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Do Now

Do Now. Write the formula or name for the following compounds: Cobalt III Chloride Nickel II Sulfide MnCl 4 ZnF 2 V(SO 4 ) 2. Non-Chemical Bonds. Metallic Bonding and Intermolecular Forces. The Nature of Metals. Metallic Bonds. Properties of Metals. Chemical Properties.

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Do Now

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  1. Do Now • Write the formula or name for the following compounds: • Cobalt III Chloride • Nickel II Sulfide • MnCl4 • ZnF2 • V(SO4)2

  2. Non-Chemical Bonds Metallic Bonding and Intermolecular Forces

  3. The Nature of Metals Metallic Bonds

  4. Properties of Metals Chemical Properties Physical Properties Lustrous - reflect light, shine when they are polished Flexible malleable – ductile – Are solids at room temperature Good conductors of heat and electricity • Tend to lose electrons easily • Have low ionization energy (energy needed to remove electrons) • Have low electronegativity (attraction for electrons) • Form positive ions when combining with other atoms can be rolled or hammered into sheets can be drawn intowires (except formercury) WRITE DOWN

  5. Metallic Bonding • The metal atoms in a piece of metal lose electrons and become positive ions (cations). • The cations are attracted to the electrons, but not strongly enough to hold onto them. • As a result, the electrons are free to move among the cations. They are mobile. • So, a metallic bond is positive ions in a “sea” of mobile electrons. WRITE DOWN

  6. Explaining Metal Properties The bond between the electrons and the cations is strong. The electrons are free to move between the cations. The electrons that bond the cations are mobile. The electrons ,which are held loosely, jump up to higher energy levels, and fall back down giving off light.

  7. Do Now • Describe how to name a compound with the stock system. • Describe how to name a binary covalent compound.

  8. The Ties that Bind Intermolecular FOrces

  9. Dipoles • Iodine monochloride (ICl) is a solid at room temperature. • What holds the crystal together? • Check the electronegativities and the shape of the iodine monochloride molecule. • The molecule is polar. • Polar molecules are often called dipoles since they have two poles. δ+ δ− 3.2 I Cl 2.7 0.5 WRITE DOWN

  10. Dipole-Dipole Attraction • There is an attraction between the oppositely charged ends of the polar iodine monochloride molecules. • dipole-dipole attraction: The attraction between the oppositely charged ends of polar molecules. • It accounts for the fact that iodine monochloride is solid. δ− δ+ ♥ ♥ Cl I δ+ δ− I Cl WRITE DOWN

  11. Van der Waals Forces • Consider iodine (I2): • It is a solid at room temperature. • It is a diatomic molecule with nonpolar bonds. • What holds the iodine crystals together? • Even though iodine consists of nonpolar, diatomic, neutral molecules, there is still a weak attraction : • Van der Waals Forces to be caused by temporary dipoles that form as the electrons move, causing temporary dipole-dipole attractions. WRITE DOWN

  12. More on Van der Waals Forces • Look at the phases of the elementsin the halogen family. • Under normal conditions (STP),fluorine and chlorine are gases,bromine is a liquid, and iodineand astatine are solids. • They are all diatomic, nonpolar molecules. What accounts for the differences in boiling points? Gas Liquid Solid WRITE DOWN As the size of an atom or molecule increases, so does the Van der Waals force of attraction.

  13. Evidence for Other Bonds • Consider the following family of compounds: • The boiling points decreasewith mass as expected. • You’d expect water to have the lowest boiling point. • But it doesn’t! Why not!? 100 80 130 1 Temperature (°C) 60 81 −40.1 40 34 −60.2 20 100 0 ? 18 −20 −40 −80 −60 0 20 40 60 80 100 120 140 Mass (amu)

  14. Hydrogen Bonds • Water’s high boiling point is evidence for intermolecular forces much stronger than Van der Waals forces. . ., but why? • Determine theelectronegativitydifferences in thefollowing compounds. • Water is a verypolar molecule. • Water molecules form a strong dipole-dipole attraction called a hydrogen bond. 2.1 2.1 0.0 2.1 2.6 0.5 2.1 2.6 0.5 2.1 3.4 1.3

  15. Forming Hydrogen Bonds • Hydrogen bonds form between molecules of nitrogen (N), oxygen (O), or fluorine (F), and hydrogen. Hydrogen bonding is FON • The slightly positive hydrogen is strongly attracted to the slightly negative nitrogen, oxygen, or fluorine. WRITE DOWN

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