1. Chapter 5: Atomic Structure & The Periodic Table
Democritus– 4th century B.C., teacher in Greece, first suggested the existence of atoms, lacked experimental support because scientific testing was unknown at the time.
2000 years after Democritus, the real nature of atoms and observable changes at the atomic level were established.
John Dalton ( )—English school teacher, performed experiments to test and correct his atomic theory.

2. Dalton’s Atomic Theory
All elements are composed of tiny indivisible particles called atoms. (*Now known to be divisible—broken down into subatomic particles)
Atoms of the same element are identical. Atoms of any one element are different from those of any other element.
Atoms of different elements can physically mix together or can chemically combine with one another in simple whole-number ratios to form compounds

3. Chemical Rxs. occur when atoms are separated, joined, or rearranged
Chemical Rxs. occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical rx.
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Atom—smallest particle of an element that retains the properties of that element
( a scanning tunneling microscope can be used to view the surface of individual atoms. Ex: pg 108 with gold atoms)

4. Subatomic Particles Electrons—negatively charged subatomic particles.
J.J. Thomson ( ) discovered electrons in 1897
performed experiments that involved passing electric current through gases at low pressure. The gases were sealed in glass tubes fitted at both ends with metal disks called electrodes. (pg 109 apparatus used)
Cathode-ray tube, the electrons travel as a ray from the cathode(-) to the anode(+)

5. Robert A. Millikan ( )
Mass of e- is 1/1840 the mass of a hydrogen atom (“proton”)
Electrons are negatively charged (1-)
E. Goldstein (1886) observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays.
(+ charge)
Protons—positivley charged particle
x that of an electron

6. James Chadwick (1891-1974) English physicist, 1932 discovered/confirmed neutrons
Neutrons—neutral particle
mass= to that of a proton

7. Properties of Subatomic Particles
Symbol
Relative electric charge
Relative mass
Actual mass (g)
Electron e- 1-
1/1840
9.11 x
Proton p+ 1+ 1
1.67 x
10-24
Neutron n0

8. Atomic Nucleus How are subatomic particles arranged in an atom:
1911- Ernest Rutherford ( ) tested the theory of Atomic Structure
Used massive alpha particles—helium atoms that have lost their 2e- and have a double + charge because of the two remaining protons

9. Rutherford’s Gold-foil Experiment
Rutherford directed a narrow beam of alpha particles at a very thin sheet of gold foil.
Alpha particles passed straight through the gold atoms without deflection
Some of the alpha particles bounced off the gold foil at very large angles.

10. Rutherford’s Theory of the Atom
Atom is mostly empty space (explaining the lack of deflection of most of the alpha particles)
All the positive charge and almost all the mass is concentrated in a small region (nucleus). (accounts for the great deflection of some of the alpha particles)
Nucleus- the central core of an atom composed of protons and neutrons. Tiny compared to the atom as a whole. Contains most of the atoms mass.

11. Atomic Number Atomic number= the number of protons in the nucleus
Identifies and element
Elements are different because they contain different number of protons
Number protons=number electrons (atoms electrically neutral

12. Mass Number Mass number= number of protons + number of neutrons
If you know the atomic number and the mass number of an atom of any element, you can determine the atom’s composition
Number of neutrons= mass number- atomic number

13. Isotopes
Atoms that have the same number of protons but different number of neutrons
Different mass numbers
Chemically alike because they have the same number of protons and electrons, which are the subatomic particles responsible for chemical behavior
Ex: Carbon-12, Carbon-14
Neon-20, Neon-21, Neon-22

14. Atomic Mass Atomic mass unit (amu)– 1/12 the mass of a Carbon-12 atom
Compare the relative masses of atoms using a reference isotope as a standard.
C-12 was assigned a mass of exactly 12 atomic mass units
He-4 atom with a mass of amu, 1/3 the mass of a C-12 atom
Ex: how many C-12 atoms would the same mass as a Nickel-60 atom?

15. Answer: 5 C-12 atoms = the mass of 1 Nickel-60 atom
Atomic masses are not in whole numbers
Ex: Cl amu=
relative abundance of the naturally occurring isotopes of the element.
In nature, most elements occur as a mixture of two or more isotopes. Each isotope of an element has a fixed mass and a natural percent abundance

16. Atomic mass—a weighted average mass of the atoms in a naturally occurring sample of the element
Weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature
Ex: Element X has two natural isotopes. The isotope with a mass of amu has a relative abundance of 19.91%. The isotope with a mass of amu has a relative abundance of 80.09%. Calculate the atomic mass of this element.
Solve: ( amu x ) + ( amu x ) = amu

17. Development of the Periodic Table
Dmitri Mendeleev
Russian Chemist
Constructed the first periodic table
Periodic table- an arrangement of the elements according to similarities in their properties.
Listed in columns by increasing atomic mass

18. Henry Moseley
British physicist
Determined the atomic number of the atoms of the elements
Arranged elements in a table by order of atomic number instead of atomic mass
Today’s periodic table
Each element is identified by its symbol placed in a square

19. Periods- horizontal rows of periodic table
number of elements/period ranges
from 2 to 32
Properties of elements within a period change as you move across from element to element
Periodic Law– when the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties

20. Group or Family– each vertical column of elements
Elements in any group or family have similar physical and chemical properties
Groups 1A—8A elements are called the representative elements (have a wide range of physical and chemical properties)
Groups 1B—8B elements in the middle of the table are called transition elements (metals)
Bottom rows of elements under the main table are called the inner transition elements (metals)
(aka rare-earth elements)
80% of all elements are metals, which are solids at room temp. except for Hg (Mercury), a liquid

21. Group 1A- alkali metals
Group 2A- alkaline earth metals
Groups 3A-8A are nonmetals
Bordering the black line that divides metals/nonmetals are metalloids
Metalloids are elements with properties that are intermediate between those of metals and nonmetals
Left side of the periodic table, except for Hydrogen are metals
Upper right corner of periodic table are nonmetals

22. Metals:
High electrical conductivity
High luster when clean
Ductile (can be drawn into wires)
Malleable (able to be beaten into thin sheets)
Nonmetals:
Poor conductors of electricity
Non lustrous
Brittle

23. Periodic Table