1. Ch. 5: Atomic Structure The Theory of the Atom
________________, a famous Greek teacher who lived in the 4th Century B.C., first suggested the idea of the atom.
________ __________ came up with his atomic theory based on the results of his experiments. (See Figure 5.1)
The Atom
The smallest particle of an ________________ is an atom.
The atom is made up of three ________________ particles.
The electron was discovered in _______ by J. J. Thomson. The electron has a _______ charge. It’s mass is much smaller than the other 2 subatomic particles, therefore it’s mass is usually ______________.
(2) The proton has a ______ charge, and it was discovered in _________ by E. Goldstein.
Democritus
John Dalton
element
subatomic
1897
(-)
ignored
(+)
1886

2. Nuclear Atomic Structure
(3) The neutron does not have a charge. In other words, it is ________. It was discovered in ____ by James Chadwick. The neutron has about the same _________ as the proton.
These three particles make up all the ____________________ in the Universe!
Nuclear Atomic Structure
The atom is made up of 2 parts/sections:
(1) The ______________ --- (in the center of the atom)
(2) The ____________ _________ --- (surrounds the nucleus)
neutral
1932
mass
visible matter
nucleus
electron cloud
nucleus (p+ & n0)
e- cloud

3. The Nucleus
Discovered by Ernest ________________ in ________.
He shot a beam of positively charged “alpha particles”, which are ___________ nuclei, at a thin sheet of ______ _____.
Rutherford
1911
helium
gold foil
99.9% of the particles went right on through to the ______________.
Some were slightly deflected. Some even ____________ ________ towards the source!
This would be like shooting a cannon ball at a piece of tissue paper and having it bounce off.
detector
bounced back

4. Rutherford’s Experiment

5. Conclusions about the Nucleus
(1) Most of the atom is more or less _________ ___________.
(2) The nucleus is very _________. (Stadium Analogy)
(3) The nucleus is very ___________. (Large Mass ÷ Small Volume)
(4) The nucleus is ______________ charged.
empty space
tiny
dense
positively
Counting Subatomic Particles in an Atom
The atomic # of an element equals the number of ____________ in the nucleus.
The mass # of an element equals the sum of the _____________ and ______________ in the nucleus.
In a neutral atom, the # of protons = # of ______________.
To calculate the # of neutrons in the nucleus, ______________ the ___________ # from the __________ #.
protons
protons
neutrons
electrons
subtract
atomic
mass

6. The element is phosphorus!
Practice Problems
Find the # of e-, p+ and n0 for sodium. (mass # = 23)
Find the # of e-, p+ and n0 for uranium. (mass # = 238)
3) What is the atomic # and mass # for the following atom? # e- = 15; # n0 = 16
Atomic # = 11 = # e- = # p+
# neutrons = = 12
Atomic # = 92 = # e- = # p+
# neutrons = = 146
Atomic # = 15 = # e- = # p+
Mass # = p+ + n0 = =31
The element is phosphorus!

7. (The # shown after the name is the mass #.)
Isotopes
An isotope refers to atoms that have the same # of ___________, but they have a different # of ___________.
Because of this, they have different _________ #’s (or simply, different ___________.)
Isotopes are the same element, but the atoms weigh a different amount because of the # of ______________.
Examples---> (1) Carbon-12 & Carbon-13
(2) Chlorine-35 & Chlorine-37
(The # shown after the name is the mass #.)
For each example, the elements have identical ___________ #’s, (# of p+) but different _________ #’s, (# of n0).
Another way to write the isotopes in shorthand is as follows:
protons
neutrons
mass
masses
neutrons
atomic
mass 12 C 35 Cl 6 17
The top number is the ________ #, and the bottom # is the __________ number. Calculating the # n0 can be found by _____________ the #’s!
mass
atomic
subtracting

8. More Practice Problems
Find the # e-, p+ and n0 for Xe-131.
Find the # e-, p+ and n0 for
3) Write a shorthand way to represent the following isotope:
# e- = # n0 = 0 # p+ = 1
Atomic # = 54 = p+ = e n0 = = 77 63 Cu 29
Atomic # = 29 = p+ = e n0 = = 34
Atomic # = p+ = e- = mass # = n0 + p+ = 1+ 0 = 1
H-1 or 1 H 1

9. Atomic Mass Weighted Average
Based on the relative mass of Carbon-12 which is exactly _______.
1 p+ ≈ __ atomic mass unit (amu) 1 n0 ≈ __ amu 1e- ≈ __ amu
The atomic masses listed in the Periodic Table are a “weighted average” of all the isotopes of the element. 12 1 1
Weighted Average
Practice Problems:
(1) Mrs. Smith’s geometry semester grades are calculated using a weighted average of three category scores:
Major Grades= 60% of your grade
Minor Grades= 30% of your grade
Semester Exam=10% of your grade
If a student had the following scores, what would they receive for the semester?
Major= 80 (B-)
Minor= 60 (D-)
Semester Exam=65 (D)

10. Weighted Average
Step (1): Multiply each score by the % that it is weighted.
Step (2): Add these products up, and that is the weighted average!
60% x 80 = 48.0
30% x 60 = 18.0
10% x 65 = 6.5
Add them up!!
A “normal average” would be calculated by simply adding the raw scores together and dividing by 3…
= 205 ÷ 3 = 68.3 = D +
72.5 (C)

11. Cl-35 (75.8% abundance) Cl-37 (24.23 % abundance)
Weighted Average
Practice Problems:
(2) In chemistry, chlorine has 2 isotopes:
Cl-35 (75.8% abundance) Cl-37 (24.23 % abundance)
What is the weighted average atomic mass of chlorine?
35 x =
37 x =
Add them up!!!
(3) Oxygen has 3 isotopes:
O-16 (99.76%) O-17 (0.037%) O-18 (0.2%)
Estimate oxygen’s average atomic mass.
Barely over 16.0 amu. +
amu

13. IONS
IONS are atoms or groups of atoms with a positive or negative charge.
Taking away an electron from an atom gives a CATION with a positive charge
Adding an electron to an atom gives an ANION with a negative charge.
To tell the difference between an atom and an ion, look to see if there is a charge in the superscript! Examples: Na+ Ca+2 I- O-2
Na Ca I O

14. Forming Cations & Anions
A CATION forms when an atom loses one or more electrons.
An ANION forms when an atom gains one or more electrons
F + e- --> F-
Mg --> Mg e-

15. PREDICTING ION CHARGES
In general
metals (Mg) lose electrons ---> cations
nonmetals (F) gain electrons ---> anions

16. Learning Check – Counting
State the number of protons, neutrons, and electrons in each of these ions.
39 K O Ca +2
#p+ ______ ______ _______
#no ______ ______ _______
#e- ______ ______ _______

17. Charges on Common Ions -3 -2 -1 +1 +2
By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom.

18. The History of Chemistry
A NOT So Very Brief Summary
Of The History of Atomic Discovery

19. Defining the Atom 460 B.C. – 370 B.C.
The Greek philosopher Democritus was among the first to suggest the existence of atoms (from the Greek word “atomos”).
He believed that atoms were indivisible and indestructible.
His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method – just philosophy

20. Dalton’s Atomic Theory – Summary based on experimental data- not philosophy.
matter is composed, indivisible particles (atoms)
all atoms of a particular element are identical
different elements have different atoms
atoms combine in certain whole-number ratios
In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements.

21. Problems with Dalton’s Atomic Theory?
1. matter is composed, indivisible particles
Atoms Can Be Divided, but only in a nuclear reaction
2. all atoms of a particular element are identical
Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)!
3. different elements have different atoms
YES!
4. atoms combine in certain whole-number ratios
YES! Called the Law of Definite Proportions
5. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements.
Yes, except for nuclear reactions that can change atoms of one element to a different element

22. How the Electron was Discovered
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron.
In 1916, Robert Millikan determined the mass of an electron
Using the Oil Drop Apparatus he found an electron contains 1/1840 the mass of a hydrogen atom and that it has one unit of negative charge.
Mass of an electron = 9.11X10-28 grams or Zero amu.

23. How the Electron was Discovered Continued:

24. Conclusions from Thomson’s and Millikan’s Experiments:
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.

25. The Discovery of Protons and Neutrons
In 1886, Eugene Goldstein observed what is now called the proton.
Protons are particles with a positive charge, and a relative mass of 1 amu(or 1840 times that of an electron).

26. The Discovery of Protons and Neutrons Continued
In 1932, James Chadwick confirmed the existence of the neutron.
A neutron is a particle with no charge with a mass nearly equal to a proton (about 1 amu).

27. Thomson’s Atomic Model
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

28. Ernest Rutherford’s Gold Foil Experiment
Rutherford did not agree with Thomson’s “Plum Pudding Model”. To disprove Thomson’s theory he designed the Gold Foil Experiment.
In this experiment he shot alpha particles – positively charged helium nuclei- at a thin sheet of gold foil.

29. Gold Foil Experiment Continued
The gold foil was surrounded by a detecting screen which allowed Rutherford to record where each of the alpha particles hit the screen.
Rutherford’s results confirmed his suspicions- Thomson’s Plum Pudding Model of the atom was wrong!

30. Summary: Rutherford’s Findings
Most of the alpha particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
Conclusions:
The nucleus is small.
The nucleus is dense.
The nucleus is positively charged.

31. The Rutherford Atomic Model
Based on his experimental evidence:
The atom is mostly empty space.
All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “nucleus”.
The nucleus is composed of protons and neutrons (they make the nucleus!)
The electrons are distributed around the nucleus and occupy most of the volume.
His model was called a “nuclear model”.

32. Neils Bohr: Creator of the Bohr Model
In 1913, Neils Bohr, a Danish physicist came up with a new atomic model.
He proposed that electrons are arranged in concentric circles, or orbits, around the nucleus of the atom.

33. Bohr Continued
To explain why the negatively charged electrons do not fall into the positively charged nucleus, Bohr suggested that electrons in a specific orbit have a specific energy.
The ENERGY LEVEL of an electron is the region around the nucleus where the electron is likely to be moving. It also keeps the electrons from falling into the nucleus.

34. The Bohr Model
In the Bohr model, the energy levels closest to the nucleus have the least amount of energy.
The energy levels furthest from the nucleus have the least amount of energy.
Think of the Bohr Model as a track. You would need MORE energy if you were to run around the outside than if you were to run around the inner circle because the outside track is much longer.

35. Subatomic Particles 1 amu Particle Charge Mass (g) Location Electron-1
9.11 x 10-28
ZERO amu
Electron cloud
Proton (p+) +1
1.67 x 10-24
1 amu
Nucleus
Neutron
(no)

36. Atomic Models Timeline