1. Atoms: The Building Blocks of Matter

2. Dalton’s Atomic Theory (1803) 1.All matter is made of atoms. 2.Atoms of the same element are identical. Atoms of different elements differ. 3.Atoms cannot be divided, created, nor destroyed. 4.Atoms combine in simple whole number ratios to form compounds. 5.Atoms are rearranged, separated, or combined in chemical reactions.

3. Dalton’s model of atom (1803) Solid, Indivisible Sphere of Matter

4. J.J. Thomson & Cathode Ray (1897)

5. Rays repelled by negatively charged object. Charge to mass ratio same regardless of type of metal or gas used in the tube Concluded that identical negative particles are in all atoms. Named them electrons. Since atoms are neutral, they must also contain a positive charge.

6. J J Thomson’s Model

7. Millikan & Oil Drop Experiment (1909)

8. Measured charge of electron Allowed mass of electron to be determined from charge and charge to mass ratio. Mass of electron much smaller than mass of atom, so the atom must contain other particles to account for most of the mass.

9. Rutherford and Gold Foil Experiment (1911)

10. Result of experiment

11. Results and Conclusions Most alpha particles went through foil while a few bounced back. Atom is mostly empty space with a dense, concentrated positive mass in the center. Electrons occupy space around the nucleus.

12. Modification of plum-pudding model Prediction of what would occur with plum-pudding model. Explanation for why a few alpha particles bounced back.

13. Rutherford’s Nuclear Model (1911)

14. James Chadwick and Neutron Experiment (1932)

15. Results Found new type of particle in the atom that was electrically neutral – the neutron. Determined mass of this particle. –Slightly larger than proton Mass of nucleus finally accounted for.

16. Chadwick Model (1932)

17. Structure of the Atom + - Atomic Number = # of protons Atomic Mass = weighted average of all isotopes measure in amu’s.

18. Subatomic Particles The proton is 1800 times bigger than the electron. ParticleSymbol Relative Electric Charge Relative Mass (amu) Location in the atom Electron e-e- 0.0005486Around nucleus Proton p+p+ +1 1.007276Part of nucleus Neutron nono 0 1.008665Part of nucleus

19. Atomic Structure Atoms are on the order of 10 -10 m. –100 million atoms lined up would be 1 cm long. The nucleus is very small compared to size of the atom.

20. If the atom was the size of a football stadium… The nucleus would be the size of a marble.

21. Closer to scale model of atom

22. Isotopes of carbon

23. Isotopes Atoms of the same element with different masses. Atomic number = number of protons Mass number = number of protons + number of neutrons

24. Isotopes Mass Number Hyphen Notation Nuclear Symbol Atomic Number

25. Atomic mass vs. molar mass Same number off of the periodic table but different units Atomic MassMolar Mass Mass of:1 atom of an element1 mole of an element Measured in:amu (atomic mass unit)grams

26. Weighted averages Find the average atomic mass of copper. Cu-63, 69.15%, 62.93 amu Cu-65, 30.85%, 64.93 amu Solution: 0.6915 x 62.93 amu = 43.52 amu 0.3085 x 64.93 amu = 20.03 amu 43.52 amu + 20.03 amu 63.55 amu