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Essential Concepts from Chem 113

As a student of General Chemistry, you are expected to know the following concepts that are crucial to your success not only in General Chemistry but also in all other upper level chemistry courses.

  1. Molecular Formula - Number of atoms of each element are represented by subscripts. Correct formula means correct symbols for elements and correct subscripts. This in turn is required for correct naming of the substance. Learn to associate symbols and names for the most common elements from the periodic table. Learn to use the periodic table to get metals and non-metals, color codes for physical states, and the common charges (aka oxidation states).
  2. Ionic and covalent compound - Ionic compounds dissociate and form ions in solution or when molten and conducts electricity. They contain at least one metal. Molecular compounds stays as a whole unit all the time. Exception are the acids which product H+ ions. Example:  HCl (g) is a molecular compound, but when dissolved in water, forms H+ and Cl- ions and can conduct electricity.
  3. Mole and Molar Mass - Mole represents 6.02E23 units of anything. In chemistry, mole represents 6.023E23 units of the formula. Molar mass is the mass of one mole of the formula units. Each atom’s molar mass is given in the periodic table. Thus, a correct molecular formula can lead to correct molar mass.
    Moles = Grams of substance/Molar Mass
  4. Mole and Molarity - Molarity stands for the number of moles of a substance dissolved in 1 L of solution. Molarity (M) is the unit for concentration of the solution. In order to know the moles that are used in a reaction, you need to know the volume and concentration of the solution. Moles = Volume x Molarity
  5. KMT - Kinetic Molecular Theory explains the behavior of ideal gases. It assumes that gas particles have no volume of their own, have no attractions to each other, and exerts pressure when pounding on the walls of the container. Temperature is a measure of average velocity (average Kinetic Energy). Students mistakenly think that all molecules move at the same speed. Be clear about the nature of distribution of velocities amongst molecules and its dependence on temperature. Higher temperature means higher velocities for all molecules and higher average velocity.
  6. Ideal Gas Law - Combining Boyle's Pressure-Volume, Charles' Pressure (or Volume)-Temp, and Avogadro's Pressure (or Volume)-number of moles relationships, the ideal gas law states PV = nRT, where 'R' is the ideal gas constant. Ideal gas law can be theoretically explained by KMT.
  7. Reaction Stoichiometry - This refers to the moles ratio with which reactants must be mixed to produce products without any wastage. A balanced equation gives the mole ratio or the reaction stoichiometry. A balanced reaction satisfies the Law of Conservation of Mass. If students master how simple ratios work and how they can be worded in different ways, they can conquer stoichiometry easily.
  8. First law of thermodynamics - is also known as the Law of conservation of energy. In a chemical reaction, the energy change is equal to heat plus work done. E = q+w
  9. Heat of Reaction - Refers to the heat energy released or absorbed when the reaction is carried out using the exact amounts as shown by the chemical equation. The key point is that a balanced equation may show simplest whole numbers as coefficients, but the actual reaction may be scaled to different amounts and accordingly the heat of reaction will be scaled up or down. This is another aspect of understanding simple ratios and thus an extension of stoichiometry problems.
  10. Atomic structure - This refers to the arrangement of protons, neutrons, and electrons in the atom. In chemistry, the arrangement of electrons and the electronic configuration is very important and is inferred from the periodic table. Atomic number, which stands for the number of protons (and electrons) is given in the periodic table.
  11. Periodic Property (From Periodic table to Electron Configuration, atomic size, IE, Electron Affinity, valence electrons) - As the name suggests, various physical and chemical properties of elements follows a trend based on their relative position in the periodic table. Chemical properties repeat itself vertically (called the group of elements) while horizontally (called the period) the properties undergo a gradual change from metallic to nonmetallic. Atomic size increases along the group while decreases across period. Pulling an electron out of the atom (Ionization energy) becomes easier going down the group, but difficult across the period. Adding an electron to the atom (electron affinity) is easier across the period but not so favorable as you go down the group.
  12. Lewis Structure - shows valence electrons as dots around the atom. In a compound, a shared pair of electrons is shown as a solid line (called a bond) while the unshared pair is called as lone pair of electrons. When writing Lewis structures, the less electronegative atom is chosen as the central atom. In more complicated cases, the atom arrangement must be shown for a beginner to assemble the valence electrons. Octet rule which says that 8 electrons around an atom confer stability is the first step in explaining bonding. Elements from the 3rd row of the periodic table start exceeding the octet so that they can avoid formal charges whenever possible.
  13. VSEPR theory – Valence shell electron-pair repulsion theory is a very useful theory to explain molecular shapes. With a proper Lewis structure, the lone pair and bonded atoms are symmetrically arranged in 3D to minimize repulsions between the electron pairs. Lone pair-Lone pair repulsion is greater than Lone pair-Bonded pair repulsion, which in-turn is greater than bonded pair-bonded pair repulsion. Thus, small deviations happen from the symmetry depending on the type of electron-pair repulsions in the molecule. For example, a water molecule has 2 lone pairs and 2 bonded pairs on the central oxygen resulting in a tetrahedral arrangement. However, the actual angle between the two bonded pairs is less than the 109.5 observed for a symmetrical tetrahedral bond angle. The shape of molecule is given with respect to atoms only. In the water molecule, the H-O-H appears bent with an angle of about 107o.
  14. VB and MO theory – Valence Bond and Molecular Orbital theories explain the observed shapes of molecules from the mixing of atomic orbitals. In VB, valence atomic orbitals are hypothesized to mix and form hybrid orbitals which are based on the observed VSEPR geometry. In MO theory, all atomic orbitals can mix to form a lower energy and higher energy molecular orbital. Electrons fill these molecular orbitals in a fashion similar to atomic orbitals. Usually in freshman chemistry courses, the MO theory is not dealt in great detail because of time constraints. It is also a difficult topic for beginners. However, a good understanding of both VB and MO theory will help you succeed in organic chemistry.
  15. Bond length, Bond energy – Bond Length is a measure of the internuclear distance between atoms. It can be obtained experimentally and is the proxy for single bond, double bond, or triple bond in molecules. As the number of bonds increases, atoms become closer and it will show up in bond length. Bond Energy represents the energy required to separate the two atoms. Thus, closer atoms are harder to separate and will have higher bond energy values. In other words, a carbon-carbon triple bond will have the shortest bond length and the largest bond energy when compared to C=C or C-C.
  16. Bond polarity Molecular Polarity: Bond polarity represents the unequal sharing of electron pair(s) between two different atoms. A more electronegative atom draws the bonded electrons closer to itself creating a dipole. This bond dipole has a partial negative charge on the electronegative atom and a partial positive charge on the other bonded atom. The extent of this dipole can be linked to the separation of the two atoms on the periodic table. It can also be predicted more accurately if you have access to the Pauling’s electronegativity values. Since these bond dipoles have magnitude and direction, the entire molecule’s polarity is the net of these bond dipoles and hence depend on the actual geometry of all the bonds.
  17. Intramolecular and Intermolecular Forces – ‘Intra’ means within and ‘Inter’ means between. For covalent compounds, intramolecular forces are the bonds that hold the atoms together. This is the attraction of electrons of one atom to nucleus of the other atom. Intermolecular forces depend on the net polarity of the molecule. These are classified as ion-ion, ion-dipole, dipole-dipole, H-bond, and London forces. Ion-Ion is present in ionic compounds. Ion-dipole is present when ionic compounds interact with polar-covalent compounds. Dipole-dipole is present in polar-covalent compounds. These interactions help explain the physical properties like density, boiling point, melting point, etc. Larger intermolecular attraction leads to larger melting and boiling points. H-bond and London forces are special cases of dipole-dipole interaction given such unique names because of their importance. H-bond happens in compounds with OH, NH, or FH bonds. The dipole is so strong and the H-atom is so small, it creates an unusually strong intermolecular attraction to another O, N, or F atom. H-bonds are the basis of the liquid state of water and double-helical structure of DNA. London Forces are instantaneous dipole formations due to distortions in the electron cloud around the nucleus which becomes significant when the size of the molecule increases. For example, chlorine, bromine, and iodine have no bond dipoles. However, their physical state at room temperature change from gas to liquid to solid. This can be explained by the increasing size (molar mass) from chlorine to bromine to iodine causing larger induced dipole forces (London forces).
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