1. Chapter 4
Atomic Theory

2. Democritus
First person to purpose that matter was not infinitely divisible
“atomos”

3. Aristotle Rejected atomic theory
did not believe in “nothingness” of space

4. John Dalton (1766-1844) Modern atomic theory – 19th century
Conservation of mass – separation, combination, or rearrangement of atoms
Was all of his theory accurate? NO!!

5. Dalton’s Atomic Theory (1808)
All matter is composed of extremely small particles called atoms
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties
Atoms cannot be subdivided, created, or destroyed
Atoms of different elements combine in simple whole-number ratios to form chemical compounds
In chemical reactions, atoms are combined, separated, or rearranged

6. Defining the Atom
The smallest particle of an element that still retains the properties of the element
How small is an atom?
World population:
Atoms in a penny:
Scanning tunneling microscope – allows individual atoms to be seen

7. Cathode Ray Tube Discovered Cathode Ray Tube
Used Cathode Ray tube to discover Electron
William Crookes
JJ Thomson

8. Discovery of the Electron
1897: J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

9. Some Modern Cathode Ray Tubes

10. Cathode Ray Results
Cathode ray deflected in a magnetic field, indicated charged particles
Deflected towards positively charged plate, indicating particles must have negative charge
Altering gas, altering material had no effect on results, so particles must be in all matter
Called…. ELECTRONS!!!!
First subatomic particles!
Meant Dalton was… wrong!!!

11. Mass and charge of the Electron
1909 – Robert Millikan determines the mass of the electron.
Mass of the electron was much smaller than that of the hydrogen atom, the smallest known atom
Meant atoms were divisible into subatomic particles
Mass = 9.1 x = 1/1840 mass of hydrogen
Charge = -1
An experiment performed by Robert Millikan in 1909 determined the size of the charge on an electron. He also determined that there was a smallest 'unit' charge, or that charge is 'quantized'. He received the Nobel Prize for his work. We're going to explain that experiment here, and show how Millikan was able to determine the size of a charge on a single electron.
What Millikan did was to put a charge on a tiny drop of oil, and measure how strong an applied electric field had to be in order to stop the oil drop from falling. Since he was able to work out the mass of the oil drop, and he could calculate the force of gravity on one drop, he could then determine the electric charge that the drop must have. By varying the charge on different drops, he noticed that the charge was always a multiple of -1.6 x C, the charge on a single electron. This meant that it was electrons carrying this unit charge. Here's how it worked. Have a look at the apparatus he used:
An atomizer sprayed a fine mist of oil droplets into the chamber. Some of these tiny droplets fell through a hole in the upper floor. Millikan first let them fall until they reached terminal velocity. Using the microscope, he measured their terminal velocity, and by use of a formula, calculated the mass of each oil drop. Next, Millikan applied a charge to the falling drops by illuminating the bottom chamber with x-rays. This caused the air to become ionized, and electrons to attach themselves to the oil drops. By attaching a battery to the plates above and below this bottom chamber, he was able to apply an electric voltage. The electric field produced in the bottom chamber by this voltage would act on the charged oil drops; if the voltage was just right, the electromagnetic force would just balance the force of gravity on a drop, and the drop would hang suspended in mid-air. Now you try it. Click here to open a simulation of Millikan's chamber. First, allow the drops to fall. Notice how they accelerate at first, due to gravity. But quickly, air resistance causes them to reach terminal velocity. Now focus on a single falling drop, and adjust the electric field upwards until the drop remains suspended in mid-air. At that instant, for that drop, the electric force on it exactly equals the force of gravity on it. Some drops have more electrons than others, so will require a higher force to stop. When you've finished playing with the apparatus, close the window and we'll continue. O.K., let's look at the calculation Millikan was now able to do. When a drop is suspended, its weight  m · g is exactly equal to the electric force applied  q · E
The values of  E, the applied electric field,  m the mass of a drop, and  g, tha acceleration due to gravity, are all known values. So you can solve for  q, the charge on the drop:
Millikan determined the charge on a drop. Then he redid the experiment numerous times, each time varying the strength of the x-rays ionizing the air, so that differing numbers of electrons would jump onto the oil molecules each time. He obtained various values for  q. The charge  q on a drop was always a multiple of -1.6 x C,  the charge on a single electron. This number was the one Millikan was looking for, and it also showed that the value was quantized; the smallest unit of charge was this amount, and it was the charge on a single electron.

12. Conclusions from the Study of the Electron
Electrons are negative.
Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.
Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons
Electrons have so little mass that atoms must contain other particles that account for most of the mass

13. Thomson’s Plum Pudding Model
Atom breakable!!
Atom has structure
Electrons suspended in a positively charged electric field
must have positive charge to balance negative charge of electrons and make the atom neutral
mass of atom due to electrons
atom mostly “empty” space
compared size of electron to size of atom
Cookie dough model 9

14. Plum Pudding Model

15. Rutherford’s Gold Foil Experiment
Alpha particles are positively charged
Particles were fired at a thin sheet of gold foil
Particle hits on the detecting screen (film) are recorded

16. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
GREATLY Deflected particles were repulsed by positive charge of nucleus
Conclusions: Nuclear Model

17. Rutherford’s Nuclear Model
The atom contains a tiny dense center called the nucleus
the volume is about 1/10 trillionth the volume of the atom
The nucleus is essentially the entire mass of the atom
The nucleus is positively charged
the amount of positive charge of the nucleus balances the negative charge of the electrons
The electrons move around in the empty space of the atom surrounding the nucleus 12

18. Protons and neutrons Rutherford - Protons Chadwick – Neutrons
Subatomic particle in nucleus
protons :+1 charge (equal, opposite of electrons)
Chadwick – Neutrons
Mass nearly equal to a proton, but carries no electrical charge

19. Atoms are indivisible! What!?!
Actually, Mr. Dalton, we have proved that part of your theory wrong.
Atoms can be divided into electrons, protons, and neutrons.

20. The Structure of the Atom
Atom – electrically neutral particle composed of protons, neutrons, electrons
Spherical shape
Atoms consist of two regions
Nucleus – 99.7% of mass
Very small, dense region in the center.
Contains protons & neutrons.
Electrons Cloud
Mainly empty space surrounding nucleus
Very large compared to the nucleus.
Contains electrons.
Subatomic particles
Protons, neutrons, and electrons

21. Atomic Particles Particle Charge Mass (kg) Location Electron -1
9.109 x 10-31
Electron cloud
Proton +1
1.673 x 10-27
Nucleus
Neutron
1.675 x 10-27

22. Atomic Number
Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Identifies the atom.
Element
# of protons
Atomic # (Z)
Carbon 6
Phosphorus 15
Gold 79

23. Isotopes Elements occur in nature as mixtures of isotopes.
Isotopes are atoms of the same element that differ in the number of neutrons

24. Isotopes…Again (must be on the test) Hydrogen-2 (deuterium)
Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Nucleus
Hydrogen–1 (protium) 1
Hydrogen-2 (deuterium)
Hydrogen-3 (tritium) 2

25. Mass Number Mass # = p+ + n0
Mass number is the number of protons and neutrons in the nucleus of an isotope.
Mass # = p+ + n0
Nuclide p+ n0 e-
Mass #
Oxygen - 10 - 33 42
- 31 15 18 8 8 18
Arsenic 75 33 75
Phosphorus 16 15 31

26. Atomic Masses Carbon = 12.0125 amu
Atomic mass is the average of all the naturally isotopes of that element.
On Periodic Table
Carbon = amu
Isotope
Symbol
nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
7 neutrons
1.11%
Carbon-14
14C
8 neutrons
<0.01%

27. Practice Problems
P. 104 # 15-17

28. Writing Nuclear Symbols3 He
Mass #
Atomic Symbol
Atomic # 2
How many protons, electrons, and neutrons?
2 protons, 2 electrons, 1 neutron
Mass # - Atomic # = # Neutrons

29. Writing Isotopes Using Hyphen Notation
Uranium-235, Helium-3, or Carbon-14
How many proton, electrons, neutrons?
Name
of atom
Mass #
92 protons, 143 neutrons, 92 electrons

30. Convert these hyphen notation to nuclear symbols.
Isotope problems
Convert these hyphen notation to nuclear symbols.
Uranium-235, Helium-3, or Carbon-14
235 U 3 He 14 C 92 2 6