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General Chemistry Name: Period: Unit 3 Note Packet – Atomic Structure, Nuclear Processes, and the Periodic Table History of Atomic Theory Carefully read pages 101 to 108 and 127 to 129 (time-line) in your textbook. Answer the following questions as you go. You are responsible for learning this material on your own; it will not be covered during lectures. Greeks(~450 BC) What was Democritus’ main contribution to atomic theory? Other important Greek philosophers disagreed with Democritus. What did they think? John Dalton (1766-1844) What are the four postulates of Dalton’s atomic theory of matter? a. b. c. d. Benjamin Franklin (1706-1790) Benjamin Franklin investigated electricity. He concluded there are two kinds of electric charges an object can have and he called them positive and negative. He concluded that like (same) charges repel each other while opposite charges attract each other. Michael Faraday (1791-1867) Early scientists wondered where these charges come from and had questions about their physical properties. Michael Faraday thought that the structure of atoms was related to electricity. Faraday and other scientists investigated this possibility using a cathode ray tube. Cathode Ray Tube (mid-1800s) Describe a cathode ray tube and draw a diagram of one: Scientists found that a kind of radiation streams from the to the of a cathode ray tube. One experiment showed that a cathode ray could spin a small paddle wheel, which suggested that it was actually a . They also discovered that a magnet deflects the cathode ray in the direction expected for charged particles. JJ Thomson (1856-1940) In 1896, JJ Thompson began to systematically study cathode rays. He measured the ratio of electricity charge to mass of the electrons flowing through a cathode ray tube by deflecting the beam of electrons with a magnetic field. Thomson discovered that magnetic and electric fields deflected the ray’s path in a mathematically predictable way. Thomson concluded that a cathode ray is composed of . This meant that atoms were not indivisible balls but instead had a substructure. Thomson named these negative 582731445 -1- particles . He determined the ratio of an electron’s electrical charge to its mass, which is . JJ Thomson also developed the “plum pudding” model of an atom (see pp. 106 and 128). In this model, negative charges (electrons) are distributed evenly throughout an atom’s positively charged interior, similar to the way chocolate chips are distributed throughout cookie dough. Robert Millikan (1868-1953) Millikan was able to determine the quantity of charge carried by the electron. He was able to balance the oil drop against the force of gravity to determine that charge. Millikan further built on Thomson’s experiments by using the values Thomson obtained for the charge-to-mass ratio. With these values he was able to successfully calculate the mass of a single electron by using his oil droplet experiment. The electron carries exactly unit of Millikan calculated the mass of an electron to be charge. . Ernest Rutherford (1871-1937) Studying the nature of radiation gave scientists further clues about the substructure of the atom. Ernest Rutherford continued the study of radiation. In 1909, Rutherford and his colleagues performed a famous experiment called the gold foil experiment. In this experiment, alpha particles (with a +2 charge) from radioactive polonium were aimed at a thin piece of gold foil. Most of the particles passed through the foil, but some were scattered in all directions. This led Rutherford to reject Thomson’s plum pudding model. Instead, he proposed that . What name, originally proposed by Rutherford, is still used to describe the core of an atom? Draw a diagram showing the apparatus used the gold foil experiment. in Draw a diagram showing the details of what went on in the gold foil experiment. 582731445 -2- Atom – the smallest part of matter that retains the characteristics of that type of matter considered to be the building blocks of matter atoms consist of a positively charged center, or nucleus, surrounded by negatively charged particles nucleus consists of protons and neutrons negatively charged particles are called electrons existing outside the nucleus in mathematical “regions of probability” particle proton neutron electron mass (g) mass (amu) relative charge location in atom Types of Atoms elements each different type of atom is an element elements (different types of atoms) are distinguished by the # of protons every atom of the same element has the same # of protons (ex: every atom of hydrogen has one proton, every atom of boron has 5 protons, etc.) the # of protons = elements are organized on the periodic table according to atomic number each element has a chemical symbol the symbol is the first letter of an element’s name capitalized and may have another letter from the name in lower case (examples: C for carbon and Al for aluminum) some symbols are derived from Latin names (example: Argentum is Latin for silver; the symbol for silver is Ag) # protons = # electrons in a neutral atom # protons ≠ (do not equal) # electrons in a charged (positive or negative) atom (called an ion) # neutrons can vary (not the same for every atom of an element) Ions the result of a neutral atom gaining or losing an electron cation anion charge of an ion = number of protons – number of electrons Isotopes atoms that have the same number of protons but different numbers of neutrons for example…..hydrogen has three isotopes: hydrogen 1 hydrogen 2 (deuterium) hydrogen 3 (tritium) Isotope notation: Mass number Charge 15N37 Atomic number 582731445 Example 7 protons 8 neutrons (15-7) 10 electrons (7+3) -3- Isotopes are named by their Mass number. In the example above the isotope is: Nitrogen 15 or it could also be written as Nitrogen-15 But if you look at your periodic table, the atomic mass of nitrogen is 14.01 atomic mass units. Rounding that to the nearest whole number, we see that the most common isotope of nitrogen is nitrogen-14. The example just happens to be talking about nitrogen-15. What is the most common isotope of Phosphorous? (round what you see on the periodic table) What is the most common isotope of Boron? Use your periodic table to write the symbols for the following: Mass number P __ protons 0 Charge B __ neutrons Atomic number Mass number Charge __ protons 0 __ neutrons Atomic number __ electrons __ electrons mass number = # protons plus (+) neutrons atomic mass or atomic weight takes into account all the different isotopes atomic mass/weight is an AVERAGE of all the different isotopes an example calculation for atomic mass is show below: 107 Ag and 48.65% 108 Example: A sample of silver is 51.35% atomic mass/weight? Ag. What is its average Answer: 0.5135 x 107 = 54.9445 0.4865 x 108 = 52.542 Total = 107.4865 amu = average atomic mass/weight (107.49 amu) CHANGES IN THE NUCLEUS chemical reactions involve interactions between outer (valence) electrons nuclear reactions change the composition of the atom’s nucleus radiation = Nuclear stability Most atoms have a stable nucleus ____________ Protons are positively charged like charges repel question: why doesn’t the nucleus fly apart since all the protons are repelling each other? Answer: another force, called the ________________________, is holding the nucleus together Strong nuclear force = the force that helps to hold the nucleus together (not a force we encounter in daily life only exists between subatomic particles that are extremely close) Protons Neutrons 582731445 - have both repulsive forces due to like charges, and the strong nuclear force - have only the strong nuclear force acting on them - help to balance the electromagnetic repulsion that the positive protons experience - help to hold the nucleus together ___________________________________ -4- Elements 1-20 _____________________________ Beyond element 20 ____________________________________________ Beyond element 83 (bismuth) _____________________________________ Interesting to note: There are “magic numbers” of protons and neutrons (2, 8, 20, 28, 50, and 82) that are stable Atoms with even numbers of protons and neutrons tend to be stable Types of Radioactive Decay Symbol(s) Explanation Alpha Emission of a helium nucleus (2 protons and 2 (α) neutrons) Penetrating ability Low, stopped by paper Beta (β) Emission of a high-speed electron; equivalent to the conversion of a neutron to a proton Medium, stopped by heavy clothing Gamma (γ) Emission of a high energy; similar to an x-ray High, stopped by lead Example of alpha emission example of beta emission example of gamma emission *notice that the superscripts and subscripts to the left of the arrow equal the superscripts and subscripts to the right! Write the nuclear equation for the alpha decay of gold-185: Decay rate Different radioactive elements decay at different rates Rate is described as a _________________, the amount of time (usually in years) for ½ of all the radioactive atoms (parent isotope) in a sample to decay to the more stable isotope (daughter product). Nuclear Fission & Fusion When nuclei of certain isotopes are bombarded with neutrons, they undergo ________________, the splitting of the nucleus into smaller fragments. There are only two fissionable isotopes, draw the symbols for each below its name (both are atoms, not ions): Uranium-235 Plutonium-239 The equation for the fission of Uranium-235 is: Notice that there are the same number of protons on each side (how many? ______) and the same number of neutrons on each side (how many? _________) 582731445 -5- In a chain reaction ____________________________________________________________________ __________________________________________________________________________________ An atomic bomb is a device that starts an _______________________ nuclear chain reaction. To be used for energy, _______________ must be controlled so that _____________ is released more _______________. Nuclear Fusion Nuclear fusion can be thought of as the opposite of nuclear fission, that is ___________ nuclei ___________ to produce a nucleus with a greater mass. This releases much ___________ energy than a fission reaction, but only happens above _____________ ˚C. Write the formula for 4 hydrogen nuclei combining to create one helium nucleus: MODELS OF THE ATOM (1900 to present) Rutherford proposed an atom with a nucleus and negative particles moving around the nucleus. The next major step forward on the model of the atom came after some important breakthroughs in our understanding of light behaves like a wave, but also like a stream of extremely tiny, fast-moving particles (wave-particle duality) light is a form of electromagnetic radiation electromagnetic radiation can be described in terms of: amplitude wavelength () frequency () speed of light (c) Notice: if is large, then v must be small since c is constant; the reverse is also true (therefore, long light is low energy while short light is high energy) visible spectrum of light is just part of the electromagnetic spectrum – the part detectable by the human eye the energy absorbed by and given off by objects was thought to be continuous Max Planck, in 1900, proposed that energy is only available in discrete packets, or (singular is ); the size of the quantum is related to the frequency of the radiation by a simple equation 582731445 -6- Bohr Model of the Atom (1913) Bohr applied Planck’s idea of quantized energy to the model of the atom. Electrons move around the nucleus in specific orbits, or energy levels. These energy levels are said to be quantized. An atom has several orbits, each representing a specific energy level. The energy levels are like steps of a ladder. Only specific energy values (steps) can exist – there are none in between. When an atom is not excited ( nucleus. The electrons are at the lowest energy level. If an atom gains energy ( ), an electron is displaced farther away from the nucleus to one of the higher energy levels. An atom emits energy when an electron falls from a higher energy level to a lower energy level in one sudden drop, or transition. The energy released is electromagnetic energy. The frequency () of the radiation emitted depends on the difference between the higher and lower energy levels involved in the transition. Remember that a specific frequency correlates with a specific wavelength (and color!) because c = . ), its electrons are in orbits close to the Bohr’s model of the hydrogen atom Bohr used this model to explain the discontinuous line spectrum of excited hydrogen atoms. When an electron is close to the nucleus of the atom Further away from the nucleus Absorption of energy increases the PE of the electron as it moves further away from the nucleus. Thus, an electron of an excited atom absorbs energy in the form of heat or electricity when it moves to a higher energy level. When the electron returns to ground state, it emits the energy it absorbed. The packet of energy emitted corresponds to a specific color on the line spectrum. Problems with Bohr’s Model: 1. it only explained the line spectrum for hydrogen 2. he couldn’t explain why the electrons, which he pictured as circling the nucleus like tiny planets, did not fall into the nucleus every time they emitted light Wave/Particle Paradox: Light sometimes acts as a particle and sometimes acts as a wave. Thomas Young noticed that when electrons are forced through a narrow slit, a pattern of wave interference emerged. He was familiar with Rutherford’s experiment that showed electrons are particles and needed to reconcile the two points of view. This led to the development of the Quantum Mechanical Model of the atom. The Quantum Mechanical Model is based on complex mathematical equations developed by Erwin Schrodinger that describe the wave nature of the atom and the locations of the electrons. 582731445 -7- Bohr Model vs. Quantum Theory Bohr Quantum Theory QUANTUM THEORY based partly on Heisenberg’s Uncertainty Principle the position and the momentum of a moving object cannot simultaneously be measured and known exactly there is an inherent limitation to knowing both where a particle is at a particular moment and how it is moving in order to predict where it will be in the future an electron is in an – probability space where an electron can be found a certain percentage of the time as defined by Schrodinger’s equations An orbital can hold _____ electrons (this means the bigger the energy level, the more orbitals it has!). Schrodinger’s equations (once all the mathematics has been done) describe where a particular electron is likely to be found We can know 4 pieces of information about the location of any electron within an atom. The following table summarizes what we can know. What we know about the location of an electron within an atom Type of information Label What it represents How many options? Energy Level n n=1 through n=7 Sublevel l Distance from the nucleus Shape Magnetic lm Orientation of Orbital (which direction it is pointing) Spin ls How electron is spinning 582731445 -8- n = 1 1 sublevel (s) n = 2 2 sublevels (s, p) n = 3 3 sublevels (s, p, d) n = 4 4 sublevels (s, p, d, f) Depends on the number of sublevels. p-orbitals follow the x, y, and z axes (one orbital per axis) Up = + ½ Down = - ½ The spinning of electrons generates an electric field. How many e- can it hold? 2n2 Level 1 = ______ Level 2 = ______ Level 3 = ______ Level 4 = ______ Level 5 = ______ (each orbital can hold 2 electrons) For 2 electrons to occupy the same orbital they must have opposite spin. Each type of orbital has a distinct shape. The same type of orbital always has the same shape, but each orbital is bigger in higher energy levels. s sublevels look like: p sublevels look like: d sublevels look like: f sublevels are weirder still… (check out http://winter.group.shef.ac.uk/orbitron/) DESCRIBING THE LOCATION OF ELECTRONS IN AN ATOM: To determine the electron configuration of an atom, 3 guidelines are followed: Pauli exclusion principle: Taken together, the four quantum numbers describe the state of a particular electron. For example, an electron may have these quantum numbers: 1,0,0,-1/2 (corresponding to n,l,lm, and ls). Think of these as a social security number for an individual, or a street address for a house. No two are identical. Aufbau principle: 1s is filled first, then 2s, 2p, 3s, 3p … Hund’s rule: equal energy orbitals are “degenerate orbitals”; for example, p x , py, and pz are degenerate orbitals Orbital Diagrams 582731445 -9- The locations of electrons in an atom are described using an orbital notation. In order to draw an orbital notation diagram, you must understand the quantum numbers (energy levels, sublevels, orientation, and spin). Summary of energy levels, sublevels, and orbitals Principal energy Sublevel Orbitals level n=1 1s 1s (one) n=2 2s, 2p 2s (one) + 2p (three – x,y,z) n=3 3s, 3p, 3d 3s (one) + 3p (three – x,y,z) + 3d (five – xy,xz, etc.) n=4 4s, 4p, 4d, 4f 4s (one) + 4p (three – x,y,z) + 4d (five – xy,xz, etc.) + 4f (seven – xyz, etc.) Number of electrons in each sublevel Sublevel Number of orbitals Maximum # of electrons (two per orbital) s p d f This diagram helps show the differences in energy for each sublevel within an electron cloud. Working from the bottom to the top, you can see how each atom’s electrons fill. Use this to write in the electrons for a neutral atom of Phosphorous. 582731445 - 10 - Assignment of electrons to an atom Please summarize the three things you know about how we assign electrons to an atom. Electrons fill orbitals in the following sequence: Electron Configuration: The electron configuration is another way of describing the positions of electrons in an atom using only letters and numbers. This is like a summary of the orbital notation. The number of electrons in each orbital is written as a superscript above the orbital designation. The electron configuration for sulfur is as follows: Write the electron configurations for the following atoms: Element H Electron configuration Element Al He O Li C Cl Mg N Ar Electron configuration At atomic number greater than 18, the sublevels begin to fill out of order. A good approximation of the order of filling can be determined using either the diagonal rule or the location of the element on the periodic table. 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p 5d 6s 6p 6d 7s 7p 7d Write the electron configurations for the following atoms. 4f 5f 6f 7f Element K Electron configuration V Co Zr Cu Looking for patterns: 582731445 - 11 - Write the electron configurations for the following elements: He Li Ne Mg Ar Ca Ti Ni Br Describe the pattern(s) in the electron configurations above: The Noble Gas Shortcut – Write the electron configurations for the following atoms using the shortcut you just learned: Be Ne Al Ca Sc Ni I Ga Valence Electrons The electrons in the outermost energy level are active in bond formation. They are only in the outermost s and p sublevels (between 1 and 8 electrons see the periodic table). Since only s and p electrons are valence electrons, the maximum number of valence electrons is eight. For the elements below, do the following: 1. Write the electron configuration 2. Determine the number of valence electrons 3. Write the number of valence electrons and circle it Example: Carbon Electron configuration: 1s22s22p2 Number of valence electrons: 4 Example: Oxygen Electron configuration: Number of valence electrons: 582731445 - 12 - For the following atoms, write the electron configuration and circle the valence electrons. 1. fluorine 7. zinc 2. phosphorus 8. carbon 3. iron 9. oxygen 4. argon 10. aluminum 5. potassium 11. hydrogen 6. magnesium 12. helium Electron Configurations for Ions Follow all the same rules that you learned in writing electron configuration for atoms. Before you begin, determine which noble gas is closest to the atom/ion in questions, then count how many electrons will be lost or gained to make the ion. Subtract or add that number of electrons to the number of electrons in the original atom. Write the electron configuration for the ion using the number of electrons in the ion. Note that it will usually be the same as the electron configuration for the nearest noble gas. For example: Magnesium atom magnesium ion: Magnesium is closest to the noble gas so a magnesium atom will gain/lose electrons when it becomes an ion. As an ion, magnesium has the same electron configuration as . Write the electron configurations for the following elements and their most common ions: Sodium sodium ion Oxygen oxide (ion) Phosphorus phosphide (ion) Beryllium beryllium ion Iodine iodide (ion) Lewis electron dot diagrams Lewis diagrams or electron dot diagrams emphasize the valence electrons. Such diagrams show: Kernel - the nucleus plus the inner electrons (all non-valence electrons) - represented by the element’s symbol Valence electrons – electrons in the outermost s and p orbitals - represented by a series of dots around the kernel (element’s symbol) - s orbital is 2 dots at 12 o’clock, px, py, and pz orbitals are 2 dots each at 3, 6, and 9 o’clock Example: sulfur Draw Lewis dot diagrams for the following atoms to show the valence electrons for each element: 1. calcium 582731445 5. bromine - 13 - 9. phosphorus 2. potassium 6. carbon 10. hydrogen 3. argon 7. helium 11. vanadium 4. aluminum 8. oxygen 12. silver The Periodic Table Mendeleev (Russian) – Moseley (English) – Periodic Law – Basic Organization Groups / Families – vertical column – Row / Period – horizontal row Periodic Characteristics: 1. atomic size 2. ionic size 3. metallic properties luster – shiny conductivity – able to transfer heat or electrons malleability – can be rolled or hammered into sheets ductility – can be drawn (pulled) into a wire explained by…. 4. ionization energy 5. electronegativity 582731445 - 14 - nonmetallic properties Explained by… Groups: Elements in the same group have more similarities than elements in the same period because . Alkali Metals (Group 1) Alkaline Earth Metals (Group 2) Transition Metals Metalloids – elements that have both metallic and nonmetallic properties Diatomic elements – HONClBrIF elements – found combined with self – Br2, I2 N2, Cl2, etc. Halogens (Group 17) Noble / Inert Gases (Group 18) 582731445 - 15 -