Download Chapter 4 Notes - DunlapChemistry

Chapter 3 Notes
History of Atomic Theory
Democritus
Greek philosopher - 460-370 B.C.
Believed the world is made up of small indivisible particles called atoms and empty
space.
Ideas based on observations.
Aristotle
Greek philosopher - ~400 B.C.
Matter is continuous with no indivisible particles.
Aristotle “won” the debate with Democritus. His idea was accepted (vs Democritus)
until the 1700s.
Lavoisier (Antoine)
French chemist in the 1780s
Observed that the mass before and mass after a reaction remained the same.
Law of Conservation of Matter (mass)
In a normal chemical reaction, matter cannot be created or destroyed.
Proust (Joseph)
French chemist ~1800
Law of Definite Proportions
Elements combine in a certain proportion or ratio (by mass) to form a compound.
Example: H2O is always 11% hydrogen and 89% oxygen
Dalton (John)
English schoolteacher and chemist in the early 1800s.
Law of Multiple Proportions.
Proposed the atomic theory of matter (aka, the Modern Atomic Theory).
Modern Atomic Theory
1.
All elements are composed of small, indivisible particles called atoms.
2.
Atoms of the same element are identical. Atoms of different elements are
different.
3.
Atoms combine in simple whole-number ratios to form compounds.
4.
Chemical reactions occur, but atoms are never changed into atoms of
another element.
Changes: Atoms of the same element may be different : isotopes
Atoms can be divisible.
Berzelius (Jons Jakob)
Swedish chemist – 1811
Established modern system of chemical symbols for elements.
Thomson (JJ)
British physicist – late 1800s/early 1900s.
Using Crookes’ tubes (cathode ray tubes), he discovered electrons.
Determined …
(1) electrons are negatively charged particles.
(2) all electrons are identical.
(3) the charge to mass ratio using deflection of rays.
Plum Pudding Model of Atoms
The atom is a sphere of positive matter in which negative electrons are embedded.
Millikan (Robert)
American physicist - ~1909
Oil drop experiment.
Discovered the charge of an electron. It has a charge of -1.
Discovered the mass of an electron. It is 1/1840 of a proton. It has a mass of 0 amu.
Goldstein (Eugen)
German physicist – 1886
Discovered canal rays – also called anode rays. This led to the development of mass
spectrometry.
Wien (Wilhelm)
German physicist – 1898
Discovered the proton. It has a +1 charge (equal and opposite to the electron). It has a
mass of 1 amu.
Rutherford (Ernest)
British chemist/physicist – 1911
Performed the Gold Foil Experiment with the help of Geiger and Marsden. Alpha
particles were “shot” at a thin sheet of gold foil.
Results of Gold Foil:
(1) Most particles (93%) went straight through the foil.
(2) Some particles were slightly deflected.
(3) About 1 in 8000 particles came straight back.
Conclusions from Gold Foil:
(1) The atom is mainly empty space.
(2) The atom contains a small, dense positively charged core called the nucleus.
Predicted the existence of neutrons.
The mass of the protons in the nucleus was only about ½ of the known mass of
the nucleus. Predicted a neutral particle must exist in the nucleus.
Chadwick (James)
English physicist – 1932
Discovered the neutron. It has a 0 charge. It has a mass of 1 amu.
Worked with Geiger and Rutherford.
Parts of the Atom
Protons
Positive (+) charge
In the nucleus
Contribute to atomic mass
Determine the atomic number
Never gained or lost
Electrons
Negative (-) charge
Outside the nucleus
In different energy levels; each row on the periodic table is a new energy level
Do not contribute to atomic mass
A proton’s mass is 1800 times greater than that of an electron
Gained, lost or shared in bonding
Neutrons
Neutral charge (no charge)
In the nucleus
Contribute to atomic mass
Never gained or lost
Different isotopes of the same element have a different number of neutrons
Atomic number
Equal to the number of protons
Identifies the element
Is the whole number on the periodic table
Examples: H-1, He-2, Li-3
At. # = # of p+ = # of e- (in the elemental state)
Mass number (also called atomic mass number)
Equal to the atomic number + the number of neutrons
Equal to the number of protons + the number of neutrons
Is a whole number; is not shown on the periodic table
Mass # = At. # + # neut.
Mass # = # p+ + # neut.
Average atomic mass
A weighted average of the masses of all the isotopes for a given element
Examples: H-1.01; Cl-35.45
Isotopes
Elements that contain the same number of protons but a different number of neutrons
Example: C-12 and C-14; both have 6 p+ and 6e-; C-12 has 6 neutron and C-14
has 8 neutrons
Determine the following for each atom: (1) element name; (2) element symbol; (3) atomic
number; (4) atomic mass number; (5) # of protons; (6) # of electrons; (7) # of neutrons.
1.
2.
3.
28 protons
30 neutrons
1. Nickel
2. Ni
3. 28
4. 58
Tungsten – 186
1. Tungsten
2. W
3. 74
4. 186
Pb
124 neutrons
1. Lead
2. Pb
3. 82
4. 206
5. 28
6. 28
7. 30
5. 74
6. 74
7. 112
5. 82
6. 82
7. 124