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Notes – Chapter 4 – The Structure of the Atom
I.
History of the Atomic Theory
A. Democritus (460 – 370 BC)
1. matter
2. movement of atoms
3. different kinds of atoms
B. Aristotle (384 – 322 BC)
1. four kinds of matter
2. transforming matter
3. idea of “atom” rejected
5. this theory was popular and easier to accept
C. Antoine Lavoisier (1770’s)
1. Experiment
2. Law of conservation of Mass
D. Joseph Proust (1779)
1. Develops Law of Definite Composition
a. example
b. elements combine in whole number ratios – WHY???
E. John Dalton (1803)
1. develops Law of Multiple Proportions
2. example of Law of Multiple Proportions
3. Dalton collected data and developed his atomic theory in 1803
4. Dalton’s Background
a.
b.
c.
d.
Dalton was a teacher at age 12
loved meteorology
studied works of Democritus, Boyle and Proust
wrote New System of Chemical Philosophy in 1808
5. Dalton’s Atomic Theory
a.
b. *
c. *
d.
e.
*modified in the Modern Atomic Theory
F. Modern Atomic Theory
1. composition of all matter
2. atoms of elements
3. mass of atoms of elements
4. combining of atoms
5. subdividing atoms
G. History of Models of the Atom
1803
1897
1909
1913
1935
Today
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solid
electron
proton
electron
neutron
quantum atom
particle
orbits
theory
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Dalton
Thomson
Rutherford
Bohr
Chadwick
Schrodinger
And others
II.
History of Atomic Structure
A. J.J. Thomson (1856-1940)
1. Experiments with cathode ray tubes
a. showed atoms have (-) charged particles – identified first subatomic particle
b. determined charge/mass ratio of the electron
2. Thomson’s Pudding Model
a. electrons (-) charge are present suspended in positive “stuff”
c. atom is like plum pudding
B. Robert Milikan (1868-1953)
1. Oil Drop Experiment (1909)
a. discovered
b. mass of the electron
C. Ernest Rutherford (1871-1937)
1. discovered proton (+) charge
2. received Nobel Prize in Chemistry
3. Gold Foil Experiment (Expectations)
a.
b.
4.
Gold Foil Experiment Results
a.
b.
c. Rutherford reasoned
5. Gold Foil Experiment Conclusions (atom has a lot of empty space with (+) protons in the
nucleus surrounded by (-) electrons)
a. the atom is mostly empty space
b.
C. Niels Bohr (1913)
1. electrons orbit nucleus in predictable paths
D. James Chadwick (1891-1974)
1. discovers neutron in 1935
2. charge of neutron
3. mass of neutron
E. The Quantum Atom Model (Theory) – Current Model
1.
2. two regions in the atom
III.
Subatomic Particles
A. Comparing Particles
Particle
Symbol
Charge
Relative Mass
Actual Mass
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0
Electron
9.11 x 10-28
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+1
Proton
1.673 x 10-24
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Neutron
n0
B. Atomic Number and Mass Number
1. Atomic number
a. identifies the element
b. no two elements have the same atomic number
1.675 x 10-24
c. examples
2. Mass number
a.
b. mass number is close to the mass of an atom in amu (atomic mass units)
c. isotopes
1) number of neutrons in isotopes (mass # - atomic # = number of neutrons )
2) examples of isotopes
C. Ions
1. Electrons and ions
a. in neutral atoms the number of electrons equals the number of protons
b. if there are more electrons than protons a negative ion forms (anion)
c. if there are fewer electrons than protons a positive ion forms (cation)
D. Changing the Number of Particles
1. You can never change the number of protons and still have the same element
2. If you change the number of neutrons in an atom, you get an isotope
3. If you change the number of electrons in an atom you get an ion
E. Nuclear Notation
2. one method to depict isotopes of an element
3. contains symbol of the element, the mass number and the atomic number
X
C
Na
C O
F. Hyphen Notation
1. element symbol or name and mass number
2. examples
a.
b
.
c.
d.
e.
IV.
Mass of Atoms
A. Atomic Mass
1 Mass of an atom
a. too small to measure in grams
b. use relative mass (amu )
O
1) 1 amu is atomic mass unit
2) 1 amu is 1/12 the mass of one C-12 atom
B. Average Atomic Mass
1. weighted average mass of all known isotopes
a. weighted means
b. mass of each isotope is multiplied by its percent occurrence in nature – than the masses
of all isotopes are added to get the average atomic mass
C. The Mole and Molar Mass
1. measures the amount of substance.
a. 1 mole = 6.02 x 1023 particles ( atoms, molecules ions, electrons)
b. Standard – 1 mole is the number of atoms in 12g of C-12
2. molar mass – mass in grams of one mole(mol) of any substance
a.
b.
V.
Radioactive Decay of Elements
A. Process
1. What is radioactive decay?
2.
Why do atoms undergo radioactive decay?
B. Comparison of alpha, beta, gamma radiation
Alpha
Beta
Gamma
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Form
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Symbol
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Mass
______________________________________________________________________________
Charge
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Notation
C. Types of Decay
1. Beta Decay (neutron  proton + electron)
a.
b.
c.
2. Alpha Decay
a.
b.
c.
D. Examples of alpha and beta decam