Attention Groups: Below I have written each group's members names and what section that group is doing. Beside each name is written what pages they are outlining and there is a spot for a title. Please find your name and complete your outline to your section.
GROUP ONE: Please put what pages you are outlining next to your name.
Introduction: Overview of Chapter 4 The atom is the smallest particle of an elelment that retains its identity in a chemical reaction. The Greek philosopher Democritus was one of the first to suggest the existence of the atom. John Dalton then transformed Democritus's ideas on atoms into a scientific theory known as Dalton's atomic theory. Atoms are very small but individual atoms are observable with instruments such as scanning tunneling micrscopes. Three kinds of subatomic particles are electrons, protons and neutrons. In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are distributed around the nucleus and take up almost all the volume of the atom. Elements differ from atoms because elements contain different numbers of protons. The number of neutrons in an atom is the difference between the mass number and atomic number. Isotopes of an element have a different number of neutrons but the same nuber of protons and it is because of this that they have different mass numbers. To caculate 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. Introduction: Overview of Chapter 5 Rutherford's atomic model could not explain the chemical properties of elements. Niels Bohr then created the Bohr model because he thought Ruthford's model needed improvement. Like the Bohr model, the quantum mechanical model of the atom restricts the energy of electrons to certain balues. The aufbau principle, the Pauli exclusion principle and Hund's rule tell you how to find the electron configuration of atoms. Light. The wavelength and frequency of light are inversely proportional to each other. When atoms absorb energy, electrons move into higher energy levels. They then return to lower energy levels when electrons lose energy by giving off light. Albert Einstein sucessfully explained experimental data by proposing that light could be described as quanta of energy. The Bohr atom.
GROUP ONE (Chapter 4.1 and 4.2 pg 101- 109)
Evan Sommerich, PJ, Brandon Boisclair, Dan McCormack, Christian Cooke, Haley Conatser Co Editor- Brandon Boisclair
Early Models of an Atom Evan Sommerich (pg-101)
Atom- the smallest partilce of an element that remains its identity in a chemical reaction.
Democraties Atomic Philosopy- he belived that atoms were indivsible and indestructable. this was later proved wrong.
Title - Sizing up the Atom
PJ (PG - 103) - This section states that you can continually break down a substance until you get to a small particle. This particle is called an atom. The atom still still has the properties of the substance but does not have the physical appearance as it did before.
The radii of most elements fall within the range of 5x10^-11 to 2 x 10^-10
Despite their small size individual atoms can be observed with instruments such as scanning tunneling microscopes (STM).
these microscopes are very powerful
Example of a STM below
example of a STM
(PJ)
Title Brandon Boisclair (pg - pg)
Daltons Atomic Theory
Dan McCormack (pg 102)
the process of recording atoms began with John Dalton (1766-1884) who was an English chemist and school teacher
by using experimental methods he changed Democritus's ideas into a scientific theory
after his experiments he studied which elements combined in chemical reactions and formed a hypothesis to explain the observations
His theory was this:
All elements are composed of tiny indivisible particles called atoms.
Atoms of the same element are identical. the atoms of any 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 ratios to form compounds
Chemical reactions occur when atoms are separated, joined, of rearranged. Atoms of one element however, are never changed into atoms of another element as a result of a chemical reaction
Title Christian Cooke (pg 107-108 pg) 1911- Rutherford and co-workers tested what was then the current theory of the atomic structure. Using alpha particles, a thin layer of them weredirected at a very thin sheet of gold. To their surpirse, the alpha particles went right thru without deflcection as they had previously thought. He then proposed a new theory - the atom was mostly empty space , which explained the lack of deflection. He proposed that the particles and mass are condensed into a small, positively charged region which was called the nucleus. THis would account for deflection because the alpha particles were also positive. The NUCLEUS is a tiny central core of an atom composed of protons and neutrons. Rutherford atomic model is known as the nuclear atom, where protons and neutrons are located in the nucleus. Electrons are distributed around the nucleus and occupy almost all of the atom. tThe nucleus is very small compared to the rest of the cell.
Subatomic Particles
Haley Conatser (pg 104- 105)
Atoms can be broken down into even smaller, more fundamental particles, called subatomic particles
Three kinds of subatomic particles are electrons, protons, and neutrons
Electrons:
- English physicist J.J. Thomson discorved the electron in 1897
- Electrons are negatively charged subatomic particles
- Thomson performed experiments that involved passing electric currents thought gases at low pressure
- He sealed the gases in glass tubes fitted at both ends with metal disks called electrodes. The elctrodes were connected to a soure of electricity. One electrote, the anode, became positively charged. The other electrode became negatively charged.
- The result was a glowing beam, or cathode ray that traveled from the cathode to the anode.
- A cathode ray is deflected by a magnet. A cathode ray is also deflected by electically chartged metal playes.
- To see a cathode ray being deflected by a magnet there is a video here: Click here
- A positively charged plate attracted the cathode ray, while a negatively charged play repels it.
- Thomson hypothesized that a cathode ray is a stream of tiny negatively charged particles moving at high speed
- Thomson originally called these particles corpuscles; later they were named electrons.
- To test his hypothesis he tried to measure the ratio of the charge of an electron to it's mass He found this ratio to be constant and in addition, the charge-to-mass ratio of electrons did not depend on the kind of gas in the cathode-ray tube or the type of metal used for the electrodes.
Robert A. Millikan (1868-1953):
- This U.S. physicist carried out experiments to find the quantity of charge carried by an electrons
- Using the value and the charge-to-mass ratio that Thomson had found he was able to found out the mass of an electron
- What Millikan told us about the charge and mass of the electron is still true today
- An electron carries exactly one unit of negative charge and its mass is 1/1840 the mass of an hydrogen atom
GROUP TWO (Chapter 4.3 and 5.1 pg 110-116, pg 127-132)
Alex Nunan, Zoey Killion, Mike Hanley, Becky Hyatt, Elizabeth Howard Co- Editor: Alex Nunan
Elizabeth Howard page 110
Atomic Number : the number of protons in the nucleus of an atom of that element
Atoms of the First Ten Elements:
Elizabeth Howard page 111
NUMBER OF NEUTRONS = MASS NUMBER - ATOMIC NUMBER
external image atomicsymbol2.gif
Isotopes
Alex Nunan (pg 112-113)
isotopes- atoms that have the same number of protons but different numbers of neutrons:
isotopes of an element have different numbers of neutrons --> different mass numbers
chemically alike because they have the same numbers of protons and electrons (subatomic particles responsible for chemical behavior)
chemical symbols for isotopes:
for example, hydrogen has three known isotopes-
second isotope has 1 neutron & mass number of 2: hydrogen-2 (or deuterium)
third isotope has 2 neutrons & mass number of 3: hydrogen-3 (or tritium)
look at link above to see symbols for hydrogen-1 and isotopes of other elements
Atomic Mass
Becky Hyatt (pg 114-117)
the actual mass of a proton or neutron is very small and cannon be measured using a mass spectrometer
an atomic mass unit (amu) is defined as one twelfth of 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 at 12 amu
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 mass of the atoms in a naturally occuring sample of the element
Conceptual Problem 4.3:
using atomic mass to determine the relative abundance of isotopes
The atomic mass of copper is 63.546 amu. Which of copper's two isotopes is more abundant: copper-63 or copper-65?
1. Analyze- Identify the relevant concepts
define what an atomic mass and a weighted average mass are
2. Solve- Apply the concepts to this problem
because the atomic mass is a weighted average of the isotopes, copper-63 must be more abundant than copper-65
Practice Problem: Boron has two isotopes: boron-10 and boron-11. Which is more abundant, given that the atomic mass of boron is 10.81?
The Development of Atomic Models Zoey Killion (pg127-129)
After discovering the atomic nuclues, Rutheford combined what was known about hte atom by proposing an atomic model in which the electrons move around the nucleus, like the planets move around the son
Rutheford's model explained only a few simple properties of atoms, his model could not explain:
why metals or complounds of metals give off characteristics colors when heated in a flame
why objects - when heated to higher and higher temperatures - first glow dull red, then yellow, then white (as shown in picture below of horseshoe)
Rutheford's atommic model also could not explain the chemical properties of elements
The Bohr Model
Niels Bohr (1885-1962), a young Danish physicist and a student of Rutheford's, believed Rutheford's model needed improvement
In 1913 Bohr changed Rutheford's model to include newer discoveries about how atoms change energy when they absorb or emit light
Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus
This time line shows the development of atomic models from 1800-1935:
1803 - John Dalton's Atom
1897 - Thomson's Model of Atom
1911 - Rutherford's Model of Atom
1913 - Bohr's Model of Atom
1926 - Electron Cloud Model of Atom
energy levels: the specific energies an electron can have
an electron can jump from one energy level to another
the electrons in an atom cannot be between energy levels
to move from one energy level to another, an electron must gain or lose just the right amount of energy
the higher an electron is on the energy ladder, the farther it is from the nucleus
quantum: the amount of energy required to move an electron from one energy level to another energy level (quantum comes from the Latin word quantus, meaning "how much")
the energy of an electron is said to be quantized
the amount of energy an electron gains or loses in an atom is not always the same
the energy levels in an atom are not equally spaced
the higher energy levels are closer together
the higher the energy level occupied by an electron, the less energy it takes to move from one energy level to the next energy level
The Quantum Mechanical Model
Mike Hanley (pg 130-131)
In 1926, the Austrian physicist Erwin Schrödinger (1887-1961) used new results to devise and solve a mathematical equation describing the behavior of the electron in the hydrogen atom.
The modern description of the electrons in the atom, the quantum mechanical model, comes from the mathematical solutions to the Schrödinger equation.
Like the Bohr model, The quantum mechanical model of the atom restricts the energy of electrons to certain values
Unlike the Bohr model, the quantum mechanical model of the atom does not involve an exact path the electron takes around the nucleus.
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.
In the quantum mechanical model, the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented by a fuzzy cloud.
The cloud is more dense when the probability of finding an electron is high.
The cloud is less dense when the probability of finding an electron is low.
There is a slight chance of finding the electron at a considerable distance from the nucleus.
Example of the quantum mechanical model
Atomic Orbitals
Olivia Richardson (pg 131-132)
By solving the Schrodinger equation, you are giving the energies an electron can have which are its energy levels
An atomic orbital is a mathematic expression that describes the probability of finding an electron at various locations around the nucleus and is also thought of as a region of space in which there is a high probability of finding an electron
the energy levels of electrons in the quantum mechanical model are labeled by principal quantum numbers (n)
the assigned values are n=1, 2, 3, 4
There could be several orbitals with different shapes at different energy levels
Energies within a principal energy level constitute energy sublevels
Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found
there are different orbitals with different shapes. s orbitals are spherical, p orbitals are dumbbell-shaped
three different kinds of p orbitals have different orientations in space.
there are 5 kinds of d orbitals, four which are shaped as a clover
the lowest principal energy sublevel (n=1) has only one sublevel called 1s
(n=2) has two sublevels, 2s and 2p, the second principal energy level has four orbitals
(n=3) has three sublevels, 3s, 3p, and 3d, the third principal energy level has nine orbitals
(n=4) has four sublevels 4s, 4p, 4d, and 4f, the fourth principal energy level has sixteen orbitals
GROUP THREE (Chapter 5.2 and 5.3 pg 133-145)
Shannon Lamy, Mitchell Martin, Andrea Luongo, James Payne, Nate Lynch,Dakota Pimentel Co-Editor: Shannon Lamy
Electron Configurations
Mitchell Martin (pg 133-135)
Aufbau Principle : electrons occupy the orbitals of the lowest energy first.
Pauli exclusion principle: Only one or two electrons can occupy a P orbital or S orbital.
Hund's Rule: electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin as large as possible.
Exceptional Electron Configurations
Shannon Lamy (pg 136-137)
Electron Configuration - A statement describing the populations of electronic energy sublevels of an atom.
You can obtain correct electron configurations for the elements up to vanadium from the aufbau diagram for orbital filling.
Orbital filling diagram
Orbital Filling Dragram
The element copper and chromium do not have electron configurations according to the aufbau diagram.
EX. The actual electron configuration for chromium is...
Chromium
Cr 1s2 2s2 2p6 3s2 3p6 3d5 4s1
^^half filled d sublevel.
EX. The actual electron configuration for copper is...
Cu 1s2 2s2 2p6 3s2 3p6 3d10 4s1Copper
^^filled bublevel.
FILLED ENERGY SUBLEVELS ARE MORE STABLE THAN PARTIALLY FILLED SUBLEVELS.
Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.
This is because there are subtle electron-electron interactions
Light
Andrea Luongo (pg 138-139)
Isaac Newton (1642-1727) studied the behavior of light by concluding that light consisted of particles.
In the 1900s, light was discovered to consist of waves, not particles.
Properties of Waves
Each complete wave cycle begins at zero, increases to its highest value, passes zero to reach its lowest value, then returns to zero.
Amplitude- the height of the wave from zero to the crest.
Wavelength- the distance between crests.
Represented by the Greek letter lambda. Lambda:
Frequency- the number of wave cycles to pass a given point per unit of time.
Usually cycles per second.
Represented by the Greek letter nu. Nu:
Hertz- the SI unit of frequency cycles per second.
Written as a reciprocal second or s^-1
As wavelength increases, frequency decreases.
Frequency & Wavelength:
Wavelength & Amplitude:
Frequency, Wavelength, & Amplitude:
Frequency * Wavelength = the speed of light (c)
Light is composed of electromagnetic waves.
Electromagnetic radiation- inlcludes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays.
All electromagnetic waves travel in a vacuum at a speed of 2.998 X 10^8 m/s.
Sunlight has continuous ranges of frequency and wavelengths.
The color of light depends on its frequency.
Spectrum of colors- wavelengths of visible light that are separated when a beam of light passes through a prism.
In the visible spectrum, red light has the longest wavelength and lowest frequency.
Electromagnectic Spectrum:
Atomic Spectra
James Payne (pg 141)
When atoms abosorb energy, electrons move into higher energy levels,
electrons lose energy by emitting light
Frecuencies of light emitted by an element seperate into discrete lines to give the atomic emission spectrum of the element
Each discrete line in an emission spectrum corresponds to exact one frecuenct of light emitted by the atom
Much knowledge of the universe comes from teh studing of atomic spectra of the stars which are glowing hot bodies of gases
Title
Dakota Pimentel (pg. 142-143)
Quantum Mechanics
Nate Lynch (pg 144-145)
-The branch of quantum physics that accounts for matter at the atomic level; an extension of statistical mechanics based on quantum theory
In quantum theories, energy and momentum have a definite relationship to wavelength. All particles have properties that are wave-like (such as interference) and other properties that are particle-like (such as localization). Whether the properties are primarily those of particles or those of waves, depends on how you observe them.
For example, photons are the quantum particles associated with electromagnetic waves. For any frequency, f, the photons each carry a definite amount of energy (E = hf)
Chapters 4 and 5 The Atom
Editor: Caroline Rubino
Attention Groups: Below I have written each group's members names and what section that group is doing. Beside each name is written what pages they are outlining and there is a spot for a title. Please find your name and complete your outline to your section.
GROUP ONE: Please put what pages you are outlining next to your name.Introduction: Overview of Chapter 4
The atom is the smallest particle of an elelment that retains its identity in a chemical reaction. The Greek philosopher Democritus was one of the first to suggest the existence of the atom. John Dalton then transformed Democritus's ideas on atoms into a scientific theory known as Dalton's atomic theory. Atoms are very small but individual atoms are observable with instruments such as scanning tunneling micrscopes. Three kinds of subatomic particles are electrons, protons and neutrons. In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are distributed around the nucleus and take up almost all the volume of the atom. Elements differ from atoms because elements contain different numbers of protons. The number of neutrons in an atom is the difference between the mass number and atomic number. Isotopes of an element have a different number of neutrons but the same nuber of protons and it is because of this that they have different mass numbers. To caculate 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.
Introduction: Overview of Chapter 5
Rutherford's atomic model could not explain the chemical properties of elements. Niels Bohr then created the Bohr model because he thought Ruthford's model needed improvement. Like the Bohr model, the quantum mechanical model of the atom restricts the energy of electrons to certain balues. The aufbau principle, the Pauli exclusion principle and Hund's rule tell you how to find the electron configuration of atoms. Light. The wavelength and frequency of light are inversely proportional to each other. When atoms absorb energy, electrons move into higher energy levels. They then return to lower energy levels when electrons lose energy by giving off light. Albert Einstein sucessfully explained experimental data by proposing that light could be described as quanta of energy.
The Bohr atom.
GROUP ONE (Chapter 4.1 and 4.2 pg 101- 109)
Evan Sommerich, PJ, Brandon Boisclair, Dan McCormack, Christian Cooke, Haley ConatserCo Editor- Brandon Boisclair
Early Models of an Atom
Evan Sommerich (pg-101)
Atom- the smallest partilce of an element that remains its identity in a chemical reaction.
Democraties Atomic Philosopy- he belived that atoms were indivsible and indestructable. this was later proved wrong.
Title - Sizing up the Atom
PJ (PG - 103) - This section states that you can continually break down a substance until you get to a small particle. This particle is called an atom. The atom still still has the properties of the substance but does not have the physical appearance as it did before.Title
Brandon Boisclair (pg - pg)
Daltons Atomic Theory
Dan McCormack (pg 102)
- the process of recording atoms began with John Dalton (1766-1884) who was an English chemist and school teacher
- by using experimental methods he changed Democritus's ideas into a scientific theory
- after his experiments he studied which elements combined in chemical reactions and formed a hypothesis to explain the observations
His theory was this:- All elements are composed of tiny indivisible particles called atoms.
- Atoms of the same element are identical. the atoms of any 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 ratios to form compounds
- Chemical reactions occur when atoms are separated, joined, of rearranged. Atoms of one element however, are never changed into atoms of another element as a result of a chemical reaction
TitleChristian Cooke (pg 107-108 pg)
1911- Rutherford and co-workers tested what was then the current theory of the atomic structure. Using alpha particles, a thin layer of them weredirected at a very thin sheet of gold. To their surpirse, the alpha particles went right thru without deflcection as they had previously thought. He then proposed a new theory - the atom was mostly empty space , which explained the lack of deflection. He proposed that the particles and mass are condensed into a small, positively charged region which was called the nucleus. THis would account for deflection because the alpha particles were also positive. The NUCLEUS is a tiny central core of an atom composed of protons and neutrons. Rutherford atomic model is known as the nuclear atom, where protons and neutrons are located in the nucleus. Electrons are distributed around the nucleus and occupy almost all of the atom. tThe nucleus is very small compared to the rest of the cell.
Subatomic Particles
Haley Conatser (pg 104- 105)Electrons:
- English physicist J.J. Thomson discorved the electron in 1897- Electrons are negatively charged subatomic particles
- Thomson performed experiments that involved passing electric currents thought gases at low pressure
- He sealed the gases in glass tubes fitted at both ends with metal disks called electrodes. The elctrodes were connected to a soure of electricity. One electrote, the anode, became positively charged. The other electrode became negatively charged.
- The result was a glowing beam, or cathode ray that traveled from the cathode to the anode.
- A cathode ray is deflected by a magnet. A cathode ray is also deflected by electically chartged metal playes.
- To see a cathode ray being deflected by a magnet there is a video here: Click here
- A positively charged plate attracted the cathode ray, while a negatively charged play repels it.
- Thomson hypothesized that a cathode ray is a stream of tiny negatively charged particles moving at high speed
- Thomson originally called these particles corpuscles; later they were named electrons.
- To test his hypothesis he tried to measure the ratio of the charge of an electron to it's mass He found this ratio to be constant and in addition, the charge-to-mass ratio of electrons did not depend on the kind of gas in the cathode-ray tube or the type of metal used for the electrodes.
Robert A. Millikan (1868-1953):
- This U.S. physicist carried out experiments to find the quantity of charge carried by an electrons- Using the value and the charge-to-mass ratio that Thomson had found he was able to found out the mass of an electron
- What Millikan told us about the charge and mass of the electron is still true today
- An electron carries exactly one unit of negative charge and its mass is 1/1840 the mass of an hydrogen atom
GROUP TWO (Chapter 4.3 and 5.1 pg 110-116, pg 127-132)
Alex Nunan, Zoey Killion, Mike Hanley, Becky Hyatt, Elizabeth HowardCo- Editor: Alex Nunan
Elizabeth Howardpage 110
Atomic Number : the number of protons in the nucleus of an atom of that element
Atoms of the First Ten Elements:
Elizabeth Howard page 111
NUMBER OF NEUTRONS = MASS NUMBER - ATOMIC NUMBERIsotopes
Alex Nunan (pg 112-113)isotopes- atoms that have the same number of protons but different numbers of neutrons:
chemical symbols for isotopes:
for example, hydrogen has three known isotopes-
Atomic Mass
Becky Hyatt (pg 114-117)- the actual mass of a proton or neutron is very small and cannon be measured using a mass spectrometer
- an atomic mass unit (amu) is defined as one twelfth of 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 at 12 amu
- 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 mass of the atoms in a naturally occuring sample of the element
Conceptual Problem 4.3:- using atomic mass to determine the relative abundance of isotopes
- The atomic mass of copper is 63.546 amu. Which of copper's two isotopes is more abundant: copper-63 or copper-65?
- 1. Analyze- Identify the relevant concepts
- define what an atomic mass and a weighted average mass are
- 2. Solve- Apply the concepts to this problem
- because the atomic mass is a weighted average of the isotopes, copper-63 must be more abundant than copper-65
Practice Problem:Boron has two isotopes: boron-10 and boron-11.
Which is more abundant, given that the atomic mass of boron is 10.81?
The Development of Atomic Models
Zoey Killion (pg127-129)
- After discovering the atomic nuclues, Rutheford combined what was known about hte atom by proposing an atomic model in which the electrons move around the nucleus, like the planets move around the son
- Rutheford's model explained only a few simple properties of atoms, his model could not explain:
- why metals or complounds of metals give off characteristics colors when heated in a flame
- why objects - when heated to higher and higher temperatures - first glow dull red, then yellow, then white (as shown in picture below of horseshoe)
- Rutheford's atommic model also could not explain the chemical properties of elements
The Bohr ModelThe Quantum Mechanical Model
Mike Hanley (pg 130-131)
Atomic Orbitals
Olivia Richardson (pg 131-132)GROUP THREE (Chapter 5.2 and 5.3 pg 133-145)
Shannon Lamy, Mitchell Martin, Andrea Luongo, James Payne, Nate Lynch,Dakota PimentelCo-Editor: Shannon Lamy
Electron Configurations
Mitchell Martin (pg 133-135)Aufbau Principle : electrons occupy the orbitals of the lowest energy first.
Pauli exclusion principle: Only one or two electrons can occupy a P orbital or S orbital.
Hund's Rule: electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin as large as possible.
Exceptional Electron Configurations
Shannon Lamy (pg 136-137)Chromium
Cr 1s2 2s2 2p6 3s2 3p6 3d5 4s1
- ^^half filled d sublevel.
- EX. The actual electron configuration for copper is...
Cu 1s2 2s2 2p6 3s2 3p6 3d10 4s1Light
Andrea Luongo (pg 138-139)
Frequency & Wavelength:
Wavelength & Amplitude:
Frequency, Wavelength, & Amplitude:
Electromagnectic Spectrum:
Atomic Spectra
James Payne (pg 141)
When atoms abosorb energy, electrons move into higher energy levels,
electrons lose energy by emitting light
Frecuencies of light emitted by an element seperate into discrete lines to give the
atomic emission spectrum of the element
Each discrete line in an emission spectrum corresponds to exact one frecuenct of light
emitted by the atom
Much knowledge of the universe comes from teh studing of atomic spectra of the stars
which are glowing hot bodies of gases
Title
Dakota Pimentel (pg. 142-143)
Quantum Mechanics
Nate Lynch (pg 144-145)
-The branch of quantum physics that accounts for matter at the atomic level; an extension of statistical mechanics based on quantum theory
In quantum theories, energy and momentum have a definite relationship to wavelength. All particles have properties that are wave-like (such as interference) and other properties that are particle-like (such as localization). Whether the properties are primarily those of particles or those of waves, depends on how you observe them.
For example, photons are the quantum particles associated with electromagnetic waves. For any frequency, f, the photons each carry a definite amount of energy (E = hf)
http://phys.educ.ksu.edu/