1. Unit 2 - Lecture 1: Structure of the Atom
Atoms, Molecules, and Ions

2. A HISTORY OF THE STRUCTURE OF THE ATOM
Unit 2 - Lecture 1: Structure of the Atom
A HISTORY OF THE STRUCTURE OF THE ATOM

3. History Greek Philosopher Democritus (460-370 B.C.):
all matter composed of small atoms
atomos = indivisible
1803, John Dalton (brit.): atoms are the fundamental building blocks of matter

4. Dalton's Postulates
Each element is composed of extremely small particles called atoms.
Figure 2.1 John Dalton ( )

5. Dalton's Postulates
All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.
Figure 2.1 John Dalton ( )

6. Dalton's Postulates
Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
Figure 2.1 John Dalton ( )

7. Dalton’s Postulates
Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

8. John Dalton’s Atomic Theory (ca 1803)
Unit 2 - Lecture 1: Structure of the Atom
John Dalton’s Atomic Theory (ca 1803)
Each element is composed of extremely small particles called atoms.
All atoms of a given element are identical. The atoms of different elements are different and have different properties (including different masses).
Atoms of an element are not changed into different types of atoms by chemical reactions. Atoms are neither created nor destroyed in chemical reactions. This is the Law of Conservation of Mass.
Compounds are formed when atoms of more than one element combine. A given compound always has the same relative number and kind of atoms. This is the Law of Constant Composition. 10

9. Unit 2 - Lecture 1: Structure of the Atom
John Dalton’s Atomic Theory
Led him to deduce the Law of Multiple Proportions:
When two or more elements combine to form more than one compound, the relative masses of the elements which combine will be in in the ratio of small whole numbers.
In carbon monoxide, CO, 12 g carbon combine with
16 g oxygen. C:O ratio is 12:16 or 3:4.
In carbon dioxide, CO2, 12 g carbon combine with
32 g oxygen. C:O ratio is 12:32 or 3:8.

10. John Dalton’s Atomic Theory
Unit 2 - Lecture 1: Structure of the Atom
John Dalton’s Atomic Theory
Almost right. A good start.
very small
Structure of the atom after Dalton (ca. 1810)

11. Unit 2 - Lecture 1: Structure of the Atom
J.J. Thomson (1897):
Cathode Rays
Atoms subjected to high voltages give off cathode rays.

12. Unit 2 - Lecture 1: Structure of the Atom
J.J. Thomson: Cathode Rays
Cathode rays can be deflected by a magnetic field.
Cathode rays are negatively charged particles (electrons).
Electrons are in atoms.

13. Unit 2 - Lecture 1: Structure of the Atom
J.J. Thomson – The Electron
“Plum pudding” model: Negative electrons are embedded in a positively charged mass.
Electrons (-)
Unlike electrical charges attract, and that is what holds the atom together.
Positively charged mass
Structure of the atom after Thomson (ca. 1900)

14. Radioactivity
Radioactivity is the spontaneous emission of radiation by an atom.
First observed by Henri Becquerel
( ).
Marie and Pierre Curie also studied it.
Nobel Prize in 1903 (physics).

15. Unit 2 - Lecture 1: Structure of the Atom
Studies of Natural Radioactivity
Some atoms naturally emit one or more of the following types of radiation:
alpha (α) radiation (later found to be He2+ - helium nucleus)
beta (β) radiation (later found to be electrons)
gamma (γ) radiation (high energy light) α
Alpha particles
Electrons (-) γ γ
Positively charged mass α
Somehow gamma radiation is in there, too.
Structure of the atom after Becquerel (early 1900s)

16. Radioactivity
Three types of radiation were discovered by Ernest Rutherford:
 particles (positive, charge 2+, mass 7400 times of e-)
 particles (negative, charge 1-)
 rays (high energy light)
Figure 2.8

17. Unit 2 - Lecture 1: Structure of the Atom
Ernest Rutherford (1910)
Scattering experiment: firing alpha particles at a gold foil

18. The Nuclear Atom
Some alpha particles bounce off the gold foil. This means the mass of the atom must be concentrated in the center and is positively charged!
Thompson’s model could not be correct.
Figure 2.11

19. Unit 2 - Lecture 1: Structure of the Atom
Ernest Rutherford
The Nucleus and the Proton
The mass is not spread evenly throughout the atom, but is concentrated in the center, the nucleus.
The positively charged particles in the nucleus are protons.
Electrons (-) are now outside the nucleus.
Structure of the atom after Rutherford (1910)

20. Unit 2 - Lecture 1: Structure of the Atom
James Chadwick – The Neutron
In the nucleus with the protons are particles of similar mass but no electrical charge called neutrons.
Electrons (-) are now outside the nucleus in quantized energy states called orbitals. (From Niels Bohr and quantum mechanics)
The positively charged particles in the nucleus are protons. + n n
Structure of the atom after Chadwick (1932)

21. Unit 2 - Lecture 1: Structure of the Atom
proton (+)
neutron
electrons - responsible for the volume and size of the atom, negatively charged
10-10 m
nucleus - responsible for the mass of the atom, positively charged
10-14 m

22. Subatomic Particles
Protons and electrons are the only particles that have a charge.
Protons and neutrons have essentially the same mass.
The mass of an electron is so small we ignore it.
Table 2.1

23. Unit 2 - Lecture 1: Structure of the Atom
Atomic Facts
Feature
Size
Mass
1 amu = 1 atomic mass unit = x g
electron (-)
10-18 m ???
amu + n n
proton (+)
10-15 m
amu
Electrons are outside the nucleus in quantized energy states called orbitals.
neutron (0)
10-15 m
amu

24. Symbols of Elements
Elements are symbolized by one or two letters.

25. Atomic Number
All atoms of the same element have the same number of protons:
The atomic number (Z)

26. Atomic Mass
The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

27. Unit 2 - Lecture 1: Structure of the Atom
Atomic Number
Carbon atom
The number of protons in the nucleus is called the atomic number Z.
Z determines the identity of an element.
Saying “the atomic number of an element is 6” is the same as saying “carbon.”
The number of electrons in the atom is also Z (because atoms have no net electric charge).
How many neutrons are in C?
- proton
- neutron

28. Unit 2 - Lecture 1: Structure of the Atom
Isotopes A
12C Z 6
The number of protons and neutrons (nucleons) in an element is called the mass number A.
A = Z + number of neutrons.
An element may have different numbers of neutrons but NOT different numbers of protons.
Atoms of an element with different numbers of neutrons are called isotopes of that element.
- proton
- neutron
How many neutrons are in C? The answer is “it depends on the isotope.”

29. Isotopes
Isotopes are atoms of the same element with different masses.
Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C

30. Unit 2 - Lecture 1: Structure of the Atom
Isotopes
number of protons (Z)
number of neutrons 6 8 92
146
mass number (A)
number of electrons
symbol 12 6
12C or C-12 6 14 6
14C or C-14 6 16 8
16O or O-16 8
238 92
238U or U-238 92

31. Unit 2 - Lecture 1: Structure of the Atom
Atomic Masses
Atomic masses are based on 12C.
The mass of 12C (or C-12) is defined to be exactly 12 amu.

32. Unit 2 - Lecture 1: Structure of the Atom
Atomic Masses
The mass (weight) shown in the periodic table is the mass of the element as its occurs naturally.
If the element has more than one isotope, the mass shown is the weighted average of the masses of the isotopes.
Mg has 3 isotopes.
24Mg % amu
25Mg % amu
26Mg % amu
weighted average of Mg: x
0.1000x
0.1101x
24.31 amu
atomic weight of Mg based on natural abundance: amu

33. Unit 2 - Lecture 1: Structure of the Atom
Ions
Atoms can gain or lose electrons to become charged particles called ions.
A chemical particle that contains a positive or negative charge
Cations are positively charged ions.
Formed when an atom loses electrons
Anions are negatively charged ions.
Formed when an atom gains electrons

34. Unit 2 - Lecture 1: Structure of the Atom
Ions
Formation of a cation 1p e- +
Hydrogen atom
1p, 0 n, 1 e-
Hydrogen ion (cation)
1p, 0 n, 0 e-
1 H
1 H+
Net charge = 0
Net charge = +1

35. Unit 2 - Lecture 1: Structure of the Atom
Ions
Formation of an anion 8p 8n
8e- +
2e- 8p 8n
10e-
Oxygen atom
8p, 8 n, 8e-
Oxygen ion (anion)
8p, 8n, 10e-
16 O2-
16 O
Net charge = 0
Net charge = -2

36. Unit 2 - Lecture 1: Structure of the Atom
Nuclear Symbols
Mass
Number Charge
Atomic
Number X
Charge = # p - # e-

37. Unit 2 - Lecture 1: Structure of the Atom
Nuclear Symbols
Using nuclear symbols to determine the number of p, n, e, and total charge O 16 8
Mass Number =
Atomic Number = 16 8
# protons = atomic number = 8
# neutrons = Mass # - Atomic # = = 8
# electrons = # protons = 8

38. Unit 2 - Lecture 1: Structure of the Atom
Nuclear Symbols O 16 8 2-
Mass Number =
Atomic Number = 16 8
# protons = atomic number = 8
# neutrons = Mass # - Atomic # = = 8
# electrons = # protons - charge = 8 - (-2) = 10

39. Unit 2 - Lecture 1: Structure of the Atom
Nuclear Symbols Ba
137 56 2+
Mass Number =
Atomic Number =
137 56
# protons = atomic number = 56
# neutrons = Mass # - Atomic # = = 81
# electrons = # protons - charge = 56 - (+2) = 54

40. Nuclear Symbols - Atoms
Unit 2 - Lecture 1: Structure of the Atom
Nuclear Symbols - Atoms
Example: Write the nuclear symbol for the following atoms:
1) 50 p, 70 n
2) 17 e-, 20 n
120Sn 50
37Cl 17

41. Unit 2 - Lecture 1: Structure of the Atom
Nuclear Symbols - Ions
Practice writing nuclear symbols from information given:
53 p, 74 n, 54 e-
53 proton (= atomic number)  I
74 neutrons + 53 proton  mass number = 127
54 electrons (one more than protons)  1-
127I1- 53

42. Fe 2) 23 e-, 30 n, net charge = +3 # protons?
23 electrons, but charge of 3+
ie 3 more protons than electrons  p= 26
 Atomic number = 26  element = Fe 56 3+ Fe 26