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CHEM CH 2 Atoms, Molecules, and IONs

Periodic Table Basics

  • The periodic table arranges elements by increasing atomic number and shows recurring chemical properties.

  • Elements are organized into horizontal rows called periods and vertical columns called groups.

  • Group numbers (old notation) include 1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A for main-group elements and 1–12 for transition metals in some schemes; newer schemes use 1–18 across the main groups.

  • Main group elements include groups 1–2 and 13–18; transition metals are groups 3–12; inner transition metals are the Lanthanide and Actinide series.

  • The periodic table also highlights regions such as Metals, Metalloids, and Nonmetals; Alkali metals and Alkaline Earth metals are specific metal families; Halogens and Noble Gases are nonmetal families.

  • The Lanthanide series includes La, Ce, Pr, Nd, Pm, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, Lu.

  • The Actinide series includes Ac, Th, Pa, U, Np, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No, Lr.

  • Markets and practical relevance: elements are chosen for technologies (e.g., smartphones), industry, and biology due to distinct properties.

Regions and Classifications of the Periodic Table

  • Metals: typically lose electrons to form cations; good conductors of heat and electricity; include alkali and alkaline earth metals, transition metals, and post-transition metals.

  • Nonmetals: typically gain electrons to form anions or share electrons in covalent bonds; poor conductors of heat and electricity (with exceptions).

  • Metalloids (semi-metals): possess properties between metals and nonmetals; useful in electronics.

  • Alkali metals (Group 1) and Alkaline earth metals (Group 2): highly reactive metals.

  • Transition Metals (Groups 3–12): variable oxidation states; form colorful compounds; used in catalysis and materials.

  • Inner Transition Metals: Lanthanides and Actinides; often shown separately at the bottom of the table.

The Smartphone as a Case Study in Element Uses

  • Indium and Indium Tin Oxide (ITO): used in touch screens as a transparent conductor (Indium tin oxide).

  • Battery components: Lithium cobalt oxide as a positive electrode in Li-ion batteries; graphite (carbon) as negative electrode; sometimes manganese is used.

  • Glass and materials: alumina and silica in glass; potassium ions help strengthen glass.

  • Rare earth elements (lanthanides) provide colors in screens and are used in magnets for speakers and vibration units; many rare earths occur in the Earth's crust at low concentrations but are economically important in electronics.

  • Wiring and microelectronics: Copper for wiring; gold and silver for microelectronics; tantalum in micro-capacitors.

  • Silicon chips: pure silicon used to manufacture integrated circuits, with dopants to create conducting regions.

  • Casings and structural materials: Magnesium alloys for some casings; plastics (carbon-based) are common; flame-retardant compounds may include bromine; nickel reduces electromagnetic interference.

Molecular vs Ionic Substances

  • Molecules are composed of two or more NONMETAL atoms joined by covalent bonds (shared electrons).

  • Examples of diatomic molecules that occur naturally and are stable:

  • Hydrogen (H2)

  • Oxygen (O2)

  • Nitrogen (N2)

  • Chlorine (Cl2)

  • Bromine (Br2)

  • Iodine (I2)

  • Identify substances that are ionic vs molecular based on bonding and composition (examples are provided in practice activities).

  • Diatomic Molecules: These elements exist naturally (and more stably) as two atoms of each element bonded with one covalent bond 7

Ions: Cations and Anions

  • A CATION is a positively charged ion (loss of electrons).

  • An ANION is a negatively charged ion (gain of electrons).

  • Monatomic cations are always metals; monatomic anions are always nonmetals.

  • Examples: number comes before the charge

    • Cations: {Na}^{+},{Ca}^{2+},{Fe}^{3+}

    • Anions: {F}^{-},{O}^{2-},{N}^{3-}

    • monatomic - one atom - monatomic cations are always metals - monatomic anions are nonmetals

Ionic Compounds

  • Ionic compounds form via ionic bonds between ions.

  • They can be monatomic (Na^+, Cl^−) or polyatomic - made of more than one atom (NO3^−, SO4^{2−}).

  • Composition: an ionic compound consists of a METAL ion and a NONMETAL (or a polyatomic ion). The compound is neutral overall; the charges balance.

  • Examples of ionic compounds:

    • MgCl2, Li2SO4, KOH, Ca(NO3)2, FePO4, KF

Predicting Ion Charge for Elements (Representative Elements)

  • Main group elements form charges related to their group numbers:

    • Group 1 = +1, Group 2 = +2, Group 3A = +3.

    • Group 5A = -3, Group 6A = -2, Group 7A = -1.

    • Group 4 = typically no charge for C, Si, Ge; Sn and Pb can behave differently (often transition-metal-like behavior in some contexts).

  • Transition metals can form multiple oxidation states; this leads to more complex naming (Roman numerals used for charge). Form more than one charge

Naming Binary Ionic Compounds (Inorganic Nomenclature)

  • If the cation is a main-group element: use the element name as the cation name.
    mg²+ = magnesium

  • Li^+ = lithium
    If the cation is a transition metal: use the element name followed by its charge in Roman numerals in parentheses.

  • Cu²+ + copper(ii)

  • Cu^+ = copper(i)

  • Fe²+ = iron(ii)

  • Fe³+ = iron (iii)

  • The old -ous/-ic nomenclature is ignored in contemporary teaching.

  • Practice: write formulas for ionic compounds from simple monatomic ions; also name common polyatomic ions and ionic compounds.

Naming the Anion and Balancing Formulas

  • Monatomic anions: names end in -ide after the root of the element:

    • N³ - = nitride

    • O²- = oxide

    • S²- = sulfide

    • Se²- = selenide

    • Te²- = telluride

  • Polyatomic anions: names and formulas are fixed (-ate)

    • Co3²- = carbonate

    • NO3^- = nitrate

    • NO2^- = nitrite

    • PO4³- = phosphate

    • PO3³- = phosphite

  • Oxoanions: memorize the -ates (e.g., NO3^-, NO2^-, SO4^{2-}, SO3^{2-}, CO3^{2-}, PO4^{3-}).

  • The approach for writing formulas: balance cations and anions so that the overall charge is neutral; determine the smallest whole-number ratio of ions to achieve neutrality (the formula unit).

Naming Polyatomic Ions and Ionic Compounds (2: Naming the anion)

  • For polyatomic anions, the name of the anion and its charge are fixed (e.g., sulfate SO4^{2-}, phosphate PO4^{3-}).

  • Example:

    • ext{Na}3 ext{PO}4 is sodium phosphate.

    • ext{K}2 ext{SO}4 is potassium sulfate.

Binary Molecular (Covalent) Compounds

  • Naming: two nonmetals form covalent bonds; use prefixes to indicate the number of atoms.

  • Rules:

    • The first element is named with its full name.

    • The second element is named with its root plus the suffix -ide.

    • Use prefixes to indicate numbers: 1{ o} ext{mono}, 2{ o} ext{di}, 3{ o} ext{tri}, 4{ o} ext{tetra}, 5{ o} ext{penta}, 6{ o} ext{hexa}, 7{ o} ext{hepta}, 8{ o} ext{octa}, 9{ o} ext{nona}, 10{ o} ext{deca}

    • For oxides, drop the 'a' or 'o' ending of the prefix before vowels (e.g., CO = carbon monoxide).

  • Examples:

    • ext{PI}_3
      ightarrow ext{phosphorus triiodide}

    • ext{N}_2 ext{O}
      ightarrow ext{dinitrogen monoxide}

    • ext{P}4 ext{O}7
      ightarrow ext tetraphosphorus heptoxide}

    • ext{BrF}_3
      ightarrow ext{bromine trifluoride}

Simple Molecular Compounds: Alkanes

  • Alkanes are saturated hydrocarbons with formula $CnH{2n+2}$ for straight-chain alkanes.

  • First ten members (straight-chain):

    • Methane: ext{CH}_4

    • Ethane: ext{C}2 ext{H}6

    • Propane: ext{C}3 ext{H}8

    • Butane: ext{C}4 ext{H}{10}

    • Pentane: ext{C}5 ext{H}{12}

    • Hexane: ext{C}6 ext{H}{14}

    • Heptane: ext{C}7 ext{H}{16}

    • Octane: ext{C}8 ext{H}{18}

    • Nonane: ext{C}9 ext{H}{20}

    • Decane: ext{C}{10} ext{H}{22}

  • Condensed structural forms accompany each formula (e.g., CH3CH3 for ethane).

Acids: Oxoacids and Acids Without Oxygen

  • Two types of acids:

    • Oxoacids (hydrogen bonded to an oxygen-containing polyatomic ion).

    • Acids that do not contain oxygen.

  • Naming Oxoacids: start from the polyatomic ion; if the polyatomic ion ends with -ite, replace with -ous acid; if it ends with -ate, replace with -ic acid.

    • NO3^- (nitrate) → nitric acid (HNO3)

    • NO2^- (nitrite) → nitrous acid (HNO2)

    • SO4^{2-} (sulfate) → sulfuric acid (H2SO4)

    • SO3^{2-} (sulfite) → sulfurous acid (H2SO3)

    • PO4^{3-} (phosphate) → phosphoric acid (H3PO4)

    • ClO4^- (perchlorate) → perchloric acid (HClO4)

    • ClO3^- (chlorate) → chloric acid (HClO3)

    • ClO2^- (chlorite) → chlorous acid (HClO2)

    • ClO^- (hypochlorite) → hypochlorous acid (HClO)

    • IO4^- (periodate) → periodic acid (HIO4)

    • BrO- → hypobromous acid (HBrO) and other related oxyacids via same pattern.

  • 3 oxygen and more get -ate - 4 or more gets prefix peri-

  • 2 oxygen or less get -ite - 1 gets prefix hypo-

  • Summary: for oxyacids, replace the -ite suffix with -ous acid and -ate suffix with -ic acid in the acid name; for hypohalos and perhalates follow the hypo- and per- prefixes with the same pattern.

Ionic compound = cation (+) + anion (-)

Laws of Definite and Multiple Proportions

  • Law of Definite Proportions (Proust): Different samples of a pure compound always contain the same proportion of elements by mass.

    • Example (given in notes): Hydrogen to Oxygen by mass is fixed at 1 part H to 8 parts O.

    • Expressed as rac{m_ ext{H}}{m_ ext{O}} = rac{1}{8} with appropriate units.

  • Law of Multiple Proportions (Dalton): If two elements form more than one compound, the weights of one element that combine with a fixed weight of the other are in ratios of small whole numbers.

    • Example: For carbon and oxygen, 12.01 g C combine with 16.00 g O to form CO; 12.01 g C combine with 32.00 g O to form CO2, giving a simple ratio of 1 : 2 for the oxygen masses relative to a fixed carbon mass.

Dalton’s Atomic Theory (Summarized)

1) Elements are composed of tiny particles called atoms.
2) Atoms of different elements have characteristic sizes, masses, and properties.
3) Compounds form when atoms bond in small whole-number ratios (Law of Definite Proportions).
4) In chemical reactions, atoms are rearranged; they are not created or destroyed.

Atomic Structure: Thomson, Millikan, Rutherford

  • J. J. Thomson: Determined the charge-to-mass ratio of the electron (−e/me).

  • Robert Millikan: Measured the elementary charge using the oil-drop experiment, enabling calculation of the electron’s mass.

  • Ernest Rutherford: Gold foil experiment revealed a dense, positively charged atomic nucleus; most of the atom is empty space.

  • Evolution of models:

    • Dalton’s solid spheres → Thomson’s plum pudding model (electrons embedded in a positive matrix) → Rutherford’s nucleus-centered model → Bohr’s planetary model (quantized orbits) → Current orbital model with electrons in probabilistic orbitals.

Nucleus, Electrons, and Atomic Size

  • The atom consists of a tiny, dense nucleus (protons and neutrons) surrounded by a large region where electrons reside.

  • Nucleus: contains protons (p+) and neutrons (n0); overall mass is concentrated in the nucleus.

  • Electron cloud: large volume, small mass compared to the nucleus; electrons contribute most of the atom’s volume.

Mass, Volume, and the Atomic Mass Unit (u)

  • 1 unified atomic mass unit (u) = 1.66054 × 10^-24 g.

  • Protons and neutrons each have a mass close to 1 u; electrons have negligible mass in comparison.

  • Modern chemistry often uses amu or atomic weight (the latter is a weighted average of isotopic masses).

Isotopes, Notation, and Atomic Structure

  • Isotopes: Atoms of the same element with different numbers of neutrons, hence different mass numbers A but the same atomic number Z (number of protons).

  • Isotope notation convention: mass number A, atomic number Z, and symbol: usually written as ^{A}_{Z} ext{X}.

  • Examples of isotopes:

    • Hydrogen: ^{1}{1} ext{H}, ext{ }^{2}{1} ext{H} (deuterium), ext{ }^{3}_{1} ext{H} (tritium)

  • Neons and other isotopes have natural abundances that determine the element’s average mass.

  • The medical isotope example: ^{75}_{34} ext{Se} (selenium-75) is mentioned as being used medically for pancreatic disorder diagnostics.

Protons, Neutrons, and Electrons in Isotopes and Ions

  • For a neutral atom: the number of electrons equals the number of protons (electrons = Z).

  • In ions, electrons differ from protons; electrons = Z − charge for cations, or Z + |charge| for anions.

  • Given an isotope symbol (e.g., ^{A}_{Z}X):

    • Protons = Z

    • Neutrons = A − Z

    • Electrons in a neutral atom = Z; in an ion, adjust by the charge.

Average Atomic Mass (Atomic Weight)

  • Definition: The weighted average mass of the isotopes of an element, taking into account their natural abundances.

  • Formula: ext{Atomic weight} = ext{Atomic mass} = ext{sum}{i} ig(mi imes piig) where mi = mass of isotope i (in u) and p_i = fractional abundance of isotope i (as a decimal).

  • Note: Atomic weight is typically not an integer due to the fractional abundances of several isotopes.

  • Example (Neon as given in the notes):

    • Isotopes and abundances:

    • ^{20}_{10} ext{Ne} with mass ~ 19.992 ext{ amu} and abundance 90.51%

    • ^{21}_{10} ext{Ne} with mass ~ 20.993 ext{ amu} and abundance 0.27%

    • ^{22}_{10} ext{Ne} with mass ~ 21.991 ext{ amu} and abundance 9.22%

    • Convert abundances to fractions: 0.9051, 0.0027, 0.0922.

    • Compute the weighted average:
      ext{Average mass} = (19.992 imes 0.9051) + (20.993 imes 0.0027) + (21.991 imes 0.0922) \
      \approx 18.0948 + 0.0566 + 2.0276 \
      \approx 20.179 ext{ amu}

    • This aligns with Neon’s standard atomic weight around 20.18 amu.

Isotope Identification and Notation Details

  • Isotopic notation example: Hydrogen isotopes (protium, deuterium, tritium) as ^{1}{1} ext{H},^{2}{1} ext{H},^{3}_{1} ext{H}. These indicate mass numbers A = 1, 2, 3 respectively.

  • Atomic number Z is read from the periodic table and defines the number of protons in the nucleus; for a neutral atom, electrons = protons.

  • The Se-75 isotope example is cited as medically used to diagnose pancreatic disorders.

Quick Reference: Typical Notations and Calculations

  • Atomic number (Z): number of protons.

  • Mass number (A): number of protons + neutrons.

  • Neutrons = A − Z.

  • Electrons for neutral atoms = Z; for ions, electrons = Z − charge (if positive) or Z + |charge| (if negative).

  • Molecular vs Ionic substances: Molecular = covalent bonds between nonmetals; Ionic = electrostatic attraction between metals and nonmetals (or polyatomic ions).

Important Equations and Concepts (LaTeX)

  • Law of Definite Proportions: rac{m_ ext{H}}{m_ ext{O}} = rac{1}{8} ext{ (by mass for a fixed compound, for example)}

  • Law of Multiple Proportions (illustrative): ratios of the second element in different compounds with a fixed first element are small whole numbers (e.g., CO vs CO2: masses of O in CO and CO2 to fixed C mass are 1:2).

  • Atomic weight (weighted average of isotopes): ext{Atomic weight} = ext{sum}i ig(mi imes piig) where mi is the isotope mass and p_i is its fractional abundance.

  • Isotope notation: ^{A}_{Z} ext{X} where A = ext{mass number} = ext{protons} + ext{neutrons}, \, Z = ext{atomic number} = ext{protons}.

  • 1 u = 1.66054 imes 10^{-24} ext{ g}.

  • For a neutral atom, electrons = protons: ext{electrons} = Z. For ions, adjust by the charge: ext{electrons} = Z - q, where q is the ionic charge.