# Atomic Structure & Elements

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1 Atomic Structure & Elements (Chapter 2: p 24-42, 46-48, Problems (all 2.--): 3, 5b, 7, 15, 17, 19, 21, 23, 25, 27, 29, 63, 69 Picture of an Atom Nucleus Electron Cloud I. Dalton s Atomic Theory 1. All matter is made up of atoms. 2. Atoms can neither be created nor destroyed. 3. Atoms of a particular element are alike. 4. Atoms of different elements are different from one another. 5. A chemical reaction involves either the union or separation of individual atoms. (~1808) II. Subatomic particles Particle Symbol Charge Relative Mass (amu) protons p neutrons n electrons e Compare charges and relative mass. An amu or atomic mass unit is a convenient relative mass unit, because a proton and neutron each have a mass of about 1 amu. 1 amu = x g. III. How p +, n, and e fit together in atoms & ions A. p + & n are bound together in the nucleus, located in the center of the atom. Note: The p + number =. What is in the nucleus of the most common form of the lithium (Li) atom? Note: 7 Li means that the sum of the p + and n in the nucleus is 7. Key proton neutron 1 7 Li nucleus

2 B. Analogy for size of nucleus relative to the whole atom: The whole atom is the size of a major league baseball park. The nucleus would be like a marble sitting out past second base. This means: 1. Nucleus: very small & dense. (Does something about the nucleus bother you?) 2. Most of the atom s space is occupied by e, which have very little mass. C. Electrons (e ) are found in orbitals located outside of the nucleus. 1.The Bohr model (planetary?) can be represented as: (Note: This drawing is not to scale.) 2. A Li atom has 3 e. 3. The circles that the e are located on are called e - e - e - 4. Electrons (e ) in orbits farther from the nucleus are less tightly bound. = the nucleus 5. Energies of e in the different orbitals were determined by observing light emission from atoms. (Like neon lights.) Based on your previous studies, what holds the e near the nucleus? A D. Symbolism: Z E Example for carbon with 6 neutrons is 12 6 C 1. E = elemental symbol, 2. A = the mass number (sum of the number of protons + neutrons.) 3. Z = E. Isotopes 1. Atoms that have the same number of protons (ie the same atomic number), but a different number of neutrons (ie different mass number. 2. Many atoms have isotopes, some of which are more stable than others. 3. Example: isotopes of carbon: 12 C, 13 C, 14 C 12 C is the most abundant isotope. 14 C is radioactive. (Used in determining age of old objects.) 12 C 13 C 14 C p + n e 2

3 4. Neutrons are thought to act as a kind of glue that holds the protons and neutrons together. Otherwise the protons would repel each other. IV. Looking at the elemental symbols on the periodic table. A. For example, look at carbon on the Periodic Table: The atomic number is The atomic weight is amu and is the weight of the average C atom on earth. Remember: The atomic mass unit (amu) is a convenient unit, defined relative to a 12 C atom. One atom of 12 C is defined to weigh amu. Are the atomic mass and mass number the same thing? B. Instruments like the mass spectrometer allow chemists to make accurate determinations of the weight and abundance of the different isotopic forms of an element. C. Carbon isotopic mass & abundance data: isotope abundance (%) mass (amu) 12 C C source: CRC Handbook, 59 th ed. 6 C D. Qualitative Atomic Mass: The weight of the average atom should be quite close to 12 (since most of the C atoms are 12 C), but a little bit above 12 (because there are some 13 C atoms which weigh more than 12 amu.) E. Quantitative Atomic Mass: Calculation of the average weight 12 C component: amu 98.89/100 = amu 13 C component: amu 1.11/100 = amu + average weight is amu This is close to the amu atomic weight value in Periodic Table. The % abundance is also called the natural abundance. Do you think the natural abundance values on earth are the same as those on other planets, meteors, asteroids, etc.? 3

4 V. Quantum mechanics (views e as waves instead of particles) describes atomic behavior better than Bohr model. A. Bohr model only works well for H atoms. B. Using a specific mathematical approach (that requires very advanced math) gives a model with much better predictive capabilities. C. We will use results from quantum mechanics. Don t sweat the math. VI. Atomic orbitals (where e hang out) A. Principal quantum numbers: 1, 2, 3, etc.. Describes levels or shells around the nucleus (correlates to period numbers on the Shells Orbitals 1 (smallest shell) one 1s 2 one 2s, three 2p 3 one 3s, three 3p, five 3d 4 one 4s, three 4p, five 4d, seven 4f B. Orbital shapes: (see the Orbitron at 1. s orbitals are one lobed and spherical 2. p orbitals are 2 lobed and are roughly dumbbell shaped. 3. d and f orbitals have relatively complicated shapes. periodic table) C. Energies You might consider that all of the orbitals of an atom (up to and beyond 7f) always exist, but only become interesting when they are occupied by e. Orbital occupancy by e - : 1. Lowest energy orbitals (those closest to nucleus) are occupied first. 2. An orbital can only contain two e. 3. When orbitals of equal energy (ex.: 2p x, 2p y, 2p z ) are being filled, put one e in each orbital first, then add start adding additional e. D. We describe the orbital occupancy of an atom (or ion) by writing its electronic configuration. In this class you will do this by direct application of the Periodic Table. 4

5 Periodic Table 1A 1 H A 3A 4A 5A 6A 7A 3 Li Na K Rb Cs Fr (223) 4 Be < Atomic number < Elemental symbol < Atomic weight 12 Mg B 4B 5B 6B 7B < B > 1B 2B 20 Ca Sr Ba Ra Sc Y La Ac (227) 22 Ti Zr Hf Rf (261) 23 V Nb Ta Ha (263) 24 Cr Mo W Sg (263) 25 Mn Tc (98) 75 Re Ns (265) 26 Fe Ru Os Hs (265) 27 Co Rh Ir Mt (266) 28 Ni Pd Pt (269) 29 Cu Ag Au (272) 30 Zn Cd Hg (277) 5 B Al Ga In Tl C Si Ge Sn Pb N P As Sb Bi O S Se Te Po (209) 9 F Cl Br I At (210) 8A 2 He Ne Ar Kr Xe Rn (222) 58 Ce Th Pr Pa Nd U Pm (145) 93 Np Sm Pu (244) 63 Eu Am (243) 64 Gd Cm (247) 65 Tb Bk (247) 66 Dy Cf (251) 67 Ho Es (252) 68 Er Fm (257) 69 Tm Md (258) 70 Yb No (259) 71 Lu Lr (260) E. Let s try a few. (Remember: Elemental identity is determined by p + number.) 1. How many electrons are there in a lithium (Li) atom, and in which orbitals are the electrons found? Draw an orbital filling diagram. Write the electronic configuration. 2. Write the electronic configuration for beryllium (Be). 3. For boron (B) 4. For carbon (C) 5. For nitrogen (N). 6. For oxygen (O). 7. For fluorine (F) 8. For neon (Ne) 5

6 9. For sodium (Na) 10. For magnesium (Mg) 11. For phosphorous (P) 12. For calcium (Ca) 13. For iron (Fe) 14. For selenium (Se) F. Important definition: The outermost s and p electrons are called valence electrons. Because they are outermost, they can be involved in sharing (covalent bond formation) or they may be lost or gained (in ion formation). Underline the valence electrons in the examples, above. Periodic Table of the Elements: Orbital filling 1A 2A 3B 4B 5B 6B 7B 8B 1B 2B 3A 4A 5A 6A 7A 8A H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Tc Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Lilac = filling s orbitals Center, Yellow = filling d orbitals Right, Green = filling p orbitals 6

7 VII. Development of the Periodic Table (Mendeleev) This discussion is important in helping you understand how science works and why humans have found it useful. A. Origin: Mendeleev organized the elements into a table based on chemical & physical properties & atomic weight (Note: Protons ( p + ) had not yet been found). 1. He arranged the table so that atoms in a column had similar chemical and physical properties: CH 4 NH 3 H 2 O HF SiH 4 PH 3 H 2 S HCl HBr 2. And so that atomic weight increased from left to right across the table. That is, F weighs more than O, O weighs more than N, etc. 3. And so that atomic weight increased from top to bottom. That is, F weighs more than Cl, Cl weighs more than Br. B. When he did this, he saw gaps in the table. 1. He thought the gaps were elements that existed that had not yet been discovered. 2. For the gap below silicon, he made very specific predictions (interpolation) of the properties of this undiscovered element. He called this undiscovered element, Eka-silicon. Periodic Table of the Elements Known in Mendeleev s Time H Li Be Na Mg B C N O F Al Si P S Cl K Ca Ti V Cr Mn Fe Co Ni Cu Zn As Se Br Sr Y Zr Nb Mo Rh Pd Ag Cd Sn Sb Te I Ba Ta W Os Ir Pt Au Hg Pb Bi 7

8 Note: Some of the lanthanides & actinides were known by 1839, but these groups have been omitted for clarity. Mendeleev was able to use the elements shown above to make some striking predictions based on periodicity. Known since ancient times (C, S Fe, Cu, Ag, Sn, Au, Hg, Pb) Discovered in the Middle Ages (P) Discovered Date source: C. About 15 years after his prediction, Eka-silicon (Ge) was discovered & found to have properties very close to those he predicted. Predicted for eka-silicon Measured for Ge Atomic wt Density (g/cm 3 ) Density Cl (g/cm 3 ) compound from Chemistry 3 rd ed., Atkins & Jones Scientists were impressed by the predictive power of Mendeleev s ideas!!! D. Why are elements in the same column in the Periodic Table similar in their chemical reactivity? 1. Representative elements in the same column have the same valence e numbers. 2. Because of this they tend: a) to form the same types of ions b) to have similar (not identical) chemical reactivities E. Why aren t elements in the same column [ex. oxygen (O) & selenium (Se)] identical in chemical reactivity? 1. nuclei are different, which influences outer e. 2. different amounts of inner e -, which cause electrostatic shielding between nuclei and outer electrons F. Another way to look at the elements in the periodic table: 1. Metals Examples: Mg, Cr, Fe, Cu, Co, Au, Zn, Ag, Ni, Ti, Na, Li, Ca, Ba a) Most tend to lose e b) Conduct heat & electricity well c) Shiny, form into thin sheets & wires, etc. 2. Non-metals Examples: C, N, O, F, S, P, Br, Cl, He, Ne 8

9 a) Most tend to gain e b) Often don t conduct heat & electricity well 3. Metalloids have intermediate properties. Examples: Si, As, Sb, Ge, Te Modern Periodic Table of the Elements * 1A 1 H A 3A 4A 5A 6A 7A 8A 2 He Li Be < < < Atomic number Elemental symbol Atomic weight 5 B C N O F Ne Na Mg B 4B 5B 6B 7B < B > 1B 2B 13 Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc (98) 44 Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po (209) 85 At (210) 86 Rn (222) 87 Fr (223) 88 Ra Ac (227) 104 Rf (261) 105 Ha (263) 106 Sg (263) 107 Ns (265) 108 Hs (265) 109 Mt (266) 110 Ds (269) 111 Rg (272) 112 Cn (277) 58 Ce Pr Nd Pm (145) 62 Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu (244) 95 Am (243) 96 Cm (247) 97 Bk (247) 98 Cf (251) 99 Es (252) 100 Fm (257) 101 Md (258) 102 No (259) 103 Lr (260) *Organized by increasing number of protons rather than atomic weight. Indicate the metals, non-metals and metalloids on the periodic table above. 9

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