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Which scientist performed the Gold Foil Experiment? What were the 2 main conclusions from this experiment about the structure of the atom? What two calculations were made as result of the Oil Drop Experiment? Over 100 subatomic particles have been discovered. As a class, they are called nucleons. The three best-known subatomic particles are represented as follows: proton Symbol Charge Mass,g p or + +1 neutron electron n 0 1.673x10-24 1.673x10-24 e-1 9.11x10-28 Elements differ from one another by the number of protons they contain in their nuclei (known as an element’s atomic number), and are arranged left to right, row by row, on the periodic table. Since only protons and neutrons are located inside the nucleus, the net charge on any nucleus is positive. Ex: Barium has an atomic number of 56, so a barium nucleus has a charge of +56. Since neutrons have about the same mass as protons, both contribute to an atom’s mass. An atom’s mass number is the sum of protons and neutrons in its nucleus. Ex: Barium- _____ has 56 protons and 81 neutrons, thus a mass number of 56 + 81 = 137. Why is atomic mass listed on the periodic table not in whole numbers? Because, atomic mass is an average of the masses of all available isotopes of the element, based on percent occurrence. Ex: Barium’s atomic mass is 137.33 amu (atomic mass units) as read from the periodic table, suggesting that barium has isotopes other than barium-137 that bring the average mass up slightly. Electron mass is so low it is considered to be negligible for the atom. Electrons do however affect charge of an atom. The overall charge of an atom is protons minus electrons. In a neutral (zero-charged) atom, the number of protons = the number of electrons. An atom is neutrally charged unless electrons have been removed or added through chemical or electrical processes. Then the atom is charged and is called an ion. A positively charged atom is called a cation; a negatively charged atom is called an anion. Examples: Element Charge state Electron movement # protons # electrons calcium neutral None 20 20 calcium +2 charge 2 lost 20 18 iodine neutral None 53 53 iodine -1 charge 1 gained 53 54 On the last test 15% of students earned 100%, 75% of students earned 85% and 10% of students earned 70%. What was the average percentage for the class on the test? Not all atoms of the same element contain the same numbers of neutrons. Isotopes are atoms of an element that contain different numbers of neutrons, thus altering mass number. Many elements exist in isotopic form. Nuclear notation (Isotopic nomenclature) for isotopes: Barium-137 (Ba-137) 137 Ba protons = neutrons = electrons = +2 56 Since atomic number for an element never changes, it is often dropped from isotopic notation and only mass number and maybe charge is represented: 137 Ba +2 137 or Ba Three important isotopes of hydrogen: protium (H-1) : most common isotope; 1 proton, 0 neutrons, 1 electron deuterium (H-2): found mostly on the sun and stars, in heavy water; 1 proton, 1 neutron, 1 electron tritium (H-3): found mostly on the sun and stars, used in many military weapons; 1 proton, 2 neutrons, 1 electron Deuterium and tritium are the primary isotopes involved in fusion on the sun. The sun converts 4 million tons of hydrogen into helium every second! AMU- Atomic Mass Unit is defined as 1/12 the mass of the carbon-12 isotope, which is assigned 12 amus. Atomic Mass- is the weighted average mass of the atoms in a naturally occurring sample of the element -reflects both mass and abundance of isotopes as they occur in nature Sample Problem: Radioactivity, the spontaneous emission of high energy particles or rays by unstable atomic nuclei, can produce subatomic particles. The transuranium elements are the most naturally radioactive. Nucleic decay produces alpha particles, beta particles, or gamma rays. Alpha particles (α) – represented by 4 2He , are high speed helium nuclei with a +2 charge and fairly low penetrating power. Beta particles (β) – represented by 0-1β , are high speed electrons with a -1 charge and high penetrating power. Gamma rays (γ) – represented by 00γ , are high energy electromagnetic waves with no charge and extremely high penetrating power. A nuclide’s radioactive stability can be predicted by its neutron to proton ratio. For smaller atoms (At#<20), the stable ratio is 1:1. As the size of the atom increases, the ratio does too, to a maximum of 1.5:1. Atoms within the band of stability, indicated by the graph on p.866 are stable and do not undergo spontaneous radioactive decay. Atoms outside the band of stability, above or below, are prone to decay to approach a more stable state. Nuclear Equations – Nuclear reactions are represented in equation form, using nuclear notation. The sum of the mass numbers (superscripts) on one side of the equation must equal the sum on the other. The same applies to atomic numbers (subscripts). Many nuclear reactions emit gamma rays, which have no mathematical effect on the equation, but a huge effect on energy released. There are numerous types of nuclear reactions. Examples include: alpha decay – the emission of alpha particles. Ex: alpha decay of uranium238 238 234 92U → 4 90Th + 0 2He + 20γ beta decay – the emission of beta particles. Ex: beta decay of iodine-131 131 53 I 131 → 54Xe 0 + -1β electron capture – results in the emission of xrays. Ex: electron capture by rubidium-81 81 0 37Rb + -1e 81 → 36Kr + x-ray photon neutron bombardment – used to initiate fission reactions in nuclear power plants 235 92U + 1 0n → 236 92U (a less stable nuclide) Nuclear Half-Lives – Each radioactive isotope has what is known as a half life, the time it takes for half the mass of an available sample to decay into another substance. Half-lives can range from microseconds to millions of years. (See p.871.) Half-life is used to calculate either how long it would take to result in a level of the isotope that could be tolerated by living things or to calculate the mass of an isotope remaining after a period of time. mfinal = minitial x .5t/T t/T is the number of half lives a sample has been through, where t = time lapsed and T = half life (must have unit agreement) Sample problems: 1. How much of a 25.0g sample of radon-222 will remain after 31 days? t/T = 31days/3.8days = 8.15 mfinal = 25.0g x (.5)8.15 = 0.0875g 2. How much of a 350g sample of a radioacitve isotope remains after 5 half lives has passed? 350g → 175g → 87.5g → 43.75g → 21.875g → 10.9375g Or: mfinal = 350g x .55 = 10.9375g 11g Fission vs. Fusion – Two nuclear processes used to produce massive amounts of energy. Fission is the splitting of large nuclei, resulting in the formation of smaller nuclei. Neutrons are used to initiate the process, which will continue on its own (self-sustaining) with each fission process producing more neutrons to initiate more fissions. This process is currently used in nuclear power plants… many issues… reaction control, used fuel rod disposal, etc. Fusion is the combining of small nuclei into a larger one. Fusion is the process that takes place on the sun and stars, a more powerful process than fission. Ex: the fusion of hydrogen isotopes to produce helium on the sun. 2 1H 2 + 1 1H 1H 3 + 1H 4 → 2He + ENERGY 4 → 2He + ENERGY Research continues today to reproduce fusion in a commercially viable way here on earth, a utopia for energy supply. The major problem is reaction containment, as fusion produces temperatures in excess of 40 million Kelvin. The quantum mechanical model of the atom is based on the study of waves and light. Important wave definitions and relationships: amplitude – the distance from the crest to the midway origin line of the wave. wavelength (, the Greek letter lambda) – the distance between similar points in a set of waves, such as crest to crest or trough to trough; normally expressed in a meter-type unit. frequency (, the Greek letter nu) – the number of waves that pass a given point per unit time; usually expressed in cycles/second (sec-1 or 1/sec), commonly known as Hertz, Hz. Frequency, wavelength, and energy are related to one another mathematically through: c = and E = h where “c” is the speed of light (3.00x108 m/s) and “h” is Planck’s constant (6.626x10-34 J/Hz). (Watch for unitary agreement!) Inspection of these relationships suggests that frequency and wavelength are inversely proportional and that energy and frequency are directly proportional. Electromagnetic Radiation (EMR) – a form of energy that travels in waves. The EMR spectrum includes all visible light and many other forms of wave energy: Fill in these blanks for the EMR spectrum: The highest energy visible light is violet light. The lowest energy waves in the EMR spectrum are radiowaves. Visible red light has shorter wavelength than infrared light. Visible green light has lower energy than Xrays. Visible indigo light has lower frequency than ultraviolet light. EMR travels at the speed of light in a vacuum, 3.00x108 m/s. Types of EMR differ from each other by wavelength, frequency, and resulting energy. All EMR is the result of electron movement between energy levels within atoms. When an electron moves to a higher energy level, farther from the nucleus, energy is consumed and the atom is referred to as “excited”. As the electron returns to a lower energy level, energy is emitted in the form of photons discrete packets of radiant energy that travel in waves (suggested by Einstein). Because of unique arrangements of electrons within atoms, each element has a characteristic light it emits when exposed to a sufficient amount of heat or electricity. When examined through a spectrophotometer, a device that breaks light into its component waves, the element’s bright-line spectrum is observed. (See p.126 for the BLS for hydrogen.) A BLS is not a continuous spectrum like the EMR spectrum, but it is distinct lines of color that correspond to very specific wavelengths, frequencies, and energies. The energies correspond to the energies required to move electrons farther from the nucleus and that emitted when they return to their original locations. Elements also have unique and specific energy requirements to completely remove an electron from an atom. Ionization energy is the amount of energy required to remove an electron from a gaseous atom of an element. The energy required is equal to the energy emitted when an electron is added to an atom from some other source. Of course, ionization energy is greater than the various energies required or emitted as electrons move within an atom. Work to relate wave behavior, atomic energy levels, and electrons by Louis deBroglie and Erwin Schrodinger led to the branch of physics called quantum mechanics. Quantum mechanics and probability, the statistical likelihood of an occurrence, are the basis for current atomic models. Electrons are particles with wave characteristics. Probability helps describe the behaviors and positions of electrons. One important assumption of today’s model is the Heisenberg Uncertainty Principle which states that it is not possible to describe both position and speed of an electron at the same point in time. An important distinction to make is that between an orbit (a pathway, as in Bohr’s Model of the atom) and an orbital. An orbital is a region of space around a nucleus in which an electron is most likely to be found. It is not a barrier or pathway, only a model describing the likely occupied area. Schrodinger developed a mathematical model for the wave behavior of an electron. The very complex equation contains four quantum numbers that are used to describe electron location and behavior… n Principle Quantum Number refers to the energy level location of the electron n is 1,2,3,4,… n (up to 7) the number of sublevels in the energy level = n (up to 4) the number of available orbitals in the energy level = n2 (up to 16) the greatest number of electrons contained in any one energy level = 2n2 (up to 32). Ex: where n=3, the maximum number of electrons is 2(3)2 = 18. l Azimuthal or Orbital Quantum Number refers to the sublevel location of the electron l is either s, p, d, or f s-sublevel orbitals are spherical in shape. They are closest to the nucleus; thus electrons located here have the lowest relative energies within the energy level. There is only 1 s-orbital per energy level. p-sublevel orbitals are dumbbell-shaped. They are a little farther from the nucleus; thus electrons located here have a little higher relative energies within the energy level. There are 3 p-orbitals (px, py, pz) per energy level, starting with energy level 2. d-sublevel orbitals are cloverleafshaped, very complex. They are even a little farther from the nucleus; thus electrons located here have even higher relative energies within the energy level. There are 5 d-orbitals per energy level, starting with energy level 3. f-sublevel orbitals are extremely complex in shape. They are farthest from the nucleus; thus electrons located here have the highest relative energies within the energy level. There are 7 f-orbitals per energy level, starting with energy level 4. ml Magnetic Quantum Number designates the number of orbitals on each sublevel and in which specific orbital the electron is likely located describes the home orbital’s geometric orientation about the x, y, and z axes ms Spin Quantum Number describes the direction of spin for the electron ms is either +1/2 (clockwise spin) or –1/2 (counterclockwise spin). The Pauli Exclusion Principle is very important here and throughout the quantum model. It states that: 1) no more than 2 electrons may occupy any one orbital; 2) if electrons occupy the same orbital, they must have opposite spin directions; 3) thus, no two electrons in one atom will have all four quantum numbers identical. Electrons are arranged from orbitals with lowest energy (closest to the nucleus) to orbitals with highest energy (farthest from the nucleus). Configurations can be written for neutral atoms and their ions. There are 3 basic rules that govern orbital filling: 1. Aufbau Principle – electrons enter lowest energy levels first. 2. Hund’s Rule – when electrons fill orbitals within a sublevel, each orbital is occupied by one electron before any orbital obtains two electrons. All electrons in singly occupied orbitals have the same spin directions. 3. Pauli Exclusion Principle – described earlier Reading e-configs and orbital(energy) diagrams from the Periodic Table: Recognize the significance of the structure of the Periodic Table: Row numbers correspond to Energy Levels (Principle Quantum Numbers, n) The PT is arranged in blocks: s-block (groups 1 & 2), p-block (groups 13-18), d-block (groups 3-12), and f-block (the lanthanide and actinide series) Read the PT from left to right, row by row, adding new electron(s) (indicated by superscript) each step of the way: 1. s-block (maximum of 2 electrons) and p-block (maximum of 6 electrons) electrons are in the energy level indicated by current row number 2. when you reach d-block (maximum of 10 electrons), drop 1 energy level below the current row number 3. when you reach f-block (maximum of 14 electrons), drop 2 energy levels below the current row number 4. when you reach p-block, return to the original energy level indicated by row number Ex: for the element arsenic (As): 1s22s22p63s23p64s23d104p3 total up the superscripts and you get 33, the atomic number of arsenic Sample orbital diagrams and electron configurations: Nitrogen (N) - ___ protons, ___ electrons electron configuration: 1s22s22p3 orbital/energy diagram: __ 1s __ 2s __ __ __ 2p Manganese (Mn) - ___ protons, ___ electrons e-config: 1s22s22p63s23p64s23d5 orbital/energy diagram: __ 1s __ 2s __ __ __ 2p __ 3s __ __ __ 3p __ 4s __ __ __ __ __ 3d Orbital/energy diagrams and electron configurations can be written for ions too. It is best to begin with the neutral diagram or configuration and make your additions to or deletions from it! Highest energy level electrons, referred to as valence electrons, are the first and often the only electrons involved in bonding. (Ions form as result of ionic bonding.) Note: Atoms tend to ionize in the easiest way for them to end with a full highest energy level. “s” and/or “p” sublevels will always be the highest energy level’s sublevels. A full highest energy level would contain 8 electrons… the octet rule. Samples: Formation of an anion → Sulfur (S) - ___ protons, ___ electrons e-config: 1s22s22p63s23p4 Sulfur ion (S-2) – ______________ electrons e-config: 1s22s22p63s23p6 Formation of a cation → Iron (Fe) - ___ protons, ___ electrons e-config: 1s22s22p63s23p64s23d6 energy diagram: __ __ 1s 2s __ __ __ __ __ __ __ __ __ __ __ __ __ 2p 3s 3p 4s 3d Remember that valence electrons are always the first and often the only electrons to become involved in bonding. It is imperative that you begin with the neutral diagram or configuration and make changes to it when dealing with cations!! Remove electrons from highest energy levels first, not necessarily the last filled. Iron ion (Fe+2) - _________________ electrons e-config: 1s22s22p63s23p64s23d6 energy diagram: __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p __ __ __ __ __ 3d Iron ion (Fe+3) - ___________________ electrons e-config: 1s22s22p63s23p64s23d65 energy diagram: __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 3s 2p 3p 4s 3d Develop the e-config for Nickel (II) ion (Ni+2): Neutral Ni - ___ protons, ___ electrons e-config: 1s22s22p63s23p64s23d8 Nickel ion (Ni+2) - ________________ electrons ion e-config: 1s22s22p63s23p63d8 As a general rule, metals tend to lose electrons, become positively charged, when they ionize. Nonmetals tend to gain electrons to become negatively charged. Metalloids, those elements on and below the stairstep on the periodic table, can do either depending on the circumstances. Based on energy diagrams and configurations, predict charge(s) for the following elements when they ionize: Nitrogen ion 1. Classify the element – nonmetal, likely to become negatively charged 2. Draw energy diagram or e-config of the neutral element: energy diagram: __ __ __ __ __ e-config: 1s22s22p3 1s 2s 2p 3. Locate valence electrons, and determine what must happen to meet the octet rule. 4. Predict charge(s) - Nitrogen is likely to gain 3 electrons to become -3 charged. You predict charge(s) for Calcium ion and Arsenic ion: Calcium – a metal, will lose e- to become + charged econfig: 1s22s22p63s23p64s2 valence electrons : level 4, 2 of them charge prediction: lose level 4 electrons to be +2 charged Arsenic – a metalloid, can act like a metal (lose e- to become + charged) and can act like a nonmetal (gain e- to become – charged). econfig: 1s22s22p63s23p64s23d104p3 valence electrons : level 4, 5 of them charge prediction: as a nonmetal – gain 3 electrons to fill level 4 electrons to be -3 charged as a metal – lose 4p-sublevel’s 3 electrons to be +3 charged, or also lose the 4s-sublevel’s 2 electrons to be +5 charged Once atomic number passes 18 or so, it becomes cumbersome to write full configurations. Thus, it is convention to use a shorthand format: 1. Locate the element to be read. 2. Find the noble gas that precedes it in atomic number and place its symbol in brackets to represent all of its electrons. 3. Begin reading at the next row of the PT, following guidelines, until you reach the target element. Ex: for the element lead (Pb) [Xe]6s24f145d106p2 total up the electrons represented by xenon and the superscripts and you get 82, the atomic number of lead You do neutral platinum, neutral selenium, and Sn+4: Normally, electron configurations are easy to predict, but there are a few exceptions in the areas of the transition metals. The most notable exceptions are chromium and copper. Their actual configurations are slightly altered from predicted in order to minimize energy. Chromium, Cr predicted: 1s22s22p63s23p64s23d4 actual: 1s22s22p63s23p64s13d5 Copper, Cu predicted: 1s22s22p63s23p64s23d9 actual: 1s22s22p63s23p64s13d10 These exceptions apply to other elements in their groups (under them) on the periodic table, but the farther down the table the less likely the exception applies. Lewis Structures depict an atom of an element with its valence electrons. Valence electrons – electrons in the highest energy level of an atom. Because of filling order, valence electrons will always be in “s” and/or “p” sublevels. Valence electrons are normally the first, and often the only, electrons to become involved in bonding.) Ex:sulfur (neutral) e-config: Lewis Structure: To construct: 1. Write the electron configuration for the atom. 2. Write the element symbol to represent the nucleus and all electrons other than valence electrons. 3. Begin placing valence electrons around the symbol in the same manner that you fill orbital diagrams. a. b. Place “s” electrons to the right of the symbol. Proceed clockwise, placing “p” electrons one at a time on the remaining three sides of the symbol. Then double up to complete placement of remaining electrons. You do Lewis diagrams for thallium, molybdenum, and phosphorus: