  β + N 6 -1 7

31.  Half-Life (t 1/2 ) 1. Time required for one-half of a radioisotope’s nuclei to decay into it products 2. Examples: Iron-59 half-life is 44.5 days. How much of a 2.000 g sample will remain after 133.5 days.

32. t 1/2 = 44.5 days half-life