Period E
Chapter 4 Wiki
The Atom
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THE ATOMSECTIONS (Here is all the info for the class that was on the cards. Feel free to edit it if you have more detail on page numbers.)Coeditor: Brendan Morrissey (Elaney Marcotte secondary) - Structure of the Atom (Pages 104 - 108)Coeditor: Chris Hart (Atomic Mass) Randall Melanson (Atomic Number + Isotopes) Marissa Chura (Mass Number) Pages 110 - 116Coeditor: Katie West (Sizing up the Atom - 103) Brett Chatfield (Early Models of the Atom 101 - 102)Coeditor: Fast Phil Royal (131 - 132) Frank Morely (130) Siri Devlin (127 - 129)Coeditor: Abbey Salvas (133 - 136) Kerry Desmond (135) Austin Burlone (136) Arrangements in AtomsCoeditor: Dan Lynch (144 - 143) Nic Cunha (138 - 140) Delia Calderon (141 - 143)
The atom is the smallest known piece of matter that retains its properties. Since the idea of it was conceived more than two thousand years ago, models of the incredibly small atom have cropped up with many changes to its shape and formation. It has evolved from a single perfect circle to a complex nucleus holding more than twenty different microscopic pieces around it in space. The atom is the most essential part of understanding chemistry and matter. The properties it holds can unlock the secrets of our world and everything in it, from the sun, to the earth, air, sea and everything we can possibly find on it. The atom, the quintessential matter of the world, is the basic building block of everything known to man, including man itself.
Section 4.1

Coeditor: Katie West
Group Members: Brett Chatfield

Early Models of the Atom

With our unaided eyes, we could never see the tiny particles that make up matter called atoms. Atoms are the smallest particles of an element that retain their identities during a chemical reaction. Though early scholars could not see the atom, they predicted them and also their shapes and behaviors.

Democritus’ Atomic Philosophy
The Greek philosopher Democritus (460 B.C. -370 B.C.) was one of the first to suggest these new things called atoms. He came up with this theory: Atoms were indivisible and indestructible. His ideas had some common things with the future scientific theory, but he never mentioned the behavior of the chemicals. He also lacked evidence from the scientific method because he never experimented on his theories.

A basic image of Democritus’ Atomic Philosophy

Dalton’s Atomic Theory
The real behavior and nature of how atoms behave and act was not concluded for more than 2,000 years after Democritus made his philosophy of them. The modern way of seeing atoms began with John Dalton (1766-1844), who was an English chemist and schoolteacher.
He was able to make variations of Democritus’ ideas using experiments to create a scientific theory. He studied the ratios of how chemicals bond during chemical reactions. He made hypotheses and theories
All elements are composed of tiny indivisible particles called atoms. to explain what he has seen. With these observations,
he created Dalton’s Atomic Theory:

  1. 1803_dalton2.jpgAtoms of the same element are identical. The atoms of any one element are different from those of any other element.
  2. Atoms of different elements can physically mix together or can chem
    ically co
  3. mbine in simple whole-number ratios to form compounds.
  4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.

A photo of what Dalton predicted. The red atoms can combine in whole number ratios with the blue atoms to form compounds.

Dalton’s Atomic Theory Clip:

Brett Chatfield 101-102

Sizing up the Atom
A pure metal, like copper, would have the same properties no matter how small you make it. When you can no longer divide it and still have the properties of copper, it is an atom. These atoms are very small – there are 4 x 10^22 times as many atoms in a pure copper coin the size of a penny as there are people on earth.

The radii of most atoms is between 5x 10^-11 and 2 x 10^-10 m. Though there are very small, individual atoms can be observed with certain instruments, like scanning tunneling microscopes. We can even move individual atoms and arrange them into patterns. This ability will lead to creating atomic-sized, or “nanoscale,” technology that could prove essential to medicine, communications, solar energy, and space exploration.
Iron atoms arranged on a copper sheet

Iron atoms arranged on a copper sheet

Kati West p. 103

Pages 104-108: Structure of the Nuclear Atom
Coeditor: Brendan Morrissey
Group Members: Elaney Marcotte

Subatomic Particles
  • One important change to Dalton's atomic theory is that atoms are now known to be divisible
  • Atoms can be broken down into smaller, more fundamental paricles called subatomic particles
  • Three kinds of subatomic particles are electrons, protons, and neutrons

  • J.J. Thomson, an English physicist, discovered the electron in 1897
  • Electrons are negatively charged subatomic particles
  • Thomson performed experiments that involved passing electric current through gases at low pressure
    • He sealed the gasses in glass tubes fitted at both ends with metal discs called electrodes
    • The electrodes were connected to a source of electricity
    • One electrode, an anode, became positively charged, while the other electrode, the cathode, became negatively charged
    • The result was a glowing beam, a cathode ray, that traveled from the cathode to the anode.
      • A cathode ray is deflected by a magnet and electrically charged metal plates
      • A postively charged plate attracts the cathode ray, while a negatively charged plate repels it
        • Thomson knew that opposite charges attract and like charges repel, so he hypothesized that a cathode ray is a stream of negatively charged particles moving at a high speed; he called these particles corpuscles, which would later be called electrons
    • To test the hypothesis, Thomson set up an experiment to measure the ratio of the charge of an electron to its mass, which would later be found as constant. The charge-to-mass ratio of electroons did not depend on the kind of gas in the cathode ray tube or the type of metal used for the electrodes
    • Thomson concluded that electrons must be parts of the atoms of all element
  • U.S. Physicist Robert A. Millikan carried out experiments to find the quantityof charge carried by an electron. Using this value and the ratio by Thomson, Millikan calculated the mass of the electron.
  • Millikan's values for electron charge and mass from 1916 are very similar to those accepted today, as it carries one unit of negative charge and its mass is 1/1840 the mass of a hydrogen atom
Brendan Morrissey (Pages 104-105)

Protons and Neutrons pg. 106Subatomic_Particle_Chart.jpg
  • Atoms have no net electric charge and are neutral
  • Electric charges are carried by particles of matter
  • Electric charges always exist in whole number multiples of a basic unit
  • A neutral particle is formed when the number of positive and negative particles are equal
  • In 1886 Eugen Goldstein used a cathode-ray tube to observe particles
    • He observed rays traveling in the opposite direction of the cathode rays
    • He called these canal rays and stated that they were composed of positive particles
      • These are now called protons
        • They have a mass 1840 times the size of an electron
  • 1932 James Chadwick confirmed the existence of the neutron
    • Neutrons
      • Particles with a mass equal to a proton
      • Have no charge
The Atomic Nucleus pg. 106

Early scientists wondered how particles were arranged in an atom. Many, including J. J. Thomson thought electrons were distributed evenly throughout the atom which was filled with positively charged material. Thomson believed in a “plum pudding model” where electrons were stuck in a positively charged area. The idea comes from raisins stuck in dough to make plum pudding. Ernest Rutherford, who was a student of Thomson, created an experiment that changed all knowledge of the atom.
Plum Pudding Model

Rutherford’s Gold-Foil Experiment pg. 107
  • In 1911 Rutherford and others at the University of Manchester, England set up an experiment to test the theories of the atomic structure
  • The experiment used massive alpha particles
    • Alpha particles are helium atoms that lose both their electrons and have a double positive charge because they have two remaining protons
    • A beam of alpha particles was fired at a thin sheet of gold foil
    • According to Thomson’s model, the alpha particles should have passed through the gold easily
      • Only a slight deflection would happen because they believed a positive charge was spread out in the gold atom
        • This would cause the positive alpha particles to be repelled from the positive charge in the gold atom
    • Many alpha particles passed through the gold without deflection
    • A small amount of the alpha particles bounced off the gold at large angles or straight back

A video about Rutherford’s experiment: Rutherford's Experiment: Nuclear Atom - YouTube

The Rutherford Atomic Model pg. 107-108
Rutherford's Gold-Foil Experiment

  • Because of the surprising results of the experiment, Rutherford came up with a new atomic structure model
  • He believed that the atom is mostly empty space
    • This explained the lack of deflection and why many of the alpha particles were able to pass easily through the gold foil
  • He found that all the mass and positive charge of the atom is concentrated in a small region
      • This has enough charge to repel some of the alpha particles and deflect them and send them straight back during the experiment
        • He named this the nucleus
          • A nucleus is the central core of the atom that is made up of protons and neutrons
  • Rutherford’s atomic model is called the nuclear atom
    • In this the protons and neutrons are in the nucleus
    • The electrons are found around the nucleus and occupy the majority of the volume of the atom
    • The nucleus is very tiny compared to the rest of the atom
      • If the atom were the size of football field, then the nucleus would be the size of a marble
  • This model tuned out to be incomplete
  • It had to be revised to explain the chemical properties of the elements
Elaney Marcotte (Pages 106-108)

Pages 110-116: Distinguishing Among Atoms
Coeditor: Chris Hart
Group Members: Randall Melanson, Marissa Chura

Atomic Mass

Chris Hart (Pages 114-116)

The mass of even the largest atom is incredibly small. Since the 1920's it has been possible to determine these tiny masses by using a mass spectrometer. With this instrument, the mass of a flourine atom was found to be 3.155 x 10^-23 g. Such data about the actual masses of individual atoms can provide useful information, but, in general, these values are inconveniently small and impractical to work with. Instead, it is more useful to compare the relative masses of atoms using a reference isotope as a
standard. The isotope chosen is carbon-12. This isotope of carbon was assigned a mass of exactly 12 atomic mass units. An atomic mass unit (Amu) is designed as one twelfth of the of a carbon-12 atom. A helium-4 atom, with a mass of 4.0026 amu, has about one third the mass of a carbon-12 atom.
A carbon-12 atom has six protons and six neutrons in its nucleus, and its mass is set as 12 amu. The six protons and six neutrons account for nearly all of this mass, so the mass of a single proton or a single neutron is about one twelfth of 12 amu, or about 1 amu. In nature, most elements occur as a mixture of two or more isotopes. Each isotope of an element has a fixed mass and a natural percent abundance. Consider the two stable isotopes of chlorine, Chlorine-35 and Chlorine-37. If you calculate the arithmetic mean of these two masses ((34.969 amu + 36.966 amu)/2), you get an average atomic mass of 35.958 amu. However, this value is higher than the actual value of 35.453. To explain this difference, you need to know the natural percent abundance of the isotopes of chlorine. Chlorine-35 accounts for 75% of the naturally occuring chlorine atoms; chlorine-37 accounts for only 25%. The atomic mass of an element is a weighted average mass of the atoms in a naturally occuring sample of the elemtn. A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature.
By knowing the atomic mass of an element is a weighted average of the mass of its isotopes, you can determine atomic mass based on relative abundance. To do this, you must know three values: the number of stable isotopes of the lement, the mass of each isotope, and the natural percent abundance of each isotope. To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. The resulting sum is the weighted average mass of the atoms of the element as they occur in nature. You can calculate atomic masses based on the given masses and natural abundances of the isotopes for each element. For example, carbon has two stable isotopes: carbon-12, which has a natural abundance of 98.89%, and carbon-13, which has natural abundance of 1.11%. The mass of carbon-12 is 12.000 amu; the mass of carbon-13 is 13.003 amu. The atomic mass is calculated as follows:
Atomic mass of carbon = (12.000 amu x 0.9889) + (13.003 amu x 0.0111) = 12.011 amu

Atomic Number

Atoms are composed of protons, nuetron and electrons. Proton and nuetrons are found within the nucleus of an atom and electrons surround the nucleus in the electron cloud. All atoms are made of a nucleus and an electron cloud, however the elements differ because they contain a different amound of protons. -The atomic number of an element is the number of protons in the nucleus. -Example: hydrogen has one proton, the mass number is one. oxygen has eight protons, the mass number is eight. lithium has three protons, the mass number is three.
external image atomic_number.jpg
external image atomic_number.jpg


All atoms of the same element have the same number of protons and electrons, but the number of nuetrons vary. -Isotopes are atroms of the same number of protons but different number of neutrons. -Because elements have different number of nuetrons, they also have different mass numbers. -Isotopes are chemically alike due to the same number of protons and neutrons which are the subatomic particles responcible for chemical behavior.
external image images?q=tbn:ANd9GcTAc9xWRBdbFAZi-3s8zy_8yKtCduOtcTxihD3AYAz6ya08xJJ22Rs7DaR7
external image images?q=tbn:ANd9GcTAc9xWRBdbFAZi-3s8zy_8yKtCduOtcTxihD3AYAz6ya08xJJ22Rs7DaR7
There are three known isotopes for hydrogen. Each isotope has one proton in the nucleus, however a different number of neutrons in the nucleus.
-Hydrogen with no neutrons= hydrogen-1 or simply hydrogen -Hydrogen with one nuetron and the mass number of two= hydrogen-2 or deuterium -Hydrogen with two nuetrons and a mass number of three= hydrogen-3 or tritium

Marissa Chura (Page 111)
Mass Number
Most of an atom's mass is in its nucleus and depends on the number of protons and neutrons. The total number of protons and neutrons in an atom is called the atom's mass number.
Helium: 2 protons, 2 neutrons. Mass number=4.
Carbon: 6 protons, 6 neutrons. Mass number =12.
If you know the atomic number and mass number, you can determine an element's composition.
Oxygen: Atomic Number=8, Mass Number=16.
Atomic number=number of protons=number of electrons. So an oxygen atom has 8 protons and 8 electrons.
Number of neutrons=mass number-atomic number
Neutrons=16-8. So an oxygen atom has 8 neutrons.
You can refer to atoms with the element name and mass number. For example, you could write oxygen-16, carbon-12, or helium-4.

Section 5.1- Models of the AtomSiri Devlin,Frank Morelyand Phillip Royal
Pages 127-132Coeditor: Phillip Royal
The Development of Atomic Models

The Development of Atomic Models~The discovery of the nucleus led to entirely new models of the atom~Rutherford used previous ideas to make his own atomic model.~His model had the electrons move around the nucleus like the planets around the sun.
~His model was not able to describe the chemical properties of any element, so new models needed to develop.
~These new models needed to better describe electron movement

The Bohr Model
external image bohrs_model.gif

- Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus
Energy levels -the fixed energies an electron can have
Quantum- the amount of energy required to move an electron from one energy level to another

  • The amount of energy an electron gains or loses in an atom is not always the same. The higher energy levels are closer together
The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level.

v The Bohr Model has failed in many ways to explain the energies absorbed and emitted by atoms with more than one electron
- Atomic orbitals are not actually circular, but come in a variety of different shapes (S,D,F,P)

The Quantum Mechanical Model
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Quantum mechanical model- modern description of the electrons on atoms, based on Schrodinger equation
- The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus

Atomic Orbitals
Atomic Orbital- is often thought of as a region of space in which there is a high probability of finding an electron
Principal energy level
Number of sublevels
Type of sublevel
N= 1
1s ( 1 orbital)
N= 2
2s ( 1 orbital), 2p ( 3 orbitals)
N= 3
3s ( 1 orbital), 3p ( 3 orbitals), 3d (5 orbitals)
N= 4
4s (1 orbital), 4p (3 orbitals), 4d ( 5 orbitals), 4f (7 orbitals)
  • Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found

external image image002.jpg

Section 5.2: Arrangement in Atoms

Coeditor: Abbey Salvas (133 - 134) Kerry Desmond (135) Austin Burlone (136)

Electron Configurations (pg. 133-134)

-in most natural phenomena, change proceeds toward the lowest possible energy
-in an atom, electrons and the nucleus interact to make most stable arrangement possible
-electron configurations: the ways in which electrons are arranged in various orbitals around the nuclei of atoms
-three rules tell how you find the electron configurations of atoms
1. aufbau principle
2. the Pauli exclusion principle
3. Hund’s rule
-Aufbau Principle
-electrons occupy orbitals of lowest energy first
-orbitals for any sublevel of a principal energy level are always of equal energy
-within a principal energy level the s sublevel is always the lowest-energy sublevel
-the range of energy levels within a principal energy level can overlap the energy levels of another principal level
-Pauli Exclusion Principle
-an atomic orbital may describe at most two electrons
-either one or two electrons can occupy an s orbital or a p orbital
-to occupy the same orbital, two electrons must have opposite spins; that is, the electron spin must be paired
-spin is a quantum mechanical property of electrons and may be thought of as clockwise and counterclockwise
-a vertical arrow indicates an electron and its direction of spin
by Abbey Salvas

Kerry Desmond (pg. 135)

Hund's Rule

-electrons occupy all orbitals of equal energy before pairing up.

-they're negatively charged; they don't want to be anywhere near each other unless it's necessary
-if there is an orbital of equal energy, they will use that, because they repel each other.
-if they are in the same orbital, they spin in opposite directions.
A video about the Aufbau Principle, Pauli Exclusion Principle, and Hund's rule

Austin Burlone (136)

Austin Burlone
P. 136
Exceptional Electron Configurations
-One is able to figure out correct electron configurations for elements up until vandium by using the aufbau diagram
-Filled energy sublevels are more stable than partially filled sublevels
-Some actual electron configurations are different from those assigned from the aufbau principle since half-filled sublevels are less stable then filled sublevels, yet more stable than other configurations
-Yet there are some exceptions to the aufbau principle (caused by subtle electron-electron interactions in orbitals with almost the same energies)
-The higher the quantum numbers are, the smaller the energy differences may be
-Other exceptions to the aufbau principle can occur as well
Clip to Electron Configuration Video
Section 5.3

Nicholas Cunha
P. 138-140


The previous sections spoke about orbitals and how they are organized in an atom, each with particular energy levels. You also learned how to write electron configurations for atoms. In the remainder of this chapter you will get a closer look into what led to Schrodinger’s equation and the quantum mechanical model of the atom. The quantum mechanical model came out of the study of light. Isaac Newton tried to explain the behavior of light through theorizing that light was made up of many particles. By 1900 however, there was enough experimental evidence to conclude that light was made up of waves. Each wave cycle starts at zero, increases to its highest value, passes through zero to its lowest value, and returns to zero again. The amplitude of a wave is the wave’s height from zero to the crest. The wavelength represented by the Greek letter lambda, λ, is the distance between crests. The frequency represented by the Greek letter nu, ν, is the number of wave cycles that pass a given point per unit of time. This is usually measured in cycles per second, which is measured with an SI unit called Hertz (Hz). A hertz can also be expressed as a reciprocal second (s-1). The product of frequency and wavelength always equals a constant c the speed of light.
c=λνWave Analysis and Frequency Comparison Charts(below is an Electromagnetic Spectrum Diagram)
wavelength_amplitude.gif wavelength_frequency.gif

The wave length and frequency of light are inversely proportional- as the wavelength of light increases, the frequency decreases for example. According to the wave model light consists of electromagnetic waves. Electromagnetic radiation includes radio waves, micro waves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays. In a vacuum they all travel at the speed of 2.998 x 108 m/s.Sunlight consists of light with a continuous range of wavelengths and frequencies. The colors of each frequency found in sunlight depend on the frequency. When light passes through a prism, the different frequencies separate into a spectrum of colors such as what happens in the atmosphere with water droplets acting as prisms creating a rainbow. Each color blends into the next in the order of red, orange, yellow, green, blue, and violet (in the visible spectrum). Red has the largest wavelength and lowest frequency in the visible spectrum.

Practice problems for Calculating the Wavelength of Light
Calculate the wavelength of the yellow light emitted by the sodium lamp shown if the frequency of the radiation is 5.10 x 1014 Hz (5.10 x 1014/s)
1. Lists the knowns and unknowns

- Frequency (v) = 5.10x1014/s
- c = 2.998 x 108 m/s
- wavelength (λ)= ?m

2. Solve for the unknown
Solve the equation c=λν for c
- λ= c

Substitute the known values to solve.
- λ= c/ν = 2.998 x 108 m/s = 5.8 x 10-7m
5.10 x 1014/s


Atomic Spectra

Pages 141-143
Delia Calderon
  • Whenatoms absorb energy, electrons, move into energy levels, and these electrons loseenergy by emitting light when they return to lower energy levels
  • Each specific frequency of visible light emitted corresponds to a particular color
  • The atomic emission spectrum of an element is given when light passes through a prism and the frequencies of light emitted by an element separate into discrete lines.
  • Each discrete line corresponds to one exact frequency of light emitted by the atom.
  • No tw elements o have the same emission spectrum; the spectra are useful for identifying elements.
    emission.gif neonspect.jpg

An Explanation of Atomic Spectra
  • The lowest possible energy of the electron is its ground state; in the ground state the electrons principle quantum number (n) is 1.Spectrum.jpg
  • Excitation of the electron by absorbing energy raises it from the ground state to an excited state.
  • A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level, a single abrupt step called electronic transition.
  • The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron.
  • Three groups of lines observed in the emission
    • Lyman Series- the lines at the ultraviolet end of the hydrogen spectrum, transition from higher energy levels to n=1
    • Balmer Series- the lines in the visible spectrum, transition from higher energy levels to n=2, involves a smaller change in electron energy than transitions to n=1
    • Paschen Series- transition from higher energy levels to n=3, lines are in the infrared range; energy changes and frequencies of emitted light are smaller still.
  • Spectral lines in each group become more closely spaced at increased values of n because the energy levels get closer together
  • Bohr’s theory of the atom was only partially satisfactory, eventually a new and better model called the quantum mechanical model displaced the Bohr model.
  • The quantum mechanical model is based on the description of the motion of material objects as waves.


Quantum Mechanics

Daniel Lynch (Co-Editor) pg. 144-145

Einstein explained that light could be described as a quanta of energy.
The quanta behave as if they were particles.
Light quanta are called photons.
Louis de Broglie reasoned a mathematical expression for the wavelength of a moving particle.
Electrons reflected from the metal surface produced curious patterns.
The patterns were like those obtained when X-rays reflect from metal surfaces.
The electrons believed to be particles were reflected as if they were waves!
Today, wave like properties of beams of electrons are useful in magnifying objects.
The electrons in an electron microscope have much smaller wavelengths than light.
This allows a much clearer enlarged image of a very small object such as a mite,
The mass of the object must be very small in order for its wavelength to be large enough to observe.
For example a 50g golf ball traveling at 40 m/s the wavelength is 3 x 10 ^-34 which is too small to observe. An electron has a mass of 9.11 x 10 ^-28 g if it were moving at a velocity of 40 m/s it would have a wavelength of 2 x 10 ^-5 which is comparable to infrared radiation and is readily measured.
Classical mechanics adequately describes the motion of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms and waves.
Heisenberg uncertainty principle - states that it is impossible to know exactly the velocity and the position of a particle at the same time. This limitation is critcal in dealing with the small particles such as electrons. It does not matter for bigger objects like planes or cars.
To understand this principle, consider how you determine the location of an object. To locate a set of keys in a dark room you can use a flashlight. You see the keys when the light bounces off them and strikes your eyes. Likewise you strike an electron with a photon of light.
The electron had such a small mass that striking it with a photon affects its motion in a way that cannot be predicted.
The act of measuring its position can change its velocity and makes it uncertain.
The discovery of matter waves paved the way for the quantum mechanical description of electron in atoms.

Quantum Mechanics videos found on Youtube

Heisenberg principle video