Chemical Reactions or Bonds Away with Valence Electrons
Chemical Reactions or “Bonds Away” with Valence Electrons Review valence electrons n Principles of “Bonds Away” n Ionic Bonds n Metallic Bonds n Covalent Bonds n Intermolecular Forces n Common Chemical Reactions n
Take Home Message n n 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 n n n 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
Chapter Deletions (No. 9) Pp 184 (A Closer look) n Pp 186 (A Closer Look) n Pp 188 (Percent Composition of Compounds) – 191 (Ion Exchange Reactions) n
Valence Electrons and Chemical Bonding Review valence electrons n Principles of “Bonds Away” n Ionic Bonds n Metallic Bonds n Covalent Bonds n Intermolecular Forces n Common Chemical Reactions n
Atoms in Proximity: “Bonds Away” n n Hypothesis: 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 Product: electrons are re-arranged into bonds n n n Give away electrons Accept electrons Share electrons
Valence Electrons and Chemical Bonding Review valence electrons n Principles of “Bonds Away” n Ionic Bonds n Metallic Bonds n Covalent Bonds n Intermolecular Forces n Common Chemical Reactions n
Ionic Bonding Atoms give away electrons whereas other atoms receive electrons n Example of lithium (Li) chloride (Cl) n 3 Li 6 + 17 35. 5 Cl = Li. Cl
Ionic Bonding n Lithium (Li) Li gives up 1 electron and is left with 2 electrons (-) and 3 protons (+); net positive (+) charge n 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
Ionic Bonding n n Atoms give away electrons while other atoms receive electrons Example of lithium chloride Li + Cl = Li. Cl n Bonding via electrical attraction between Li+ and Cln n n 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: Na. Cl is based on ionic bonds and is salt
Ionic Bonding: Salt
Valence Electrons and Chemical Bonding Review valence electrons n Principles of “Bonds Away” n Ionic Bonds n Metallic Bonds n Covalent Bonds n Intermolecular Forces n Common Chemical Reactions n
Metallic Bonds n n n Elements that 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 n n n Malleability (easily worked or pounded) Conductive of electricity (good conductors) Examples n n Gold jewelry Copper wire
Valence Electrons and Chemical Bonding Review valence electrons n Principles of “Bonds Away” n Ionic Bonds n Metallic Bonds n Covalent Bonds n Intermolecular Forces n Common Chemical Reactions n
Covalent Bonds n Extremes of behavior in bonding n n n Accept or give away electrons (ionic bonds) No tendency to share (noble gases) Intermediate between these two extremes but n n n Do not form ionic bonds Do not form metallic bonds Yet share 1, 2, 3 and 4 electrons in unique arrangement called covalent bonds n Key: orbits of valence electrons are shared so that electrons are shared (and move) between valence shells of adjacent atoms
Covalent Bonds n n n Example of hydrogen fluoride (HF) 1 H and 9 F 1 19 Note: Valence shell for both atoms are full Single bond shared n Double bond n
Covalent Bonds: Carbon n n 6 C 12 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
Behavior of Valence Electrons: Five Options n n n 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)
This Week’s Lab: Evaporation and Chemical Structure n Vaporization and chemical properties of molecules Liquid to gas state change n State change has energy cost: endothermic (temperature decrease) n n Temperature change is a function of chemical structure of molecule n Bonding and polarity
Evaporation and Chemical Structure n Organic compounds n Carbon based or hydrocarbons bond with other elements via covalent bonds) n Alkanes: C and H only n n Alcohols: C, H and OH (hydroxyl group) n n Pentane (C 5 H 12) Ethanol (C 2 H 5 OH) 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
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 n Method n Measure temperature change electronically n Record for 6 hydrocarbons n Analyze data (graphically) based on understanding of the bonds for each molecule n
Valence Electrons and Chemical Bonding Review valence electrons n Principles of “Bonds Away” n Ionic Bonds n Metallic Bonds n Covalent Bonds n Intermolecular Forces n Common Chemical Reactions n
Intermolecular Forces: Polarization & Hydrogen Bonding n n n Example of water (H 2 O) +H H+ OWhen one molecule’s distribution of atoms results in one side of the molecule having either a + or – charge Resulting distribution of charges causes adjoining H 2 O 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)
Intermolecular Forces: Van der Waal Forces n n n 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 n n Graphite pencil lead Stack of paper
Valence Electrons and Chemical Bonding Review valence electrons n Principles of “Bonds Away” n Ionic Bonds n Metallic Bonds n Covalent Bonds n Intermolecular Forces n Common Chemical Reactions (p. H) n
Acid – Base Reaction: Measurement n n p. H 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 n n n Rainwater of 5. 6 means what? Cell p. H value of 6 -8 means what? Importance to biological systems and buffering
Valence Electrons and Chemical Bonding Review valence electrons n Principles of “Bonds Away” n Ionic Bonds n Metallic Bonds n Covalent Bonds n Intermolecular Forces n Common Chemical Reactions (p. H) n
Take Home Message n n 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 n n n 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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