Download Atomic Theory

Learning Target:
Atomic Theory timeline
Let’s Take a Trip Through Time!
Democritus: 460 – 370 B.C.
First Concept of an Atom
 There are various basic
elements from which all
matter is made
 Everything is composed
of small atoms moving in
a void
 Some atoms are round,
pointy, oily, have hooks,
etc. to account for their
properties
 Ideas rejected by leading
philosophers because
void = no existence
John Dalton
1766-1844
 Introduced his ideas in 1803
 Each element is composed of
extremely small particles called
atoms
 All the atoms of a given element
are identical, but they differ
from those of any other element
 Atoms are neither created nor
destroyed in any chemical
reaction
 A given compound always has
the same relative numbers and
kinds of atoms
J.J. Thompson
1856-1940
 Discovered electron
1897 – Cathode Ray
Experiment
 Plum Pudding model
1904
 Electrons in a soup of
positive charges
 Discovered isotopes
1913
Cathode Ray Tube Experiment
Ernest Rutherford
1871-1937
 Nucleus Theory 1910
 alpha particle gold foil
experiment
 An atom’s mass is
mostly in the nucleus
 The nucleus has a
positive charge
(Moseley)
 Electrons orbit around
the nucleus
Alpha Particle Experiment
A beam of alpha particles bombarded a thin gold foil.
Most of the alpha particles went through, however, a few
of them bounced back, some at large angles.
Learning Targets:
1.
Compare wave and particle nature of light.
2.
Define a quantum of energy, and explain how it is related to an
energy change of matter.
Main Idea
 Light, a form of electromagnetic radiation, has
characteristics of both wave and a particle.
 Rutherford’s model: the atom’s mass is concentrated in
the nucleus and electrons move around it.
 Drawbacks:
 Does not explain how the electrons were arranged
around the nucleus.
 Does not explain why negatively charged electrons aren’t
pulled into the positively charged nucleus.
 In the early 1900’s, scientists observed certain elements
emitted visible light when heated in a flame.
 Analysis of the emitted light revealed that an element’s
chemical behavior is related to the arrangement of the
electrons in its atoms.
The wave nature of light
 Visible light is a type of electromagnetic radiation, a
form of energy that exhibits wave-like behavior as it
travels through space.
 All waves can be described by several characteristics.
 The wavelength (λ) is the shortest distance between
equivalent points on a continuous wave.
 The frequency (v) is the number of waves that pass a
given point per second.
 The amplitude is the wave’s height from the origin to a
crest.
 The speed of light (3.00 x 108 m/s) is the product of it’s
wavelength and frequency, c = λv
 Sunlight contains a continuous range of
wavelengths and frequencies.
 A prism separates sunlight into a continuous
spectrum of colors.
 The electromagnetic spectrum includes all
forms of electromagnetic radiation.
Electromagnetic Wave Relationship:
 c = λv
 c = speed of light in vacuum
 λ = wavelength
 v = frequency
 C = 3.00 x 108 m/s = constant
Example:
 What is the wavelength of a microwave that has a
frequency of 3.44 x 109 Hz? (Hz = 1/s or s-1)
 Ans:
 λ= c/v
 λ = 3.00 x 108 m/s
3.44 x 109 1/s
 λ = 8.72 x 10-2 m
Shortcomings of the wave model
 Fails to explain why some objects emit only certain
frequencies of light at a given temperature.
 In other words, why some metals emit electrons when
light of a specific frequency shines on them –
emissions of different wavelengths.
The
Particle
Nature
of
Light
 Planck’s Conclusion:
matter can gain or lose energy in small, specific
amounts called quanta.
 Quantum: the minimum amount of energy that can be
gained or lost by an atom
 Planck’s constant has a value of 6.626 x 10-34 J●s
 Energy emitted by hot objects was quantized
 relation exists between the energy of a quantum and the
frequency of the emitted radiation.
 The photoelectric effect is when electrons are
emitted from a metal’s surface when light of a
certain frequency shines on it.
 For a given frequency, v, matter can emit or absorb
energy only in whole number multiples.
 Albert Einstein proposed in 1905 that light has a
dual nature.
 Light’s dual nature: a beam of light has a
wavelike and particlelike properties.
 Think of it as bundle of energy called photons.
 Photon: massless particle that carries a quantum of
energy
 Mathematically shown as:
 Ephoton = hv
 Ephoton represents energy
 h is Planck’s constant = 6.626 x 10-34 Js
 v represents frequency
Example:
 What is the energy of a photon from the violet portion
of the Sun’s light if it has a frequency of 7.230 x 1014 s-1?
 Answer:
 E = h●v, v= 7.230 x 1014 s-1
h = 6.626 x 10-34 J•s
 E = (6.626 x 10-34 Js)(2.230 x 1014 s-1)
 E = 4.791 x 10-19 J
Practice…
 Calculate the energy possessed by a single photon of
each of the following types of electromagnetic
radiation. Also, identify the type of radiations.
6.32 x 1020 s-1
2. 9.50 x 1013 Hz
1.
Ans:
1. 4.19 x 10-13 J, gamma radiation
2. 6.29 x 10-20 J, infrared radiation
Assessment …
 What is the smallest amount of energy that can be
gained or lost by an atom?
 Ans: quanta
 What is a particle of electromagnetic radiation with no
mass called?
 Ans: photon
Atomic emission spectra
 Light in a neon sign is produced when electricity
is passed through a tube filled with neon gas and
excites the neon atoms.
 The excited atoms emit light to release energy.
 The atomic emission spectrum of an element is
the set of frequencies of the electromagnetic
waves emitted by the atoms of the element.
 Each element’s atomic spectrum is unique.
Learning Targets:
Compare the Bohr and quantum mechanical models of the atom.
De Broglie’s wave article duality and the Heisenberg uncertainty
principle.
Identity the relationships among energy levels, sublevels and atomic
orbitals.
 Main Idea:
Wavelike properties of electrons help relate atomic
emission spectra, energy states of atoms, and atomic
orbitals.
Niels Bohr
1885-1962
 Planetary Model 1913
 Nucleus surrounded by
orbiting electrons at
different energy levels
 Electrons have definite
orbits
 Utilized Planck’s Quantum
Energy theory
 Bohr correctly predicted the frequency lines in
hydrogen’s atomic emission spectrum.
 The lowest allowable energy state of an atom is called
its ground state.
 When an atom gains energy, it is in an excited state.
 Bohr suggested that an electron moves around the
nucleus only in certain allowed circular orbits.
 Each orbit was given a number, called the quantum
number. Hydrogen’s single electron is in the n = 1 orbit
in the ground state.
 When energy is added, the electron moves to the n = 2
orbit.
Bohr’s Model
 Bohr’s model explained the hydrogen’s spectral lines,
but failed to explain any other element’s lines.
 The behavior of electrons is still not fully understood,
but it is known they do not move around the nucleus
in circular orbits.
The Quantum Mechanical model
 Louis de Broglie (1892–1987) hypothesized that
particles, including electrons, could also have wavelike
behaviors.
 The de Broglie equation predicts that all moving
particles have wave characteristics

represents wavelengths
h is Planck's constant.
m represents mass of the particle.
represents frequency.
Werber Heisenberg 1901-1976
 Heisenberg showed it is impossible to take any
measurement of an object without disturbing it.
 The Heisenberg uncertainty principle states that it is
fundamentally impossible to know precisely both the
velocity and position of a particle at the same time.
 The only quantity that can be known is the probability
for an electron to occupy a certain region around the
nucleus.
Schrödinger
 treated electrons as waves in a model called the
quantum mechanical model of the atom.
 Schrödinger’s equation applied equally well to
elements other than hydrogen.
 The wave function predicts a three-dimensional
region around the nucleus called the atomic orbital.
 Principal quantum number (n) indicates the relative
size and energy of atomic orbitals.
 n specifies the atom’s major energy levels, called the
principal energy levels.
 Energy sublevels are contained within the principal
energy levels.
 Each energy sublevel relates to orbitals of different shape.
 4 types of orbitals:
 s orbital – spherical shape
 p orbital – dumbbell shape
 d orbital
 f orbital
Orbitals
Atomic Theory
JJ
Thomson
Democratus
400 BC
1803
John
Dalton
1904
Niels
Bohr
1910
Ernest
Rutherford
1913
1926
Schroedinger
/ Heisenberg