Chemistry: Isotopes, Atomic Structure, and Bonding - Practice Flashcards

Course Announcements

• Lecture – PowerPoint slides should be watched in slide-show mode to view animations and read purple pop-up text boxes for clarifications. Online resources available in the chapter module in D2L.
• “Lecture Slide” Quizzes are NOT for credit; they are for personal knowledge checks.
• ALEKS – practice material for mastery learning; fill in the pie; assignments due Wednesdays at 11:59pm CST.
• Group Learning – on D2L in Assessment-Assignments; due Fridays at 11:59pm CST; check the schedule.
• Lab – not meeting in person; virtual and/or dry labs; check D2L Content-Materials-Laboratory Materials-submodule per lab.
• Pre-lab Quiz on D2L in Assessment-Quizzes; due the week of the lab on Tuesdays at 11:59pm CST; 2 attempts.
• Laboratory Assignment on D2L in Assessment-Assignments; due the following week on Mondays at 11:59pm CST; must be completed to release Post-lab quiz.
• Post-lab Quiz on D2L in Assessment-Quizzes; due the following week on Thursdays at 11:59pm CST; 1 attempt.
• Office hours – set up via Zoom; check D2L for invites.
• Visit the Course D2L page for announcements, syllabus, schedule, lecture notes, links to videos, pre-/post-lab quizzes, group learning assignment, drop box, extra credit assignments, grades, discussion board, classlist, announcements.

Atomic Structure: Protons, Neutrons, and Electrons

• Most of the atom is empty space; mass is concentrated in the nucleus.
• Mass of the atom is primarily in the nucleus.
• Proton charge is opposite in sign but equal in magnitude to the electron charge.
• Protons and neutrons reside in the nucleus; electrons orbit around it.

Atomic Masses and the Mole

• The mass of 1 atom of carbon-12 is defined to be 12 amu.
• Atomic mass A is the total number of protons and neutrons: A = Np + Nn.
• Atomic number Z equals the number of protons: Z = N_p.
• In a neutral atom, the number of electrons equals the number of protons: N_e = Z.
• Note: A and Z are integers (non-decimal); atomic masses are decimals (amu).

Atomic Number, Mass Number, and Isotopes

• Isotopes have the same atomic number Z but different mass numbers A.
• This means they have the same number of protons but different numbers of neutrons.
• A convenient format uses the element symbol with Z and A noted; for example, hydrogen isotopes illustrate different A values.

Protons, Neutrons, and Electrons: Quick Determination

• Given species with Z and A, determine:
  • Protons: Z
  • Neutrons: N = A - Z
  • Electrons: e = Z (for a neutral species)
• Example (from slide): If Z = 17 and A = 35, then
  • Protons = 17
  • Neutrons = 35 - 17 = 18
  • Electrons = 17

Atomic Mass Scale and Mass Spectrometry

• Atomic mass unit (amu) is defined as 1/12 the mass of a carbon-12 atom:
  1 ext{ amu} = rac{1}{12} m(^{12} ext{C}).
• Mass spectrometry is used to determine atomic masses precisely.

The Periodic Table

• Horizontal rows are periods; vertical columns are groups (families).
• The periodic table is a practical reference (sometimes called a “cheat sheet”), so learn to use it.
• Chemical formula notation indicates the composition and can reveal metal–nonmetal, nonmetal–nonmetal, or metal–metal combinations.

Isotope Abundance, Isotope Mass, and Atomic Weight

• Isotope data table typically lists: Isotope, Symbol, Atomic Mass, Mass Number, Isotope Mass, Natural Abundance %.
• Example observations: Oxygen has three stable isotopes with varying abundances: ^16O, ^17O, ^18O with abundances ~99.757%, 0.038%, and 0.205% respectively (numbers may vary slightly by source).
• How abundance relates to atomic weight: greater natural abundance of a heavier isotope raises the average atomic weight of the element.

Percent Abundance per Isotope

• For isotopes of an element, percent abundance is calculated as:
  ext{Percent abundance} = rac{ ext{number of atoms of an individual isotope}}{ ext{total number of atoms of all isotopes of the element}} imes 100.
• This provides the fractional contribution of each isotope to the element’s overall presence.
• Self-check: the sum of all isotope abundances must equal 100%.

Atomic Weight (Average Atomic Mass)

• Atomic weight (the decimal on the periodic table) equals the average mass of all naturally occurring isotopes of the element.
• It accounts for the relative abundances of isotopes:
  ar{m} = rac{ ext{sum of (fractional abundance × isotope mass)}}{1} = ar{m} =
  igg(rac{ ext{abundance}1}{100}igg) m1 + igg(rac{ ext{abundance}2}{100}igg) m2 + \, ext{…}
• Alternate compact formula:
  ar{A} = rac{ ext{sum of (abundance × mass)}}{100} where abundances are given as percentages.
• Example: Oxygen isotopes with masses 15.9949, 16.9991, 17.9992 amu and abundances 99.757%, 0.038%, 0.205% yield an average atomic mass very close to 16.000 amu (approximately 15.999 to four significant figures).

Isotope Example Calculation (Oxygen)

• Given: ^{16}O mass = 15.9949 amu (99.757%), ^{17}O mass = 16.9991 amu (0.038%), ^{18}O mass = 17.9992 amu (0.205%).
• Calculation steps:
  • Fractional abundances: 0.99757, 0.00038, 0.00205
  • Weighted masses: 0.99757 imes 15.9949 \,+
    0.00038 imes 16.9991 \,+
    0.00205 imes 17.9992
  • Sum ≈ 15.9994 amu
• To four significant figures: 16.00 amu (matching standard periodic table values around 15.999).

Two Stable Isotopes and Abundance (Concept Check)

• If the atomic mass is closer to the lighter isotope’s mass, the lighter isotope is more abundant.
• Given: two isotopic masses 10.0129 amu and 11.0093 amu with an average atomic mass of 10.81 amu, the lighter isotope (10.0129 amu) is more abundant.
• Answer: 3. Isotope with mass 10.0129 amu is more abundant.

Polyatomic Compounds and Ions

• Polyatomic Ions consist of two or more atoms covalently bonded and possessing an overall ionic charge; memorize common polyatomic ions.
• Quiz on polyatomic ions may be administered in Lab.

Oxoanions: -ate and -ite Endings

• Two forms exist for oxoanions:
  • Higher oxygen content uses the suffix -ate.
  • Lower oxygen content uses the suffix -ite.
• Memorize: More O means -ate; Less O means -ite.
• Examples:
  • Sulfate: ext{SO}_4^{2-}
  • Sulfite: ext{SO}_3^{2-}
  • Nitrate: ext{NO}_3^{-}
  • Nitrite: ext{NO}_2^{-}
• If the anion contains hydrogen, prefix with hydrogen (e.g., hydrogen sulfate: ext{HSO}4^{-} a.k.a. bisulfate; hydrogen carbonate: ext{HCO}3^{-} a.k.a. bicarbonate).

Oxoanions Containing Four Forms

• When four forms exist, add prefixes and suffixes: add "per_ate" and "hypoite" forms to indicate very high or very low oxygen content.
• Memory aid presented (though specific examples in transcript):
  • Per____ate, fluorate, fluorite, hypofluorite (illustrative; follow actual chemistry nomenclature in formal contexts).

Naming Ionic Compounds

• Binary Ionic Compounds contain two different elements (one metal, one nonmetal).
• Naming rule:
  • Name the element that appears first in the formula (usually the metal).
  • Name the second element with its root + ending -ide (for monatomic anions).
  • Example patterns: HCl → hydrogen chloride; NaCl → sodium chloride; SiC → silicon carbide (the -ide indicates two elements total).
• Key hint: -ide usually signals a binary compound with a metal and a nonmetal.

Determining Charges on the Elements to Make Ions

• Charge correlates with group position relative to a noble gas; determine charge by how many electrons are gained or lost to attain a noble gas configuration.
• Examples:
  • Li (group 1) loses 1 electron to form Li⁺; nearest noble gas is He.
  • Oxygen (group 16) tends to gain 2 electrons to form O²⁻ to reach Ne.
• The goal is to minimize the number of electrons gained or lost to achieve a noble gas configuration; the resulting charges balance in ionic compounds.

Transition Metals Ion Charges

• Transition metals can have multiple oxidation states; you cannot assign a fixed charge from the group number alone.
• The actual charge of a transition metal cation in a binary ionic compound is determined by the charge of the accompanying anion.

Determining Ionic Compound Formulas (Cross-and-Drop)

• Ionic compounds are neutral overall; use cross-and-drop to balance charges.
• Example: Aluminum oxide.
  • Al³⁺ and O²⁻ combine so that the total charges cancel: 2 Al³⁺ (total +6) with 3 O²⁻ (total -6) → Al₂O₃.
• Calcium phosphate example:
  • PO₄³⁻ (3−) and Ca²⁺ (2+) require a combination that balances to zero; Ca₃(PO₄)₂ is the balanced formula.

Naming Ionic Compounds with Transition Metals

• If a transition metal cation can have more than one charge, use Roman numerals in parentheses to indicate the charge.
  • Example: Fe can be Fe²⁺ or Fe³⁺; Cu can be Cu⁺ or Cu²⁺.
• For cations with two possible charges, the suffix -ous denotes the smaller positive charge and -ic denotes the larger positive charge (e.g., Cu⁺ = cuprous, Cu²⁺ = cupric).
• Mn, Fe, Cu, and other transition metals commonly display multiple oxidation states; use the anion to deduce the metal’s charge.

Flow Chart for Naming Ions (McGraw-Hill Connect)

• Ionic naming involves two branches:
  • Cation with a single possible charge: name metal first; anion name follows; add -ide if monatomic.
  • Cation with multiple possible charges: name metal first; specify charge with Roman numeral; use monatomic/polyatomic anion name as appropriate.
• This chart helps distinguish between Binary Ionic Compounds vs. those with transition metals.

Which Formula is Represented by the Correct Formula and Chemical Name?

• Example choices:
  1) Iron(III) nitrate; Fe(NO₃)₃
  2) Calcium hydroxide; CaOH
  3) Lithium nitride; LiN
  4) Potassium sulfate; KSO₄
• Correct answer: 1. Iron(III) nitrate is Fe(NO₃)₃. The others are unbalanced or incorrect (Ca(OH)₂, Li₃N, K₂SO₄).

Hydrated Compounds

• Hydrated ionic compounds contain a defined amount of water trapped within the solid lattice.
• Water of hydration is the water associated with the compound.
• The hydrated formula is written as MX·nH₂O, where n is the number of water molecules per formula unit.

Molecules and Molecular Compounds

• A molecule is a combination of at least two non-metal atoms in a specific arrangement and ratio held together by covalent bonds.
• Homonuclear diatomic molecules: two atoms of the same element in a molecule (e.g., H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂).
• Heteronuclear diatomic molecules contain atoms of different elements (e.g., CO, NO).
• Prefixes: di, tri, tetra, etc. Prefixes indicate the number of atoms in molecular compounds; poly- means many.

The Seven Diatomic Molecules

• Elements that naturally exist as diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂.
• Mnemonic and visual patterns can help remember the seven diatomic molecules.

Prefixes Used in Naming Binary Nonmetal (Molecular) Compounds

• Numerical prefixes (used only for molecular compounds):
  • 1: mono-, 2: di-, 3: tri-, 4: tetra-, 5: penta-, 6: hexa-, 7: hepta-, 8: octa-, 9: nona-, 10: deca-, 12: dodeca-.
• Important note: The prefix mono- is generally omitted for the first element except in carbon monoxide (CO).

Naming Binary Nonmetal Molecular Compounds: Examples

• Water: H₂O
• Sulfur dioxide: SO₂
• Hydrogen peroxide: H₂O₂
• Sulfur trioxide: SO₃
• Ammonia: NH₃
• Carbon monoxide: CO
• Nitric oxide: NO
• Carbon dioxide: CO₂
• Nitrogen dioxide: NO₂
• Hydrazine: N₂H₄
• Disulfur decafluoride: S₂F₁₀
• Note on naming: Use prefixes to indicate the number of each element; omit prefix for the first element when appropriate; retain the ending of the second element as -ide when applicable.

Predict the Correct Name of a Molecule (Diagram-Based Question)

• Given a diagram where nitrogen (N) is the central atom and oxygen (O) are surrounding, the correct name for the depicted molecule is:
  • 3) Dinitrogen oxide (dinitrogen monoxide) [often written as N₂O].

Empirical vs Molecular Formulas

• Empirical formula: the simplest whole-number ratio of elements in a compound.
• Molecular formula: the actual number of each type of atom in a molecule; may be a multiple of the empirical formula.
• The integer n represents how many empirical units are in a molecule.
• If a compound has the empirical formula CH₂O and the molecular formula is C₆H₁₂O₆, then n = 6 (molecular formula = (CH₂O)₆).

Examples: Empirical vs Molecular Formulas

• Empirical formula: smallest whole-number ratio of atoms (n may be 1, 2, 6, etc.).
• Molecular formula: can be a multiple of the empirical formula to reflect the actual number of atoms in a molecule.

Which Polyatomic Formula Is Not Correct?

• Given examples (identify not correct):
  • Mg(NO₃)₃ (not correct due to charge balance; nitrate is NO₃⁻, total 3×−1 = −3; Mg²⁺ cannot balance −3).
  • NH₄CO₃ (not correct; should be (NH₄)₂CO₃ to balance charges).
  • NaClO₄ (correct; chlorate with Na⁺ forms NaClO₄).
  • Al₂(SO₄)₃ (correct; aluminum sulfate with Al³⁺ and SO₄²⁻).

Which Kind of Change is Shown?

• Question: What kind of change is shown in the accompanying diagram or scenario?
  • Options: 1) Chemical Change 2) Physical Change 3) Both Chemical and Physical Change 4) No change
• The transcript does not provide the diagram or explicit answer; use context from the example in class to determine the change type.

Connections to Foundational Principles and Real-World Relevance

• Isotopes and atomic weight tie into how mass is measured in labs and in natural samples; mass spectrometry is a key analytical technique.
• Periodic table organization reflects recurring chemical properties (periodicity) and informs predictions about reactivity and bonding.
• Ionic vs molecular compounds highlight how electron transfer vs sharing governs bonding, properties, and naming conventions used in chemistry, biology, medicine, and environmental science.
• Understanding polyatomic ions and oxyanion patterns (ate/ite) helps in balancing reactions and writing correct formulas across inorganic chemistry and biochemistry.

Key Formulas and Notation Recap (LaTeX)

• Atomic number and composition:

  • Protons: Z = N_p
  • Neutrons: N = A - Z
  • Electrons (neutral): N_e = Z

• Isotope abundance and atomic weight:

  • ar{m} = rac{ ext{sum of (abundance}i imes ext{mass}i)}{100}
  • Or, with fractional abundances: ar{m} = rac{ ext{sum of (f}i mi)}{1} where fi = rac{ ext{abundance}i}{100}

• Atomic mass unit: 1 ext{ amu} = frac{1}{12} m(^{12} ext{C})

• Cross-and-Drop balancing for ionic compounds (example):

  • For Al₂O₃: total positive = total negative; 2 imes (+3) + 3 imes (-2) = 0

• Diatomic molecules (seven): ext{H}2, ext{N}2, ext{O}2, ext{F}2, ext{Cl}2, ext{Br}2, ext{I}_2

• Naming prefixes for binary molecular compounds: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-

• Oxoanions suffix rules:

  • Higher oxygen: -ate; lower oxygen: -ite; more oxygen: -ate; hydrogen-containing forms use hydrogen- prefix (e.g.,

  • hydrogen sulfate: ext{HSO}4^{-}; hydrogen carbonate: ext{HCO}3^{-})

End of Notes