1. CHE 106: General Chemistry
CHAPTER TWO
Copyright © James T. Spencer
All Rights Reserved

2. Chapter 2: Atoms, Molecules and Ions
What is Chemistry
Logic
Magic
Chapt. 2.1

3. Atoms, Molecules and Ions
Science: Atomic Theory
“The strength of a science is that its conclusions are derived by logical arguments from facts that result from well-designed experiments. Science has produced a picture of the microscopic structure of the atom so detailed and subtle of something so far removed from our immediate experience that it is difficult to see how its many features were constructed. This is because so many experiments have contributed to our ideas about the atom.”
B. Mahan from University Chemistry
Chapt. 2.1

4. “Seeing “Atoms
STM image showing single-atom defect in iodine adsorbate lattice on platinum.

5. Atoms, Molecules and Ions
Science: Atomic Theory
from a fundamental understanding of the macroscopic behavior of substances comes an understanding the microscopic behavior of atoms and molecules (Baseball rules from Baseball Game?)
Macroscopic Microscopic
Substances Atomic theory
Mixtures
Physical Properties and Changes
Question: Can matter be infinitely divided?
Most Greek Philosophers - Yes
Democritus (460 BC) and John Dalton (1800s) - No (“atomos”means indivisible”)
Chapt. 2.1

6. Atoms, Molecules and Ions
History of Atomic Theory and Scientific Inquiry
Aristotle - “metaphysics”, thought experiments and no experimental observations necessary to substantiate ideas.
Archimedes ( BC) - Scientific Method, determined composition of the King of Syracuse’s crown by measuring density through water displacement.
Roger Bacon ( ) - Experimental Science “ It is the credo of free men - the opportunity to try, the privilege to err, the courage to experiment anew. ...experiment, experiment, ever experiment”.
Chapt. 2.1

7. Aristotle ( BC)
All of the sciences (epistêmai, literally "knowledges") can be divided into three branches: theoretical, practical, and productive. Whereas practical sciences, such as ethics and politics, are concerned with human action, and productive sciences with making things, theoretical sciences, such as theology, mathematics, and the natural sciences, aim at truth and are pursued for their own sake.

8. Archimedes ( BC)
Archimedes was a native of Syracuse (not NY). Stories from Plutarch, Livy, and others describe machines invented by Archimedes for the defence of Syracuse (These include the catapult, the compound pulley and a burning-mirror).
Archimedes discovered fundamental theorems concerning the centre of gravity of plane figures and solids. His most famous theorem gives the weight of a body immersed in a liquid, called Archimedes' principal.
His methods anticipated integral calculus 2,000 years before Newton and Leibniz.

9. Archimedes ( BC)

10. Archimedes ( BC)
Suspecting that a goldsmith might have replaced some of the gold by silver in making a crown, Hiero II, the king of Syracuse, asked Archimedes to determine whether the wreath was pure gold. The wreath could not be harmed since it was a holy object.
The solution which occurred when he stepped into his bath and caused it to overflow was to put a weight of gold equal to the crown, and known to be pure, into a bowl which was filled with water to the brim. Then the gold would be removed and the king's crown put in, in its place. An alloy of lighter silver would increase the bulk of the crown and cause the bowl to overflow.
Pure Gold?
Equal Weight of Gold
Crown Displaced More Water

11. Greek Philosophers Air Fire Water Earth
Greek “Elements”
Water
Earth
Democratus - First to say that all matter is NOT infinately divisible. [But the Greeks did not test their ideas]
Alchemy - Pseudoscience by fakes and mystics devoted to turning base metals to gold BUT they did make (by accident) many ground breaking discoveries of nature (chemical reactions).

12. Scientific Measurement
Robert Boyle - Robert Boyle ( ) was born in Ireland. He became especially interested in experiments involving air and developed an air pump with which he produced evacuated cylinders. He used these cylinders to show that a feather and a lump of lead fall at the same rate in the absence of air resistance. In his book “The Sceptical Chemist” (1661), Boyle urged that the ancient view of elements as mystical substances should be abandoned and that an element should instead be defined as anything that cannot be broken down into simpler substances.

13. Scientific Measurement
Antoine Lavoisier ( ) - Furthered measurement as basis for scientific reasoning.
“Je Veux Parler Des Faits” - Do Not Rely Upon Speculation But Build Upon Facts.
More on Lavoisier on Next Slide

14. Antoine Lavoisier
Antoine Lavoisier was born in Paris, and although Lavoisier's father wanted him to be a lawyer, Lavoisier was fascinated by science. From the beginning of his scientific career, Lavoisier recognized the importance of accurate measurements. He wrote the first modern chemistry (1789) textbook so that it is not surprising that Lavoisier is often called the father of modern chemistry. To help support his scientific work, Lavoisier invested in a private tax-collecting firm and married the daughter of one of the company executives. Guillotined for his tax work in 1794.

15. Atoms, Molecules and Ions
Earth
History Atomic Theory and Scientific Inquiry
Lavoisier ( ) - founder of “modern chemistry”, not to rely on speculation but to build upon facts, ended the “time of alchemy”.
Alchemy
Water
Fire
pure water
“earth”
evaporate out water from
dust sealed container
alchemists said that the water
was “transmuted” to earth
Law of Conservation of Mass
Lavoisier showed that the amount of “earth” found at the
end of the experiment was equal to the weight the container
lost, therefore, the water was not “transmuted” to earth.
Chapt. 2.1

16. Scientific Method Theory Form and test hypothesis Patterns and Trends
Observations
and Experiments
Chapt. 2.1

17. John Dalton ( )
John Dalton ( ), an Englishman, began teaching school when he was 12. He was fascinated with meteorology (keeping daily weather records for 46 years), which led to an interest in gases and their components, atoms. He switched to chemistry when he saw applications in chemistry for his ideas about the atmosphere. He proposed the Atomic Theory in Dalton was a humble man with several apparent handicaps: he was poor; he was not articulate; he was not a skilled experimentalist, and he was color-blind (a terrible problem for a chemist). In spite of these disadvantages he did great things.

18. Atomic Theory John Dalton’s Atomic Theory
Designed a theory to account for a variety of experimental observations:
Each element is composed of extremely small particles (called atoms).
All atoms of a given element are identical (therefore, atoms of different elements are different and have different properties).
Chapt. 2.1

19. Atomic Theory (Continued)
John Dalton’s Atomic Theory
Atoms of an element are not changed into different types of atoms by chemical reactions and atoms are neither created nor destroyed in chemical reactions.
Compounds are formed when atoms combine and a given compound always has the same relative number and kind of atoms.
Chapt. 2.1

20. Atomic Theory Dalton’s Atomic Theory Atoms are the building blocks:
Elements are composed of only one kind of atom.
Compounds are made by mixing atoms in definite proportions
Mixtures do not involve the type of “small scale” (but strong) interactions found in Elements and Compounds
Chapt. 2.1

21. Atomic Theory; Dalton’s Theories
Law of Constant Composition (or Definite Proportion, first proposed by Joseph Proust):
In any given compound, the relative number and kind of atoms are constant (same proportion of elements by mass).
implies that atoms interact in a specific way when they form a compound.
the elements making up a particular compound combine in the same proportions regardless of the manner in which the compound was prepared.
Chapt. 2.1

22. Atomic Theory; Dalton’s Theories
Law of Constant Composition (or Definite Proportion):
Copper Carbonate ALWAYS contains 5.3 parts Copper to 4 parts Oxygen and 1 part Carbon (by Weight).
Carbon Dioxide ALWAYS contains 1.00 parts Carbon to 2.67 parts Oxygen
Chapt. 2.1

23. Atomic Theory; Dalton’s Theories
Law of Conservation of Mass:
the total amount of material present after a chemical reaction is the same as the amount present before the reaction.
Matter (elements, etc...) cannot be created nor destroyed during chemical reactions.
Total Mass Before Chemical Reaction
Total Mass After
Chemical Reaction =
Chapt. 2.1

24. Atomic Theory; Dalton’s Theories
Law of Multiple Proportions:
If two elements form more than one compound, then the ratios of the masses of a second element that combine with 1 g of the first elements can always be reduced to small whole numbers
Mass of O Comb. w/ 1 g C
I 1.33 g
II 2.66 g
III 3.99 g
Chapt. 2.1

25. Atomic Theory; Dalton’s Theories
Law of Multiple Proportions:
If two elements form more than one compound, then the ratios of the masses of a second element that combine with 1 g of the first elements can always be reduced to small whole numbers C
Mass of O Comb. w/ 1 g C
I 1.33 g
II 2.66 g
III 3.99 g I
III II O C O O C O C O O
1: : :1
Chapt. 2.1

26. Atomic Theory; Dalton’s Theories
Law of Multiple Proportion: Another Example
Oxygen combines with hydrogen to form 2 compounds
Compound 1
8 grams of oxygen combines with 1 gram of hydrogen
[H2O]
Compound 2
16 grams of oxygen combines with 1 gram of hydrogen
[H2O2]
Chapt. 2.1

27. Atomic Theory; Dalton’s Theories
Law of Multiple Proportion: Yet Another Example
Chlorine combined with oxygen to form four binary compounds [A, B, C, D].
Compound Mass of O combined Div. by
with g Cl
A g 1.00
B g 4.00
C g 6.00
D g 7.00
Allowed Dalton to prepare the first atomic mass table
Chapt. 2.1

28. Guy-Lussac
Joseph Guy-Lussac ( ) found that (at the same temperatures and pressures):
2 volumes of hydrogen reacts with 1 volume of oxygen to yield 1 volume of water vapor + =
Water O H
Amedeo Avogadro ( ) proposed that (at the same temperatures and pressures), equal volumes of different gases contain the same number of particles:
2 molecules of H + 1 molecule of O yield 1 molecule of water

29. Experiments in Atomic Theory
Dalton’s Laws Set Groundwork for Atomic Theory but Important Experiments Lead to Our Modern Understanding
Faraday - Electrodeposition
Millikan - Oil Drop Experiment
Roetgen - Radioactivity
Curie - Radioactive Particles
Rutherford - Gold Foil Experiment

30. Michael Faraday ( )
Experiments in electro-magnetism, electrical power conversion, etc...
Humble scientist rose from very poor background to become one of the most influential of his age. Believed that careful observations were most important.
“Try desperately to succeed - and do not hope for success”

31. Atomic Structure Electrodeposition Cell Electrical Nature
Michael Faraday (1833) (first ideas about the nature of electricity
The weight of a material deposited at an electrode by a given amount of electricity is always the same.
The weights of various materials deposited by fixed amounts of electricity are proportional to their equivalent weights. [remember equivalent weights]
electrodes
deposition
electrolyte + -
Electrodeposition
Cell
If equivalent weights of a substance contain the same number or an integral multiple of molecules or atoms then a fixed number of atoms reacts with a fixed amount of electricity. This suggests that electricity is also composed of particles.
Chapt. 2.1

32. Sir J. J. Thomson
British physicist who worked with electrical currents and fields.
Appointed Prof. of Physics at Cambridge when he was 27 and Received the Nobel Proze in for his characterization of the electron.

33. Atomic Structure
J. J. Thomson: Cathode Ray Tube (CRT) Experiment
Set up a large electrical potential between a pair of electrodes in a glass tube and an electrical current will flow between the elctrodes.
The current will flow even when all the air is pumped out of the tube. The invisible charge carriers were called “cathode rays”.
Cathode rays travel in straight lines and form a luminious spot when they hit a glass tube.
(-)
(+)
Cathode Ray Tube
[evacuated glass tube]
Chapt. 2.1

34. Atomic Structure: CRT Electric Field Magnetic Field
The cathode rays are deflected by an electric field.
(-)
(+)
Electric Field
The cathode rays are deflected by an magnetic field.
(-)
(+)
The same effect was observed regardless of what gas was used in the discharge tube. Therefore, electricity must be a universal fragment.
Magnetic Field
Chapt. 2.1

35. Electricity: Thomson’s charge to mass
CRT
(-)
(+) 1 2 3
(-)
(+)
Magnetic Field
Electric Field
Spot mag field elec. field
1 On Off
3 Off On
2 Off Off
On On
Chapt. 2.1

36. Thomson’s charge to mass
Ee = Electrical Field He = Magnetic Field
[where e = electric charge (unk) and  = velocity]
Set up experiment such that;
Electrical Field = Magnetic Field
Ee = He or
E / H
Now, turn off the mag. field and measure deflection of beam ()
Using Newton’s 2nd Law can calculate e/m
CRT
(-)
(+) 1 2 3
(-)
(+)
Magnetic Field

37. Thomson’s charge to mass
calculated charge to mass ratio (e/m)
for electron = 1.76 x 108 coulombs/g
found;
(1) e/m was 1000x greater than for any known ion
(2) e/m of independent of gas in tube [Universal Fragment]
(3) Not electrified atoms but fragments (called electrons)

38. Robert Millikan ( )
Nobel Prize, 1923; for his work on the elementary charge of electricity and on the photoelectric effect.
Robert Millikan was one of the first American scientists to be recognized in Europe. In 1909 he performed the first of a series of experiments to measure the fundamental charge of an electron, the Millikan Oil Drop Experiment. The value determined by this experiment was used in Bohr's formula for the energy of the Hydrogen line spectrum as a first confirmation of the quantized atom. He named and studied "cosmic rays" as well.

39. Electricity: Millikan’s electron mass
Oil Drop Experiment (1909)
Goal: to measure the electrical charge on each oil droplet
Procedure: measure the velocity of the falling oil drop both with and without the high voltage plates urned on
Found: charges were always multiples of 1.60 x C
Postulate: charge of one electron was 1.60 x C
atomizer -
high
voltage
viewer +
Ionization by radiation causes
the oil to pick up “extra” electrons
Chapt. 2.1

40. Electricity: electron mass
charge = e = x 108 coul g-1
mass m
Thomson
charge = e = x coul
Millikan
Combine and Solve
mass = charge = x C = x g
1.76 x 108 coul g x 108 C g-1
mass of the electron was 2000x smaller than the
lightest atom (hydrogen)
Chapt. 2.1

41. Wilhelm Conrad Roentgen
Wilhelm Conrad Roentgen was born in Lennep, Germany, on 27 March He obtained a degree in mechanical engineering and, in 1869, was awarded a degree in physics. While working as a professor of physics at Wurzburg University, he made his famous discovery. He called the unknown radiation "X rays," since "X" frequently stands for an unknown quantity in mathematics. His unique discovery truly changed the world and immediately became a useful tool for medical science.
Wilhelm Conrad
Roentgen

42. Radioactivity: Wilhelm Roetgen and Henri Becquerel
CRT
metal
target
e beam
X-rays
- not affected by magnetic fields
- passed thru many materials
-produced images on film
(ionized Ag emulsions)
invisible
radiation
(X-rays) U
glowed in dark (phosphorescence)
emitted high energy
radiation in the dark (radioactivity)
Chapt. 2.1

43. 1903 Nobel Prize for Radioactivity
Pierre and Marie Curie
Henri Becquerel

44. Marie Sklodawaska Curie
The most famous of all women scientists, Marie Sklodowska-Curie is notable for many firsts.
In 1903, she became the first woman to win a Nobel Prize for Physics (Pierre Curie and Henri Becquerel, for the discovery of radioactivity.
She was also a professor at the Sorbonne University in Paris (1906).
In 1911, she won an unprecedented second Nobel Prize (in chemistry for her discovery radium. She was the first person ever to receive two Nobel Prizes.)
She was the first mother of a Nobel Prize Laureate; daughter- Nobel Prize 1932.
Marie Sklodowska-Curie
In 1934, Maria Curie died of leukemia

45. Radioactivity: Marie Curie and Ernest Rutheford
Marie Curie ( ) - separated the pure radioactive material (Uranium) which was spontaneously radioactive (from the mineral pitchblende)
Ernest Rutheford ( ) - found radiation from uranium was of three types (, , and )  - U  +
slits 
 - heavy particles with +2 charge, combines with electrons
to form helium, 4He
 - electrons with -1 charge
 - high energy electromagnetic radiation
Chapt. 2.1

46. Nuclear Atom: Thomson’s Model (ca. 1900)
Since the electron made up only a small amount of an atom’s mass it was proposed that it must similarly make up a small amount of the atoms volume.
“Plum-pudding” model
positive charge spread
over sphere
= electron

47. Ernest Rutherford
Ernest Rutherford ( ) was born on a farm in New Zealand. In he placed second in a scholarship competition to attend Cambridge University, but was awarded the scholarship when the winner decided to stay home and get married. As a scientist in England, Rutherford did much of the early work on characterizing radioactivity. He also invented the name proton for the nucleus of the hydrogen atom. He received the Nobel Prize in chemistry in 1908.

48. Nuclear Atom: Rutheford and the Gold Foil
experiment - fired
heavy  particles
at a thin gold foil
and looked
for deflections 
4He particles 
thin
gold
foil
slits
detector
found - most particles passed straight through foil, some had
deflections thru small angles BUT some had VERY
large deflections ( = 180°)
“...as if you fired a 15-inch
cannon shell at a piece of tissue
paper and it came back and hit you...”

49. Nuclear Atom: Rutheford and the Gold Foil
A:C around 13,000:1
Gold Foil A
 Beam A C B B A A

50. Rutheford’s Atom
Based on gold foil experiment and previous work with electrical and nuclear particles, proposed a nuclear theory;
(1) atoms are mostly empty space with very dense (pos. charged) nuclear core (<10-12 cm dia.)
(2) atoms are highly “non-uniform”
(3 ) atomic nucleus must contain large electrical forces of considerable mass (since small electron cannot be responsible for such large deflections)

51. Nature’s Basic Forces + - + + + +
Electromagnetic - force between charged or magnetic particles (electrical and magnetic forces are very closely related) DRIVES MOST OF CHEMICAL BEHAVIOR (Coulomb’s Law; F = kQ1Q2/d2)
Gravitational - force between objects proportional to their masses.
Strong Nuclear - force keeping like charged nucleons (such as protons) together (very strong but very short range).
Weak Nuclear - nuclear force observed in some radioactive behavior (weaker than electromagnetic but stronger than gravitational). + - m m + + + +
Strong Nucl. > Electromagnetic > Weak Nucl. > Gravitational

52. Modern Atomic Structure
atomic dimensions; nucleus 10-4 Å and atom Å (1 Å = m) “... if a nucleus were 2 cm (ca. 1 in.) then the atom would be 200 m (ca. 200 yds)”
atom composed of many “subatomic” particles but only three of these are important to chemists
atomic mass (1 amu = 4 x g), charge (1 esc = 1.60 x coul), density (1014 g/cm3)
atom = dense nucleus with mostly empty space; electrons of most chemical import. (matchbox of nucl. = 2.5 billion tons)
particle charge (esu) mass (amu)
proton
neutron
electron x 10-4

53. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
A) express this diameter in picometers

54. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
A) express this diameter in picometers
1.5 angstroms

55. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
A) express this diameter in picometers
1.5 angstroms meters angstrom

56. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
A) express this diameter in picometers
1.5 angstroms meters pm angstrom meter

57. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
A) express this diameter in picometers
1.5 angstroms meters pm angstrom meter
150 pm

58. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
B) How many carbon atoms could be aligned side by side in a straight line across the width of a pencil line that is 0.10 mm wide?
0.10 mm 1 meter angstrom 1 carbon mm meter angs.

59. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
B) How many carbon atoms could be aligned side by side in a straight line across the width of a pencil line that is 0.10 mm wide?
0.10 mm

60. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
B) How many carbon atoms could be aligned side by side in a straight line across the width of a pencil line that is 0.10 mm wide?
0.10 mm 1 meter mm

61. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
B) How many carbon atoms could be aligned side by side in a straight line across the width of a pencil line that is 0.10 mm wide?
0.10 mm 1 meter angstrom mm meter

62. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
B) How many carbon atoms could be aligned side by side in a straight line across the width of a pencil line that is 0.10 mm wide?
0.10 mm 1 meter angstrom 1 carbon mm meter angs.

63. Modern Atomic Structure
Sample exercise: The diameter of a carbon atom is 1.5 angstroms.
B) How many carbon atoms could be aligned side by side in a straight line across the width of a pencil line that is 0.10 mm wide?
0.10 mm 1 meter angstrom 1 atom mm meter angs.
= 6.7 x 105 atoms

64. Atomic Theory: Isotopes
differences/similarities between atoms of an element;
all atoms of an given element have the same number of protons (and therefore the same number of electrons to balance charge)
atoms of an element may have different numbers of neutrons - called isotopes
AE C 12C 13C 14C Z
atomic number (Z) - number of protons
mass number (A) - number of protons + number of neutrons
nuclide - atoms of a specific elemental isotope

65. Atomic Theory: Isotopes
14N
7 electrons, 7 protons, 7 neutrons
8 electrons, 8 protons, 9 neutrons
17 electrons, 17 protons, 18 neutrons
92 electrons, 92 protons, 146 neutrons 7
17O 8
35Cl 17
238U 92

66. Atomic Theory: Isotopes
Sample exercise:How many protons, neutrons, and electrons are in a 39K atom?

67. Atomic Theory: Isotopes
Sample exercise:How many protons, neutrons, and electrons are in a 39K atom?
Atomic# = 19

68. Atomic Theory: Isotopes
Sample exercise:How many protons, neutrons, and electrons are in a 39K atom?
Atomic# = 19 # of protons = 19
# of electrons = 19

69. Atomic Theory: Isotopes
Sample exercise:How many protons, neutrons, and electrons are in a 39K atom?
Atomic# = 19 # of protons = 19
# of electrons = 19
Mass # = 39

70. Atomic Theory: Isotopes
Sample exercise:How many protons, neutrons, and electrons are in a 39K atom?
Atomic# = 19 # of protons = 19
# of electrons = 19
Mass # = = 20 neutrons

71. Atomic Theory: Isotopes
Sample exercise:Give the complete chemical symbol for the nuclide that contains 18 protons, 18 electrons, and 22 neutrons.

72. Atomic Theory: Isotopes
Sample exercise:Give the complete chemical symbol for the nuclide that contains 18 protons, 18 electrons, and 22 neutrons.
Atomic # = 18 , element is Argon

73. Atomic Theory: Isotopes
Sample exercise:Give the complete chemical symbol for the nuclide that contains 18 protons, 18 electrons, and 22 neutrons.
Atomic # = 18 , element is Argon
40Ar 18

74. Atomic Theory: Isotopes
Allotropes - Different chemical forms of the same element existing in the same physical state.
Fullerene
Graphite
Diamond

75. Periodic Table; Dmitri Mendeleev (1869)
Displays chemical reactivity trends and relationships and constructed to account for (and predict) chemical reactivity of the elements.
For example:
Li, Na, K soft metals, v. reactive w/ water
He, Ne, Ar gases and not reactive
F, Cl, Br reactive with many other elements in a similar fashion
Cu, Ag, Au Metal w/ similar reactivity

76. Periodic Table; Dmitri Mendeleev

77. Periodic Table Group or Family Row 1 Alkali metals Li, Na, K,...
2 Alkaline earth metals Be, Mg, Ca,...
16 Chalcogens (chalk formers) O, S, Se,...
17 Halogens (salt formers) F, Cl, Br,...
18 Noble Gases (inert gases) He, Ne, Ar,...
Group
or Family
Row

78. Periodic Table alkali metals non-metals alkaline earth metals
metalloids
non-metals
noble gases
rare earth

79. Periodic Table (1869) metals non-metals
conductors insulators
shiny dull
high thermal conductivity thermal insulators
solids at RT freq. non-solids at RT
ductile brittle
Metalloids (along line in table) have properties
between metals and non-metals

80. Molecules and Ions
Molecule - “assembly” of two or more atoms (with properties different from constituent types of atoms (see “Law of Multiple Proportions”). i.e., H2O, H2O2, CaCO3, HNO3, H2SO4,...
some elements found in nature as molecules (i.e., O2, N2, etc... [diatomic])
Formulas
Molecular - actual numbers and types of atoms in a molecule
Empirical - smallest whole number ratio of constituentStructural - “picture” showing how the atoms are attached to one another

81. Molecules Molecular Empirical Structural Formula Formula Formula
H2O (water) H2O
H2O2 (hydr. peroxide) HO
C2H4 (ethylene) CH2
C6H12O6 (glucose) CH2O

82. Formulas What is its empirical formula?
Ethylene is a gas at room temperature and is the starting material for for many plastics. Its molecular formula is C2H4.
What is its empirical formula?
What other molecular formulas are possible for this same empirical formula?

83. Formulas
Ethylene is a gas at room temperature and is the starting material for for many plastics. Its molecular formula is C2H4.
What is its empirical formula?
CH2
What other molecular formulas are possible for this same empirical formula?
C2H4 , C3H6 , C4H8 , C5H10 , ...

84. Formulas What is its empirical formula?
Cucurbituril is a compound with cage-like molecules big enough to surround and loosely trap smaller molecules. It has the molecular formula C36H36N24O12.
What is its empirical formula?

85. Formulas What is its empirical formula?
Cucurbituril is a compound with cage-like molecules big enough to surround and loosely trap smaller molecules. It has the molecular formula C36H36N24O12.
What is its empirical formula?
C3H3N2O

86. Formulas
Sample exercise: Give the empirical formula for the substance whose molecular formula is Si2H6.

87. Formulas
Sample exercise: Give the empirical formula for the substance whose molecular formula is Si2H6.
SiH3

88. Ions atoms can gain or lose electrons to become charged (called ions)
positive ion = cation negative ion = anion
Na (neutral has 11 electrons) can easily lose 1 electron to become a cation (Na+1)

89. Ions
Polyatomic ions; molecules with charges i.e., NO3-1, SO4-2, PO4-3, etc...
chemical properties of ions may be VERY different from similar neutral species
Predicting charges on ions - use periodic table (gain or lose electrons to end up with the same number as the nearest noble gas)

90. Ions

91. Ions
Sample exercise: How many protons and electrons does the Se2- ion possess?

92. Ions
Sample exercise: How many protons and electrons does the Se2- ion possess?
Se atomic number = 34

93. Ions
Sample exercise: How many protons and electrons does the Se2- ion possess?
Se atomic number = 34
# of protons = 34
# of electrons = = 36

94. Ionic Compounds
transfer of electrons between atoms, Na + Cl = [Na]+[Cl]-
ionic compounds contain anions and cations, typically combinations of metals and non-metals (molecular compounds, in which electrons are shared, are usually result from the combination of non-metals only); FeS, LiBr, CuSO4, TiO4, etc...
total charge is neutral; total (+) = total (-)
ionic compounds are arranged in a 3D array (packing of ping-pong balls)
usually only empirical formulas can be written for ionic compounds (because no real molecular unit in solid phase but “extended” lattice)
usually solids but soluble in water insol. in organic sols.

95. Ionic Compounds total charge is neutral; total (+) = total (-)
Cation Anion Charges Empirical Formula
sodium (Na) chlorine (Cl) Na Cl NaCl
magnesium(Mg) nitrogen (N) Mg N Mg3N2
aluminum (Al) bromine (Br) Al Br AlBr3
barium (Ba) sulfate (SO4) Ba SO BaSO4
lithium (Li) carbonate (CO3) Li CO Li2CO3
nickel (Ni) chloride (Cl) Ni Cl NiCl2
Ni Cl NiCl3

96. Ionic Compounds Cell Face Unit Cell - + - + - + - + - - + - + - + - +

97. Ionic Compounds
Sample exercise: Which of the following compounds are molecular?
CI4
FeS
P4O6
PbF2

98. Ionic Compounds
Sample exercise: Which of the following compounds are molecular?
CI4
FeS
P4O6
PbF2

99. Ionic Compounds
Sample exercise: Write the empirical formulas for the compounds formed by the following ions:
a) Na+ and PO43-

100. Ionic Compounds
Sample exercise: Write the empirical formulas for the compounds formed by the following ions:
a) Na+ and PO43-
Na3PO4

101. Ionic Compounds
Sample exercise: Write the empirical formulas for the compounds formed by the following ions:
b) Zn2+ and SO42-

102. Ionic Compounds
Sample exercise: Write the empirical formulas for the compounds formed by the following ions:
b) Zn2+ and SO42-
ZnSO4

103. Ionic Compounds
Sample exercise: Write the empirical formulas for the compounds formed by the following ions:
c) Fe3+ and CO32-

104. Ionic Compounds
Sample exercise: Write the empirical formulas for the compounds formed by the following ions:
c) Fe3+ and CO32-
Fe2(CO3)3

105. Nomenclature: naming inorganic compounds
Method for unambiguously referring to the a. 15 million known molecules)
Organic compounds - containing C combined typically with H, O, N, and S (originally associated with living organisms but no longer relevant definition)
Inorganic compounds - all other compounds

106. Nomenclature: naming inorganic compounds
Traditional names for compounds long known (ammonia [NH3], water [H2O], Zeise’s salt [Pt(C2H4)Cl3]-1], Muriatic Acid [HCl], etc...)
common names (somewhat systematic, ferrous chloride, cupric chloride, etc...)
International Union of Pure and Applied Chemistry rules (IUPAC)

107. Nomenclature: naming ionic compounds
Ionic compounds are names based upon the component ions.
Positive ion (cation) named and written first
Negative ion (anion) named and written last
Solve ambiguity in charge by using Roman numerals
Cation Anion Compound Name
Na+ Cl- NaCl sodium chloride
Al+3 O-2 Al2O3 aluminum oxide
Fe+2 O-2 FeO iron(II) oxide
(ferrous oxide)
Fe+3 O-2 Fe2O3 iron(III) oxide
(ferric oxide)

108. Nomenclature: naming cations
Monoatomic - take the name from the element
Li+1 lithium ion Sr+3 strontium ion
Ca+2 calcium ion
Polyatomic - only one common polyatomic cation
NH4+1 ammonium ion
Multiple Cationic Charge Possible - specify charge with Roman numerals to be unambiguous
Fe+2 iron(II) ion Fe+3 iron(III) ion
Cr+6 chromium(VI) ion Cr+5 chromium(V) ion
For metals, older method used to distinguish between ions differing by one charge unit by adding suffix (-ous for lower charge, -ic for higher charge)
Fe+2 ferrous ion Fe+3 ferric ion
Co+2 cobaltous ion Co+3 cobaltic ion

109. Nomenclature: naming anions
Monoatomic - add -ide suffix
F-1 fluoride ion P-3 phosphide ion
O-2 oxide ion B-5 boride ion
Polyatomic - some common use -ide suffix
OH-1 hydroxide ion CN-1 cyanide ion
N3-1 azide ion O2-2 peroxide ion
Oxyanions - (1) when only two, the one with less oxygen ends in -ite and the one with more oxygen ends with -ate
NO2-1 nitrite ion NO3-1 nitrate ion
SO3-2 sulfite ion SO4-2 sulfate ion
Oxyanions- for species with more than two members use prefixes (hypo- less oxygen and per- more oxygen) ClO-1 ClO2-1 ClO3-1 ClO hypochlorite chlorite chlorate perchlorate

110. Nomenclature: acids
Acid - compound which yields H+ when dissolved in water
write hydrogen first; HCl, H2SO4, H3PO4, etc...
anions which end in -ide use hydro- as prefix and -ic as suffix
Anion Acid
Cl- (chloride) HCl (hydrochloric acid)
F- (fluoride) HF (hydrofluoric acid)
oxyacids - replace -ate suffix of anion with -ic,
replace -ite suffix of anion with -ous (leave prefixes!)
Anion Acid
ClO2- (chlorite) HClO2 (chlorous acid)
ClO3- (chlorate) HClO3 (chloric acid)
ClO4-1 (perchloric) HClO4 (perchloric acid)

111. Nomenclature: molecular compounds
Similar to ionic compounds
More positive element (left and down on periodic table) named first (first in formula also)
Second element name ends with -ide
Use numbering prefixes if necessary
Prefix Number
Mono- 1
Di- 2
Tri- 3
Tetra- 4
Penta- 5
Hexa- 6
Hepta- 7
Octa- 8
Nona- 9
Deca- 10
Formula Name (text prob. 2.45)
N2O5 dinitrogen pentoxide
IF7 iodine heptafluoride
XeO3 xeon trioxide
SiCl4 silicon tetrachloride
H2Se dihydrogen selenide
P4O6 tetraphosphorus hexoxide

112. Nomenclature: examples
Formula Name
ZnCl2
(NH4)2SO4
FeF3
HBr
HBrO4
SF6
HCN
zinc(II) chloride
ammonium sulfate
iron(III) fluoride
hydrobromic acid
perbromic acid
sulfur hexafluoride
hydrogen cyanide

113. End Chapter 2 Atomic Theory
Experiments leading to the discovery of atomic structure
The Periodic Table
Molecules and Ions
Nomenclature