1. Unit 3: “Atomic Structure”
Chemistry: Mr. Blake

2. Section 4.1 Defining the Atom
The Greek philosopher Democritus (460 B.C. – 370 B.C.) 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 – but just philosophy

3. Dalton’s Atomic Theory (experiment based!)
All elements are composed of tiny indivisible particles called atoms
Atoms of the same element are identical. Atoms of any one element are different from those of any other element.
John Dalton
(1766 – 1844)
Atoms of different elements combine in simple whole-number ratios to form chemical compounds
In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.

4. Problems with Dalton’s Atomic Theory?
1. matter is composed of 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)! Different elements have different atoms. YES!
3. atoms combine in certain whole-number ratios
YES! Called the Law of Definite Proportions
4. 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

5. Sizing up the Atom
Elements are able to be subdivided into smaller and smaller particles – these are the atoms, and they still have properties of that element
If you could line up 100,000,000 copper atoms in a single file, they would be approximately 1 cm long
Despite their small size, individual atoms are observable with instruments such as scanning tunneling (electron) microscopes

6. Section 4.2 Structure of the Nuclear Atom
One change to Dalton’s atomic theory is that atoms are divisible into subatomic particles:
Electrons, protons, and neutrons are examples of these fundamental particles
There are many other types of particles, but we will study these three

7. Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron

8. Modern Cathode Ray Tubes
Television
Computer Monitor
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

9. Mass of the Electron
Mass of the electron is
9.11 x g
The oil drop apparatus
1916 – Robert Millikan determines the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge

10. Conclusions from the Study of the Electron:
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

11. Conclusions from the Study of the Electron:
Eugen Goldstein in 1886 observed what is now called the “proton” - particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron)
1932 – James Chadwick confirmed the existence of the “neutron” – a particle with no charge, but a mass nearly equal to a proton

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

13. Electron cloud
Nucleus

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

15. Ernest Rutherford’s Gold Foil Experiment - 1911
Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil
Particles that hit on the detecting screen (film) are recorded

16. Rutherford’s problem:
In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?
Target #2
Target #1

17. The Answers:
Target #1
Target #2

18. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The nucleus is small
The nucleus is dense
The nucleus is positively charged

19. 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 distributed around the nucleus, and occupy most of the volume
His model was called a “nuclear model”

20. # protons in an atom = # electrons
Atomic Number
Atoms are composed of identical protons, neutrons, and electrons
How then are atoms of one element different from another element?
Elements are different because they contain different numbers of PROTONS
The “atomic number” of an element is the number of protons in the nucleus
# protons in an atom = # electrons

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

22. Mass Number
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 75
Arsenic 75 33
Phosphorus 16 15 31

23. Complete Symbols
Contain the symbol of the element, the mass number and the atomic number. X
Mass
number
Atomic
Subscript →
Superscript →

24. Br Symbols 80 35 Find each of these: number of protons
number of neutrons
number of electrons
Atomic number
Mass Number Br 80 35

25. Symbols
If an element has an atomic number of 34 and a mass number of 78, what is the:
number of protons
number of neutrons
number of electrons
complete symbol

26. Symbols
If an element has 91 protons and 140 neutrons what is the
Atomic number
Mass number
number of electrons
complete symbol

27. Symbols
If an element has 78 electrons and 117 neutrons what is the
Atomic number
Mass number
number of protons
complete symbol

28. Isotopes
Dalton was wrong about all elements of the same type being identical
Atoms of the same element can have different numbers of neutrons.
Thus, different mass numbers.
These are called isotopes.

29. Isotopes
Frederick Soddy ( ) proposed the idea of isotopes in 1912
Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons.
Soddy won the Nobel Prize in Chemistry in 1921 for his work with isotopes and radioactive materials.

30. We can also put the mass number after the name of the element:
Naming Isotopes
We can also put the mass number after the name of the element:
carbon-12
carbon-14
uranium-235

31. Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons.
Protons
Electrons
Neutrons
Nucleus
Hydrogen–1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2

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

33. 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

34. An anion forms when an atom gains one or more electrons
A cation forms when an atom loses one or more electrons.
F + e- --> F-
Mg --> Mg e-

35. metals (Mg) lose electrons ---> cations
In general
metals (Mg) lose electrons ---> cations
nonmetals (F) gain electrons ---> anions

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

37. One Last Learning Check
Write the nuclear symbol form for the following atoms or ions:
A. 8 p+, 8 n, 8 e- ___________
B. 17p+, 20n, 17e- ___________
C. 47p+, 60 n, 46 e- ___________

38. 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.

39. Atomic Mass How heavy is an atom of oxygen?
It depends, because there are different kinds of oxygen atoms.
We are more concerned with the average atomic mass.
This is based on the abundance (percentage) of each variety of that element in nature.
We don’t use grams for this mass because the numbers would be too small.

40. Measuring Atomic Mass
Instead of grams, the unit we use is the Atomic Mass Unit (amu)
It is defined as one-twelfth the mass of a carbon-12 atom.
Carbon-12 chosen because of its isotope purity.
Each isotope has its own atomic mass, thus we determine the average from percent abundance.

41. To calculate the average:
Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results.
If not told otherwise, the mass of the isotope is expressed in atomic mass units (amu)

42. Composition of the nucleus
Atomic Masses
Atomic mass is the average of all the naturally occurring isotopes of that element.
Isotope
Symbol
Composition of the 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%
Carbon = 12.01

43. - Page 117
Question
Knowns and Unknown
Solution
Answer

44. The Periodic Table: A Preview
A “periodic table” is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties
The periodic table allows you to easily compare the properties of one element to another

45. The Periodic Table: A Preview
Each horizontal row (there are 7 of them) is called a period
Each vertical column is called a group, or family
Elements in a group have similar chemical and physical properties
Identified with a number and either an “A” or “B”

46. Section 6.1 Organizing the Elements
A few elements, such as gold and copper, have been known for thousands of years - since ancient times
Yet, only about 13 had been identified by the year 1700.
As more were discovered, chemists realized they needed a way to organize the elements.

47. Section 6.1 Organizing the Elements
Chemists used the properties of elements to sort them into groups.
In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties
One element in each triad had properties intermediate of the other two elements

48. Mendeleev’s Periodic Table
By the mid-1800s, about 70 elements were known to exist
Dmitri Mendeleev – a Russian chemist and teacher
Arranged elements in order of increasing atomic mass
Thus, the first “Periodic Table”

49. He left blanks for yet undiscovered elements
Mendeleev
He left blanks for yet undiscovered elements
When they were discovered, he had made good predictions
But, there were problems:
Such as Co and Ni; Ar and K; Te and I

50. A Better Arrangement
In 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number
The arrangement used today
The symbol, atomic number & mass are basic items included-textbook page 162 and 163

52. The Periodic Law
When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
Horizontal rows = periods
There are 7 periods
Vertical column = group (or family)
Similar physical & chemical prop.
Identified by number & letter (IA, IIA)

53. Areas of the Periodic Table
Three classes of elements are: ) metals, 2) nonmetals, and ) metalloids
Metals: electrical conductors, have luster, ductile, malleable
Nonmetals: generally brittle and non-lustrous, poor conductors of heat and electricity

54. Areas of the Periodic Table
Some nonmetals are gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br)
Notice the heavy, stair-step line?
Metalloids: border the line-2 sides
Properties are intermediate between metals and nonmetals

55. Squares in the Periodic Table
The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms:
Atomic number and atomic mass
Black symbol = solid; red = gas; blue = liquid (from the Periodic Table on our classroom wall)

57. Groups of Elements - Family Names
Group IA (1) – alkali metals
Forms a “base” (or alkali) when reacting with water (not just dissolved!)
Group 2A (2)– alkaline earth metals
Also form bases with water; do not dissolve well, hence “earth metals”
Group 7A (17) – halogens
Means “salt-forming”
Group 8A (18) – noble gases
Nonreactive because of their electron configuration

58. ELEMENTS THAT EXIST AS DIATOMIC MOLECULES
Remember:
HOFBrINCl
These elements only exist as PAIRS. Note that when they combine to make compounds, they are no longer elements so they are no longer in pairs!