2. MODELS OF THE ATOM A HISTORICAL PERSPECTIVE

3.  Anything that has mass and takes up space  If you did not know this definition, how would you describe matter?

4.  460 B.C. - Democritus thought matter could not be divided indefinitely. Democritus This led to the idea of atoms in a void. The term “atomos” means uncuttable or indivisible He speculated that all matter was composed of atoms

5. fire air water earth  350 B.C - Aristotle modified an earlier theory that matter was made of four “elements”: earth, fire, water, air.  Aristotle was wrong. However, his theory persisted for 2000 years. Aristotle

6.  1800 -Dalton proposed a modern atomic model based on experimentation not on pure reason. All matter is made of atoms. Atoms of an element are identical. Each element has different atoms. Atoms of different elements combine in constant ratios to form compounds. Atoms are rearranged in reactions. His ideas account for the law of conservation of mass (atoms are neither created nor destroyed) and the law of constant composition (elements combine in fixed ratios).

7.  Every atom was a spherical hard sphere  Each element was different and therefore had a different symbol

8.  First to provide existence of electrons  Used cathode ray tubes (remember Grade 9 science???)  cathode ray tube cathode ray tube  Reasoned that electrons could be produced from electrodes of various metals; therefore all atoms contained electrons  since atoms are electrically neutral they must also contain a positive charge

9.  Thomson postulated that an atom consists of a cloud of positive charge with negative charged electrons embedded randomly inside of it  Chocolate chips are electrons embedded in a positive sphere

10.  Rutherford shot alpha (  ) particles at gold foil. Most particles passed through. So, atoms are mostly empty space Some positive  -particles deflected or bounced back! Thus, a “nucleus” is positive & holds most of an atom’s mass. Radioactive substance path of invisible  -particles Lead block Zinc sulfide screenThin gold foil

11.  Rutherford’s reaction to the experiment was like shooting a 15 inch shell at a piece of paper and having it bounce back at you.

12.  Because of the observations through the gold foil experiment, Rutherford proposed a model where electrons “orbited” the nucleus.  Rutherford named the positive charges in the nucleus as protons

13.  Worked with Rutherford to determine the masses of nuclei of different elements  In experiments found that masses of nuclei were different than the sum of the masses of protons  Concluded that nucleus also contains neutral charged particles called neutrons

14.  Atoms consist of a 1. Nucleus – protons (+) (identifies atom) and neutrons (0) 2. Orbiting electrons (-) (determines chemical properties of atom)  Size of nucleus is small compared to atom (page 137)  Some atoms are different. They have different number of neutrons but same number of protons and electrons - isotopes

15.  isotope notation indicates the chemical symbol of the element symbol used and the isotope of the element  Z - represents atomic number  A - represents mass number (total number of protons and neutrons  X - represents the chemical symbol  Occur naturally in nature

16.  Can also be written as element name and mass number  Why don’t we indicate the atomic number?  Ex. carbon-12 and carbon-14  Some isotopes are unstable and spontaneously decay releasing radioactive material such as alpha and beta particles etc.  Homework page 142 #1-4,7

17.  tour of super collider tour of super collider

18.  Contradicted laws of physics where electron in motion must continuously release energy, thereby creating smaller orbits and eventually crashing into nucleus

19. Visible light is part of a broader spectrum called the electromagnetic spectrum. Type of radiation is dependant on the wavelength. The shorter the wavelength, the higher the frequency, resulting in more energy.

20.  All radiation travels at the speed of light =3.00x10 8 m/s in a vacuum

21. Electrons orbit the nucleus in “shells” Electrons can be bumped up to a higher shell if hit by an electron or a photon of light.

23.  There are 2 types of spectra: continuous spectra & line spectra. It’s when electrons fall back down that they release a photon (unit of light energy). These jumps down from “shell” to “shell” account for the line spectra seen in gas discharge tubes (through spectroscopes).

24.  Atoms can absorb specific amounts of energy and release the same amount of energy  Forms an emission spectrum, not a continuous spectrum

25.  Atomic spectra of elements Atomic spectra of elements

26.  Will find emission spectra in IR region, UV regions etc. depending on the amount of energy released

27.  Bohr incorporated Rutherford’s planetary model but made some restrictions based on the spectra he observed  1) atoms have specific energy levels called stationary states (fixed circular orbit)  2) while in a specific energy state, the electrons do not emit energy  3) electrons can change orbits by emitting or absorbing specific quantities of energy

28.  Model only successfully explained one- electron systems such as hydrogen, He +, Li 2+,etc.  Unable to explain emission spectra of atoms with two or more electrons

29.  Atomic spectra of elements Atomic spectra of elements  Sublevels within an energy level  Large gaps and small gaps