1. The Structure of the Atom
Chapter 4
4.1 - Early Theories of the Atom
4.2 - Subatomic Particles
4.3 - How Atoms Differ
4.4 - Unstable Nuclei & Radioactivity
– CHEMISTRY 112

2. 4.1 Early Theories of Matter
Dalton’s Atomic Theory
All matter is composed of extremely small particles called atoms.
All atoms of a given element are identical. Atoms of different elements are different from one another.
Atoms cannot be created or divided into smaller particles or destroyed.
Different atoms combine in simple whole number ratios to form compounds.
In a chemical reaction, atoms are separated, combined, or rearranged.
Dalton’s Atomic Model
Atom - the smallest particle of an element that retains the properties of the element.

3. 4.1 Early Theories of Matter
E. Goldstein discovered the proton in 1886.
J.J. Thomson discovered the electron in 1897 during cathode ray tube experiments in the late 1890’s.

4. 4.1 Early Theories of Matter
Robert A. Millikan determined the mass and charge of the electron in 1916.
one unit of negative charge
mass is 1/1840 of a hydrogen atom

5. 4.1 Early Theories of Matter
In 1911 Ernest Rutherford discovered the nucleus during his gold foil experiment.

6. 4.1 Early Theories of Matter
Ernest Rutherford’s gold foil experiment.

7. 4.1 Early Theories of Matter
Neils Bohr developed the planetary model of the atom
Electrons are in a particular path have a fixed energy
Energy level-region around a nucleus where the electron is likely to be moving

8. 4.1 Early Theories of Matter
Erwin Schrodinger developed the Quantum Mechanical Model
Describes the electronic structure of the atom as the probability of finding electrons within certain regions of space

9. 4.1 Early Theories of Matter
James Chadwick discovered the neutron in 1932.
In 1913 Henry Mosley used X-rays to count the number of protons in the atomic nuclei of different atoms.

10. 4.2 Subatomic Particles & the Nuclear Atom
Located within the Nucleus
Proton (p+)
Positively charged particle (each carries a charge of +1)
Relative mass = 1 amu
Actual mass = X kg
Neutron (n0)
Neutrally charged particle
Actual mass = X kg
Serves as the glue that holds the nucleus together as well as a buffer between the charges of protons and electrons

11. Subatomic Particles Located outside the nucleus in the electron cloud
Negatively charged particle (each carries a charge of -1)
Relative mass = 1/1840 amu
Actual mass = 9.11 X kg
The electron is the part of the atom that will function in bonding and reactions

12. 4.3 How Atoms Differ Atomic Number O Mg
the number of protons in the nucleus of an atom
indicated at the top of the element’s block on the periodic table 8 O
15.999
Oxygen has an atomic number of 8
There are 8 protons in an atom of Oxygen 12 Mg
24.305
Magnesium has an atomic number of 12
There are 12 protons in an atom of Magnesium

13. Isotopes
Atoms of the same element with the same number of protons, but different numbers of neutrons
Since the atoms have different numbers of neutrons, they also have different mass numbers
Mass number = # of protons + # of neutrons

14. Abbreviating Isotopes
Hyphen Notation
Simply write the Name of the atom, put a hyphen, and then write the mass number
Carbon-12 vs. Carbon-14
Carbon 12 has 6 protons and 6 neutrons
Carbon 14 has 6 protons and 8 neutrons
Nuclear Designation
Element symbol is written in the center
Mass number goes in the upper left corner
Atomic number goes in the lower left corner 12 C 6

15. Different Isotopes Br Boron-10 Boron 11 35Cl 66Zn 17 30
Identify the number of protons, neutrons, and electrons each of the following have.
Boron-10
Boron 11
35Cl Zn 35 Br
79.904
p+: ________ no: ________ e-: ________
p+: ________ no: ________ e-: ________
p+: ________ no: ________ e-: ________
p+: ________
no: ________
e-: ________
p+: ________
no: ________
e-: ________

16. Calculating Atomic Mass
Mass Number
the number of protons + neutrons in a given isotope
Atomic Mass
The weighted average mass of all of the isotopes of that element
[(Mass of isotope A)(percent abundance )] + [(Mass of isotope B)(percent abundance)]

17. Practice Calculating Atomic Mass
Calculate the atomic mass of helium given the following information:
There are two naturally occurring isotopes of helium:
Isotope % Abundance Mass
helium
helium

18. Practice Calculating Atomic Mass
There are two naturally occurring isotopes of helium:
Isotope % Abundance Mass
helium
helium
( x ) + ( x ) = =

19. Practice Calculating Atomic Mass
There are three naturally existing isotopes of silicon: silicon-28, silicon-29, and silicon-30. Their percents of natural abundance is listed respectfully: %, 4.70 %, and 3.09 %.
Calculate the average atomic mass of silicon and express your answer in 4 significant digits.

20. 4.4 Unstable Nuclei & Radioactive Decay
Nuclear Reactions
reactions that involve a change in the nucleus of an atom.
Radioactivity
the spontaneous release of radiation.
Radiation
rays and particles emitted by radioactive materials
Radioactive atoms emit radiation because their nuclei are unstable.
There are three main types of radiation
Alpha decay
Beta decay
Gamma decay

21. Alpha (α) radiation 2 2 two protons and 2 neutrons Positive charge
Symbols:  4 4He
reduces the atomic number by 2
reduces the mass by 4

22. Beta (β) radiation Fast moving electron Negative charge Symbols:  0-1
increases the atomic number by 1
does not change the mass

23. Gamma (γ) radiation high energy radiation
released with alpha and beta radiation
symbol:  
does not change the mass or atomic number

24. So, it takes 6 years for the 60g sample to decay into 7.5g.
Half lives
The time it takes for 1/2 of the mass of the isotope to be decayed.
If I have a 60g sample and the half life is 2 years, how long will it take for there to be 7.5g left of the sample?
60g  30g  15g  7.5g
2 years years years
So, it takes 6 years for the 60g sample to decay into 7.5g.