CHAPTER 4: ATOMIC STRUCTURE. Intro Video! (nothing about Bohr, I promise)

Transcription

1 CHAPTER 4: ATOMIC STRUCTURE Intro Video! (nothing about Bohr, I promise)

2 I. HISTORY OF ATOMIC THEORY A. Highlights: 1. Democritus: suggested matter was made of tiny indivisible particles 2. Aristotle: rejected idea of the atom & suggested that all matter was composed of one continuous substance called hyle (similar to clay); accepted as true until 17 th century 3. Antoine Lavoisier (1770s): explained burning & law of conservation of matter (mass remains constant) 4. Joseph Proust (1799): experimentally showed that matter does not combine randomly; law of definite proportions

3 5. John Dalton (1803): experimented with many compounds, carefully massing the products formed and comparing them with the original reactants; led to law of multiple proportions 6. Dalton revived the particle theory of matter and made the following statements based on experimental evidence: a. All elements are composed of atoms which are tiny indivisible particles. b. All atoms of the same element are exactly alike. c. Atoms of different elements are different. d. Compounds are formed by joining of atoms from two or more elements and always in whole number ratios. 7. Corrections to Dalton s theory: a. Atoms can, to a small degree, be altered by chemical change. b. Atoms are not indivisible and can be changed by nuclear reactions. c. Atoms of the same element can be different, isotopes.

4 Copy into notes: Figures 4-2, 4-3, and 4-5 (4-2 & 4-3: do not draw pictures, just copy text; 4-5: draw pictures, copy text & copy caption) Classwork: p. 91 #1-5 (will check tomorrow)

5 II. DISCOVERING SUBATOMIC PARTICLES A. William Crookes (English physicist, 1870 s): worked with cathode rays in discharge tubes & found that they could be deflected by a magnet; concluded that the rays consisted of tiny electrically charged particles 1. Due to Crookes work & further research, the scientific community determined that these particles contained a negative charge, but there was no concrete evidence of this

6 B. J.J Thomson (English scientist, 1897): proved Crookes was right. 1. Wanted to determine the ratio of this particle s negative charge to its mass 2. Worked with cathode rays to show that both electric and magnetic fields caused deflections & measured these deflections 3. Found a charge-to-mass ratio of the charged particle & compared it to other known charge-to-mass ratios; discovered that the mass of this particle was less than the mass of the smallest atom (hydrogen) 4. Determined that this charged particle must be smaller than an atom & therefore part of the atom; called these electrons 5. Calculated an e/m (charge/mass) ratio from data e/m = 1.76 x 10 8 coulomb/g.

7 C. Robert Millikan (American physicist, ) 1. Performed the famous oil drop experiment & determined the exact charge of the electron 2. All the charges were multiples of a smaller charge; he calculated the smallest charge to be 1.60 x coulomb 3. He used Thomson s e/m ratio & solved for the mass of an electron; 9.11 x g 4. The discovery of the electron & its mass led to J.J. Thomson s plum pudding model small negative charges distributed throughout a uniform positive charge

8 D. Lord Ernest Rutherford (English scientist, 1912) & squad (Niels Bohr, Hans Geiger, Ernest Marsden) 1. Performed the famous gold foil experiment shot alpha particles (+) from a radioactive source at a thin sheet of gold foil. Most passed right through, some were deflected, and a few bounced back! 2. Made the following conclusions: a. The foil was mostly empty space b. The deflections were caused by a positive charge c. The positive charge was in a very small area in the center of the atom (nucleus) d. The atom then consisted of a small dense positive core (nucleus) orbited by electrons e. The opposite charge of the positive core and the negative electrons canceled each other out, resulting in a neutral atom.

9 3. He further surmised that the nucleus of hydrogen, the least massive atom, is a single particle with a positive charge, the proton. 4. Nuclei of larger atoms then contained a larger number of protons. 5. Problem: he found that if he set the mass of the proton at 1.00, then the protons that account for the positive charge do not account for the total mass only part of it. 6. Therefore, the atom had more mass than the amount contributed by the protons. 7. Suggested tightly bound proton-electron pairs which he called neutrons to account for the missing mass no real evidence for this.

10 E. James Chadwick - English scientist (1932). 1. Bombarded the element beryllium with alpha particles. 2. A new type of particle was given off which had no charge and about the same mass as the proton. 3. Concluded that these particles were Rutherford s proton-electron pairs and kept the name neutron for these particles. Copy Figures 4-9, 4-12, and 4-13 (include all text), & Table 4-1

11 Work in pairs or groups of 3. Create a poster addressing the following models of the atom: John Dalton's solid sphere model J.J. Thomson's plum pudding model Ernest Rutherford's nuclear model Poster must include: Who Person(s) responsible What A description and an illustration of the model; make sure to address how protons, neutrons, and electrons are included in the model Where Country, state, university, etc. When Date(s) Why The prior knowledge at the time How A summary of the technology or evidence used to develop the model References Page numbers from our textbook and at least two other helpful resources

12 III. DIFFERENCES IN ATOMS A. Atom: smallest amount of a chemical element that can still take part in a chemical reaction. 3 parts: 1. Proton (p + ): positive particle found in the nucleus Mass: amu Charge: x coulomb 2. Neutron (n o ): neutral particle found in the nucleus Mass: amu (about the same as p + ) Charge: 0 3. Electron (e - ): negative particle found orbiting the nucleus Mass = amu (about 2000x less massive than p + & n o Charge = x coulomb sketch & label this

13 B. Atomic number: number of protons in an atom 1. Elements are composed of atoms which all have the same number of protons 2. atomic number = number of protons = number of electrons (in a neutral atom) 3. Henry Moseley (English physicist, 1913) found that the wavelength of x-rays from a source depended on the number of protons present; performed experiments to determine the atomic numbers of many elements

14 C. Isotopes: atoms of the same element (therefore, same # of protons) with differing numbers of neutrons 1. Isotopes of an element usually have the same chemical behavior because chemical behavior is determined by electrons 2. Example: Isotopes of hydrogen (copy this diagram exactly, words included)

15 D. Mass number: number of protons and neutrons in the nucleus of a particular atom 1. mass number = #protons + #neutrons 2. Example: Find the number of neutrons in potassium-40 Since it is called potassium-40, that means that the mass number is 40 Find # of protons (from periodic table -- # of protons is the same as the atomic number) potassium s atomic number = 19 mass number = #protons + #neutrons 40 = 19 + x 21 = x Potassium-40 contains 21 neutrons

16 Copy Figure 4-14 (p. 98) Bookwork: p. 97 #6-10, p. 99 #11-13, p. 101 #14