Chapter 2

Atomic Theory

400 B.C.E. atomists

1804 C.E. Dalton

1903 Thompson

1911 Rutherford1913

Bohr1932

Chadwick

A History of Atomic Models

Some early philosophers believed that matter had ultimate, tiny, indivisible particles

Leucippus and Democritus

Other philosophers believed that matter was infinitely divisible

Plato and Aristotle

Because there was no experimental way of proving who was correct, the best debater was the person assumed

correct, i.e., Aristotle

Early Philosophy of Matter

In the late 17th century, the scientific approach to understanding nature became established.

For the next 150+ years, observations about nature were made that could not easily be explained by the infinitely

divisible matter concept.

Scientific Revolution

Lavoisier Proust Dalton

A Summary of the Nature of Electrical Charge19th Century

Positive (red) and negative (yellow) charges attract.

Positive-positive and negative-negative charges repel.

Positive and negative charges cancel.

1+ 1- 0

J.J. Thomson and Cathode Rays

Examined the electrically charge “particles” in cathode rays Measured the amount of force necessary to deflect the paths of the particles

The particles have a negative charge. The amount of deflection was related to two factors, the charge and mass of

the particles.

Every material tested contained these same particles.

+++++++++++

-------------

Power Supply - +

Cathode

(-) (+)

Thomson’s Conclusions

If the particle has the same amount of charge as a hydrogen ion, then it must have a mass almost 2000 x

smaller than hydrogen atoms!

The only way for this to be true is if these particles were pieces of atoms.

Apparently, the atom is not unbreakable.

Thomson believed that these particles were the ultimate building blocks of matter.

These cathode ray particles became known as electrons.

A New Theory of the Atom

Thomson proposed that instead of being a hard, marble-like unbreakable sphere, the way Dalton

described it, the atom actually had an inner structure.

The structure of the atom contains many negatively charged electrons. These electrons are held in the atom by their attraction for a positively charged

electric field within the atom.

Predictions of the Plum Pudding AtomThe mass of the atom is due to the mass of the electrons within it. Electrons are the only particles in Plum Pudding atoms, therefore the only source of mass.

The atom is mostly empty space and should not have a bunch of negatively charged particles near each other as they would repel.

Radioactivity

In the late 1800s, Henri Becquerel and Marie Curie discovered that certain elements would constantly

emit small, energetic particles and rays.

These energetic particles could penetrate matter.

Ernest Rutherford discovered that there were three different kinds of emissions:

alpha rays made of particles with a mass 4 x H atom and + charge

beta rays made of particles with a mass ~1/2000th H atom and – charge

gamma rays that are energy rays, not particles

Rutherford, Marsden and Geiger

Detector

alpha rays

Gold foil

Rutherford’s Experiment

To minimize alpha loss by scattering from air molecules, the experiment was carried out in a fairly good vacuum, the metal box being evacuated through a tube T (see below). The alphas came from a few milligrams of radium (to be precise, its decay product radon 222) at R in the figure below, from the original paper, which goes on: "By means of a diaphragm placed at D, a pencil of alpha particles was directed normally on to the scattering foil F. By rotating the microscope [M] the alpha particles scattered in different directions could be observed on the screen S." Actually, this was more difficult than it sounds. A single alpha caused a slight fluorescence on the zinc sulphide screen S at the end of the microscope. This could only be reliably seen by dark-adapted eyes (after half an hour in complete darkness) and one person could only count the flashes accurately for one minute before needing a break, and counts above 90 per minute were too fast for reliability. The experiment accumulated data from hundreds of thousands of flashes.

Rutherford’s Experiment-Expected Result

Rutherford’s Experiment-Actual Result

Rutherford’s Results

Over 98% of the α particles went straight through

About 2% of the α particles went through but were deflected by large angles

About 0.005% of the α particles bounced off the gold foil

Rutherford's partner in the initial phase of this work was Hans Geiger, who later developed the Geiger counter to detect and count fast particles. Many hours of staring at the tiny zinc sulphide screen in the dark must have focused his mind on finding a better way!

In 1909, an undergraduate, Ernest Marsden, was being trained by Geiger. To quote Rutherford (a lecture he gave much later):

"I had observed the scattering of alpha-particles, and Dr. Geiger in my laboratory had examined it in detail. He found, in thin pieces of heavy metal, that the scattering was usually small, of the order of one degree. One day Geiger came to me and said, "Don't you think that young Marsden, whom I am training in radioactive methods, ought to begin a small research?" Now I had thought that, too, so I said, " Why not let him see if any alpha-particles can be scattered through a large angle?" I may tell you in confidence that I did not believe that they would be, since we knew the alpha-particle was a very fast, massive particle with a great deal of energy, and you could show that if the scattering was due to the accumulated effect of a number of small scatterings, the chance of an alpha-particle's being scattered backward was very small. Then I remember two or three days later Geiger coming to me in great excitement and saying "We have been able to get some of

the alpha-particles coming backward …" It was quite the most incredible event that ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you."

Rutherford developed the nuclear theory of the atom.

1. Most of the atom’s mass and all of its positive charge are contained in a small core called the nucleus.

2. Most of the volume of the atom is empty space through which the tiny, negatively charged electrons are dispersed.

3. The number of negatively charged electrons outside the nucleus is equal to the number of positively charged particles (protons) inside the nucleus, so that the atom is electrically neutral.

Nucleus: protons and neutrons

electrons

Structure of the Nucleus

Rutherford proposed that the nucleus had a particle that had the same amount of charge as an electron but

opposite sign – the proton.

Because protons and electrons have the same amount of charge, for the atom to be neutral there must be equal

numbers of protons and electrons.

Some Problems

How could beryllium have four protons stuck together in the nucleus?

If a beryllium atom has four protons, then it should weigh 4 amu; but it actually weighs 9.01 amu! Where is the extra mass coming from?

Protons - subatomic particles found in the nucleus

The proton has a charge of +1.60 x 1019 C. The proton has a mass of 1.67262 x 10−24 g.

Electrons - negatively charged particles found in all atoms. Cathode rays are made of streams of electrons. The electron has a charge of −1.60 x 1019 C. The electron has a mass of 9.1 x 10−28 g.

There Must Be Something Else!

To answer these questions, Rutherford and Chadwick proposed that there was another particle in the nucleus

– it is called a neutron.

Neutrons are subatomic particles with a mass = 1.67493 x 10−24 g and no charge, and are found in the

nucleus.

Subatomic Particle

Mass (g)

Mass (amu) Charge Symbol

Proton 1.67262 x 10-24 1.00727 +1 p, p+, H+

Electron 0.00091 x 10-24 0.00055 -1 e, e-

Neutron 1.67493 x 10-24 1.00866 0 n, n0

1 amu = 1 atomic mass unit = 1/12 the mass of a carbon-12 atom

The Properties of Protons, Neutrons, and Electrons

Quantum Model

approximately 10-10 m

Atom

approximately 10-14 m

Nucleus

nucleus electrons

protonneutron

The Nuclear Model

The Number of Protons Defines the Element

Helium nucleus: 2 protons, 2 neutrons

Carbon nucleus: 6 protons, 6 neutrons

Structure of the Nucleus Isotopes and Ions

Soddy (1913) discovered that the same element could have atoms with different masses, which he called isotopes.

The observed atomic mass is a weighted average of the weights of all the naturally occurring atoms.

The percentage of an element that is one isotope is called the isotope’s natural abundance.

Isotopes Atoms of the same element which differ in the

number of neutrons in their nuclei. Consider carbon.

C-12

6p,6n

6e

98.89 %C-13

6e

6p,7n

1.11 %C-14

6e

6p,8n

0.0000000001 %

Isotopes

All isotopes of an element are chemically identical and undergo the exact same chemical reactions.

All isotopes of an element have the same number of protons.

Isotopes of an element have different masses.

Isotopes of an element have different numbers of neutrons.

Isotopes are identified by their mass numbers, which is the sum of all the protons and neutrons in the nucleus.

Atomic Mass (Weight)

Carbon 98.99 % C-12

1.11 % C-13 Weighted average = 12.01 amu trace % C-14 }

Chlorine 75.5 % Cl-35 Weighted average = 35.45 amu 24.5 % Cl 37 }

Isotopes

Atomic number Number of protons = Z

Mass Number Protons + neutrons = whole number = A

Relative Abundance = relative amount found in a sample

XZ

A

# neutrons = A-Z

25p+

30n0

25e-

18p+

20n0

18e-

47p+

62n0

47e-

6p+

7n0

6e-

40p+

50n0

40e-

28p+

33n0

28e-

Ar1838

Mn2555

Ag47109

Zr4090

C613

Ni2861

a) b) c)

d) e) f)

1

Al27

13

C13

6

Mo96

42

Cs133

55

2. Complete the table:

6 6 13

42 54 42

13 14 13 13 27

55 78 55

Review : What is a weighted average ?

The atomic weight of the element bromine (Br) is listed as 79.904 amu. Which would predict to be the most

abundant isotope of Br ?

Br 80 35 X ☜ Does not exist

Naturally occurring Bromine is actually 51% Br-79 and 49% Br-81.

Atomic Mass

We previously learned that not all atoms of an element have the same mass - isotopes

Isotopes have identical chemical properties.

In calculations, we generally use the average mass of all an element’s atoms found in a sample.

We call the average mass the atomic mass.

Atomic Mass = Σ (fractional abundance of isotope)n x (mass of isotope)n

Calculating Atomic Mass

3. Gallium has two naturally occurring isotopes: Ga-69, with mass 68.9256 amu and a natural abundance of 60.11%, and Ga-71, with mass 70.9247 amu and a natural abundance of 39.89%. Calculate the atomic mass of gallium.Convert the percent natural abundances into decimal form by dividing by 100.

Fraction Ga-69 = 60.11/100 = 0.6011

Fraction Ga-71 = 39.89/100 = 0.3989

Atomic mass = (0.6011 × 68.9256 amu) + (0.3989 × 70.9247 amu)

= 41.4321 amu + 28.2919 amu

= 69.7231 = 69.72 amu

A gaseous sample is ionized by bombardment

with electrons.

The positive ions thus formed are subjected to an electrical force by the electrically charged velocity selector plates and a magnetic

force by a perpendicular magnetic field.

Ions with a particular velocity pass through and are deflected into

circular paths by the magnetic field.

Ions with different masses strike the detector in different regions.

The more ions of a given type, the greater the response of the detector .

Mass Spectrometry

The response of an ion detector converted to a relative scale

The percent natural abundances of the mercury isotopes are

196Hg, 0.146% 198Hg, 10.02% 199Hg, 16.84% 200Hg, 23.13%

201Hg, 13.22% 202Hg, 29.80% 204Hg, 6.85%

80Hg200.59

Mass Spectrum of Chlorine

Atomic Mass of Cl = 35.45 amu

Chapter 2 The Periodic Table

A. “Metals” vs.”Nonmetals” (Before 1800)

1. Metals - Solids, Lustrous, Malleable, Ductile, Conductors

2. Nonmetals - Solids, Liquids, Gases, Poor Conductors

B. In early 1800’s, 50-60 Elements Known

1. 1829 - Doebereiner - A System of “Triads”

Cl, Br, I Li, Na, K Fe, Co, Ni

S, Se, Te Ca, Sr, Ba

Classification of the Elements

Classification of the Elements2. 1865 - Newlands - “Law of Octaves”

a. Arranged lightest of known elements by atomic weight

b. Observed a periodicity of eight

H Li Be B C N O

F Na Mg Al Si P S

Cl K Ca Cr Ti Mn Fe

c. Note: He, Ne, Ar, and Kr were undiscovered

d. Introduced the idea of periodicity

Mendeleev

3. 1869 - Mendeleev and Meyer

a. Arranged Elements by Atomic Weight

b. Also Considered Chemical Properties of Elements

c. The Result - A Table in Which the “Families” of Elements Got Larger

d. Additional Triumph of Mendeleev Spaces left for Undiscovered Elements

Classification of the Elements

“I began to look about and write down the elements with their atomic weights and typical properties, analogous elements and like atomic weights on separate cards, and this soon convinced me that the properties of elements are in periodic dependence upon their atomic weights.” --Mendeleev, Principles of Chemistry, 1905, Vol. II

Mendeleev:

Predicted Properties of Eka-aluminum and Eka-Silicon

Features of the Modern Periodic Table A. The Common Form of the Table is One of Many

B. Periods = Rows Groups = Columns

C. “A” Groups = “Representative” “B” Groups = “Transition”

Lanthanides and Actinides = “Inner Transition”

D. Periodic Law - Elements within a given group have similar chemical and physical properties. These properties change gradually with an increase in atomic number.

E. The terms metals, nonmetals, and metalloids are still used.

Mendeleev didn’t get it all correct !!

Chapter 2

The Periodic Table

Group Behavior

Hydrogen

Alkali Metals-Group IA Li, Na, K, Rb, Cs, Fr

Alkali Earth Metals-Group IIA Mg, Ca, Sr, Ba, Ra

Halogens-Group VIIA F, Cl, Br, I, At

Noble Gases-Group VIIIA He, Ne, Ar, Kr, Xe, Rn

Important Groups