Atomic Structure

1. Electromagnetic Radiation

Gamma Rays

are emitted as the result of transitions within the atomic nucleus.  They originate from both natural and artificial radioactive materials, nuclear explosions, and a variety of other sources in outer space. they are able to pass through three meters of concrete! .

X-rays, and is energetic enough to pass easily through many materials (including the soft tissues of animals).  The high penetration depths of these powerful waves, coupled with their ability to expose photographic emulsions, has led to their extensive use in medicine to investigate textures in the human body, as a therapy, and as a surgical tool.  Uncontrolled exposure to X-rays, however, can lead to mutations, chromosomal aberrations, and other forms of cell damage.

Ultraviolet Radiation

Ultraviolet light wave energies can ionize the atoms of a number of gas molecules present in the Earth’s atmosphere, and this is the process by which the ionosphere is created and sustained.  Although small doses of this relatively high-energy light can promote the production of vitamin D in the body and cause minimal tanning of the skin, too much ultraviolet radiation can lead to serious sunburn, permanent retinal damage, and the promotion of skin cancer.It is used to kill micro-organisms in water and food, as a photocatalyst for caged compounds, and hardens casts in medical treatments

Infrared Radiation

The human body does not emit visible light, but it does emit weak infrared radiation (often abbreviated as IR), which is felt and can be recorded as heat.  Molecules of all objects existing above absolute zero (-273 degrees Celsius) emit infrared rays, the amount of emission generally increasing with temperature.  Approximately half of the sun's electromagnetic energy is emitted in the infrared region, and household items such as heaters and lamps also produce large quantities. Common tools that rely on detection of infrared radiation include night vision scopes, electronic detectors, sensors in satellites and airplanes, and astronomical instrumentation. The “heat-seeking missiles” used by the military are also guided by infrared detectors.  In outer space, infrared wavelengths of radiation map the celestial dust between stars, as evidenced by the large dark patches visible from Earth when viewing the Milky Way Galaxy.  In the household, infrared radiation plays a familiar role in heating and drying clothes, as well as allowing the remote control operation of garage doors and home entertainment systems.

Microwaves

constitute the highest frequency radio waves, and are emitted by the Earth, buildings, cars, airplanes, and other large objects.  In addition, low-level microwave radiation permeates space, where it is speculated to have been released by the Big Bang during the creation of the universe.  Higher frequency microwaves are the basis for RADAR, an acronym that stands for RAdio Detecting And Ranging, a transmission and reception technique used in tracking large objects and calculating their speed and distance.  Astronomers utilize extraterrestrial microwave radiation to study the Milky Way and other nearby galaxies.  Microwaves are also employed for transporting information from Earth to orbiting satellites in vast communications networks, for relaying information from ground-based stations over long distances, and in terrain mapping.

Radio Waves

have very little energy, so they appear to transfer energy in a smooth,  continuous fashion.  That is why they are used in the transmission of signals for radio and television stations.  Important in magnetic resonance imaging (MRI), radio waves are also produced by stars in distant galaxies, and can be detected by astronomers using specialized radiotelescopes.

Visible light

(the part that our eyes see) represents only about 2.5% of the entire (spectrum of electromagnetic radiation.  The visible region of the electromagnetic spectrum is perceived as colours ranging from low-energy, long-wavelength deep red (wavelength of 780 nanometers) to relatively high-energy, short-wavelength deep violet (400 nanometers).

2. Spectroscopy

Emission Spectroscopy

When a sample of hydrogen gas receives an electrical spark, some of the H-H bonds  in the sample are broken and the resulting hydrogen atoms become “excited.”  In order to get rid of the extra energy, light of various wavelengths is emitted, producing what is called an emission (line) spectrum.

Absorption Spectroscopy

absorption spectroscopy is the study of electromagnetic radiation absorbed by atoms or molecules that changes energy levels (i.e. dark lines within a continuous spectrum represent the wavelengths of light that have been absorbed by a sample).  Often, absorption spectroscopy is used to identify the elements present in stars.  It can also be used to determine the concentration of chemical compounds in samples. 

Fireworks

Energy is supplied to a fireworks system when a fuse ignites the chemicals.  The high 

temperature of the explosion vaporizes the atoms or molecules, allowing the electrons to move to an excited energy state.  As the energy that they have absorbed isn’t constant, a line emission spectrum is produced by excited atoms as they quickly return to the ground state, shedding the extra energy as light of one or more characteristic wavelengths.  Falling electrons release photons or light with an energy equivalent to the distance of the drop between orbitals.  Blue is the product of a large drop in energy level, while red is produced as a result of a smaller drop.

The following is a selection of the colours of elements that are used to generate colourful 

fireworks:

3. Atomic Models

Daltons Atomic theory

1. All elements are composed of tiny, invisible particles called atoms.

2. Atoms of the same element are identical, and those of different elements are different.

3. Atoms of different elements can be physically mixed or chemically combined in simple whole-number ratios to form compounds.

4. Well, chemical reactions occur when Adam gets separated, joined or rearranged, atoms of one element are never changed into atoms of another element as a result of chemical reactions.

Discovered electrons.  Realizing that the accepted atomic model did not include electrons, Thomson proposed the “plum pudding” model.  Here, negatively charged electrons were stuck in a lump of positively charged material, similar to raisins stuck in dough.  

Ernest Rutherford (1871-1937) 

Given the small size of atoms, this question proved difficult to answer.  In his 1911 gold foil experiment, he and his co-workers at the University of Manchester (England) aimed a beam of alpha particles at a sheet of gold foil surrounded by a fluorescent screen.  The team found that most of the particles passed through the foil without deflection.  A few, however, were greatly deflected.Rutherford concluded that the atom is mostly empty space (which allowed the alpha particles to easily pass through the gold foil), with a positive charge concentrated in the center, or nucleus (which caused the deflection of some of the alpha particles).While Rutherford’s model of protons making up a nucleus surrounded by electrons worked very well to explain a few simple properties of atoms, it did not explain, for example, why metals give off characteristic flame colours.

Johannes Rydberg (1854-1919)

he proposed a simple relationship relating the various lines in the spectra of the elements.  His expression included a constant that later became known as the Rydberg constant.

Neils Bohr (1885-1962) 

he proposed that electrons are arranged in orbits (concentric circular paths) around the nucleus, similar to the planetary model.In this model, electrons following a particular path avoided falling into the nucleus by having fixed energies (i.e. if they didn’t lose energy, they wouldn’t fall into the nucleus).  The energy level of an electron was said to be the region around the nucleus where it was most likely to be found.Bohr proposed that when radiation with a specific amount of energy was absorbed by a hydrogen atom in the ground state (the state of lowest possible energy), its electron would jump from the ground state to a higher, unstable energy level (called the excited state).  The electron would eventually lose energy and change to a lower energy level by emitting energy in the form of light.  (Remember that the light emitted by each electron transition would have a characteristic frequency or color based on the amount of energy that was released, resulting in a line spectrum.)The fixed energy levels of electrons may be compared to the rungs of a ladder; the lowest rung corresponds to the lowest energy level, and one can climb up or down the ladder by moving from rung to rung.  Just as a person on a ladder cannot stand between the rungs, an electron cannot exist between energy levels.  To move from one energy level to another, an electron must gain or lose just the right amount of energy, just as someone moving from one ladder rung to another must move just the right distance.

The amount of energy required to move a hydrogen atom’s electron from its present state to the next higher state is called a quantum of energy.

Max Planck (1858-1947) 

(The phrase “quantum leap”--used to describe an abrupt change--comes from this concept).  This idea contradicted the assumption at the turn of the century that matter could absorb or emit any quantity of energy.  Planck’s theory that energy could occur only in small packets of energy was the basis of Bohr’s atomic theory.In general, the higher an electron is on the energy “ladder,” the greater its energy and the farther it is from the nucleus.  Unlike the rungs of a ladder, however, the energy levels in an atom are not equally spaced, so the amount of energy lost or gained by an electron is not always the same.  Similar to a set of steps that become closer together as you climb higher, the energy levels get closer together as you move away from the nucleus.

Louis de Broglie (1892-1987)

He reasoned that if light waves can behave like a stream of particles, then perhaps particles like electrons could similarly behave like waves (a.k.a. “wave-particle duality”).  In his discussions, he related the circumference of an atomic orbit to whole number wavelengths of an electron travelling around the nucleus. 

Heisenberg’s Uncertainty principle

states that “It is impossible to know simultaneously both the momentum and position of a particle with certainty.”  In other words, the more accurately we know a particle’s position, the less accurately we can know its momentum (and vice versa).  Applied to the electron, this principle implies that we cannot assume that electrons move in a well-defined orbit (as in the Bohr model). 

Erwin Schrödinger

used the concept of wave-particle duality and complex differential calculus to develop an equation that described the location and energy of an electron in a hydrogen atom.  Unlike the Bohr model which defines an exact path for the movement of an electron, the quantum mechanical model estimates the statistical probability of finding an electron in a particular space in the atom (called an orbital).  In illustrations, the probability of finding an electron within a certain space around the nucleus is represented by a fuzzy cloud.  Where the probability of finding an electron is high, the cloud is more dense, and vice versa

4. Periodic Trends/ Electron Configurations

Orbitals 

Regions in space in which it is most probable to find electrons with a specific amount of energy, not to be confused with the term "orbit"

Principal quantum number (n)

can take the values 1, 2, 3, 4, 5, 6, or 7 (just like sections in a theatre).  It indicates the energy level in which an electron is located.  The larger the value, the further the orbital is from the nucleus and the higher its energy.

l quantum number

the quantum number identifies the sublevel within an energy level where an electron can be found

m quantum number

represents the orbital within the sublevel where the electron is located

Electronegativity

It may be seen from the preceding table that electronegativity increases from left to right across a period.  For the representative elements (the s and p blocks, a.k.a. the main group elements), electronegativity decreases as you go down a group.  Unfortunately, the electronegativities of the transition metals (the d group) are not as predictable

Atomic Radius

• Defined as half the distance between adjacent nuclei in a crystal for metals

• Defined as half the distance between chemically bonded identical atoms for non-metals

• Decreases from left to right across the periodic table

• Increases down a group

Ionic Radii

The size of both positive and negative ions decreases as you move across a period, but increases as you move down a group (because electrons are in higher principal energy levels).

Ionization Energy

• The energy required to remove an electron from an atom

• High electronegativity leads to high ionization energy

• Decreases down a group

• Increases from left to right across the periodic table