1 / 24

Valence Electrons and Chemical Bonds 06 and 08 October 2015

This lecture provides an overview of the principles of valence electrons and chemical bonds, including ionic, metallic, and covalent bonds. It also covers intermolecular forces and water's unique properties. Examples and explanations are given to illustrate these concepts.

kimberlymurray
Download Presentation

Valence Electrons and Chemical Bonds 06 and 08 October 2015

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Valence Electrons and Chemical Bonds06 and 08 October 2015 • Principles of Valence Electrons and Bonds • Ionic Bonds • Metallic Bonds • Covalent Bonds • Intermolecular Forces • Water’s Unique Properties

  2. Periodic Table

  3. Two Atoms in Proximity • Observation: when two atoms are brought together, electrons re-arrange to the lowest energy state (i.e., where valence electrons are most stable) • Consequence: distribution of electrons among atoms re-arranged into bonds • Give away electrons • Accept electrons • Share electrons

  4. Orbits, Shells and Electrons http://www.colorado.edu/physics/2000/applets/a2.html

  5. Ionic Bonding • Some atoms give away electrons, whereas other atoms receive electrons • Example of water (two atoms of hydrogen + one atom of oxygen) – H2O 11H + 816O + 11H = H2O • Example of lithium (Li) chloride (Cl) 36Li + 1735.5Cl = LiCl

  6. 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

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

  8. Ionic Bonding: Salt

  9. Metallic Bonds • Elements that do not give or take electrons (no ionic bonding) but share electrons • Valence electrons move freely between adjacent 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

  10. Covalent Bonds (Remember These) • Extremes of behavior in bonding • Accept or give away electrons (ionic bonds) • No tendency to share (noble gases) • Intermediate between these two extremes • 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 (contrast with metallic bonds)

  11. Simplest Example: Hydrogen Gas and Covalent Bonds

  12. Covalent Bonds: Carbon • 612C is a special case (important in biology & chemistry) • 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 (2 orbits) • Unique behavior of C: four simultaneous C - C bonds C C C C C

  13. Periodic Table

  14. 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/shell (metallic bond) • Share one or more electrons plus their orbits with another atom (covalent bond)

  15. Valence Electrons and Chemical Bonds06 and 08 October 2015 • Principles of Valence Electrons and Bonds • Ionic Bonds • Metallic Bonds • Covalent Bonds • Inter-molecular Forces • Water’s Unique Properties

  16. Intermolecular Forces: Polarization & Hydrogen Bonding • Example of water (H2O) • When one molecule’s distribution of atoms results in one side of the moleculehaving either a + or - charge • Resulting distribution of charges causes adjoining H2O molecule to align with + and - charges to be most stable • Called “polarity” of molecule (e.g., magnet)

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

  18. Valence Electrons and Chemical Bonds06 and 08 October 2015 • Principles of Valence Electrons and Bonds • Ionic Bonds • Metallic Bonds • Covalent Bonds • Intermolecular Forces • Water’s Unique Properties

  19. Water and Its Properties: Liquid State • Water is liquid over broad range of temperatures 0oC to 100oC • Comparison with other compounds CompoundChemical FormulaFreeze VaporRange(oC) (oC) (oC) Water H2O 0 100 100 Ammonia NH3 -78 -33 45 Methane CH4 -182 -164 18

  20. Water and Its Properties: pH • pH scale: ionization of H2O H+ and OH- • Increase in H+ results in more acid solution from 7 to 0 • Increase in OH- results in more basic solution from 7 to 14 • Examples • Rainwater of 5.6 means what? • Cell pH value of 7-9 means what? • Importance to biological systems and buffering See text for more discussion of pH

  21. Water and Its Properties: Freezing • Water is unusual in that H2O is less dense as a solid than as a liquid (it floats!) • Mechanism • H2O expands when it solidifies • Due to hydrogen bonding • Consequence • Ponds, lakes and ocean freeze from the top down

  22. Hierarchy Theory and Emergent Properties of H2O • Principle of hierarchy theory • Principle of emergent properties • a priori: combine one atom of O with two atoms of H and what would you expect? • Emergent properties • Liquid • Hydrogen bonding and polarity • H+ and OH- in solution • Solvent • Range of temperature at which liquid • Three phases (gas, liquid and solid)

  23. Importance of Chemical Bonds • Interactions between atoms (i.e., bonding) responsible for creation of myriad of molecules and compounds (natural and human made) • Interactions between atoms (i.e., bonding) responsible for all chemical reactions that effect a change in energy (i.e., kinetic and potential) • Granola bar for lunch • Role of covalent bonds in living systems is basis for most of the uniqueness of life • Water molecule’s unique properties and importance in sciences (physical, chemical and biological)

More Related