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HNRS 227 Lecture 11 Chapter 8 and Chapter 9

HNRS 227 Lecture 11 Chapter 8 and Chapter 9. The Periodic Table and Chemical Reactions presented by Prof. Geller with materials from Prof. Taylor. Recall from Chapter 8. History of atomic theory Electron and nucleus (protons and neutrons in nucleus both made up of quarks)

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HNRS 227 Lecture 11 Chapter 8 and Chapter 9

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  1. HNRS 227 Lecture 11 Chapter 8 and Chapter 9 The Periodic Table and Chemical Reactions presented by Prof. Geller with materials from Prof. Taylor

  2. Recall from Chapter 8 • History of atomic theory • Electron and nucleus (protons and neutrons in nucleus both made up of quarks) • The Bohr Model of the Atom • The Quantum Concept • Bohr’s Theory • Quantum Mechanics • Quantum numbers • Principal, angular momentum, magnetic, and spin • Electron Configuration • Pauli Exclusion Principle • Filling of orbital shells

  3. Periodic Table of Elements • An underlying principle • most stable state for an atom is one in which the outermost shell is filled with the maximum number of electrons • 1st Shell (1 orbit; 2 electrons) • Hydrogen (11H; 1 electron; stable ?) • Helium (24He; 2 electrons; stable ?) • Periodic Table’s 1st Row • Hydrogen and Helium

  4. Periodic Table of Elements • 2nd shell has 4 orbits with 2 electrons (maximum) per orbit (total of 8 electrons/shell) • Most stable configuration is the following: • 1st shell filled with 2 electrons • 2nd shell filled with 8 electrons • Total of 10 electrons (1020Ne) • 2nd row of Periodic Table • 8 elements (list and relate to the above)

  5. Periodic Table of Elements • 3nd shell has 4 orbits with 2 electrons maximum per orbit (total of 8 electrons/shell) • Most stable configuration is the following: • 1st shell filled with 2 electrons • 2nd shell filled with 8 electrons • 3rd shell filled with 8 electrons • Total of ___ electrons (1840Ar) • 3nd row of Periodic Table • 8 elements (list and relate to the above)

  6. Periodic Table of Elements • Rows • Number of elements in a row is not chance but reflects the maximum number of electrons in the outermost shell • Row 1 = 2 • Row 2 = 8 • Row 3 = 8 • Row 4 = 18 • etc

  7. Periodic Table of Elements Columns • Elements in a given column have similar chemical properties • All elements in column have the same number of valence electrons • Column IA has 1 electron in outer shell • Column IIA has 2 electrons in outer shell • Column IIIA has 3 electrons in outer shell • Column IVA has 4 electrons in outer shell • Column VA has 5 electrons in outer shell

  8. Periodic Table of Elements

  9. Taylor’s Take Home Message • Atoms are the chemical building blocks of all matter • Structure of atoms (electrons, neutrons, protons and their arrangement) determine the unique behavior/attributes of the elements • Of the above (No. 2), the “place” and “pairing” of the electrons are the most critical in chemical reactions • Electrons reside in defined shells (orbits) surrounding the nucleus of the atom and the electrons in the outermost shell (valence electrons) determine an atom’s chemical reactivity • Utility and periodicity of the Periodic Table of Elements is a function of the distribution of all electrons in shells, the valence electrons in the outermost shell, and the mass of the element

  10. Chemical Reactions and BondsChapter 9 • Review valence electrons • Principles of “Bonds Away” • Ionic Bonds • Metallic Bonds • Covalent Bonds • Intermolecular Forces • Common Chemical Reactions

  11. Taylor’s Take Home Message • When atoms combine to produce molecules and compounds, expect the chemical properties of the molecules/compounds to be far different than that of the constituent atoms (hierarchy theory) • Atoms bind together by re-arranging and sharing electrons • Ionic bonds • Metallic bonds • Covalent Bonds • Intermolecular forces (e.g., hydrogen bond) • Chemical interactions make and break bonds between atoms and in so doing effect a change in energy (potential and kinetic) • Weak chemical bonds (e.g., covalent bonds) play a very important role in the chemistry of life

  12. Chapter Items that won’t be emphasized from Chapter 9 • p. 184 • A Closer look • p. 186 • A Closer Look • pp. 188-191 • Percent Composition of Compounds • Ion Exchange Reactions

  13. Atoms in Proximity:Chemical Bonds • Chemical Action • when two atoms are brought together, electrons will tend to re-arrange themselves to the lowest energy state where the valence electrons are most stable • Chemical Reaction • electrons are re-arranged into bonds • Give away electrons • Accept electrons • Share electrons

  14. Ionic Bonds • Some atoms give away electrons whereas other atoms receive electrons • Example of lithium (Li) chloride (Cl) 36Li + 1735.5Cl = LiCl

  15. Ionic Bonding • Lithium (Li) Li gives up 1 electron and is left with 2 electrons (-) and 3 protons (+); net positive (+) charge • Chlorine (Cl) Cl has 1 unpaired electron in valence shell, so Cl tends to accept an electron and is left with 18 electrons (-) and 17 protons; net negative (-) charge

  16. Ionic Bonding Summary • Some atoms give away electrons while other atoms receive electrons • Example of lithium chloride Li + Cl = LiCl • Bonding via electrical attraction between Li+ and Cl- • Li+ + Cl - = Li+Cl- • Consequence: ionic bonds are underpinned by charged ions and tend to form crystals of very specific and repeating geometry (very rigid) • Example: NaCl is based on ionic bonds and is salt

  17. Ionic Bond Example: Salt

  18. Metallic Bonds • Some elements do not give or take electrons (ionic bonds) BUT share electrons • Valence electrons tend to move freely between both atoms (contrast with ionic bonds) • Significance of sharing electrons: compounds tend to show two features • Malleability (easily worked or pounded) • Conductive of electricity (good conductors) • Examples • Gold jewelry • Copper wire

  19. Covalent Bonds • Extremes of behavior in bonding • Accept or give away electrons (ionic bonds) • No tendency to share (noble gases) • Intermediate between these two extremes but • Do not form ionic bonds • Do not form metallic bonds • Yet share 1, 2, 3 and 4 electrons in unique arrangement called covalent bonds • Key: orbits of valence electrons are shared so that electrons are shared (and move) between valence shells of adjacent atoms

  20. Covalent Bond Example • Example of hydrogen fluoride (HF) • 11H and 919F • Note: Valence shell for both atoms are full • Single bond shared • Double bond

  21. Covalent Bonds with Carbon • 612C is a special case (profoundly important) • Valence electrons for C are 4 (1 in each orbit) and intermediate between giving and accepting • C - C single covalent bond (1 orbit) C • C - C two covalent bonds involving 2 orbits • Unique behavior of C C C-C-C (or H or N or __) C

  22. Behavior of Valence Electrons • Five Options • No action (e.g., inert gases) • Give away one or more electrons in valence state (positive ion leading to ionic bond) • Accept one or more electrons to valence state (negative ion leading to ionic bond) • Share an electron with many other atoms without respect to an orbit (metallic bond) • Share one or more electrons plus their orbits with another atom (covalent bond)

  23. Regarding Next Week’s Lab: Evaporation and Chemical Structure • Vaporization and chemical properties of molecules • Liquid to gas state change • State change has energy cost: endothermic (temperature decrease) • Temperature change is a function of chemical structure of molecule • Bonding and polarity

  24. Evaporation and Chemical Structure • Organic compounds • Carbon based or hydrocarbons bond with other elements via covalent bonds) • Alkanes: C and H only • Pentane (C5H12) • Alcohols: C, H and OH (hydroxyl group) • Ethanol (C2H5OH) • Structural formula • Hydrogen bonding: H bonded to N, O or F (tight bond) • Process: as chemical vaporizes, temperature change is chemical specific and is a “window” onto the chemical structure of molecule

  25. Evaporation and Chemical Structure • Hypothesis • temperature changes with vaporization in a manner that is predictable, based on the bonding among atoms involving C, H and OH • Method • Measure temperature change electronically • Record for 6 hydrocarbons • Analyze data (graphically) based on understanding of the bonds for each molecule

  26. Intermolecular Forces: Polarization & Hydrogen Bonding • Example of water (H2O) +H H+ O- • When one molecule’s distribution of atoms results in one side of the molecule having either a + or – charge • Resulting distribution of charges causes adjoining H2O molecule to align itself with + and – charges to be most stable • Called “polarity” of molecule (e.g., magnet) • Relate to lab exercise: greater polarity, greater bonding and less evaporation (less temperature change)

  27. Intermolecular Forces: Van der Waal Forces • In polarity, specific and rigid + and – fields on each molecule that does not change over time • When molecules converge, inevitable that electrons shift and re-distribute (e.g., planar compound) • In re-distribution, small net attraction between molecules arise and two molecules for weak bond • Graphite pencil lead • Stack of paper

  28. Acid – Base Reaction: Measurement • pH scale • Any increase in H+ results in more acid solution from 7 to 0 • Any increase in OH- results in more basic solution from 7 to 14 • Examples • Rainwater of 5.6 means what? • Cell pH value of 6-8 means what? • Importance to biological systems and buffering

  29. Taylor’s Take Home Message • When atoms combine to produce molecules and compounds, expect the chemical properties of the molecules/compounds to be far different than that of the constituent atoms (hierarchy theory) • Atoms bind together by re-arranging and sharing their electrons • Ionic bonds • Metallic bonds • Covalent Bonds • Intermolecular forces (e.g., hydrogen bond) • Chemical interactions make and break bonds between atoms and in so doing effect a change in energy (potential and kinetic) • Weak chemical bonds (e.g., covalent bonds) play a very important role in the chemistry of life

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