Download Chapter 4 Atomic Structure Notes

Chapter 4
I
Atomic Structure
Atoms
A. Democritus (@ 4BC Greece)
1. matter is composed of tiny indivisible particles called atoms
2. indivisible, indestructible, fundamental units of matter
3. philosophized but did not experiment
B. John Dalton( 1766-1844 England)
1. Experimented to arrive at this theory
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all elements are composed of submicroscopic, indivisible particles called
atoms
atoms of the same element are the same
atoms of different elements are different
atoms of different elements can combine in simple whole numbers to form
compounds
in chemical reactions atoms are separated, joined or rearranged, BUT are
not changed into different atoms
C. Atoms are the smallest particle of an element that retains the properties of that
element.
II
Electron, Protons, Neutrons
A. Electrons - negatively charged particle
- discovered in 1897 by an English physicist Sir J.J. Thomson
- American scientist Robert A. Milikan determined the value of the charge
and the mass (1916)
B. Neutron – the existence of a neutral particle was confirmed in 1932 by the
English physicist Sir James Chadwick
- its mass nearly equals that of the proton
C. Proton – was first observed in 1886 by E. Goldstein
- A particle with one unit of a positive charge
*** Atoms are electrically neutral***
III
Structure of the Atom
In 1911, Ernest Rutherford performs the gold foil experiment which determined:
A. Atoms consist of a small positively charged dense nuclei
B. Almost all of the mass is in the nucleus
C. The negatively charged electrons are in the area that creates most of the volume
of the atom.
IV
Atomic No. is the number of protons in the nucleus of an atom.
*** Since atoms are electrically neutral, the number of positive protons, must equal
the number of negative electrons. ****
V
Mass No. is the total number of protons and neutrons in the nucleus of an atom.
*** # of Neutrons = Mass No. – Atomic No.
# of Neutrons = Mass No. - # of Protons
Shorthand
Superscript
12
C
Mass No.
6
Subscript
Atomic No.
Ex. Calculate protons, neutrons & electrons for carbon & sodium
12
C
6
p=6
n=6
e=6
VI
Isotopes are atoms of the same element ( therefore the same number of protons)
But differ in the number of neutrons
Ex.
Carbon - 12
and
p=6
n=6
e=6
Carbon – 14
p=6
n=8
e=6
VII
Atomic mass is the average mass of all naturally occurring samples of an element.
(It reflects the mass and relative abundance)
VIII
Calculating Atomic Mass
A. Steps:
1. You must know: a. the number (no.) of stable isotopes
b. the mass of each isotope
c. the natural abundance of each isotope (%)
2. Find the mass that each isotope contributes to the weighted sample by
multiplying the mass by its relative abundance.
3. Add the products.
IX
The Periodic Table
A. Dimitri Mendeleev (1869) (Russian) – publishing the 1st periodic table based on
increasing atomic mass no.
1. The elements fell into 7 columns based on chemical & physical properties
2. He left spaces for undiscovered elements
B. Henry Mosely (1913) (British) publishes the “modern” periodic table based on
increasing atomic no..
1. Periodic Law – states that the properties of the elements repeat in a regular
way if the elements are arranged by increasing atomic no..
(columns)2. Groups (Families) – elements have similar chemical and physical properties
therefore they have the same no. of outer e-(s)
(Rows) 3. Periods – have the same principal energy level
C. Glenn Seaborg (1944) (American) moves the heaviest element (of the time) to the
bottom of the table.
D. Families
1/IA = Alkali Earth Metals
2/IIA = Alkaline Earth Metals
3-12 /B’s = Transition Metals***
13/IIIA = Boron’s Family
14/IVA = Carbon’s Family
15/VA = Nitrogen’s Family
16/VIA = Oxygen’s Family
17/VIIA = Halogen’s Family
18/VIIIA = Nobel Gases
(* Lanthanide & Actinide = Inner Transition
Metals)
E. Stepline – divides the table
1. Left – Metals (EXCEPT for Hydrogen)
2. Right – Nonmetals
3. Bordered on Two sides – Metalloids