History and Structure of an Atom

1. Atoms

2. History of an atom • Comes from the Greek word atomos • Democritus - Greek philosopher around the year 400 BC. • Democritus concluded that matter could not be divided into smaller and smaller pieces forever. Eventually, the smallest piece of matter would be found. Which he called atomos

3. The Dalton Model - 1803 1. All elements are composed of indivisible particles. 2. Atoms of the same element are exactly alike. 3. Atoms of different elements are different. 4. Compounds are formed in definite ratios Ex: H2O is water, and H2O2 is hydrogen peroxide

4. Atoms • Atoms are so small that, even today, direct visual inspection is all but impossible • As you can see, details of the atoms cannot be seen.

5. Discovered electrons in 1891. Cathode ray tubes J. J. Thomson (1856-1940)

6. Thomson’s Plum Pudding Model

7. Ernest Rutherford • British physicist who discovered the nucleus in 1911 • Gold Foil or Alpha Scattering experiment

8. Discovery of the Nucleus

9. Rutherford vs. Thomson • Rutherford explained the scattering of the alpha particles by saying there must be a nucleus (small, dense, positive center) • Thomson model spread the entire mass of the atom throughout that space • Rutherford put most of the mass of the atom at the center of the atom, in a space much, much smaller than the atom itself

10. Neils Bohr • Electrons move in definite orbits around the nucleus, like planets moving around the sun. • Bohr proposed that each electron moves in a specific energy level.

11. Parts of an Atom • Protons: positively (+) charged, located in the nucleus • Neutrons: no charge (0), located in nucleus • Electrons: negatively (-) charged, located around the nucleus

12. Size of an atom • If this were the proton in a hydrogen atom, it would take a screen 1 mile across to display the electron's orbit.

13. Electrons vs. Protons • If an electron weighed the same as a dime, a proton would weigh the same as a gallon of milk.

14. Atomic Number • Always expressed as a whole number on the periodic table. • The atomic number equals the number of protons in an atom.

15. Atomic Mass • Almost always expressed as a decimal fraction on the periodic table • The mass number of an atom is equal to the sum of the protons and neutrons in its nucleus.

16. Ions • Particles that have a positive or negative charge. • Formed by gaining or losing electrons • Metals tend to lose electrons (positive ions) and nonmetals tend to gain electrons (negative ions)

17. How do I find the number of protons? • Always equal to the atomic number. • If the number of protons changes, a different element is formed

18. Number of electrons? • Basically, the number of electrons in an atom is equal to the atomic number. • The numbers of electrons are easily changed within an atom. If the chemical symbol is written with a charge, representing an ion, the charge indicates the number of electrons that have been added or removed.

19. # of neutrons? • Subtract the atomic number from the mass number to find the number of neutrons. • Number of neutrons in an atom = mass number (as a whole) - atomic number

20. How many protons are in each of these atoms? carbon iron bromine gold uranium How many electrons are in each of these atoms? helium chlorine sodium mercury barium How many neutrons are in each of these atoms? copper lead silver iodine tin Practice