Atomic Theory A Brief History Atoms are made up of subatomic particles called protons, neutrons and...
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Transcript of Atomic Theory A Brief History Atoms are made up of subatomic particles called protons, neutrons and...
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.