AtomicTheory

A Brief History

Atoms are made up of subatomic particles called protons, neutrons and electrons

How do we know that?

Vocabulary Atom: The smallest unit of an element,

having all the characteristics of that element and consisting of a dense, central, positively charged nucleus surrounded by a system of electrons.

Molecule: The smallest particle of a substance that retains the chemical and physical properties of the substance and is composed of two or more atoms.

Compound: A compound is a substance made up of atoms representing more than one element bonded together and exhibiting distinct physical and chemical characteristics Example: H2O, H2SO4

Background

Law of Conservation of Mass (Lavoisier, 1789) During a chemical reaction, the total mass of the

reactants is equal to the total mass of the products.

Law of Definite Proportions (Proust, 1799) When atoms combine to form compounds, they

always combine in the same simple, small whole number proportions.

Example: Water is always H2O Example: Sulfuric Acid is always H2SO4

Aristotle(circa. 400 B.C.)

Matter is not made of particles, but rather is continuous.

The continuous matter is called “hyle.” There were only four elements

Earth, Air, Fire, Water

Democritus(circa. 400 B.C.)

Matter is made of empty space and tiny particles called “atoms.”

Atoms are indivisible. There are different types of atoms for

each material in the world.

Because the early Greek philosophers did not experiment and because Aristotle was an established teacher and because the church was opposed to “soul atoms”, the views of Democritus were not accepted until the 19th

century.

Why was Democritus Ignored?

Pre-Atomic Theory Postulates

Law of Conservation of Mass During a chemical reaction, the total mass of the

reactants is equal to the total mass of the products.

Law of Definite Proportions When atoms combine to form compounds, they

always combine in the same simple, small whole number proportions.

Example: Water is always H2O Example: Sulfuric Acid is always H2SO4

John Dalton(early 1803)

Matter consists of tiny particles called atoms which are indivisible and indestructible.

All atoms of a particular element are identical. Atoms of different elements differ in mass and

properties. Atoms combine in whole number ratios to form

compound atoms. In chemical reactions, atoms are combined,

separated, or rearranged but are never created, destroyed, or changed

Why were Dalton’s views accepted?

The scientific method is now the proper way to “do science.”

Dalton’s theory was based on experimental observations: the law of Conservation of Mass and the law of Definite Proportions.

Dalton’s theory correctly predicted the outcome of future experiments. These predictions became the law of Multiple Proportions.

The Dalton Atom

John Dalton examined the empirical proportions of elements that made up chemical compounds.

At this stage, the atom was still seen as an indivisible object, with no internal structure.

Amedo Avogadro

Avogadro, among other achievements, was able to explain the existence of diatomic molecules. Avogadro’s Law: Equal volumes of any gas at

the same temperature and pressure, have the same number of particles.

1 mole = 22.4 Liters

J.J. Thomson set up a crookes tube with a anodic and cathodic ends

When electricity was applied to the tube, a beam was emitted from the cathodic (-) plate

Thomson then assumed the particles emitted were negative

To test this theory, he applied a magnetic field to the tube and “bent” the beam

What happens with like charges?

He tested the tube further by applying an electrical field to the tube using paddles

The tube turned around

Thomson determined that the tube turned as tiny particles hit the paddles

Demonstration

Molecular Expressions: Electricity and Magnetism - Interactive Java Tutorials: Crookes Tube

He concluded that the particles in the tube were negatively charged and had mass

mass = 9.109 x 10-31kg

Since these particles are negatively charged, but the atoms are neutral, there must be other particles in an atom

Problem: This requires too many electrons!

Thomson Model

The discovery of the electron by J. J. Thomson showed that atoms did have some kind of internal structure.

The Thomson model of the atom described the atom as a "pudding" of positive charge, with negatively charged electrons embedded

J.J. Thomson’s Plum Pudding Model

Positively charged “pudding”

Negatively charged particles later named electrons

Thomson movie

Milliken and the Oil Droplet

In 1909, Robert Milliken performed an experiment using droplets of oil to determine the charge of an electron. electrons, e, e-, -1.602 x 10-19C

Ernest Rutherford conducted experiments to test the Thomson model

He directed alpha particles through a thin gold foil and measured them with a film

Most particles went through the foil

•But, some were deflected, Why?

Rutherford’s HypothesisEngland, 1911

Rutherford hypothesized that the particles were travelling through a void and occasionally bouncing off a concentrated positive charge.

Conclusion There must be a dense region with positive

charges surrounded by the electrons An atom is mostly empty space with a dense

region in the middle. This dense region is called the “nucleus” He measured the number of particles deflected

and the angles and calculated that the radius of the nucleus was 1/10,000 of the whole atom

Problem: Electrons should spiral into the nucleus.

Let there be protons!

The discovery was made and protons were recognized

The mass of a proton is 2000x the mass of an electron

1.673 x 10-27 kg

We’re not done yet ...

30 years later, Irene Curie, the daughter of the great Madame Curie, produced a beam of particles that could go through almost anything

And James Chadwick determined this beam was not affected by a magnetic field (no charge!)

Neutrons were given credit

Since like charges repel, how can the nucleus be stable with protons (+) and neutrons (0)?

Coulomb’s Law: the closer two charges are, the greater the force between them

As the distance between like charges decreases, the force between them increases.

Try it!

Coulomb’s Law

Problems with Rutherford’s model

According to classical physics, an electron in orbit around an atomic nucleus should emit photons continuously as they are accelerating in a curved path.

The loss of energy should cause the electron to collide with the nucleus and collapse the atom.

Elemental Quandary

The Rutherford model was unable to explain the difference in the visible spectrum for each element.

Visible-line Spectrum

When an elemental gas is excited by electricity, it emits a distinct visible light pattern.

The color of each spectral line is identified by the wavelength ()

Electromagnetic Spectrum

All of the frequencies or wavelengths of electromagnetic radiation.

Wavelength

The wavelength is the distance between repeating units of a wave pattern (λ) and measured in nm

Frequency

Frequency is the measurement of the number of times that a repeated event occurs per unit of time (Hz)

The blue wave has the greatest frequency.

Hydrogen

Carbon

Oxygen

Xenon

Compare these spectrum

•Hydrogen, Carbon, Oxygen and Xenon

In comes Niels BohrDenmark, 1913

In 1913, Bohr proposed that electrons were restricted to certain fixed circular orbits.

Orbits are energy levels Electrons can jump from ground state to an

excited state by absorbing energy or a photon with the precise wavelength.

Neils Bohr(early 1900’s)

Electrons travel around the nucleus in specific energy levels.

Electrons have a ground state and an excited state Electrons do not radiate energy in their normal

energy level called the ground state. Electrons absorb energy and move to energy

levels further from the nucleus called excited states.

Electrons lose energy (light) as they return to lower energy levels.

The Bohr Atom

NucleusGround State

Excited States

-

+

The Bohr Planetary Atomic Model

The Bohr Atom

In the Bohr Model the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun

The Modern Atom

The modern atom is further defined by the works of these scientists:

de Broglie Max Plank Albert Einstein Heisenberg Erwin Schrodinger

Problems with the Planetary Model

This model only works for Hydrogen

Max PlankGermany, 1918

Energy is gained or lost in discrete “packets” called quanta

Calculated the amount of energy and determined that it is a constant Plank’s Constant hv

Founded quantum mechanics theory He was also an accomplished

musician!

de Broglie, 1924

Electrons move like waves and so have properties of waves.

Albert Einstein Einstein was simultaneously working on the

photoelectric effect, the theory of relativity and the energy-mass relationship.

Heisenberg, 1925 Heisenberg proposed that it is not possible to

know the position and momentum of an electron at the same time. Heisenberg Uncertainty Principle

Erwin SchrödingerAustria, 1920’s

Electrons have characteristics associated with waves and particles; wave-particle duality.

Electrons are located around the nucleus in “orbitals” An orbital is a probability that an electron

will be there 4 quantum numbers indicate the

probable location of the electron wave.

Schrödinger Wave Equation

2/x2 + 2/y2 + 2/z2 + 82m/h2(E-V)=0

(E-V) = 2 2me4/h2n2

The equation predicts the orbital

The Modern Atomic ViewThe Wave-Mechanical Model

Another View

The Theory

No two electrons can have the same quantum number (Pauli Exclusion Principle) No two electrons can occupy the same space

at the same time A quantum number is an address of the

electron Electrons exist in orbitals around the nucleus

Let’s Review

The Dalton Atom

John Dalton examined the empirical proportions of elements that made up chemical compounds.

At this stage, the atom was still seen as an indivisible object, with no internal structure.

Thomson Model

The discovery of the electron by J. J. Thomson showed that atoms did have some kind of internal structure.

The Thomson model of the atom described the atom as a "pudding" of positive charge, with negatively charged electrons embedded

Rutherford Model

The Rutherford model described the atom made up of a dense nucleus of approximately containing positively charged particles, surrounded by an electron cloud of approximately.

“Nuclear Model”

Niels Bohr

The Bohr Model is probably familar as the "planetary model" of the atom, the figure is used as a symbol for atomic energy

The neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun

•Father of Quantum Physics

•Electrons absorb and emit energy in discrete “packets” called quanta

Max Plank

Erwin Schrödinger

Electrons exist in specific orbitals and are assigned separate quantum numbers

Summary

The model of the atom changed over time. How? What? When? Where? Why? Get into your study groups and each student

answer a different question. Write your responses on the bottom of your

notes page.