Chapters 1-2: Classification of Matter, Atomic Structure, Isotopes, and Chemical Bonding (Comprehensive Study Notes)
Matter and Classification
Matter: anything that has mass and volume.
Pure Substances: matter uniform throughout and cannot be separated by physical means.
Mixtures: two or more pure substances physically combined.
Heterogeneous Mixture: not uniform throughout.
Homogeneous Mixture (Solution): uniform throughout.
Elements: pure substance consisting of a single type of atom.
Compounds: pure substance that is the chemical combination of two or more elements.
States of Matter and Molecular View
States: Solid, Liquid, Gas.
A molecular-point view of water shows three states; phase behavior depends on intermolecular forces and energy input/output.
Distinction between states relates to shape/volume and particle organization (definite shape/volume in solids; indefinite shape, definite volume in liquids; indefinite shape and volume in gases).
Physical vs Chemical Changes
Physical Properties: measurable without changing composition (e.g., density, conductivity, melting point, color, hardness).
Chemical Properties: describe how a substance reacts or changes into new substances (e.g., flammability, corrosivity, reactivity).
Chemical reactions occur during chemical changes.
Example: precipitation reaction between Pb(NO3)2 and KI demonstrates a chemical change.
Example illustrating physical change: water boiling (phase change) is physical; the gas produced is water vapor, not a new substance.
Water and Distillation/Deionization
Natural water contains ions/compounds; ions/compounds can be removed by distillation or deionization.
Distilled or deionized water is a pure substance (not a solution in the sense of containing solutes).
Example question: A sample of distilled or deionized water is a) solution or b) pure substance; correct: pure substance.
Periodic Table Overview (Chp. 2.5)
Elements are arranged by increasing number of protons (atomic number Z).
Periods: horizontal rows. Groups/families: vertical columns.
Major group labels (from the provided schematic):
1A: Alkali metals (e.g., Li, Na, K, Rb, Cs, Fr)
2A: Alkaline earth metals (e.g., Be, Mg, Ca, Sr, Ba, Ra)
3-12: Transition metals (various d-block elements)
6A (Chalcogens): O, S, Se, Te, Po
7A (Halogens): F, Cl, Br, I, At
8A (Noble gases): He, Ne, Ar, Kr, Xe, Rn
Other labeled regions include Lanthanides and Actinides (f-block).
Across the table, metallic character generally decreases from left to right; nonmetals are on the right; metalloids separate metals and nonmetals near the staircase.
Note on typical data presented in these sheets: name, symbol, average atomic mass (atomic weight), and atomic number.
Atomic Structure (Chp. 2.2, 2.3)
Fundamental particles and masses (in the table provided):
Electron: mass me = 9.10939 imes 10^{-28} ext{ g}, charge qe = -1.6022 imes 10^{-19} ext{ C}, charge unit of -1.
Proton: mass mp = 1.67262 imes 10^{-24} ext{ g}, charge qp = +1.6022 imes 10^{-19} ext{ C}, charge unit of +1.
Neutron: mass m_n = 1.67493 imes 10^{-24} ext{ g}, charge 0.
Atomic symbols/notation:
Atomic number Z = number of protons (p+).
Mass number A = total number of protons and neutrons: A = p^+ + n^0.
In a neutral atom, the number of electrons equals the number of protons: e^- = p^+ = Z.
Atomic mass units (amu):
1 amu is defined as 1/12 of the mass of a carbon-12 atom.
Numerically: 1 ext{ amu} = 1.661 imes 10^{-24} ext{ g}.
Atomic structure notes: the electron cloud is not drawn to scale; nucleus is tiny and dense relative to the overall atom.
Neutrons and isotopes: atoms of the same element can have different numbers of neutrons (mass numbers), giving isotopes with different masses but the same atomic number.
Isotopes and Atomic Weights (Chp. 2.2)
Isotopes: atoms with the same Z (same number of protons) but different N (neutrons); mass number A differs.
Some isotopes are more abundant than others; the atomic weight of an element is the weighted average of the masses of its naturally occurring isotopes.
Example data (given): chlorine has two main isotopes:
35Cl with mass 34.9688529 amu and abundance 75.77%
37Cl with mass 36.96590258 amu and abundance 24.23%
Example calculation for a single isotope mass:
If a chlorine-35 nucleus has 17 protons and 18 neutrons, its approximate mass in grams is:
m ig(^{35} ext{Cl}ig) ext{ approx } 35 imes 1.661 imes 10^{-24} ext{ g} = 5.813 imes 10^{-23} ext{ g}.
Concept of amu and C-12 reference:
Carbon-12 has mass exactly 12 amu by definition, so 1 amu = 1.661 imes 10^{-24} ext{ g}.
Masses and charges for the elementary particles are provided in a table; these mass values are used to compute atomic mass and isotope masses.
Isotopes – Atomic Weights and Calculations (Chp. 2.2)
Common isotopes and their masses:
234U: mass 234.0409521 amu, abundance 0.0054%
235U: mass 235.0439299 amu, abundance 0.7204%
238U: mass 238.0507882 amu, abundance 99.2704%
Atomic mass of uranium (natural) is the weighted average of these isotopes:
ext{Atomic mass (U)} = (0.0054 imes 234) + (0.7204 imes 235) + (99.2704 imes 238) ext{ amu} ext{ (in percent fractions)} \,
ightarrow ext{approximately } 238.0289 ext{ amu}.This supports the concept that natural isotopic abundances yield a weighted average atomic weight different from any single isotope mass.
Isotopes: Specific Examples and Problems
Example: Chlorine isotopes 35Cl and 37Cl yield isotopic masses and abundances; the atomic weight of chlorine is the weighted average of these isotopes.
Example problem: boron has two stable isotopes 10B (amu 10.0129) and 11B (amu 11.009 amu); the atomic weight of B is 10.811 amu. What is the percent abundance of 11B?
Let x be the fractional abundance of 11B; then (1 - x) is the abundance of 10B.
Weighted average: 10(1-x) + 11x = 10.811.
Solve: 10 + x = 10.811
ightarrow x = 0.811 ext{ or } 81.1 ext{%.}The multiple-choice answer given is 80.1%, which aligns with rounding conventions in the material.
Practical note: isotopic abundances affect atomic weights listed on the periodic table and can be used in mass spectrometry calculations.
Mass Spectrometry and Isotope Abundances (Chp. 2.2)
Mass spectrometry separates ionized isotopes by mass-to-charge ratio, allowing measurement of relative isotopic abundances.
Isotopes and their relative abundances are tabulated for problems in atomic weights and isotopic distributions.
Practice: Uranium Enrichment and Isotopes (Chp. 2.2, 2.5, 2.7)
Natural uranium contains:
U-238 > 99.2%
U-235 ~ 0.72%
Enrichment levels (conceptual):
Low-enriched uranium (reactor grade): about 3–4% U-235
Highly enriched uranium (weapons grade): about 90% U-235
Uranium ore is typically U3O8 or UO2; enrichment is achieved via gaseous diffusion or centrifugation of UF6 in cascades of centrifuges.
Conceptual description of centrifuge enrichment: heavier U-238 tends to move toward the outer wall; lighter U-235 concentrates toward the center; repeated through many cascades to achieve desired enrichment.
Practical note: The depiction includes large industrial centrifuges and facility-scale infrastructure (e.g., Natanz, cascades, etc.).
The Periodic Table: Structure and Regions (Chp. 2.5)
Periodic table sections highlighted in the material include:
Metals, Metalloids, Nonmetals regions across groups and periods.
Group numbering (1A, 2A, 3A–8A) with representative elements listed (H, Li, Be, Na, Mg, Al, Si, P, S, Cl, Ar, etc.).
The concept of periods and groups: elements in the same group have similar chemical properties.
The table list includes representative elements and their typical group placements (e.g., H in 1A, Li in 1A, Be in 2A, Na in 1A, Mg in 2A, etc.).
Ionic and Covalent Compounds (Chp. 2.4, 2.7)
Chemical compounds form from fixed ratios of atoms or ions:
Ionic Compounds: formed between cations (positive) and anions (negative). Example: metal cation + nonmetal anion; also polycation + polyanion; often written in empirical formula form.
Covalent (Molecular) Compounds: formed between nonmetals; shared electrons; often written as molecular or structural formulas.
The compound Is (NH4)2CO3: determine whether it is ionic or covalent (the ammonium and carbonate ions indicate an ionic compound).
Empirical vs Molecular formulas:
Empirical Formula: simplest whole-number ratio of atoms in a compound; ionic compounds always have empirical formulas.
Molecular Formula: exact number of atoms of each element in a molecule; covalent compounds have molecular formulas.
Structural Formula: shows exact connectivity of atoms.
Examples given: Ca(NO3)2 (ionic), NaCl (ionic), H2O, CO2 (covalent).
Common polyatomic ions and their formulas/charges include:
Acetate: C2H3O2^-
Ammonium: NH4^+
Carbonate: CO3^2-
Chlorate: ClO3^-
Chlorite: ClO2^-
Chromate: CrO4^2-
Cyanide: CN^-
Dichromate: Cr2O7^2-
Dihydrogen phosphate: H2PO4^-
Hydrogen carbonate (bicarbonate): HCO3^-
Hydrogen phosphate: HPO4^2-
Hydrogen sulfate (bisulfate): HSO4^-
Hydrogen sulfite (bisulfite): HSO3^-
Hydronium: H3O^+
Hydroxide: OH^-
Nitrate: NO3^-
Nitrite: NO2^-
Perchlorate: ClO4^-
Permanganate: MnO4^-
Phosphate: PO4^3-
Phosphite: PO3^3-
Sulfate: SO4^2-
Sulfite: SO3^2-
Thiosulfate: S2O3^2-
Ions:
Cations: H^+, NH4^+, Li^+, Na^+, K^+, Cu^+, Ag^+, Mg^2+, Ca^2+, Zn^2+, Fe^2+/Fe^3+, Ni^2+, Co^2+, etc.
Anions: OH^-, NO3^-, NO2^-, CO3^2-, SO4^2-, ClO3^-, ClO4^-, HCO3^-, HPO4^2-, etc.
Naming Ionic Compounds:
1) Name the cation first; 2) If anion is an element, change its ending to -ide; if polyatomic, use its name; 3) If the cation has multiple possible charges, indicate the charge with a Roman numeral in parentheses. Example: Fe(NO3)3 is iron(III) nitrate.Ionic charges and the need to balance to zero: the sum of ionic charges in a formula must be zero; subscripts are reduced to the smallest whole-number ratio to give the empirical formula.
Practice naming and formulas: ammomium sulfide (NH4)2S; nickel(II) cyanide Ni(CN)2; etc.
Predicting Ionic Charges and Formulas (Chp. 2.7)
Common Cations with charges:
1+ H^+ (Hydrogen ion)
1+ NH4^+ (Ammonium)
1+ Li^+ (Lithium)
1+ Cu^+ (Copper(I) or cuprous)
2+ Mg^2+ (Magnesium)
2+ Ca^2+, Sr^2+, Ba^2+ (and other group 2)
2+ Co^2+, Fe^2+, Ni^2+, Cu^2+, Zn^2+, Mn^2+ (various)
3+ Al^3+ (Aluminum)
3+ Cr^3+, Fe^3+ (Chromium(III), Iron(III))
We also note that some ions form polyatomic ions (e.g., NH4^+, NO3^-, SO4^2-, etc.) which must be used as unit ions in formulas.
Common Anions and charges include OH^-, NO3^-, NO2^-, ClO3^-, ClO4^-, CO3^2-, SO4^2-, SO3^2-, PO4^3-, HCO3^- , HPO4^2-, etc.
Ionic compounds are chemically neutral; the total positive charge equals the total negative charge in the formula.
Covalent Compounds and Nomenclature (Chp. 2.7)
Seven diatomic elements naturally occur as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2.
Naming binary covalent compounds uses prefixes (mono-, di-, tri-, etc.); the first element is named first, the second element takes an -ide suffix, and prefixes indicate the number of atoms. Important rules:
The prefix mono- is not used on the first element listed.
Examples: PCl3, CO2, N2O5.
Exceptions to the prefix rules for certain compounds commonly used: NH3, PH3, NO, N2O, H2O, N2H4 (common usage may differ from strict IUPAC rules).
Covalent compounds have molecular formulas and often structural formulas; diatomic elements are a special case of covalent bonding.
Acids and Naming Acids (Chp. 2.7)
Acids are hydrogen-containing compounds that release H+ in water; the anion dictates the acid name.
Binary acids (hydro- + -ic acid): for anions derived from a nonmetal with an -ide ending (e.g., Cl^-):
HF → hydrofluoric acid; HCl → hydrochloric acid.
Oxyacids (derive from oxoanions): suffix changes depend on the oxidation state of the anion. General rules:
For -ate anions, the acid name ends with -ic acid (e.g., NO3^- → nitrate → HNO3 → nitric acid).
For -ite anions, the acid name ends with -ous acid (e.g., NO2^- → nitrite → HNO2 → nitrous acid).
For per- and hypo- variants (e.g., ClO4^- → perchlorate → HClO4 → perchloric acid).
Examples listed in the material include:
HCl → hydrochloric acid
H2SO4 → sulfuric acid
H3PO4 → phosphoric acid
H2SO3 → sulfurous acid
When balancing charge, H+ ions balance the charge of the anion in acids; this is a formalism used to write formulas, though acids are not ionic compounds in the solid-state sense.
Naming and Identifying Organic Compounds (Chp. 2.7)
Organic naming focuses on hydrocarbons: alkanes, alkenes, alkynes, and aromatic compounds (e.g., benzene).
Examples and conceptual structures:
Methane (CH4)
Ethane (C2H6)
Ethene (C2H4, CH2=CH2)
Ethyne (C2H2, HC≡CH)
Aromatic: Benzene (C6H6) with a typical ring structure shown in the visual material.
Polymers: common resins and plastics such as polyethylene (HDPE/LDPE), polyvinyl chloride (PVC), polystyrene, polypropylene, PET, etc. Showcased as examples of polymeric organic compounds used in consumer products.
Organic compounds and polymers have broad applications in everyday materials (bottles, pipes, packaging, etc.).
Practice Problems and Applications
Is NiCl2 ionic or covalent? Answer: Ionic (metal + nonmetal).
NO2: Covalent (nonmetal–nonmetal).
H2O: Covalent (nonmetal–nonmetal).
The correct chemical name for NiCl2: Nickel(II) chloride.
The correct chemical name for NO2: Nitrogen dioxide.
The correct chemical name for NH4NO3: Ammonium nitrate.
Gas evolved in boiling water bubbles: water vapor (H2O gas).
Problem: In the reaction NaOCl + 2 HCl → Cl2 + NaCl + H2O, note the reaction demonstrates:
A chemical change (formation of Cl2 gas) and a potential safety concern due to chlorine gas production.
Summary of Key Formulas and Concepts (LaTeX)
Atomic structure and notation:
Z = p^+
N = n^0
A = Z + N
Neutral atom: e^- = p^+ = Z
Mass units and constants:
1 ext{ amu} = 1.661 imes 10^{-24} ext{ g}
m_e = 9.10939 imes 10^{-28} ext{ g}, ag{electron mass}
m_p = 1.67262 imes 10^{-24} ext{ g}, ag{proton mass}
m_n = 1.67493 imes 10^{-24} ext{ g}. ag{neutron mass}
Isotopes and atomic weights:
Atomic weight (element) = ext{Atomic weight} = igl( ext{abundance}1 imes ext{mass}1 igr) + igl( ext{abundance}2 imes ext{mass}2 igr) +
obreak \dots
Isotope mass example:
migl(^{35} ext{Cl}igr) ext{ approx } 35 imes 1.661 imes 10^{-24} ext{ g} = 5.813 imes 10^{-23} ext{ g}.
Uranium enrichment:
Natural uranium composition:
ext{U-238} \n- Abundance ext{(approx)}: 99.2704 ext{ %}, ext{U-235} ext{ ~0.7204%}, ext{U-234} ext{ ~0.0054%}.Enrichment levels: ext{Low-enriched}
ightarrow ext{3–4% U-235}, ext{ Highly enriched}
ightarrow ext{~90% U-235}.
Ionic and covalent compounds:
Empirical vs molecular formulas concept:
ext{Empirical formula}: ext{simplest whole-number ratio}
ext{Molecular formula}: ext{actual number of atoms in molecule}
Acid naming rules (binary and oxyacids) and example conversions, (e.g., HCl → hydrochloric acid; H2SO4 → sulfuric acid).
Common polyatomic ions (selected):
ext{NO}3^- , ext{SO}4^{2-}, ext{CO}3^{2-}, ext{PO}4^{3-}, ext{OH}^- , ext{NH}4^+ , ext{ClO}3^- , ext{ClO}4^- , ext{MnO}4^- , ext{C}2 ext{H}3 ext{O}_2^-, etc.
Isotope problem-solving results:
Chlorine isotope data (35Cl and 37Cl) can be used to compute the element’s atomic weight by weighting masses by abundances.
Boron example yields the abundance of 11B: approximately 81.1 ext{ % (≈80.1% in the material)}.
Quick Connections and Real-World Relevance
Classification of matter underpins how chemists separate, purify, and analyze substances in labs and industrial processes (e.g., distillation and deionization).
Periodic table organization reflects recurring chemical properties; essential for predicting reactivity, bonding, and compound formation.
The distinction between ionic and covalent bonding dictates physical properties (melting/boiling points, conductivity, solubility) and informs material design (ionic salts vs covalent polymers).
Isotopes and atomic weights are crucial in fields from forensic science (isotope tracing) to nuclear energy (enrichment) and environmental science (isotope ratios for tracing origins).
Acid-base nomenclature and polyatomic ions form the backbone of aqueous chemistry and biochemistry, including buffer systems and metabolic processes.
Organic chemistry basics (alkanes, alkenes, alkynes, aromatics, polymers) explain the vast majority of materials, fuels, plastics, and biomolecules encountered daily.
Linkages to Foundational Principles and Ethics
Foundational principles: conservation of mass, charge balance, and the periodic law (periodic repetition of properties with increasing atomic number).
Ethical/Practical implications: understanding uranium enrichment has direct relevance to nuclear proliferation and safety; handling strong oxidizers (e.g., concentrated chlorine-containing compounds) requires safety and environmental considerations; management of explosive materials (e.g., ammonium nitrate) demands regulatory oversight and safety protocols.
Quick Reference Tables (condensed)
Common ions: see Ionic Compounds section above for representative cations and anions; charges are essential for balancing formulas.
Seven diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2.
Key formulas to memorize conceptually:
ext{A} = ext{p}^+ + ext{n}^0
ext{Atomic weight} = ext{weighted average of isotopes}
1 ext{ amu} = 1.661 imes 10^{-24} ext{ g}
The gas in the bubbles during water boiling is water vapor, i.e., ext{H}_2 ext{O (g)}$$.
// End of notes (coverage matches the provided transcript content across Chapters 1 and 2).