1. Early Models of the Atom

2. Defining the Atom
An atom is the smallest particle of an element that retains its identity in a chemical reaction.
Philosophers and scientists have proposed many ideas on the structure of atoms.

3. Defining the Atom Democritus’s Atomic Philosophy
How did Democritus describe atoms?
Democritus

4. Democritus (460 B.C B.C.)
first to suggest the existence of atoms
proposed all life and matter are made up of small particles called atoms
believed that atoms were indivisible and indestructible

5. Democritus (460 B.C. - 370 B.C.) dismissed by Aristotle and Plato
atom = Greek word “atomos” meaning indivisible
dismissed by Aristotle and Plato
idea was largely abandoned for 2000
years

6. John Dalton 1800 proposed modern atomic model
transformed Democritus’ ideas on atoms into a scientific theory through using experimental methods

7. John Dalton
based on his results, he came up with his atomic theory

8. Dalton’s Atomic Theory
Elements are composed atoms.
Atoms of the same element are identical.
Atoms of different elements different.
Atoms of different elements combine in simple, whole-number ratios to form compounds.
In a chemical reaction, atoms are combined, separated or rearranged but never created, destroyed or changed.

9. Law of Multiple Proportions
Dalton’s Atomic Theory (#4)
whenever 2 elements combine to form two or more compounds, the mass of one element that combines with a given mass of the other is in the ratio of small whole numbers
Water (H2O)
Peroxide (H2O2)

10. Law of Conservation of Mass
Dalton’s Atomic Theory (#5)
mass can not be created nor destroyed in a physical change or
chemical reaction

11. Law of Definite Proportions
Dalton’s Atomic Theory
in any chemical compound, the masses of the elements are always in the same proportions

12. Law of Definite Proportions
mass ratio does not change no matter how much of the sample or how it was formed

13. Parts of the Atom Proton: positive charge Neutron: Neutral charge
Electron: Negative charge
Objectives
Slide Menu

14. Discovery of the Electron
1897, discovered electrons
Cathode Ray Tube
stream of negatively charged particles
J. J. Thomson
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Website

15. Discovery of the Electron
charge to mass ratio constant
therefore electrons are part of all elements
Millikan in 1916
electron = -1 charge
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Website
Click on Picture to Access
Website

16. Protons discovered by Goldstein
observed in cathode-ray tube (ray traveling in opposite direction)
has a positive charge

17. Neutrons discovered by Chadwick in 1932
mass is equal to that of proton
neutral charge

18. Plum Pudding Model J.J. Thomson
electrons evenly distributed throughout positively charged material
resembles a fruit cake
disproved in 1909 by Rutherford

19. Gold-Foil Experiment

20. Gold-Foil Experiment 1911, Rutherford
beam of alpha particles shot at Au foil
should have passed through foil with slight deflection
majority not deflected
deflection at large angles
Discovery of NUCLEUS (add in)

21. Rutherford’s Atomic Model
atom is mostly empty space
positive charge and most of mass is in a small central region called the nucleus
nucleus contains both protons and neutrons

22. Rutherford’s Atomic Model
electrons are around the nucleus and make up the bulk of the atom
nucleus is tiny compared to size of atom

23. Size of Nucleus
If an atom is the size of a football stadium, the nucleus would be the size of a marble on the 50 yard line. The rest of the stadium is the area where an electron can be found.

24. Atomic Number Elements differ because of the number of protons
Atomic Number = number of protons
Atomic Number identifies an element

25. Atomic Number Atoms are electrically neutral
Number of protons must equal the number of electrons

26. Mass Number
Mass number = number of protons + number of neutrons

27. Mass Number Example: Helium Atom Atomic Number = 2 therefore,
Protons = 2
Neutrons = 2
Mass Number = 4

28. Mass Number Example: Oxygen Atom Atomic Number = 8 Mass Number = 16
Protons = 8 (= atomic number)
Electrons = 8
Neutrons = mass number – atomic no.
= 16 – 8 = 8

29. Nuclear Symbols

30. Isotopes same number of protons but different number of neutrons
chemically alike
different mass numbers

31. Isotopes Hydrogen 1 proton 1 electron Deuterium 1 proton 1 electron
1 neutron
Tritium
1 proton
1 electron
2 neutron

32. Atomic Mass Unit 1/12 mass of a carbon-12 atom
Carbon-12 chosen as reference
abbreviated by amu

33. Atomic Mass
weighted average mass of the atoms in a naturally occurring sample of the element
determined based on relative abundance of isotopes
mass of isotope x abundance = atomic mass

34. Atomic Mass
10X = amu and relative abundance of 19.91% : 11X = amu and relative abundance of 80.09%. Calculate the atomic mass of this element.
10X = amu x = amu
11X = amu x = amu
atomic mass = amu

36. Rutherford’s Atomic Model
The one thing that this model could not explain is…….
the chemical properties of elements.