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Comprehensive Study Notes: Periodic Trends and Bonding (OpenStax-based Transcript)

Page 1: Periodic Table Trends

  • Topic: Periodic Table Trends and related foundational concepts in chemistry.

  • Emphasis on patterns in element properties as you move through the periodic table.

  • OpenStax-based framework referenced throughout (textbook origin).

Page 2: Dmitri Mendeleev

  • Dmitri Mendeleev: Russian chemist and inventor; Born in 1907 St. Petersburg (as per transcript).

  • Published the first organized periodic table of elements based on atomic weight.

  • Predicted two new elements (gallium Ga and scandium Sc) and was correct in both predictions.

  • Predicted trends in elements based on their place in the table, which aligned with later discoveries.

  • Significance: Pioneered a periodic organization that allowed prediction of unknown elements and their properties.

Page 3: Henry Moseley

  • Henry Moseley concluded that periodic table should be arranged by the number of protons (the atomic number).

  • This arrangement became the standard and is the version used today.

  • Significance: Reconciled inconsistencies in Mendeleev’s table and provided a more fundamental basis for periodicity.

Page 4: The Modern Periodic Table Structure

  • Periodic table organization concepts:

    • MAIN-GROUP ELEMENTS (also called representative elements)

    • TRANSITION ELEMENTS

    • INNER TRANSITION ELEMENTS (lanthanides and actinides)

  • Group labeling in older notation: 1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A; corresponding to columns on the modern table.

  • Visual categories on the page (examples shown in the transcript):

    • 1A: Metals (main-group)

    • 2A: Metals (main-group)

    • 8A: Noble gases

    • 3A, 4A, 5A, 6A, 7A: Other main-group elements

    • 8A: Noble gases at the far right

  • Blocks and series include: Hydrogen (H) and helium (He) among the lighter elements; transition metals in the center; lanthanides and actinides shown as inner-transition elements.

  • Key takeaway: The modern table is arranged by atomic number with clear groupings (main-group, transition, inner-transition) and includes separate lanthanide and actinide rows.

Page 5: Metals, Metalloids, and Nonmetals

  • Figure 2.11 illustrates which elements are metals, metalloids, and nonmetals.

  • Examples listed in the page include:

    • Chromium (Cr), Copper (Cu), Cadmium (Cd), Lead (Pb), Bismuth (Bi)

    • Boron (B), Silicon (Si), Arsenic (As), Antimony (Sb), Tellurium (Te)

    • Carbon (graphite), Sulfur (S), Chlorine (Cl), Bromine (Br), Iodine (I)

  • Takeaway: The table visually separates metals, metalloids, and nonmetals with representative examples.

Page 6: Molecules in Elements

  • Figure 2.16 shows elements that occur as molecules in nature.

  • 1A–8A groups contain diatomic and small molecules: H2, N2, O2, F2, Cl2, Br2, I2, etc.

  • P4, S8, Se8 are examples of molecules composed of multiple atoms in elemental form.

  • Terminology:

    • Diatomic molecules: two atoms (H2, N2, O2, F2, Cl2, Br2, I2)

    • Tetratomic molecules: four atoms (e.g., P4)

    • Octatomic molecules: eight atoms (e.g., S8)

  • Significance: Explains why some elements exist naturally as diatomic or polyatomic molecules.

Page 7: A Biological Periodic Table

  • A specialization of the periodic table for biology:

    • Building-block elements (essential for life)

    • Major minerals (macroelements important for biological processes)

    • Trace elements (needed in smaller amounts)

  • Examples mentioned in the transcript include:

    • Building-block: Na, Mg, Ca, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Mo

    • Elements in the “BCNOF” region: B, C, N, O, F, Si, P, S, Cl, As, Se, Sn, I (illustrative layout of essential life elements)

  • This framing highlights how biology emphasizes a subset of periodic table elements due to physiological roles.

Page 8: Periodicity and Common Ions (Monatomic Ions)

  • The periodic table prompts discussion of common monatomic ions and expected charges.

  • Patterns to notice: metals tend to form cations; nonmetals tend to form anions.

  • Examples of common monatomic ions include:

    • H+, Li+, Na+, K+ (group 1 cations)

    • Mg2+, Ca2+ (group 2 cations)

    • Cr2+, Cr3+, Mn2+, Fe2+, Co2+, Cu+, Zn2+, etc. (various oxidation states)

  • These ions are often shown in basic ion-charge patterns for teaching purposes.

Page 9: Common Monoatomic Ions (Table 2.3)

  • Cations (positive charges): H+, Li+, Na+, K+, Ag+ (silver), Mg2+, Ca2+, Zn2+, Sr2+, Ba2+, Cd2+, Al3+ (examples listed on the page)

  • Anions (negative charges): F-, Cl-, Br-, I-, O2- (oxide), S2- (sulfide), N3- (nitride)

  • Note: The table emphasizes that many common ions have charges of -1, -2, or -3 and that many are blue in the original figure to indicate common ions.

  • Concept: The charge of common ions follows predictable patterns based on group/family membership.

Page 10: Periodic Trends in a Nutshell

  • Key trends to remember:

    • Ionization energy (IE)

    • Electron affinity (EA)

    • Atomic radius

    • Nonmetallic character

    • Metallic character

  • These trends reflect how atoms attract/hold electrons and how their sizes change across the table.

  • The page summarizes these trends to be revisited with more detail in subsequent sections.

Page 11: Atomic Radius – Size

  • Trend 1: Down a group, atomic radius increases due to the addition of electron shells (increase in principal quantum number n).

  • Trend 2: Across a period, atomic radius decreases due to increasing nuclear charge (more protons) pulling electrons closer to the nucleus.

  • Core idea: Radius grows with more shells; shrinks with greater effective nuclear charge across a period.

Page 12: Ionization Energy – Removing an Electron

  • Definition: Ionization energy is the energy required to remove an electron from a neutral atom.

  • Result of sufficient energy: The atom is ionized and becomes charged; protons and electrons are no longer equal.

  • General trend: The larger the atom, the easier it is to remove an outer electron.

  • Specific trend: First ionization energies decrease down a group, and increase across a period.

  • Implication: Elements high up and to the right tend to have high IE; elements down and to the left tend to have low IE.

Page 13: Electron Affinity – Adding an Electron

  • Electron affinity (EA) is the energy change when a gaseous atom accepts an electron.

  • Opposite of ionization energy: energy release or absorption when gaining an electron.

  • Some elements release energy when gaining an electron (negative EA); others require energy (positive EA).

  • Trend: Electron affinities generally become more negative across a period (EA decreases numerically / becomes more exothermic).

Page 14–15: Ionic vs Covalent Bonding

  • Ionic bonding: transfer of electrons from a metal to a nonmetal, forming ions and an ionic compound.

  • Covalent bonding: sharing of electrons between nonmetals, forming covalent molecules.

  • Presentation in figures: diagrams contrasting ionic vs covalent bonds and whether compounds are ionic or covalent.

  • Key takeaway: Bond type depends on electronegativity differences and element types involved.

Page 16: Formation of Ionic Compounds

  • Concept: Transferring electrons from one element to another yields an ionic compound.

  • Result: Electrostatic attraction between oppositely charged ions forms the crystal lattice of an ionic solid.

Page 17: Covalent Bonding – H2 Molecule

  • Covalent bonds form when two nonmetals share electrons.

  • Example: Formation of a covalent bond between two hydrogen atoms to produce

    • H–H (molecule) with shared electron pair(s).

  • Significance: Explains the bonding in most organic and many inorganic molecules.

Page 18: Predicting Common Oxidation States (Charge) by Group

  • Concept: Elements have predicted common oxidation states/charges.

  • Representative patterns discussed (from groups):

    • 1A elements tend to +1 (e.g., H+, Li+)

    • 2A elements tend to +2 (e.g., Mg2+, Ca2+)

    • 7A elements tend to -1 (halides, e.g., F-, Cl-, Br-, I-)

    • 6A elements tend to -2 (e.g., O2-, S2-)

    • 3A elements can form +3 (and sometimes -3 in certain contexts)

  • Examples listed in the transcript include:

    • H+, Li+, Na+, Mg2+, K+ (and other common cations listed in the original figure)

    • Anions: F-, Cl-, O2-, N3-, etc.

  • Practical note: Oxidation states guide naming and formula construction for ionic compounds.

Page 19: Metals with Several Oxidation States

  • Table 2.4 (partial) shows metals that exhibit multiple oxidation states:

    • Copper: Cu+1 (copper(I), cuprous) and Cu+2 (copper(II), cupric)

    • Cobalt: Co+2 (cobalt(II)) and Co+3 (cobalt(III))

    • Iron: Fe+2 (iron(II), ferrous) and Fe+3 (iron(III), ferric)

    • Manganese: Mn+2 (manganese(II)) and Mn+3 (manganese(III))

    • Tin: Sn+2 (tin(II), stannous) and Sn+4 (tin(IV), stannic)

  • Takeaway: Some metals form more than one common oxidation state, which affects compound naming and properties.

Page 20: Polyatomic Ions (Partial List)

  • Polyatomic anions and related acids (partial list):

    • Ammonium: NH4+ (cation)

    • Nitrate: NO3-; Nitrite: NO2-; Peroxide: O2^2-;

    • Sulfate: SO4^2-; Sulfite: SO3^2-;

    • Hydrogen sulfate: HSO4-; Acetate: CH3COO-;

    • Cyanide: CN-; Hydroxide: OH-;

    • Carbonate: CO3^2-; Hydrogen carbonate: HCO3-;

    • Phosphate: PO4^3-; Dihydrogen phosphate: H2PO4-;

    • Chlorate: ClO3-; Perchlorate: ClO4-; Chlorite: ClO2-; Hypochlorite: ClO-;

    • Chromate: CrO4^2-; Dichromate: Cr2O7^2-;

    • Permanganate: MnO4-.

  • These ions are widely used in aqueous chemistry and acid-base/naming contexts.

Page 21: Polyatomic Ions to Learn for Class

  • Consolidated list of important polyatomic ions:

    • Cations: NH4+ (ammonium)

    • Anions: NO3-, NO2-, SO4^2-, CO3^2-, PO4^3-, MnO4-, CN-, OH-, CH3COO-, ClO3-, CrO4^2-, Cr2O7^2-

  • Use: Recognize these ions and their formulas for naming and balancing reactions.

Page 22: Hydrates

  • Definition: A hydrate is a compound with a specific number of water molecules integrated into its solid structure.

  • Example: Copper(II) sulfate pentahydrate, CuSO4 · 5H2O.

  • Other examples listed: BaCl2 · 2H2O, LiCl · H2O, MgSO4 · 7H2O, Sr(NO3)2 · 4H2O.

  • Significance: Hydrates illustrate water of crystallization and how water content affects formula naming.

Page 23: Ionic Neutrality and Aluminum Oxide

  • Example: Aluminum oxide, Al2O3, formed from Al3+ and O2- ions.

  • Charge balance check: 2(+3) + 3(-2) = 0, illustrating why Al2O3 is electronically neutral as a compound.

  • General lesson: Ionic compounds must be electrically neutral; charges from cations and anions must sum to zero in each formula unit.

Page 24: Ionic Compound Naming Rules

  • If the compound contains a metal and a nonmetal, the compound is Ionic.

  • Metal cation with a single possible charge (fixed charge):

    • Examples: Alkali metals, alkaline earth metals, Ag+, Al3+, Cd2+, Zn2+

    • Naming: Name metal first; add -ide to the root of the nonmetal name.

  • Metal cation with variable charges (multivalent metals):

    • Examples: Fe, Cu, Cr, Mn, Sn, Co, Ni, etc.

    • Naming: Name metal first; specify the charge of the metal cation with Roman numeral in parentheses; add -ide to the nonmetal root.

  • Practical rule: Roman numerals indicate the oxidation state of multivalent metals.

Page 25: Molecular vs Empirical Formulas

  • Molecular formula shows the actual number of atoms in a molecule (true formula).

  • Empirical formula shows the simplest whole-number ratio of elements in a compound.

  • Example: For

    • Molecular formula:

    • Empirical formula: NH2 for N2H4 (the two may be the same or differ depending on the compound).

  • Connection: For covalent compounds, both descriptions are used to convey composition.

Page 26: Naming Molecular (Covalent) Compounds

  • Binary molecular compounds consist of two different elements.

  • Naming rules:
    1) Name the first element that appears in the formula.
    2) Name the second element with its ending changed to -ide.

  • Examples from the page:

    • HCl → hydrogen chloride

    • HI → hydrogen iodide

  • Important: Quantitative prefixes are not shown in these two examples but are used when more than one atom of each element is present.

Page 27: Greek Prefixes for Molecular Compounds

  • Greek prefixes denote the number of atoms of each element in a molecule:

    • Mono- (1), Di- (2), Tri- (3), Tetra- (4), Penta- (5), Hexa- (6), Hepta- (7), Octa- (8), Nona- (9), Deca- (10), and so on.

  • Common prefixes include: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-

  • Use: Essential for naming molecular compounds with more than one atom of an element.

Page 28: Mono- Prefix Omission and Oxide Naming Convention

  • The prefix mono- is generally omitted for the first element in a binary molecular compound (e.g., NO not mononitrogen oxide).

  • For ease of pronunciation, the last letter of prefixes ending in o or a may be dropped when naming oxides (e.g., N2O5 becomes dinitrogen pentoxide, not dinitrogen pentaoxide).

  • Practical rule: Follow standard conventions to avoid redundancy and ensure readable names.

Notes:

  • Throughout these pages, examples and conventions are presented as OpenStax-style teaching material, emphasizing periodic trends, bonding types, ionic/covalent naming, and common ions.

  • The content blends historical context (Mendeleev, Moseley) with modern periodic table organization and practical nomenclature rules for ionic and molecular compounds.