unit 4: atomic theory structure of the atom (& radioactivity)
TRANSCRIPT
Unit 4: Atomic Theory
Structure of the Atom (& Radioactivity)
Early Atomic TheoriesModels of the Atom
Date scientist discovery________________ 100 BC Democritus/Greeks concept of the atom 1770 Antoine Lavosier Law of conservation of mass 1800 Joseph Proust Law of definite proportions 1803 John Dalton Law of multiple proportions Atomic Model I 1880 William Crookes Cathode Rays (electrons) 1885 Goldstein Canal Rays (protons) 1900 J.J. Thomson Plum Pudding Model
Electron Atomic Model II
1909 Ernest Rutherford nucleus of atom Atomic Model III 1913 Niels Bohr Planetary Model Atomic Model IV 1920- Schroedinger/Planck/ Modern or Wave Model Present DeBroglie/Einstein/etc. Atomic Model V
Early Atomic Theories Atomists and Democritus
Greeks approx. 2,500 years ago Matter was made up of atoms
“atomos” or “Indivisible” particles Seashell experiment—broken into
smaller & smaller pieces
Early Atomic Theories John Dalton
1766-1844; returned to theory of atoms Atoms are like billiard balls (solid spheres)
which can’t be broken down further 4 major postulates
1. All elements are composed of atoms2. Atoms of the same element are identical3. Atoms can physically mix or chemically
combine in simple whole number ratios4. Reactions occur when atoms separate, join,
or rearrange
Dalton’s Model of the Atom
No subatomic particles!
Early Atomic Theories
William Crookes Developed Crookes tube (a.k.a cathode
ray tube) in 1870’s First evidence for existence of electrons
because you could “see” electrons flow and confirm their existence.
Tube is precursor to today’s TV picture tubes
Building the Atom – The Electron J.J. Thomson
Discovered electron in 1897 Discovered positively charged particles
surrounded by electrons Found the ratio of the charge of an
electron to its mass to be 1/1837
Thomson’s Cathode Ray Tube Experiment J.J. Thomson
Video: Cathode Ray Tube Demo
Building the Atom – The Electron J.J. Thomson
Cathode ray tube experiments – advancement of Crookes tube
“plum-pudding model”
Thomson’s Model of the Atom (Plum Pudding Model)
Millikan’s Oil Drop Experiment Robert Millikan
Oil drop experiment Determined the charge and mass of an
electron
Video: Millikan's Oil Drop Experiment
Building the Atom – The Nucleus Ernest Rutherford
Discovered nucleus (dense core of atom) in 1911
Famous Gold foil experiment
Quote from E.R.’s Lab Notebook “It is about as incredible as if you had fired
a 15-inch shell at a piece of tissue paper and it came back and hit you.” -ER
Rutherford’s Gold Foil Experiment
Video Clip: Rutherford Gold Foil Experiment
Rutherford’s Model of the Atom
Building the Atom – The Neutron James Chadwick
Discovered the neutron (no charge, but same mass as proton)
Neutrons help disperse the strong repulsion of positive charges
Nucleus diameter = 10-5 nm Atom diameter = 10-1 nm If Nucleus = basketball -->
then Atom = 6 miles wide!
Building the Atom Niels Bohr
Improved on Rutherford’s work “Planetary model”- positive center is
surrounded by electrons in defined orbits circling the center
Bohr Model of the Atom (Planetary Model)
Bohr Model of the Atom Vocab. Energy level – the location where an
electron is found at a set distance from the nucleus dependent on the amount of energy it has
Ground state – the typical energy level where an electron is found; lowest energy
Excited state – an energy level higher than the ground state for an electron; temporary condition
Schrödinger Model (Quantum Mechanical Model) Quantum Mechanical Model
Erwin Schrödinger; Mathematical model Electron locations are based on probability Electrons are not particles, but waves!
Interactive Simulation: try it!
Defined: Orbital – region where an electron is likely to
be found 90% of the time
Schrödinger Model of the Atom(Quantum Mechanical Model)
Atomic TheoryAtom – the smallest particle of
matter that retains its properties. Smallest individual unit of an element
One atom of hydrogen is different from one atom of carbon.
Subatomic particles – the component parts of an atom: proton, neutron, and electron
Atomic Theory Ion - atom with the same number of
protons but a different number of electrons i.e. an atom with a charge!
If the atom has a (+) charge it has more protons than electrons.
If the atom has a (-) charge it has more electrons than protons.
Subatomic ParticlesSubatomic
ParticleMass and
Abbreviation Charge Location Discoverer
Protonp+
Mass =1 amu+1 Nucleus ----
Neutron nMass =1 amu 0 Nucleus Chadwick
in 1932
Electrone-
Mass ≈ zero amu
-1
Electron cloud
(outside nucleus)
Thomson in 1897
Atomic SymbolsAtomic symbol – the letter or
letters that represent an element.
13
Al
26.981
Atomic number
Atom symbol
Atomic mass or weight
Atomic NumberAtomic number = the number of protons in the nucleus.
(same for every atom of that element)
13Al
26.981
Atomic number
Atom symbol
Atomic mass or weight
Mass Number
Mass Number = # protons + # neutrons
A Boron atom can have:5 p + 5 n = 10 amu
Named as boron-10
Mass number
Atomic number
Calculations w/ Subatomic ParticlesAtomic number = # of protonsMass number = # of protons + # of
neutrons
(For a neutral atom): # of protons = # of electrons
(For a charged ion): Charge = #p+ - #e-
Isotope Notation Isotope (Isotopic Notation)
Mass #
Atomic #
Atomic Symbol
Z
AX
Example: Uranium-238
Example
Example
Sample ProblemWrite the atomic symbols for the following:
The isotope of carbon with a mass of 13
The nuclear symbol when A = 92 and the number of neutrons = 146.
31
Isotopes Isotope – atoms of the same element with
different numbers of neutrons (different mass numbers) Example: Carbon-12 Carbon-14
Atomic mass – weighted average of the masses of all the isotopes of an element
Atomic Mass
The weighted average is the addition of the contributions from each isotope.
Isotopic Abundance is the percent or fraction of each isotope found in nature.
33
Most Abundant Isotope
34
13
Al
26.981
Atomic number
Atom symbol
Atomic mass or weight
Usually can round atomic mass on the periodic table to nearest whole number
(but not always!!)
Example: Determine the average atomic mass of magnesium which has three isotopes with the following masses: 23.98 amu (78.6%), 24.98 amu (10.1%), 25.98 amu (11.3%).
1) Multiply the mass number of the isotope by the
decimal value of the percent for that isotope.
2) Add the relative masses of all of the isotopes to
get the atomic mass of the element.
35
Example: Determine the average atomic mass of magnesium which has three isotopes with the following masses: 23.98 amu (78.6%), 24.98 amu (10.1%), 25.98 amu (11.3%).
36
Now You Try!If 90% of the beryllium in the world has a mass number of 9 and only 10% has a mass number of 10, what is the atomic mass of beryllium?
Radioactivity - Vocabulary
Radioactivity - the spontaneous emission of radiation from a substance
Radiation - rays and/or particles emitted from radioactive material
Nuclear reactions - reactions involving changes in an atom’s nucleus
Radioactivity Radioactive isotopes are unstable
These isotopes decay over time by emitting particles and are transformed into other elements
Particles emitted: Alpha (α) particles: helium nuclei Beta (β) particles: High speed electrons Gamma (γ) rays: high energy light
39
Types of Radiation – α particles
Alpha radiation - stream of high energy alpha particles Alpha particles consist of 2 protons and 2
neutrons and are identical to helium-4 nucleus.
Symbol: 4He 2+
2
Not much penetrating power, travel a few centimeters, stopped by paper, no health hazard
Types of Radiation – β particles
Beta radiation - High speed electrons To form beta radiation a neutron splits into
a proton and an electron The proton stays in nucleus and the
electron propels out at high speed.
Symbol: 0e- 0β
-1 -1
100 times more penetrating then alpha radiation, pass through clothing to damage skin
Types of Radiation – Radiation
Gamma radiation Similar to X rays Doesn’t consist of particles (instead, high
energy light) Symbol:
0 0
Penetrates deeply into solid material, body tissue, stopped by Pb or concrete, dangerous
End of Unit 4!