Chemistry: Week 17

This week of the Chemistry course from the Ron Paul Curriculum covered a lot of concepts, mainly focusing on covalent bond formations, hydrogen, coordinate covalent bonds, multiple chemical bonds, and structural formula.

Lesson number one started off describing covalent bonds. Unlike ionic bonds, covalent bonds share valence electrons instead of transferring them. We quickly reviewed the basic processes for chemical bond formations, and then focused on the diatomic gas, hydrogen, as a basic example of covalent bonding. To help understand the concept we drew the electron dot structure and Lewis structural formula for the formation of H2. The lesson concluded with the comparison of ionic to covalent bonds.

Ionic Bonds:

  • tend to solid crystalline structures at room temperature
  • very high melting points
  • tend to form electrolyte solutions when dissolved in water

Covalent Bonds:

  • can be solid, liquid, or gas at room temperature
  • lower melting points
  • do not form electrolyte solutions when dissolved in water

  • electrolyte solutions when dissolved in water
  • Lesson two continued describing covalent bonds, including the concepts of the potential energy of two atoms, the bond length, and bond dissociation. Using the example of H and Cl forming the covalent compound HCl, we learned how this bond shows the bonding pair of electrons and the non binding pairs. HCl has 1 bonding pair, where the Hydrogen is sharing two electrons with Chlorine, and 3 non binding pairs. Coordinate covalent bonds is a bond that is formed when both electrons are donated by a single atom. A good example of this is NH3.

    All of the molecules and covalent bonds that have been looked at through the past lessons have obtained a noble gas electron configuration through sharing pairs of electrons. The driving force behind this is called the octet rule, where it is the tendency of the atoms to have 8 electrons in their valence shells. This is the same for ionic bonds and covalent bonds. Covalent bonds are capable of having single, double, and triple bonds. The greater number of electrons being shared the larger the energy input is needed to break that bond apart.

    Lesson three focused on binary compounds and the naming of covalent compounds. Greek prefixes are used to indicate the number of nonmetals that are present:

    Rules:

    • The name of the compound, generally, has the elements in the order given in the formula.
    • The first element is given its exact name.
    • The second element is given the suffix -ide.
    • Prefixes are used to denote how many atoms of each element are present.

    The rest of the lesson we worked on learning formulas and naming them.

    Lesson four was about structural formulas and how to write them. We started by writing the skeletal structure of the compound H2O. Next you write out all the valence electrons of the elements in a “electron pot” and then determine the number of covalent bonds using the N – A = S equation. Once you have figured that out, place the electrons around the Lewis structure to account for chemical bonds and distribute the remainder of the valence electrons around the structure so that each atom has its valence energy level completely filled.

    And that was week 17 of Chemistry.

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