Download Chapter 4: The Structure of the Atom Early Ideas about Matter Name

Chapter 4: The Structure of the Atom
Early Ideas about Matter
Name
Date
Democritus
Aristotle
John Dalton
400 BC atomos – smallest particle
350 BC hyle – continuous matter
1803 AD billiard ball – indivisible,
smallest particle
JJ Thomson
1897 AD plum pudding – electrons
in positive pudding
mass
(
/ charge) ratio e–
1910 AD discovered the charge of
the electron
1911 AD nuclear model – small,
dense nucleus with esurrounding empty
space
Robert Millikan
Ernest Rutherford
Theory
Experiment
none, philosophical argument
none, philosophical argument
laws of: conservation of matter
definite proportion
multiple proportion
Cathode ray tube / e- beam
(a Crookes tube)
Oil drop experiment (gravity,
e– charge, and charged plates)
α – particle / gold foil
Dalton’s Atomic Theory
1. Matter is composed of extremely small particles called atoms
2. Atoms are indivisible and indestructible
3. Atoms of a given element are identical in size, mass, and chemical properties
4. Atoms of a specific element are different from those of another element
5. Different atoms combine in simple whole-number ratios to form compounds
6. In a chemical reaction, atoms are separated, combined, or rearranged
Atomic structure
Atoms are made of three particles
1. protons (p+)
2. neutrons (n0)
3. electrons (e–)
}
nucleus
Isotopes: atoms of the same element with different numbers of neutrons
Examples: hydrogen isotopes are protium, deuterium, and tritium
Hyphen notation: protium = hydrogen-1, deuterium = hydrogen-2, tritium = hydrogen-3
Nuclear symbols: protium = H , deuterium = H, tritium = H
from Wikipedia: http://en.wikipedia.org/wiki/Isotope
Calculating nucleons in isotopes
Mass number
= number of p+ + number of n0
Atomic number = number of p+
= mass number – atomic number
Number of n0
(top number in a nuclear symbol)
(bottom number in a nuclear symbol)
Examples: (for neutral atoms)
H
Li
n0 = 2
n0 = 4
p+ = 1
p+ = 3
e– = 1
e– = 3
Relative atomic masses
The atomic mass unit (amu) is defined as of a carbon – 12 nuclide
Mass defect explains why
He = 4.002 603 254 15(6) but
W = 181.948 204 2(9)
+
The more p squeezed into a nucleus, the more mass they must convert to binding energy
(by Einstein’s famous equation E = mc2)
Atomic masses on the Periodic Table are not whole numbers because they are the weighted
average of all the naturally occurring nuclides of that isotope
Example: magnesium has three nuclides: 24Mg at 78.99% and 23.985 041 7 amu,
25
Mg at 10.00% and 24.985 836 92 amu, and 26Mg at 11.01% and 25.982 592 93 amu.
24
Mg: 23.985 041 7 amu x 0.789 9 = 18.945 784 amu
Mg: 24.985 836 92 amu x 0.100 0 = 2.498 584 amu
26
Mg: 25.982 592 93 amu x 0.110 1 = 2.860 683 amu
25
24.305 amu
Unstable Nuclei and Radioactive Decay
Three common types of radioactive decay
Alpha (α) = He
Beta (β)
= e (or e–)
Gamma (γ) = EMR (electromagnetic radiation or very high energy light)
Examples:
α:
Ra →
Rn + He
or
Ra →
Rn + α
β: C → N + e
or
Ra →
Rn + e–
Th + He + 2 γ
or
U→
Th + α + 2 γ
γ:
U→
Nuclear stability: The ratio of the number of p+ to n0 is the primary factor in causing the
instability that results in nuclear decay
Stable nuclei for small atomic numbers have a ratio of 1 p+ to 1 n0
Stable nuclei for high atomic numbers have a ratio of 1 p+ to 1.5 n0
Unstable nuclei will undergo a series of α and/or β decays until they reach a stable p+ to n0
ratio