Atomic Structure, Atomic Mass Calculations, and Chemical Nomenclature

Exam Structure and Scope

• Exam 1 Content Scope:
  • Covers material from Lecture 2 through Lecture 7.
  • The only topic from Lecture 7 present on Exam 1 is atomic mass.
  • All subsequent topics covered in Lecture 7 beyond atomic mass will be tested on Exam 2.
• Exam 1 Format:
  • Consists of $20$ core questions.
  • Contains $1$ optional bonus question (Question 21).
  • Leaving the bonus question blank carries no penalty.
  • The bonus question can cover any content up through Monday of the following week, including unexamined topics.
• Question Types and Timing:
  • Core questions focus on elemental identification and fundamental subatomic particle counts (e.g., identifying element symbols like $\text{Li}$, defining protons, determining proton counts for elements like cobalt).
  • Completion time ranges from $15$\,minutes to an average of $25$\,minutes.

Subatomic Particles and Atomic Structure

• Protons:
  • Positively charged subatomic particles with a charge of $+1$.
  • Located in the nucleus of the atom.
  • The number of protons determines the atomic number ($Z$) of an element.
• Electrons:
  • Negatively charged subatomic particles with a charge of $-1$.
  • Located outside the nucleus within atomic orbitals.
  • In a neutral atom, the number of electrons equals the number of protons.
• Neutrons:
  • Electrically neutral subatomic particles with a charge of $0$.
  • Located inside the nucleus alongside protons.
  • Contribute to the mass and overall nuclear stability of the atom without affecting its charge.

Determining Subatomic Particles: Uranium Case Study

• Periodic Table Identification:
  • Uranium ($\text{U}$) is located at the bottom of the periodic table in the actinides series.
• Atomic Number ($Z$):
  • Represented by the top number on an element's periodic table entry square.
  • Directly equals the total number of protons in the nucleus.
  • For Uranium ($\text{U}$), the atomic number is $92$, meaning it contains $92\,\text{protons}$.
  • Protons contribute a total positive charge of +92$.\n* **Electron Count in Neutral Atoms**:\n * Neutral atoms contain no net positive or negative charge notation.\n * To balance a positive proton charge of +92$, a neutral Uranium atom must contain$92\,\text{electrons}$(each contributing$-1 charge).\n * Gaining or losing electrons transforms a neutral atom into an ion.\n* **Mass Number (A)**:\n * Represents the sum of protons and neutrons:\n    \text{Mass Number} = \text{Protons} + \text{Neutrons}\n * For Uranium-238 (^{238}\text{U}$), the mass number is$238$.
• Calculating Neutron Count:
  • Rearranging the mass number equation gives:     $\text{Neutrons} = \text{Mass Number} - \text{Protons}$
  • Calculation for Uranium-238:     $\text{Neutrons} = 238 - 92 = 146\,\text{neutrons}$
• Atomic Mass vs. Mass Number:
  • The bottom number on the Uranium periodic table square is 238.029$.\n * While close to the mass number (238), the bottom number is the **atomic mass**, which is a calculated weighted average of all naturally occurring isotopes of Uranium.\n * Uranium exists naturally as Uranium-238, Uranium-235, and approximately 11 other unstable, radioactive isotopes.\n\n# Calculating Weighted Atomic Mass\n\n* **Definition of Atomic Mass**:\n * A calculated, weighted average of the atomic masses of all naturally occurring isotopes of an element, incorporating fractional natural abundances.\n* **General Equation**:\n  \text{Atomic Mass} = (\text{Mass}{\text{isotope 1}} \times \text{Fraction}{\text{isotope 1}}) + (\text{Mass}{\text{isotope 2}} \times \text{Fraction}{\text{isotope 2}})\n* **Converting Percent Abundance to Fractional Abundance**:\n * Percent abundance values must be converted to decimals by dividing by 100 prior to multiplication:\n    \text{Fractional Abundance} = \frac{\text{Percent Abundance}}{100}\n* **Chlorine Isotope Calculation Example**:\n * **Isotope 1 (Chlorine-35)**:\n * Mass = 34.969\,\text{amu} (atomic mass units).\n * Percent Abundance = 75.78\%\n * Fractional Abundance = \frac{75.78}{100} = 0.7578\n * **Isotope 2 (Chlorine-37)**:\n * Mass = 36.966\,\text{amu}.\n * Percent Abundance = 24.22\%\n * Fractional Abundance = \frac{24.22}{100} = 0.2422\n * **Calculation Steps**:\n    \text{Contribution}{\text{Cl-35}} = 34.969\,\text{amu} \times 0.7578 = 26.4995082\,\text{amu}\n    \text{Contribution}{\text{Cl-37}} = 36.966\,\text{amu} \times 0.2422 = 8.9531652\,\text{amu}\n    \text{Calculated Atomic Mass} = 26.4995082\,\text{amu} + 8.9531652\,\text{amu} = 35.4526734\,\text{amu}\n* **Verification Against Periodic Table**:\n * Chlorine (\text{Cl}$) on the periodic table displays an atomic mass of$35.453\,\text{amu}.\n * The calculated value (35.4526734\,\text{amu}$, or rounded to$35.45) matches the value printed on the periodic table, serving as a check for calculation accuracy.\n* **Common Error Sources**:\n * Forgetting to divide percentages by 100$.
  • Mismatching/flip-flopping an isotope's mass with the wrong isotope's percent abundance.
  • Omitting required parentheses when entering calculation strings into a calculator.
• Practical Application:
  • Weighted average calculations follow the exact same mathematical procedure used to calculate final course grades based on weighted syllabus categories.

Covalent Compounds and Nomenclature

• Covalent Bonds:
  • Formed exclusively when two atoms share electrons.
  • Electron sharing occurs due to orbital overlap.
  • Formed only between nonmetal atoms.
• Location of Nonmetals:
  • Located primarily on the top-right portion of the periodic table.
  • Hydrogen ($\text{H}$) is the single major exception, located at the top left of the periodic table while functioning as a nonmetal.
• Rules for Naming Covalent Compounds:
  1. First Element:
  • State the full name of the element listed first in the formula.
  • Add a Greek prefix corresponding to its subscript count.
  • Exception: If the first element has a subscript of $1$ (only one atom), omit the prefix "mono-" entirely.
  1. Second Element:
  • Take the root name of the element and replace its suffix with "-ide".
  • Add a Greek prefix corresponding to its subscript count (always include the prefix, including "mono-" for single atoms).
  1. Combine: Join the modified terms into a two-word compound name.
• Greek Prefixes for Subscripts:
  • $1$: Mono-
  • $2$: Di-
  • $3$: Tri-
  • $4$: Tetra-
  • $6$: Hexa-
  • $7$: Hepta-
  • $9$: Nona- (or Nana-)
  • $10$: Deca-
• Covalent Naming Examples:
  • $\text{H}_2\text{O}$:
  • Common name: Water.
  • First element: Hydrogen with subscript $2$ $\rightarrow$ Dihydrogen.
  • Second element: Oxygen with subscript $1$ $\rightarrow$ modified to "-ide" (Oxide) with prefix "mono-" $\rightarrow$ Monoxide.
  • Systematic scientific name: Dihydrogen monoxide.
  • $\text{CO}_2$:
  • First element: Carbon with subscript $1$ $\rightarrow$ apply single-atom first-element exception $\rightarrow$ Carbon.
  • Second element: Oxygen with subscript $2$ $\rightarrow$ modified to "-ide" (Oxide) with prefix "di-" $\rightarrow$ Dioxide.
  • Systematic scientific name: Carbon dioxide.

Ionic Compounds and Nomenclature

• Ionic Bonds:
  • Formed when electrons are completely transferred from one atom to another, producing distinct positive and negative charges.
• Cations:
  • Positively charged ions ($+$).
  • Formed when an atom loses one or more electrons.
  • Mnemonic: The letter 't' in "cation" resembles a positive sign ($+$).
  • By scientific convention, the cation is always written first in chemical formulas and names.
• Anions:
  • Negatively charged ions ($-$).
  • Formed when an atom gains one or more electrons.
  • By scientific convention, the anion is always written second in chemical formulas and names.
• Rules for Naming Ionic Compounds:
  1. Name the cation (first element) directly by its elemental name.
  2. Name the anion (second element) by taking its elemental root name and changing the suffix to "-ide".
  3. Do NOT use any prefixes (e.g., no "mono-", "di-", "tri-").
• Ionic Naming Examples:
  • $\text{NaCl}$:
  • Cation: Sodium ($\text{Na}$).
  • Anion: Chlorine ($\text{Cl}$) $\rightarrow$ modified to Chloride.
  • Systematic name: Sodium chloride (common table salt).
  • $\text{K}_2\text{O}$:
  • Cation: Potassium ($\text{K}$).
  • Anion: Oxygen ($\text{O}$) $\rightarrow$ modified to Oxide.
  • Systematic name: Potassium oxide (Note: no prefix used for the two potassium atoms).

Periodic Table Trends for Ionic Charges

• Group 1 (Column 1) - Alkali Metals:
  • Located in the far left column of the periodic table.
  • Share similar chemical behaviors (e.g., reactive with water and oxygen in identical manners).
  • Consistently form ions with a $+1$ charge.
• Group 2 (Column 2) - Alkaline Earth Metals:
  • Located adjacent to alkali metals in column 2.
  • Consistently form ions with a $+2$ charge.
• Group 17 (Column 17) - Halogens:
  • Located in column 17 on the right side of the periodic table.
  • Consistently form ions with a $-1$ charge.
• Group 16 (Column 16):
  • Consistently form ions with a $-2$ charge.
• Charge Directionality Trend:
  • Ionic charge becomes increasingly negative moving from right to left across the nonmetal columns.
• Specific Metal Ion Exception:
  • Aluminum ($\text{Al}$) forms an ion with a $+3$ charge.

Questions & Discussion

• Question: What is the format and length of Exam 1?
  • Answer: Exam 1 officially consists of 20 core questions, plus 1 optional bonus question (Question 21). Leaving the bonus blank does not harm the grade. It takes students between 15 and 25 minutes on average.
• Question: How is the isotopic abundance of chlorine across the universe measured?
  • Answer: It involves complex spectroscopic instruments, space probes, and mathematical physics estimations.
• Question: Why calculate atomic mass manually when the value is already on the periodic table?
  • Answer: The exercise establishes mastery of calculating weighted averages, a core analytical skill used broadly in science and for determining overall course grades.