1. CHAPTER 4 – ATOMIC STRUCTURE:

2. 4.1 DEFINING THE ATOM
Atom – the smallest particle of an element that retains the identity of the element in a chemical reaction.

3. Early Models:
a.) Democritus – atoms are invisible and indestructible his theory was conceptual not experimental.

4. b. Dalton – studied the ratios in which elements combine in chemical reactions. His theories reflected the thoughts of Democritus but were based on the results of his experiments

5. Dalton’s Atomic Theory
All elements are composed of tiny particles called atoms
b. Atoms of the same atom are identical and are different from all other elements.
c. Atoms of different elements can physically or chemically combine in simple whole number ratios to form compounds
d. Chemical reactions may separate, join, or rearrange atoms but never change the properties of the individual atoms

6. Sizing Up Atoms:
Copper is an example of an element – if you were to grind a pure copper penny down to it’s individual atoms you would have approximately 2.4 x 1022 atom. The earth has only about 6.9 x 109 people.
There are about 100,000,000 million copper atoms in 1cm
The radii of atoms varies but most are around 5 x to 2 x 10-10m
Individual atoms can be observed using a scanning tunneling microscope.
The ability to manipulate atoms in a nanoscale world opens science to many opportunities.

8. STRUCTURE OF THE NUCLEAR ATOM Subatomic particles:
Atoms are divisible!
The three subatomic particles are protons, neutrons, and electrons
Atomic Part
Location
Charge
Mass
Symbol
Change
Proton
Nucleus
Positive
1 a.m.u. p+
New atom
Neutron
None n0
Isotope
Electron
Orbit
Negative
1/1840 a.m.u. e-

9. J.J. Thomson ( ) – discovered the electron by passing electric current through gases at low pressure. The cathode ray, a glowing beam, travels from the negatively charged cathode to the positively charged anode. He then deflected the beam using a magnet.
How did he know the charge of the beam?

10. What did Thompson discover about the ratio of the charge of an electron to its mass?
He found the ratio to be constant and that the ratio did not depend on the type of gas or metal in the electrodes. He summarized then that electrons must be part of all atoms

11. Robert A Millikan – ( ) used the charge to mass ratio of an electron to calculate the mass of an electron. His value for charge and mass are similar to today’s accepted value.

12. Consider these four presuppositions;
Atoms are neutral, no net electric charge
Electric charges are carried by particles of matter
Charges exist in whole-number multiples, there are no fractions of a charge
A given number of negatively charged particles must combine with an equal number of positively charged particle to keep the particle neutral

13. Eugen Goldstein – ( ) observed rays in a cathode ray tube traveling in a direction opposite to that of the cathode ray. He called them canal rays

14. James Chadwick – (1891 – 1974) – confirmed the existence of the neutron

15. Ernest Rutherford – (1871-1937) a former student of Thompson gave shape to today’s image of an atom.
In 1911 his gold foil experiment yielded evidence of an atomic nucleus with a large area of space around surrounding that nucleus.
How did he come to this conclusion?

17. 4.3 DISTINGUISHING AMONG ATOMS
Atomic Number – is the number of protons in the nucleus of an atom.
Mass Number – equal the sum of the protons and neutrons in an atom

18. Give the number of protons, neutrons, and electrons for the following atoms

19. ISOTOPE: Atoms that have the same number of protons but different number of neutrons
Isotopes of Hydrogen

20. ATOMIC MASS: Although the actual masses of individual atoms are useful, it is more useful to compare the relative masses of atoms using a reference isotope as a standard. The isotope that was chosen is Carbon – 12 and was assigned a mass of exactly 12 atomic mass units.

21. Atomic Mass Unit: (amu) Is defined as one twelfth of the mass of a carbon – 12 atom.
Magnesium has three isotopes % magnesium 24 with a mass of amu, 10.00% magnesium 25 with a mass of amu, and the rest magnesium 26 with a mass of amu. What is the weighted average atomic mass of magnesium?
The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element and reflects both the mass and relative abundance of t isotopes of that atom as they occur in nature.
Weighted average mass = x amu x amu x amu
= 18.95amu amu amu
= 24.31amu

22. PERIODIC TABLE: Is an arrangement of elements organized in groups or rows with respect to similar traits shared by different elements.
Groups – Or families, are vertical columns of elements that have similar chemical and physical properties
Periods – There are 7 horizontal rows known as periods in which the properties of the atoms vary as you move across the table. Periods are associated with the energy levels of the element.