Atomic Theory CP

1. Atomic Theory CP History of the Atom; Modern Atomic Theory, Subatomic Particles

2. Reasons to look at history • Shows that present theory is the work of many scientists over many years • To become aware of the logic used in obtaining evidence to support the theory • To gain a better understanding of the properties of atoms

3. Greek Philosophers • Democritus (460 – 371 B.C.) • Stated that the world was made of tiny particles that are mostly empty space and that they are the building blocks of all matter • Tiny particles could not be divided • Termed them atomos, or atoms • Means “uncut” or “indivisible” • Aristotle (384 – 322 B.C.) • Proposed that matter was continuous (i.e. came from only four elements – earth, air, fire, and water)

4. Dalton’s Atomic Theory • John Dalton (1766 – 1844) • Work marks the beginning of the development of the modern atomic theory • Revived Democritus’ ideas and backed them up with scientific experimentation and data! • Dalton’s Atomic Theory: • Matter is composed of extremely small particles called atoms • Atoms are indivisible and indestructible • Atoms of a given element are identical in size, mass, and chemical properties • Atoms of a specific element are different from those of another element • Different atoms combine in simple whole-number ratios to form compounds • In a chemical reaction, atoms are separated, combined, and rearranged

5. Dalton’s Atomic Theory • Dalton’s theory successfully explained: • The law of conservation of mass (proposed by Lavoisier) • States that mass is conserved in any process, such as a chemical reaction. • Conservation is the result of the separation, combination, or rearrangement of atoms • The law of multiple proportions (proposed by Proust) • Certain elements can combine to form two or more different chemical compounds • Experimental evidence led to general acceptance of his atomic theory! • Not all of the theory was accurate. • It has had to be revised with new information

6. Modern Atomic Theory • All matter is made up of a very small particles called atoms • Atoms of the same element have the same chemical properties, atoms of different elements have different properties • Any natural sample of an element has an average mass characteristic of that element • Compounds are formed when atoms of two or more elements unite • Atoms are not subdivided in physical or chemical reactions

7. Protons, Neutrons, and Electrons Defining the atom

8. The Atom • Atom: • Smallest particle of an element that retains the properties of the element • Size: • Consider this: • In 2006 the population of Earth was about 6.5 x 109 people • A solid copper penny contains 2.9 x 1022 atoms • This is 5 trillion times larger than the worlds population!!!!!

9. Electrons • J.J. Thomson • 1879 Cathode Ray Tube • Provided the first evidence that atoms are made of even smaller particles • Cathode-Ray Tube: • Vacuum tube, all gases removed • Metal electrodes at each end • Electricity runs from the negative cathode to the positive anode • Using a magnet, the negative charged cathode ray bent towards a positively charged plate

10. Electrons • Observations of the cathode-ray behavior led to the hypothesis that cathode rays consist of identical negatively charged particles • Thomson proposed the “Plum Pudding” model of the atom

11. Thomson’s Model • Found the electron • 1 unit of negative charge • Mass 1/2000 of hydrogen atom • Later refined by Millikan to 1/1840 • Concluded that there must be a positive charge since atom was neutral • Atom was like plum pudding • A bunch of positive stuff, with electrons able to be removed

12. Nucleus • Ernest Rutherford – 1908 • Discovered the nucleus • The core of the atom and contains protons and neutrons • Rutherford’s nuclear model proposed that an atom is mostly empty space • This is based on studies of the bombardment of thin metal foils with fast-moving positively charged particles

13. Rutherford’s Experiment

14. Rutherford’s Experiment • His experiment demonstrated that: • An atom is mostly empty space • There is a small dense, positive piece at the center • Alpha particles are deflected if they got close enough to the positive center

15. Protons and Neutrons • Protons: • Rutherford also concluded that the nucleus also contained positively charged particles call protons. • Subatomic particle with a positive charge equal in magnitude to the negative charge of an electron • The number of protons determines the identity of an atom!! • Neutrons: • Discovered by James Chadwick (Rutherford’s coworker) • Showed the nucleus also contained a neutral subatomic particle, the neutron • Electrically neutral subatomic particle • The number of neutrons can vary within an element  Isotopes!

16. Rutherford Activity

17. Structure of the Atom • There are two regions • The nucleus: • Protons and Neutrons • Positive charge • Almost all of the mass • Electron Cloud: • Most of the volume of an atom • Region were electrons can be found

18. Comparing Subatomic Particles

19. Atomic Number • Atoms of each element contain a unique positive charge in their nuclei • Therefore, the number of protons in an atom identifies that particular atom! • The number of protons in an atom is referred to as the atomic number • Because all atoms are neutral: • The number of protons = the number of electrons • If you know the atomic number of an element you now know both the number of protons and the number of electrons! Atomic Number

20. Mass Number • The total number of protons and neutrons in the nucleus of an atom • Mass # = p + n • Usually found at the bottom of the atomic symbol • How to find the number of neutrons: • Mass number – atomic number = # neutrons • Example: • Nitrogen: • Mass number: 14 • Atomic number: 7 • Number of Neutrons: 7

21. Fill in the Gaps

22. Mass Number Continued • Although a given type of atom will usually contain a certain number of neutrons in the nucleus, a small percentage will not • Ex: most hydrogen atoms contain no neutrons • A small percentage contain one neutron and a smaller percentage two neutrons • What do we call atoms with a different number of neutrons?

23. Isotopes • In nature, most elements are found as a mixture of isotopes. • Isotope: • Atoms that have the same number of protons but different numbers of neutrons • This means that the number of neutrons in a atom can change! • Protons MUST stay the same • Different isotopes of the same element have different mass numbers (i.e. different neutrons) • A given element has at least two isotopes, and they all have different mass numbers • Example: • Gadolinium (Gd) has 6 stable isotopes

24. Isotopes • Isotopes are chemically alike because they have identical numbers of protons and electrons • It is useful to compare the relative masses of atoms to a standard reference isotope • Carbon-12 is the standard reference isotope • Carbon-12 has a mass of exactly 12 atomic mass units (amu)

25. Atomic Mass • A weighted average mass of the atoms in a naturally occurring sample of the element • Mass of an atom in atomic mass units (amu) • Equal to 1/12th of the mass of carbon • Weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature • This is the number which appears on the periodic table

26. Atomic Mass Number • Example: • 99% of all carbon atoms are the isotope containing 6 neutrons, the remaining 1% is the heavier isotope containing 7 neutrons, which raises the aver mass of carbon from 12.000 to 12.011

27. Calculations • Carbon has two major, stable isotopes. • Carbon-12 (98.93% abundance) • Carbon-13 (1.07% abundance) • Carbon-14 is in trace amounts and is radioactive • Calculate the average atomic mass of carbon • Carbon 12: (0.9893 x 12) = 11.8716 • Carbon 13: (0.0107 x 13) = 0.1391 • AAM: 12.0107

28. Calculations • Rubidium has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is 72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium? • Average atomic mass = (0.722 x 85) + (0.278 x 87) • Average atomic mass = (61.37) + (24.186) • Average atomic mass = 85.56 amu