Chemical Reactions or “Bonds Away” with Valence Electrons

Chemical Reactions or “Bonds Away” with Valence Electrons • Review valence electrons • Principles of “Bonds Away” • Ionic Bonds • Metallic Bonds • Covalent Bonds • Intermolecular Forces • Common Chemical Reactions

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

Chapter Deletions (No. 9) • Pp 184 (A Closer look) • Pp 186 (A Closer Look) • Pp 188 (Percent Composition of Compounds) – 191 (Ion Exchange Reactions)

Valence Electrons and Chemical Bonding • Review valence electrons • Principles of “Bonds Away” • Ionic Bonds • Metallic Bonds • Covalent Bonds • Intermolecular Forces • Common Chemical Reactions

Atoms in Proximity: “Bonds Away” • 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 • Give away electrons • Accept electrons • Share electrons

Ionic Bonding • Atoms give away electrons whereas other atoms receive electrons • Example of lithium (Li) chloride (Cl) 36Li + 1735.5Cl = LiCl

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

Ionic Bonding • 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

Ionic Bonding: Salt

Metallic Bonds • 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 • Malleability (easily worked or pounded) • Conductive of electricity (good conductors) • Examples • Gold jewelry • Copper wire

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

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

Covalent Bonds: 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

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)

This 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

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

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

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)

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

Valence Electrons and Chemical Bonding • Review valence electrons • Principles of “Bonds Away” • Ionic Bonds • Metallic Bonds • Covalent Bonds • Intermolecular Forces • Common Chemical Reactions (pH)

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

Valence Electrons and Chemical Bonding • Review valence electrons • Principles of “Bonds Away” • Ionic Bonds • Metallic Bonds • Covalent Bonds • Intermolecular Forces • Common Chemical Reactions (pH)

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

Chemical Reactions or “Bonds Away” with Valence Electrons