Chapter 10: Modern Atomic Theory and the Periodic Table. How does atomic structure relate to the periodic table? 10.1 Electromagnetic Radiation

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1 Chapter 10: Modern Atomic Theory and the Periodic Table How does atomic structure relate to the periodic table? 10.1 Electromagnetic Radiation Electromagnetic (EM) radiation is a form of energy that exhibits wave-like properties. Types of EM radiation vary in wavelength and frequency, and taken all together form the EM spectrum The Bohr Atom Up until about 100 years ago, scientists believed that energy could be absorbed or emitted in any quantity. o The spectrum of energy was continuous, like the colors of a rainbow. Chem Spring 2018

2 This view of EM energy could not explain line spectra observed when an excited gas was observed with a special instrument (spectroscope). o These colors arise when electrons in an atom absorb energy ( excite ), and then release it ( relax ) in discrete amounts. Energy is not released continuosuly, but only at specific wavelengths. Planck proposed that energy was quantized, or came in discrete amounts, while studying black body radiation. o Einstein used Planck s ideas to explain the photoelectric effect, and in doing so supported the particulate nature of light, or photons. Bohr applied the new quantum theory to propose an organization of electrons around the nucleus, and explain the line spectrum of hydrogen. o Electrons were constrained to specific orbits. While Bohr s model worked surprisingly well for the hydrogen atom, it did not succeed in explaining the spectral lines of any multielectron species, such as a helium atom. Chem Spring 2018

3 o Quantum theory had to be refined to explain the behavior of electrons in multielectron atoms. De Broglie proposed that electrons could behave as waves. Schrödinger tied all these ideas together and proposed an equation for the wave mechanics of electrons in atoms. One result was a probabilistic view of nature: electrons were not limited to discrete orbits, but rather were found in regions or orbitals around the nucleus with certain probabilities. Solutions to the Schrödinger equation provide a hierarchy for organizing electrons around the nucleus in terms of 3 quantum numbers which define these properties: o Distance from the nucleus Principal energy level or shell; n = 1, 2, 3, o Shape of the orbitals Subshell; s, p, d, f, o Orientation of the orbitals; How orbitals are positioned on the xyz-coordinates centered on the nucleus (for example, p x, p y, p z ) Chem Spring 2018

4 10.3 Energy Levels of Electrons Chem Spring 2018

5 10.4 Atomic Structures of the First 18 Elements The organization of electrons plays a defining role in the properties and reactivities of the elements. o Knowing this organization helps us understand properties and reactivities, and understand the layout of the periodic table. The Aufbau principle states that electrons fill the lowest energy orbitals first. This will result in the ground state electron configuration. Writing the electron configuration of an element. Chem Spring 2018

6 Another type of diagram, the orbital diagram, can be drawn showing more details of the electron organization. Some additional details and rules will help explain this diagram. o A fourth quantum number tells us the spin of the electron, spin up ( ) or spin down ( ). o Hund s rule states that when orbitals of the same energy are available, electrons fill these orbitals unpaired first with the same spin, and then as more electrons are added they will pair up with the opposite spin. o The Pauli Exclusion Principle states that an orbital can hold at most two electrons, and they must have opposite spins. We reduce the vertical energy diagram to a horizontal series of lines (or boxes). A larg gap between the lines indicates an increase in energy, while a small gap indicates orbitals of the same energy (within the same subshell). o Label the lines with the shell-subshell underneath, like 2p. Chem Spring 2018

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8 10.5 Electron Structures and the Periodic and the Periodic Table Valence electrons are the electrons in the outermost (highest numbered) energy level of an atom. o Only valence electrons are involved in chemical bonding and chemical reactions. o The remaining, underlying electrons are called core electrons. Chem Spring 2018

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10 Ch 11: Chemical Bonds: The Formation of Compounds from Atoms 11.1 Periodic Trends in Atomic Properties When a property gradually changes across a period, and there is a tendency for this change to repeat at regular intervals, we say that the property exhibits a Periodic Trend. o Elements get less metallic from left to right in a period. o Atomic radius decreases from left to right in a periodic (more protons attracting electrons in the same shell). o (First) Ionization energy is the energy required to remove the first electron from a neutral atom. Ionization energy increases from left to right in a period. After removing valence electrons, subsequent IE s jump as electrons are removed from the core Lewis Structures of Atoms The Lewis Structure (electron dot picture) of an element is a representation that shows the valence electrons of the element. o Valence electrons are drawn as dots around element symbol. o Electrons in the Lewis structure can be paired or unpaired. o Number of unpaired electrons shows covalent bonding capacity. Elements react to achieve eight valence electrons by transferring or sharing electrons; this is the octet rule. o Eight valence electrons is stable, as in the Noble gases. Chem Spring 2018

11 11.3 Ionic Bond: Transfer of Electrons from One Atom to Another Ions are formed when a neutral atoms gains or loses one or more electrons. o Metals lose electrons to form positive ions (cations). Metals lose just enough electrons to reach the electron configuration of the previous Noble gas in the Periodic Table. o Nonmetals gain electrons to form negative ions (anions). Nonmetals gain just enough electrons to reach the electron configuration of the next Noble gas in the Periodic Table. Chem Spring 2018

12 An ionic bond is formed when a metal transfers electrons to a nonmetal and the resulting ions attract each other. o An ionic bond may also be formed with polyatomic ions. Ionic bonds are strong, and ionic compounds are solids at room temperature. o Anions and cations alternate in a crystal lattice. Chem Spring 2018

13 Comparing the Radii of Isoelectronic Species Isoelectronic species have the same number of electrons, and for us, the same electron configuration. Protons Electrons 11.4 Predicting Formulas of Ionic Compounds We ve seen this already! o The current chapter provides the rationale for why the simple ions have the charges they do. o The text connects how you can use the periodic table to predict ionic charge as atoms lose or gain electrons to achieve a Nobel gas configuration. Then knowing the charges, you can predict the formulas of ionic compounds. You can go this route, or use the information about charges memorized earlier in the course. Chem Spring 2018

14 11.5 The Covalent Bond: Sharing Electrons Nonmetals bond to each other by sharing unpaired electrons to complete octets forming a covalent bond. o Hydrogen forms a duet, and an element in groups 1 to 13 may not be able to form an octet because they don t have enough unpaired electrons to begin with! Elements will form as many covalent bonds as they have unpaired electrons. o Elements in the same group will form the same number of covalent bonds because they have the same number of unpaired electrons. o Elements with more than one unpaired electrons may form multiple bonds to achieve their octet Electronegativity This property is used to predict the polarity of molecules. Electronegativity is the attraction that an atom of an element has for shared electrons in a molecule or polyatomic ion. Chem Spring 2018

15 o The shared electron pair if the covalent bond is drawn closer to the more electronegative element. o The unequal sharing of the bonding electrons leads to a charge imbalance and a polarized bond. o Electronegativity exhibits a periodic trend. H Electronegativities 2.1 Li Be B C N O F Na Mg Al Si P S Cl K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At The type of bond between two elements depends on the difference in their electronegativities. In the future, you will use bond polarity and VSEPR geometry to predict if molecules are polar! Chem Spring 2018

16 11.7 Lewis Structures of Compounds (and polyatomic ions) Here is a strategy for writing the Lewis structure of a molecule or ion. o We will only work with Main Group elements (1-2 and 13-18). o This is a major topic on the last test. o The Models lab will provide OJT! 0. If the compound is ionic, write separate Lewis structures for the anion and cation. # of VE = last digit of group number Calculate 6N+2 (N = # of non-h) 6N+2 # VE = 0, all single > 0, multiple Distribute as lone pairs to terminal atoms first Shift lone pairs, not bonding pairs Chem Spring 2018

17 11.10 Molecular Shape The shape around a central atom is determined by the number of bonded atoms and the number of lone pairs. o The arrangement of the central atom and the atoms bonded to it is given by the shape. o The lone pairs on the central atom are not part of the shape, but they do repel the bonding atoms into the proper orientation. The VSEPR model provides the rationale for the shape. o Valence Shell Electron Pair Repulsion o Put simply, bonding pairs and lone pairs repel each other to get as far apart as possible (the phrase that pays: minimize repulsion ). The following table will be provided on the last chapter test. Molecular Shape Around A Central Atom Bonded Lone Atoms Pairs Shape 2 0 Linear 3 0 Trigonal planar 2 1 Bent 4 0 Tetrahedral 3 1 Trigonal pyramidal 2 2 Bent Chem Spring 2018

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19 Let s Work On Some Examples Chem Spring 2018

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