WAVE NATURE OF LIGHT - Brim's Science Stuff
SC3. Students will use the modern atomic theory to explain the characteristics of atoms.
a. Discriminate between the relative size, charge, and position of protons, neutrons, and electrons in the atom.
b. Explain the relationship of the proton number to the element’s identity.
c. Explain the relationship of isotopes to the relative abundance of atoms of a particular element.
The Development of the Atomic Theory
THE INDIVISIBLE ATOM
Greek - Democritus (450 BC)
- Atom was indivisible
- Theorized the existence of the atom
- Also, theorized that there were just four 'elements' - fire, water, air, earth
Benjamin Franklin ( 1760’s)
- discovers that an object can have an electrical charge.
- Charge can be positive or negative
VIDEODISC - Electroscope
1. Can you explain what is happening?
John Dalton (1803)
- Atom was indivisible
- All elements are composed of atoms
- The same atoms for one element are exactly alike
- Atoms are neither created or destroyed in a chemical reaction
- In a chemical reaction, atoms are separated, combined, or rearranged
- Different atoms combine in simple whole number ratios to form compounds (Law of Definite Proportions)
THE DIVISIBLE ATOM
Sir William Crookes (1880)
- determined that rays were traveling from one end to another in the cathode ray tube
- the cathode ray tube led to the invention of the television
- others experimented with the cathode ray tube and discovered that the type of gas had no effect so the ray must be a part of all matter. A magnet deflected the ray, so it must be composed of charged particles, and it deflected toward the positive, so the charged particles must be negative.
J. J. Thomson (late 1890's)
- discovered the electron using the cathode ray tube
- determined that the electron was smaller than a hydrogen atom. This was a shocking discovery. Many thought Dalton was wrong.
- Knew the atom was neutral and the electron was negative, so there must be positive material with a lot more mass. Said the atom was a positive pudding-like material throughout which negatively charged electrons were scattered - Plum Pudding or Chocolate Chip Cookie Model ( Page 94, figure 4-9).
Robert Millikan (1909)
- Knew the electron was negative, but actually determined the charge (as a measurement). Was able to use this measurement to determine the mass of an electron. It was 1/1840 of a hydrogen atom
VIDEODISC – Rutherford’s Experiment
1. What is an alpha particle?
2. What did Rutherford expect to see? What did he actually see?
3. What is Rutherford’s model of the atom?
Ernest Rutherford (1909) (1919)
- Did a famous gold foil experiment (the alpha scattering experiment)
- Calculated that the atom was mostly empty space through which electrons move.
- Concluded that the atom has a small, dense, positively charged, centrally-located nucleus surrounded by negatively charged electrons
- By 1919 he had refined the concept of the nucleus. Called the positive particles protons.
Rutherford and James Chadwick (1932)
- showed the nucleus also had a neutron.
- The neutron was basically equal in mass to the proton but had no electrical charge.
SUBATOMIC PARTICLES:
Fill in the blanks with the appropriate information.
Relative How to determine
Particle Location Charge Mass (amu) the number in a neutral atom
PROTON
____________________________________________________________________________________________________
NEUTRON
____________________________________________________________________________________________________
ELECTRON
The atom at this point: is spherically shaped with a dense, centrally located, positively charged nucleus surrounded by one or more negatively charged electrons in an electron cloud. Most of the atom consists of fast moving electrons traveling through empty space surrounding the nucleus. Electrons are held within the atom by an attraction to the positive nucleus. The nucleus has neutral neutrons and positive protons and 99.9% of the mass. Since atoms are electrically neutral, the number of protons must equal the number of electrons.
Henry Moseley (1912)
- Discovered that atoms of each element contain a unique positive charge.
- The number of protons in an atom identifies it.
- ATOMIC NUMBER = number of protons (and number of electrons). It is unique for each element
The MASS NUMBER is the number of neutrons and protons (called nucleons) in the nucleus. This means that you can calculate the number of neutrons in an atom by taking the mass number and subtracting the number of protons.
ISOTOPES
Isotopes are atoms of the same element that have a different number of neutrons. They differ in mass, but the atom’s chemical behavior are the same. To identify an isotope, you write the element’s name and follow it with its mass number. Examples: Carbon-12, Uranium-238. It can also be identified by writing
C U
The mass of individual atoms is too small to measure so the mass of an atom is compared to a standard, Carbon-12 – Carbon-12 = 12 amu. AMU stands for atomic mass unit. The ATOMIC MASS is the weighted average of the isotopes of an element and can be calculated if you know the isotope’s mass numbers and the percentage abundance of each. For more practice, page 103.
PRACTICE
The atomic mass of iridium-191 is 191.0 amu and iridium-193 is 193.0 amu. The percentage abundance for each is 37.58% (iridium-191) and 62.62% (iridium-193). Calculate the average atomic mass of iridium.
SC1. Students will analyze the nature of matter and its classifications.
a. Relate the role of nuclear fusion in producing essentially all elements heavier than hydrogen.
NUCLEAR CHANGE
Review: Remember that a chemical reaction involves a change of one or more substances into new substances. Chemical reactions involve only an atom’s electrons. The nucleus remains unchanged.
Characteristics of Chemical and Nuclear Reactions
Chemical Reactions
1. Occur when bonds are broken and formed
2. Atoms remain unchanged, although they may be rearranged
3. Involve only valence electrons
4. Associated with small energy changes
5. Reaction rate is influenced by pressure, temperature, concentration, and catalysts.
Nuclear Reactions
1. Occur when nuclei emit particles and/or rays
2. Atoms are often converted into atoms of another element
3. May involve protons, neutrons, and electrons
4. Associated with large energy changes
5. Reaction rate is not normally influenced by pressure, temperature, and catalysts.
Nuclear reactions involve a change in the nucleus. Some substances spontaneously emit radiation called RADIOACTIVITY. They do this because their nuclei are unstable. The rays and particles emitted are called RADIATION. Unstable nuclei lose energy by emitting radiation in a spontaneous (does not require energy) process called radioactive decay. The atom undergoes decay until it becomes stable. By emitting radiation, atoms of one element change into atoms of another element. Elements with an atomic number greater than 83 are radioactive. Elements with #1-82 have isotopes that may be radioactive.
VIDEODISC – Hot Pocket Change
1. What type of radiation does this emit?
2. What is your hypothesis on why lead stops the radiation and wood does not?
HISTORY
1896 Henri Becquerel
Discovered mysterious rays coming from uranium. Called it RADIATION. Radioactivity is the spontaneous emission of radiation from an element.
1898 Marie and Pierre Curie discovered these rays in Radium and Polonium
TYPES OF RADIATION
1. Alpha radiation, (
Made of alpha particles. Is composed of two protons and two neutrons. Has a 2+ charge and a mass of 4 amu. Has the least amount of energy of any of the radiation – is stopped by paper. A new element is created when alpha decay happens. The mass number and the atomic number change.
Alpha decay: Ra ( Rn + He Radium-226 decays to radon-222 and an alpha
particle
2. Beta radiation, (
made of fast moving electrons. Has a –1 charge and a mass of 1/1840. An electron is emitted during beta decay because it has been removed from a neutron, leaving behind a proton. Is stopped by foil and has more energy than alpha radiation.
Beta decay: Na ( Mg + ( Sodium-22 decays into magnesium-22 and a beta
particle
3. Gamma radiation, (
high energy radiation that possesses no mass. Has no charge. Usually accompanies alpha or beta radiation. Is slowed down by lead or concrete. Accounts for most of the energy lost in a radioactive decay process.
Gamma decay: U ( Th + He + ( Uranium-238 decays into thorium-234 and an alpha
particle and two gamma rays of different frequencies
NUCLEAR STABILITY
The ratio of neutrons and protons determines nuclear stability. Too few or too many neutrons make the nucleus unstable. It needs to lose energy to become stable. So it decays. Few radioactive elements are found in nature – they have already decayed into stable atoms.
The STRONG NUCLEAR FORCE acts only on subatomic particles that are extremely close together. It overcomes the electrostatic repulsion between positive protons. Neutrons add an attractive force within the nucleus because they are not positive or negative. They are also subjected to the strong nuclear force. They help hold the nucleus together.
To determine stability, you need to calculate the neutron to proton ratio. If you were to plot these on a graph (neutrons on the y-axis and protons on the x-axis), you would get a band of stability for elements 1-82. All non-radioactive elements are within this band. Radioactive elements are found outside the band (page 811, figure 25-11).
After radioactive decay, the new atom is now positioned more closely, if not within, the band of stability.
RADIOACTIVE SERIES
A series of nuclear reactions that begins with an unstable nucleus and results in the formation of a stable nucleus is called a radioactive series. Uranium-238 goes through a radioactive series and eventually become lead-206 (page 814).
RADIOACTIVE DECAY RATES
Naturally occurring (radioactive isotopes) radioisotopes are not uncommon on earth. Even though radioisotopes have been decaying for 15 billion years (lifespan of earth), there are still some left to decay. The difference in different isotopes decay rates provides the explanation. Radioactive decay rates are measured in half-lives. A HALF-LIFE is the time it takes for one half of the radioactive sample to decay. Different isotopes have different half-lives. Examples include uranium-238 (4.46 x 109 years), carbon-14 (5730 years), radon-222 (3.8 days) and polonium-214 (163.7 microseconds).
If you were given a 10.0g sample of carbon-14, how much would be stable after 3 half-lives?
radioactive
stable
original sample 1st half-life 2nd half-life 3rd half-life
5.0g stable 7.5g stable 8.75g stable
THREE TYPES OF NUCLEAR REACTIONS
1. SPONTANEOUS - a single nucleus releases energy and particles of matter
2. FISSION - nucleus is split in two
3. FUSION - two nuclei join together – this is how all elements are made.
*ALWAYS involve a change in the nucleus and a release of energy
WRITING AND BALANCING NUCLEAR EQUATIONS PRACTICE
Remember when balancing these equations that mass number and atomic number are conserved.
1. Write the nuclear equation for the alpha decay of polonium-208.
2. Write the nuclear equation for the beta decay of copper-66.
SC3. Students will use the modern atomic theory to explain the characteristics of atoms.
b. Relate light emission and the movement of electrons to element identification.
WAVE NATURE OF LIGHT
Rutherford’s model of the atom did not explain how the electrons of an atom are arranged in the space around the nucleus. His model did not take into account the chemical behavior among various elements. In the early 1900’s it was observed that certain elements emitted visible light when heated in a flame. Analysis of the emitted light revealed that an element’s chemical behavior was related to the electron arrangement in the atom.
Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space. Radiant energy, which includes visible light, is shown in the electromagnetic spectrum.
Lowest Highest
Frequency frequency
104 1022
Radio Microwave Radar Infrared Visible Light Ultraviolet X-ray Gamma
longest ROYGBIV shortest
wavelength ( increasing energy ( wavelength
104 10-12
Wave Characteristics:
1. wavelength, ( measured from crest to crest, measured in m, cm, nm
2. frequency, ( measured by number of waves per second, hertz, Hz, waves/s
3. amplitude height of a wave crest or trough
4. speed, c 3.00 x 108 m/s in a vacuum. Speed is always the same. Waves can
have different wavelength and frequency. Wavelength and frequency are inversely related.
speed = wavelength x frequency c = ((
[pic]
White light can be refracted into its component colors and each color has its own wavelength.
PRACTICE:
1. Calculate the speed of a wave whose wavelength is 1.5 meters and whose frequency is 280 hertz.
2. Find the wavelength of a wave whose speed is 5.0 m/s and whose frequency is 2.5 hertz.
3. The speed of light is 3.00 x 108 m/s. Red light has a wavelength of 7.0 x 10-7m. What is its frequency?
Particle Nature of Light
Wave model cannot explain why heated objects emit only certain frequencies of light at a given temperature or why some metals emit electrons when colored light of a specific frequency shines on them.
1900. Max Planck – Matter can gain or lose electrons only in small specific amounts called quanta. A quantum is the minimum amount of energy that can be gained or lost. Glowing objects emit light, which is a form of energy.
Energyquantum = h( h = Planck’s constant = 6.626 x 10-34 J·s
The energy of the radiation increases as the frequency increases. This explains why ultraviolet light has more energy than violet light.
PRACTICE:
1. Calculate the energy of a gamma ray photon whose frequency is 5.02 x 1020 Hz.
2. What is the difference in energy between a photon of violet light with a frequency of 6.8 x 1014 Hz and a photon of red light with a frequency of 4.3 x 1014 Hz?
Photoelectric Effect
Electrons (photoelectrons) are emitted from a metals’ surface when light of a certain frequency shines on the surface. Increased intensity results in more electrons being ejected from the metal’s surface. (Example: a solar calculator converts the energy of light shining on it to electrical energy.) In 1905, Einstein proposed that electromagnetic radiation had wavelike and particle-like natures. The photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.
Ephoton = h( h = Planck’s constant
For the photoelectric effect to occur, a photon must possess, at a minimum, the energy required to free an electron from an atom of metal.
GLOW IN THE DARK STARS
1. laser light - 670 nm wavelength stars glow? ____________________________
2. flashlight - 500 nm wavelength stars glow? ____________________________
3. ultraviolet light - 254 nm wavelength stars glow? ____________________________
4. glowing star light - 520 nm wavelength
On a separate piece of paper, do the following:
A. Convert wavelength in nm to meters for each of the four colors of light. Use the conversion 1 x 10-9m = 1 nm
B. Using the speed of light formula and Planck’s energy formula, determine the energy for each of the four colors of light. That means there will be two calculations for each color of light in Part B.
C. Answer the questions:
1. Would infrared light cause the stars to phosphoresce?
2. Would microwave light cause the star to phosphoresce?
3. Glow in the Dark Stars are an example of an object that has a minimum requirement for photons. Give another example of minimum energy photons.
SC3. Students will use the modern atomic theory to explain the characteristics of atoms.
c. Relate light emission and the movement of electrons to element identification.
ATOMIC EMISSION SPECTRA
A neon light works because electricity is passed through the neon gas in the tube. The gas absorbs energy and becomes excited. Excited and unstable atoms then release energy by emitting light.
Atomic emission spectrum is a set of frequencies of electromagnetic waves emitted by atoms of the element. They are usually distinct color lines. Each element’s atomic emission spectrum is unique, can be used to identify the element, and can be used to determine if the element is part of an unknown compound.
VIDEODISC:
1. Why is it necessary to create a vacuum?
2. What would happen if common gases were used?
3. Can you think of ways that the energy of emitted photons with frequencies outside the visible range can be used?
Page 144, Fig. 5-8 shows the emission spectrum of hydrogen. Notice that it is discontinuous – it is made up of only certain frequencies of light.
DEMO: Atomic Emission Spectra examples: Air, argon, bromine, carbon dioxide, chlorine gas, helium, hydrogen, iodine vapor, krypton, mercury vapor, neon, nitrogen, oxygen, water vapor, xenon.
QUANTUM THEORY
In 1913, Neils Bohr came up with the quantum model of the hydrogen atom. He also correctly predicted the frequencies of the spectral lines in the hydrogen atomic emission spectrum.
The hydrogen atom has only certain allowable energy states. The lowest is called the GROUND STATE. When atoms gain energy, they are said to be in an EXCITED STATE. Hydrogen has one electron, but can have different excited states.
Bohr said that the electron moves around the nucleus in only certain allowable circular orbits. The smaller the orbit, the lower the atom’s energy state or energy level. Each orbit is assigned a quantum number, n.
Bohr said that hydrogen’s electron in the ground state did not radiate energy. When excited, the electron moves up to another energy level. Only certain atomic energies are possible and so only certain frequencies of electromagnetic radiation can be emitted.
n = 1 excites to n = 3 which then drops to n = 2 (red light) and then drops to n = 1 (UV light). See page 147, Figure 5-11.
In 1928, Louis de Broglie stated that Bohr’s quantized electron orbits had characteristics similar to waves. Particles could have wavelike behaviors. The energy had wavelike characteristics. (You do not need to know the equation.)
Werner Heisenberg came along and concluded that it is impossible to make any measurement on an object without disturbing the object. It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.
Erwin Schödinger, in 1926, developed an atomic model in which the electrons are treated as waves. This model is called the quantum mechanical or wave mechanical model. This model does not describe the path of the electron around the nucleus. The three dimensional region around the nucleus called an atomic orbital describes the electron’s probable location. It is a fuzzy cloud in which the density of the cloud at a given point is proportional to the probability of finding the electron at that point. The electron cloud has no definite boundary – it is arbitrarily drawn at 90%. See page 152, figure 5-15.
QUANTUM MECHANICAL MODEL OF THE ATOM
1. The atom has a dense, centrally located, positively charged nucleus full of protons and neutrons surrounded by mostly empty space called an electron cloud where the electrons are. (Rutherford)
2. The energy of electrons is quantized (has only specific amounts of energy). (Bohr)
3. Electrons exhibit both wave and particle behaviors. (de Broglie)
4. The absolute location of an electron is impossible to determine – its location and velocity cannot be determined at the same time. (Heisenberg)
5. The electrons travel in orbitals that have characteristic sizes, shapes, and energies, but do not describe how the electrons move. (Schödinger)
QUANTUM NUMBERS
These numbers describe the most probable location of an electron in an atom.
1. Principle quantum number, n
Is the energy level number. Gives information about the relative size and energy of the atomic orbitals. It can have values of 1, 2, 3, 4…. The greatest number of electrons possible in any one level is 2n2.
Example: The maximum number of electrons that can occupy the first level is 2(1)2 = 2; the fourth level is 2(4)2 = 32.
2. Angular Momentum Quantum Number, l
Is the energy sublevel number. It gives information as to the shape of the orbitals. The first four levels are s, p, d, and f.
Example: The first energy level has only an s sublevel. The second energy level has an s and p sublevel. This third energy level has s, p, and d sublevel.
VIDEODISC:
1. What is the Heisenberg uncertainty principle?
2. How is the principle quantum number related to the size and energy of an atomic orbital?
3. What is the electron configuration of iron?
3. Magnetic Quantum Number, m
gives information about the orientation in space of an orbital. The s sublevel has one orbital, the p sublevel has three orbitals, the d sublevel has 5 orbitals, and the f sublevel has 7 orbitals. This number determines which p, d, or f orbital the electrons are in.
4. Electron Spin Quantum number, s
indicates the direction of the electron spin. Spin is either clockwise or counterclockwise and is designated with a +1/2 or a –1/2 .
ELECTRON CONFIGURATION
Electron configuration tells us how electrons are distributed among the various atomic orbitals. It is the arrangement with the lowest possible energy. It is a simple way of keeping track of electrons. Electrons fill the levels according to a set of rules:
1. Aufbau Principle
Each electron occupies the lowest energy orbital available. Electrons are added from the ground state up. Electrons fill in increasing energy order.
2. Pauli Exclusion Principle
Each orbital can hold a maximum of only two electrons – one spinning clockwise and one spinning counter clockwise.
3. Hund’s Rule
The most stable arrangement of electrons in orbitals is to fill singly and then go back and double up.
BLANK PERIODIC TABLE
Examples:
1 Hydrogen 1s1
2 Helium 1s2
3 Lithium 1s22s1
4 Beryllium 1s22s2
5 Boron 1s22s22p1
6 Carbon 1s22s22p2
7 Nitrogen 1s22s22p3
8 Oxygen 1s22s22p4
9 Fluorine 1s22s22p5
10 Neon 1s22s22p6
11 Sodium 1s22s22p63s1
12 Magnesium 1s22s22p63s2
13 Aluminum 1s22s22p63s23p1
PRACTICE - The Electron Hotel
The order that the levels and sublevels fill is based on energy. They fill in the following order:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
Can use filling diagram, page 156. Can also use the periodic table.
ODDBALL ELECTRON CONFIGURATION
FOR ODDBALLS AND IONS
The transition elements do not always follow the rules for electron configuration. Notice the electron configurations for elements 21-30. The ones in bold are oddballs:
Scandium 1s22s22p63s23p64s23d1
Titanium 1s22s22p63s23p64s23d2
Vanadium 1s22s22p63s23p64s23d3
Chromium 1s22s22p63s23p64s13d5
Manganese 1s22s22p63s23p64s23d5
Iron 1s22s22p63s23p64s23d6
Cobalt 1s22s22p63s23p64s23d7
Nickel 1s22s22p63s23p64s23d8
Copper 1s22s22p63s23p64s13d10
Zinc 1s22s22p63s23p64s23d10
The chromium and copper do not follow the rules. The “d” sublevel holds ten and is more stable when is it half full then the “s” when it is full. So it steals an electron from “s” orbital to be more stable. The same thing happens as it nears ten electrons to be full.
Now, you write the electron configurations for some elements listed below.
42-Molybdenum _______________________________________________________________
43-Technetium _______________________________________________________________
44-Ruthenium _______________________________________________________________
45-Rhodium _______________________________________________________________
46-Palladium _______________________________________________________________
47-Silver _______________________________________________________________
48-Cadmium _______________________________________________________________
Name the transition elements that will be oddballs that you must remember:
ELECTRON CONFIGURATION FOR IONS
An ion is an atom that has a charge. It is not neutral and its number of protons does not equal its electrons. If it has lost electrons, it is positive and called a CATION. If it has gained electrons, it is negative and called an ANION. For atoms that occur as ions, the electron configuration can also be written.
1. First, write the regular electron configuration for the element.
2. If the ion is positive, take away electrons. If the ion is negative, add electrons to the highest partially filled energy level.
EXAMPLES:
Na 1s22s22p63s1 Al 1s22s22p63s23p1
Na+1 1s22s22p6 Al+3 1s22s22p6
F 1s22s22p5 H 1s1
F-1 1s22s22p6 H-1 1s2
PRACTICE:
Magnesium, Mg Magnesium ion, Mg+2
Oxygen, O Oxygen ion, O-2
Lithium, Li Lithium ion, Li+1
Iron, Fe Iron ion, Fe+3
Atoms that have absorbed energy and are in an excited state have electrons that have moved to a shell level that is higher than what is normal. The electrons that move come from the valence electrons (outermost electrons).
ORBITAL DIAGRAMS
This is another means of symbolizing where electrons are in energy levels and sublevels. Involves three basis symbols:
Unoccupied orbital orbital with one electron orbital with two electrons
Can also be drawn with circles or lines instead of squares.
_____
PRACTICE – handout
ELECTRON CONFIGURATION SHORTHAND
Electron configurations can be written in a shorthand form. You can take the noble gas (ONLY) that occurs before the element in question and then tack on the remaining configuration.
Example: Sodium 1s22s22p63s1 Shorthand: [Ne] 3s1
Copper 1s22s22p63s23p64s13d10 Shorthand: [Ar] 4s13d10
VALENCE ELECTRONS
Valence electrons are the electrons in the outermost energy levels. Using the two examples from above, sodium and copper, the number of valence electrons for each are:
Sodium 1s22s22p63s1 Shorthand: [Ne] 3s1 valence electrons: 1
Copper 1s22s22p63s23p64s13d10 Shorthand: [Ar] 4s13d10 valence electrons: 1
Other examples:
Magnesium 1s22s22p63s2 valence electrons: 2
Chlorine 1s22s22p63s23p5 valence electrons: 7
You can also use the periodic table itself to determine the valence electrons for any atom.
ELECTRON DOT DIAGRAMS
Usually we are only concerned with the electrons in the outermost energy level. Those are the electrons involved in chemical reactions – the VALENCE electrons. We can symbolize these electrons using the Lewis Electron Dot Diagram. To write an electron dot diagram, follow these steps:
1. Write the symbol for the element (symbolizes the nucleus and all inner energy levels)
2. Write the electron configuration for the element. Select the electrons in the outermost energy level. Use the HIGHEST principle quantum number regardless of highest energy.
3. Each side of the symbol represents an orbital. Draw the dots on the appropriate sides to represent the electrons in that orbital. It is important to remember which electrons are paired and which are not.
s s
p p
p p
p p
Examples:
Sodium 1s22s22p63s1 valence electrons: 1 Na
Copper 1s22s22p63s23p64s13d10 valence electrons: 1 Cu
Magnesium 1s22s22p63s2 valence electrons: 2 Mg
Chlorine 1s22s22p63s23p5 valence electrons: 7 Cl
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