Chapter 4 and 5

Atomic Structure and Electrons in Atoms

Editor: Shannon Leavey

Group 1 pg 101-108

Co-editor: Nicole Sheehan

Kayla Anghinetti, Bryan Dextradeur, Elena Conroy, and Lindsey Chou

Kayla Anghinetti (Pg 101-102)
Early Models of the Atom
- An atom is the smallest particle of an element that retains its identity in a chemical reaction.
- Democritus’s Atomic Philosophy
-Greek philosopher Democritus was among the first to suggest existence of atoms.
-He believed that atoms were indivisible and indestructible.
- Dalton’s Atomic Theory
-Modern process of discovery regarding atoms began with John Dalton.
-By using experimental methods, Dalton transformed Democritus’s ideas on atoms into a scientific theory.
-Dalton studied ratios in which elements combine in chemical reactions.
-He formulated hypotheses and theories to explain his observations.
-Result was Dalton’s Atomic Theory
-All elements are composed of tiny indivisible particles called atoms
-Atoms of the same element are identical. Atoms of one element are different from those of any other element.
-Atoms of different elements can physically mix together or can chemically combine in simple whole-number rations to form compounds.
-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.

Bryan Dextradeur (Pg 103)

Sizing Up The Atom

-Despite their small size, individual atoms are observable with instruments such as scanning tunnel microscopes.
-An atom is the smallest possible particle of an element that still holds its chemical properties.
-Atoms cannot be divided and still hold the chemical properties of an element.
-Atoms are extremely small.
-There are 4x1012 more atoms in a penny than there are people on Earth.
-The radii of most atoms fall within the range of 5x10-11m to 2x10-10m.
-Individual atoms can be moved around and arranged into patterns.
-The ability move around and arrange other atoms holds future promise for the creation of atomic-sized electronic devices, such as circuits and computer chips.
-This atomic-scale, or "nanoscale" technology, could become essential to future applications in medicine, communications, solar energy, and space exploration.


Bryan Dextradeur

Elena Conroy (pg.104-105)
Subatomic Particles
-most of Dalton's atomic theory is accepted however atoms are now known to be divisible
-atom can be broken down to subatomic particles
-three kinds of subatomic particles are electrons, protons, and neutrons

-electrons are negatively charged subatomic particles
-discovered in 1897 by English physicist J.J. Thomson
-Thomson performed experiments by putting a type of gas at low pressure into a cathode tube and put metal dishes at the end called electrodes
-one was positively charged, called anode
-one was negatively charged called cathode
-the result was a glowing beam that traveled from the cathode to the anode, creating the cathode ray
-by using a magnet, he was able to conclude that the particles in a cathode ray had a negative charge because they were attached to the positive charged magnet/plate
-Thomson found the ratio of the charge of the electron to its mass to remain constant, no matter what the type of gas was used in the cathode tube or what type of metal the plates were
-U.S. physicist Robert A. Millikan experimented to find the quantity of the charge carried by an electron, by using the charge-to-mass ratio
-His findings were close to what’s accepted today-electrons carry one unit of negative charge
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Example of cathode ray: Elena Conroy

Nikki Sheehan (page 106)
Protons and Neutrons:

What remains of atoms after electrons have been lost?

-This problem can be thought through using four simple ideas
-1. atoms have no net electric charge so they are electrically neutral
-something is electrically neutral when you do not recieve a shock when touching it
-2. electric charges are carried by paricles of matter
-3. electric charges always exist in whole-number multiples of a single basic unit (there are no fractions of charges)
-4. when a given number of negitively charged particles combines with an equal number of positively charged particles, an electrically neutral particle is formed
-considering these four points, it shows that a particle with one unit of positive charge should remain when an atom looses an electron
-evidence for a positively charged particle was found by Eugen Goldstein in 1886
-he observed a cathode ray tube and discovered rays traveling in the direction opposite of the cathode rays
-he then called these rays "canal rays" and said they were composed of positive particles called protons
-later, the exist ence of another subatomic particle (the neutron) was proven by James Chadwick
-neutrons have no charge while electrons have a negative charge and protons have a positive charge
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Goldstein's cathode ray tube Nikki Sheehan

Properties of subatomic particles Nikki Sheehan

Lindsey Chou (pg 106-108)
The Atomic Nucleus

Early theories about subatomic particles were inaccurate.
-The small size of atoms made them difficult to study.
- J. J. Thomson’s atomic model (aka “plum-pudding model”) depicted the electrons as being “stuck into a lump of positive charge”
-Thomson believed electrons were evenly dispersed within an atom full of uniform and positively charged material
-A student of Thomson’s, Rutherford, did work that challenged these ideas

Rutherford’s Gold-Foil Experiment
-Rutherford tested the current theory (Thomson’s model) of atomic theory in 1911 at the University of Manchester, England, where he worked
-The experiment used alpha particles = helium atoms that have a double positive charge because they have lost two electrons
-A beam of alpha particles was shot at a thin sheet of gold foil
-Hypothesized results = alpha particles would pass through the foil with a small deflection from the positive charge spread out in the gold atoms

(Rutherford's Gold-Foil Experiment, Lindsey Chou)
Results and Implications:
-Many particles passed through foil without deflection so the atom must have a great amount of empty space
-Some particles were deflected by the foil at large angles, and some even came back to the source so nearly all the positive charge and mass are concentrated in a small area within the atom, giving that area a positive enough charge to deflect particles

The Rutherford Atomic Model
-Atom is mostly empty space
-Nearly all the positive charge and mass are concentrated in a small region, which Rutherford named the nucleus
-Nucleus is the central core, composed of protons and neutrons, of an atom
-Electrons are distributed around the nucleus (See diagram)
-Rutherford atomic model is known as the nuclear atom
-This model turned out to be incomplete and later had to be revised

(Rutherford Atomic Model, Lindsey Chou)
Group 2

Co-editor: Hannah Kumlin

Ellie Kawa, Jessen Foster, Emily Stewart, and Chris Kelly

Ellie Kawa (Pgs 110-113)
Atomic Number
-Atoms are composed of protons, neutrons, and electrons
-Protons and neutrons make up the nucleus
(Ellie Kawa)
(Ellie Kawa)

-Electrons surround the nucleus
-Elements are different because they contain different numbers of protons
-Atomic number- the number of protons in the nucleus of an atom of that element
-Hydrogen atoms have one proton, the atomic number of hydrogen is 1
-Oxygen atoms have eight protons, the atomic number of oxygen is 8
-Atomic number identifies an element
-Atoms are electrically neutral- the number of electrons must equal the number of protons

Mass Number
(Ellie Kawa)
(Ellie Kawa)

-Mass of an atom is concentrated in its nucleus and depends on the number of protons and neutrons
-Mass number- the total number of protons and neutrons in an atom
-Helium atom has 2 protons and two neutrons, the mass number is 4
-Carbon atom has six protons and six neutrons, the mass number is 6
-If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition
-Number of neutrons in an atom is the difference between the mass number and the atomic number
-Composition of an atom can be represented in shorthand notation using atomic number and mass number
-Atomic number-subscript; mass number- superscript
-Can also refer to atoms by using the mass number and the name of the element
-Isotopes- atoms that have the same number of protons but different numbers of neutrons
-Because isotopes of an element have different numbers of neutrons, they also have different mass numbers
-Isotopes are chemically alike because they have identical numbers of protons and neutrons, which are the subatomic particles responsible for chemical behavior

(Ellie Kawa)
(Ellie Kawa)

Atomic Number
Mass Number
Number of Electrons

Hannah Kumlin (Pg 114-116)

Atomic Mass

-The actual mass of a proton or neutron is very small (1.67 X 10^-24 g.) The mass of an electron is 9.11 X 10^-28 g which is minute in comparison. These values show that the mass of even the largest atom is extremely small.
-Since the 1920s, it has become possible to determine these tiny masses by using a mass spectrometer. Although this machine is quite useful when it comes to gathering data about the actual masses of individual atoms and provides useful information, the values that this machine shows us are extremely small and are therefore impractical to work with. Instead, it is more useful to compare the relative masses of atoms using a reference isotope as a standard.


-Say the isotope chosen is a carbon-12 atom. This isotope of carbon was assigned a mass of exactly 12 atomic mass units. Using these units, a helium-4 atom, with a mass of 4.0026amu, has about about one-third the mass of a carbon-12 atom. On the other hand, a nickel-60 atom has about five times the mass of a carbon-12 atom.

*An atomic mass unit is defined as one twelfth of the mass of a carbon-12 atom.*

Because the mass of any single atom depends mainly on the number of protons and neutrons in the nucleus of an atom, you might predict that the atomic mass of an element should be a whole number. However, this is usually not the case because 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.
The atomic mass of an element is a weighted average of the masses of its isotopes. By knowing this, we are able to determine atomic mass based on relative abundance. In order to do this, we need to know three values:
1) the number of stable isotopes of the element
2) the mass of each isotope
3) 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.


Carbon has two stable isotopes: carbon-12, which has a natural abundance of 98.89%, and carbon-13 which has a 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 = (12.000amu X 0.9889) + (13.003 amu X 0.0111) = 12.011 amu)

You can also visit THIS website for another example and more information.
Group 3

Co-editor: Allie Chabot

Neal McGovern, David Monti, Lauren Murphy, and Teresa Lynch

Allie Chabot(pg 133-135)
Electron Configurations

-The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations.
-The three rules that tell you how to find the electron configurations of atoms are:
-Aufbau Principle
-Pauli Exclusion Principle
-Hund’s Rule
(Allie Chabot)
(Allie Chabot)

Aufbau Principle
-The Aufbau Principle states that electrons occupy the orbitals of lowest energy first.
-The orbitals for any sublevel of a principle energy level are always of equal energy.
-Within a principle energy level, the S sublevel is always the lowest energy level.
-The energy levels within a principle energy level can overlap the energy levels of another principle level.

Pauli Exclusion Principle
-The Pauli Exclusion Principle states that an atomic orbital may describe at most two electrons.
- A spin is a quantum mechanical property of electrons. (Allie Chabot)
- Either one or two electrons can occupy an S or P orbital; in order to occupy the same orbital, two electrons must have opposite spins.
-The vertical arrow indicates an electron and its direction of spin.

Hund’s Rule
-Hund’s Rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.
- A short cut method for showing the electron configuration of an atom involves writing the energy level and the symbol for every sublevel occupied by an electron.
- The number of electrons occupying that sublevel is indicated with a superscript.
Ex.) Helium: two electrons in one S orbital= 1S2
Ex.) Oxygen: two electrons in one S, two electrons in two S, and four electrons in two P orbitals= 1S22S22P4
- The sum of the superscripts always equals the number of electrons in the atom.
- When the configurations are written, the sublevels within the same principle energy level are generally written together.

Teresa Lynch (pg. 136)

Exceptional Electron Configurations

-You can find the electron configurations for elements up to vanadium, which has an atomic number of 23.
-If you continue you would assign incorrect configurations to chromium and copper. Their correct configurations give chromium a half-filled d sublevel and copper a filled one.
-Some electron configurations differ from those assigned using aufbau principle (half-filled sublevels are less stable than filled ones, but are more stable than other configurations).
-Exceptions to this principle are due to: electron-electron interactions in similar energy orbitals. (At higher quantum numbers differences between energy in sublevels is smaller than copper and chromium= more exceptions to the principle).

Copper is an exception to the aufbau principle. (Teresa Lynch)

David Monti (pg 138-139)
- The quantum mechanical model grew out of the study of light
- Isaac Newton believed that light consists of particles
- However, in the early 1900’s, it was discovered that light actually exists in waves
- These waves start at zero, increases to the highest value, passes through zero to reach its lowest value, and then returns to zero
- Amplitude: the waves height from zero to the crest
-Wavelength: the distance between crests, represented by 
- Frequency–:the number of wave cycles to pass by a given point per unit of time, represented by v, measured in cycles per second
-Hertz : the SI unit of cycles per second
- The product of frequency and wavelength always equals a constant (c), the speed of light

c = lv
- The wavelength and frequency of light are inversely proportional to each other. Ex (As the wavelength of light increases, for example, the frequency decreases).
- Light consists of electromagnetic waves
- Electromagnetic radiation: includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays
- The color of light for each frequency found in sunlight depends on its frequency
- When sunlight passes through a prism, the different frequencies separate into a spectrum of colors
- The spectrum colors blend to the next: red, orange, yellow, green, blue, and violet
- Red has the longest wavelength and violet has the shortest

Teresa Lynch (pg 141)

Atomic Spectra
·-Electrons move into higher energy levels when atoms absorb energy. These electrons lose energy by emitting light when they return to low energy levels.
-Ordinary light= mixture of all the wavelengths of light
-Light emitted by atoms= mixture of specific frequencies
-Atomic Emission Spectrum is the pattern formed when light goes through a prism (or diffraction grating) to separate the different frequencies of light it contains. (Each line corresponds with one frequency)
-No two elements have the same atomic emission spectrum and they are often used to identify the element.


Diagram of the atomic emission spectrum (hydrogen). (Teresa Lynch)

Lauren Murphy(pg.142-143)

An Explanation of Atomic Spectra:
Key concept: 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.
-The lowest possible energy of the electron is its ground state
-In this state the electrons principle quantum number (n) is 1
-Excitation of the electron by absorbing energy raises this number
-A quantum of energy in the form of light is emitted when the electron drops to a lower energy level
-This is a single abrupt step called an electronic transition
-Each transition produces a line of a specific frequency in the spectrum


-Above is an image depicting the three groups of lines observed in the emission spectrum of hydrogen atoms
-The lines in the Lyman series match unexpected values for the emission due to the transition of electrons from higher energy levels to n= 1
-The lines in the Balmer series result from transitions from higher energy levels to n = 2
-This usually is a smaller change in electrons then those from the Lyman series
-Transitions to n= 3 is the Paschen series
-The energy changes are still generally smaller
-Spectral lines in each group get closer together as n increases
-There is an upper limit to the frequency of emitted light for each set of lines
-The upper limit exists because an electron with enough energy completely escapes the atom.

Neal McGovern(pg. 144-145)

Quantum Mechanics
-In 1905, Albert Einstien explained experimental data byproposing that light could be described as a quanta of energy called photons
-This dual wave-particle behavior of light was hard for classical scientists to accept
-In 1924, Louise de Broglie wondered if particles of matter behave as waves as well as light?
-He came up with a mathematical expression for the wavelenghth of a moving partivle, but was unable to prove his theory
-Clinton Davisson and Lester Germer were able to prove it in 1927, though, by shooting beams of electrons at metals.
-The pair noticed that when the electrons reflected off the beams they reflected in a wavelike pattern, and it proved that de Broglie was right.
-Today, the wave-like properties of electrons are used in microscopes to examine extremely small things such as dust mights.
-De Broglie's equation states that every moving object has a wavelike behavior, but the mass of the object must be extremely small for it to be observed.
-German physicist Werner Heisenberg states that it is impossible to know both the position and the velocity of a particle at the same time in the Heisenberg uncertainty principle.
-For example, you use a flashlight to find a set of keys in a dark room, and the light reflects off the keys into your eyes.
-To find an electron, you must strike it with a photon of light. Because an electron has such a small mass, the photon striking it affects it far more than the keys.
-The effect the photon has on the electon cannot be predicted exactly, so the exact and velocity of the electron is affected when it is hit with light, hence, the velocity is uncertain.
-This discovery paved the way for Schrödingers theory, which included the concept of electron orbitals and configurations, the wavelike pattern of matter, and the uncertainty principle

external image 200px-Heisenberg_gamma_ray_microscope.svg.png
Heisenberg's gamma-ray microscope for locating an electron (shown in blue). The incoming gamma ray (shown in green) is scattered by the electron up into the microscope's aperture angle θ. The scattered gamma-ray is shown in red. Classical optics shows that the electron position can be resolved only up to an uncertainty Δx that depends on θ and the wavelength λ of the incoming light.