Chapter 2 Notes: The Structure of Matter and Chemical Elements

States of Matter

  • Solid: fixed volume and fixed shape.
  • Liquid: fixed volume but shape changes with the container.
  • Gas: indefinite volume and shape; fills the container; compressible.
  • Note: Gases are generally compressible; solids and liquids are not (in everyday conditions).

The Chemical Elements

  • Salt water electrolysis demonstration: separation of salt water into substances with an accompanying note that these substances are elements because they cannot be broken down into simpler substances by chemical means.
  • Salt water example illustrates how elements and compounds can be identified or separated under certain processes.
  • Common elements and symbols included: Copper (Cu), Gold (Au), Carbon (C), Aluminum (Al), Neon (Ne), Mercury (Hg), Iron (Fe).

The Periodic Table and Its Visual Cues

  • Elements are shown with color-coded states at room temperature in some diagrams:
    • Black symbols: solids
    • Blue symbols: liquids
    • Red symbols: gases
  • Metals, metalloids/semimetals, and nonmetals are arranged in the table by increasing atomic number and grouped by similar chemical properties.
  • Example row elements shown in the slide snippet include Na, Mg, K, Ca, etc., illustrating representative alkali and alkaline earth metals.

The Periodic Table: Definition and History

  • Periodic: occurring at regular intervals; recurring from time to time; consisting of a series of repeated stages or digits.
  • Dmitri Mendeleev (1869) created an early version of the periodic table.
  • The periodic table organizes elements due to similar characteristics, and its roughly 150+ year history reflects evolving understanding of atomic structure.

Elements with Symbols and Atomic Masses (Representative Sample)

  • 1 H — Hydrogen — 1.008
  • 2 He — Helium — 4.003
  • 3 Li — Lithium — 6.941
  • 4 Be — Beryllium — 9.012
  • 5 B — Boron — 10.81
  • 6 C — Carbon — 12.01
  • 7 N — Nitrogen — 14.01
  • 8 O — Oxygen — 16.00
  • 9 F — Fluorine — 19.00
  • 10 Ne — Neon — 20.18
  • 11 Na — Sodium — 22.99
  • 12 Mg — Magnesium — 24.31
  • 13 Al — Aluminum — 26.98
  • 14 Si — Silicon — 28.09
  • 15 P — Phosphorus — 30.97
  • 16 S — Sulfur — 32.07
  • 17 Cl — Chlorine — 35.45
  • 18 Ar — Argon — 39.95
  • 19 K — Potassium — 39.10
  • 20 Ca — Calcium — 40.08
  • 21 Sc — Scandium — 44.96
  • 22 Ti — Titanium — 47.88
  • 23 V — Vanadium — 50.94
  • 24 Cr — Chromium — 52.00
  • 25 Mn — Manganese — 54.94
  • 26 Fe — Iron — 55.85
  • 27 Co — Cobalt — 58.93
  • 28 Ni — Nickel — 58.69
  • 29 Cu — Copper — 63.55
  • 30 Zn — Zinc — 65.41
  • 31 Ga — Gallium — 69.72
  • 32 Ge — Germanium — 72.61
  • 33 As — Arsenic — 74.92
  • 34 Se — Selenium — 78.96
  • 35 Br — Bromine — 79.90
  • 36 Kr — Krypton — 83.80
  • 37 Rb — Rubidium — 85.47
  • 38 Sr — Strontium — 87.62
  • 39 Y — Yttrium — 88.91
  • 40 Zr — Zirconium — 91.22
  • 41 Nb — Niobium — 92.91
  • 42 Mo — Molybdenum — 95.94
  • 43 Tc — Technetium — 98.00
  • 44 Ru — Ruthenium — 101.1
  • 45 Rh — Rhodium — 102.9
  • 46 Pd — Palladium — 106.4
  • 47 Ag — Silver — 107.9
  • 48 Cd — Cadmium — 112.4
  • 49 In — Indium — 114.8
  • 50 Sn — Tin — 118.7
  • 51 Sb — Antimony — 121.8
  • 52 Te — Tellurium — 127.6
  • 53 I — Iodine — 126.9
  • 54 Xe — Xenon — 131.3
  • 55 Cs — Cesium — 132.9
  • 56 Ba — Barium — 137.3
  • 57 La — Lanthanum — 138.9
  • 58 Ce — Cerium — 140.1
  • 59 Pr — Praseodymium — 140.9
  • 60 Nd — Neodymium — 144.2
  • 61 Pm — Promethium — 145
  • 62 Sm — Samarium — 150.4
  • 63 Eu — Europium — 152.0
  • 64 Gd — Gadolinium — 157.3
  • 65 Tb — Terbium — 158.9
  • 66 Dy — Dysprosium — 162.5
  • 67 Ho — Holmium — 164.9
  • 68 Er — Erbium — 167.3
  • 69 Tm — Thulium — 168.9
  • 70 Yb — Ytterbium — 173.0
  • 71 Lu — Lutetium — 175.0
  • 72 Hf — Hafnium — 178.5
  • 73 Ta — Tantalum — 180.9
  • 74 W — Tungsten — 183.9
  • 75 Re — Rhenium — 186.2
  • 76 Os — Osmium — 190.2
  • 77 Ir — Iridium — 192.2
  • 78 Pt — Platinum — 195.1
  • 79 Au — Gold — 197.0
  • 80 Hg — Mercury — 200.6
  • 81 Tl — Thallium — 204.4
  • 82 Pb — Lead — 207.2
  • 83 Bi — Bismuth — 209.0
  • 84 Po — Polonium — 209
  • 85 At — Astatine — 210
  • 86 Rn — Radon — 222
  • 87 Fr — Francium — 223
  • 88 Ra — Radium — 226
  • 89 Ac — Actinium — 227
  • 90 Th — Thorium — 232.0
  • 91 Pa — Protactinium — (231)
  • 92 U — Uranium — 238.0
  • 93 Np — Neptunium — (237)
  • 94 Pu — Plutonium — (242)
  • 95 Am — Americium — (243)
  • 96 Cm — Curium — (247)
  • 97 Bk — Berkelium — (247)
  • 98 Cf — Californium — (251)
  • 99 Es — Einsteinium — (252)
  • 100 Fm — Fermium — (257)
  • 101 Md — Mendelevium — (258)
  • 102 No — Nobelium — (259)
  • 103 Lr — Lawrencium — (260)
  • 104 Rf — Rutherfordium — (263)
  • 105 Db — Dubnium — (262)
  • 106 Sg — Seaborgium — (266)
  • 107 Bh — Bohrium — (267)
  • 108 Hs — Hassium — (277)
  • 109 Mt — Meitnerium — (268)
  • 110 Ds — Darmstadtium — (281)
  • 111 Rg — Roentgenium — (272)
  • 112 Cn — Copernicium — (285)
  • 113 Nh — Nihonium — (284)
  • 114 Fl — Flerovium — (289)
  • 115 Mc — Moscovium — (288)
  • 116 Lv — Livermorium — (293)
  • 117 Ts — Tennessine — (294)
  • 118 Og — Oganesson — (294)
  • Lanthanides: Ce, Pr, Nd, Pm, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, Lu
  • Actinides: Th, Pa, U, Np, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No, Lr

Atomic Structure: Subatomic Particles and Atomic Mass Units

  • Fundamental subatomic particles:
    • Proton: charge +1, mass ~ 1.00728 u
    • Neutron: charge 0, mass ~ 1.00867 u
    • Electron: charge -1, mass ~ 0.000549 u
  • Nucleus contains protons and neutrons; electrons occupy surrounding cloud (modern quantum view).
  • Neutral atoms have equal numbers of protons and electrons, yielding a net charge of 0.
  • Atomic notation and convention:
    • Atomic mass unit (amu) is used for masses of atoms and subatomic particles.
    • Atomic number Z = number of protons = number of electrons in a neutral atom.
    • Mass number A = Z + N (where N is the number of neutrons).
  • Isotopes differ in neutron number while having the same Z.
  • Isotope notation example: ^{A}{Z}\mathrm{X}, e.g., ^{12}{6}\mathrm{C} for Carbon-12.

Electrical Charge, Protons, Neutrons, and Electrons

  • Opposite charges attract; like charges repel.
  • Net charge of the atom is zero for neutral atoms unless ionized.
  • Key relationships:
    • Proton: +1 charge
    • Electron: -1 charge
    • Neutron: 0 charge
  • Example: Charged species arise from loss or gain of electrons; cations are positively charged and anions are negatively charged.

Ion Formation and Ions

  • Ions arise from loss or gain of electrons to achieve a noble-gas electron configuration.
  • Ionization equations:
    • ext{M}
      ightarrow ext{M}^{+} + e^{-}
    • ext{N} + e^{-}
      ightarrow ext{N}^{-}
  • Main-group elements:
    • Metals tend to LOSE electrons to reach the nearest noble-gas configuration.
    • Nonmetals tend to GAIN electrons to reach the nearest noble-gas configuration.
    • Transition metals: common oxidation states are not strictly predictable from position in the table.
    • Noble gases: do not form ions under normal conditions.

Common Ions and Charge Patterns

  • Typical cation and anion charges (illustrative range):
    • Cations: +1, +2, +3, +4
    • Anions: -1, -2, -3, -4
  • Common ions and their formation patterns:
    • ext{M}
      ightarrow ext{M}^{+} + e^{-}
    • ext{N} + e^{-}
      ightarrow ext{N}^{-}
  • Examples: +2, +3, +4, -4, -3, -2, -1, +1 as typical integer charges.

Ion Examples: Practice Problems (Worked Through)

  • Potassium (K):
    • Potassium has Z = 19 protons and 19 electrons in the neutral atom.
    • It loses one electron to form ext{K}^{+}. Reaction: ext{K}
      ightarrow ext{K}^{+} + e^{-}
  • Gallium (Ga):
    • If Ga^{3+} is formed, it would have 31 - 3 = 28 electrons remaining.
  • Oxygen (O):
    • Oxygen has Z = 8; it gains two electrons to form ext{O}^{2-}.
    • Reaction: ext{O} + 2e^{-}
      ightarrow ext{O}^{2-}
  • Chlorine (Cl):
    • If Cl^{-} is formed, it has one more electron than protons:
    • Reaction: ext{Cl} + e^{-}
      ightarrow ext{Cl}^{-}

Isotopes and Isotopic Abundance

  • Isotopes share the same number of protons (Z) but differ in the number of neutrons (N).
  • Mass number A = Z + N.
  • Natural isotopes occur in different relative abundances (percent natural abundance).
  • Example: Carbon isotopes
    • ^{12}_{6}\mathrm{C} (Carbon-12) — 6 neutrons
    • ^{13}_{6}\mathrm{C} (Carbon-13) — 7 neutrons
    • ^{14}_{6}\mathrm{C} (Carbon-14) — 8 neutrons
  • Isotope notation and neutron counts:
    • For a given element X, N = A - Z.
  • Isotopes differ in neutron count but leave Z unchanged.
  • Heavier isotopes may be radioactive (e.g., ^{14}_{6}\mathrm{C} has a radioactive half-life).

Atomic Mass, Isotopic Abundance, and Atomic Weight

  • Atomic mass value (atomic weight) is the weighted average of the masses of an element’s isotopes, based on their natural abundances:

    • ar{m} =

    \sumi fi mi where fi is the fractional abundance of isotope i and m_i its mass.

  • Example seed calculation ( Mg ): isotopes Mg-24 (78.99%), Mg-25 (10.00%), Mg-26 (11.01%) yield an atomic mass approximately 24.305 u (as shown in the example).

  • Isotopes frequently cited: Carbon-12, Carbon-13, Carbon-14; Radiocarbon dating uses ^{14}_{6}\mathrm{C} due to its known half-life.

Diagrams of Isotopes and Radiocarbon Dating Mention

  • Carbon-12, Carbon-13, and Carbon-14: different mass numbers but same Z; mass differences reflect different numbers of neutrons.
  • Radiocarbon dating uses the predictable decay of ^{14}_{6}\mathrm{C} to estimate the age of organic materials.
  • Uranium isotopes (e.g., ^{235}{92}\mathrm{U} and ^{238}{92}\mathrm{U}) illustrate how isotopes can have different neutron counts yet same Z.

Diatomic Molecules and Monoatomic Elements

  • Elements exist as monoatomic species in their pure form: e.g., noble gases like Ar, or metals in their standard state.
  • Diatomic molecules: two atoms of the same element bonded together in the pure form, represented as X2. Examples include: \mathrm{H{2}}, \mathrm{N{2}}, \mathrm{O{2}}, \mathrm{F{2}}, \mathrm{Cl{2}}, \mathrm{Br{2}}, \mathrm{I_{2}}.
  • Note: Br2 is a liquid at room temperature; many other diatomic molecules are gases under standard conditions.

Quick Review: Practice Problems (Page 37 Questions)

  • 1) What element has the chemical symbol Hg? → Mercury
  • 2) How many protons are in Tin? → Tin (Sn) has Z = 50 protons
  • 3) How many electrons are in Neon? → Neon (Ne) has 10 electrons (neutral atom)
  • 4) In an atom of Pt how many neutrons are in the nucleus? → Pt has Z = 78; atomic mass ~ 195.1; neutrons ≈ 195 − 78 ≈ 117
  • 5) What is the chemical symbol for Lawrencium? → Lr
  • 6) What are the chemical symbol and atomic number for Radon? → Symbol: Rn, Atomic number: 86
  • 7) What ion does phosphorus form? → Phosphide: \text{P}^{3-} (gain of 3 electrons to reach a noble-gas configuration)
  • 8) If a Ba ion has 54 electrons, what is the charge on the ion? → Barium (Z = 56) with 54 electrons means it has lost 2 electrons, so charge = +2 → \text{Ba}^{2+}
  • 9) Which is monoatomic vs diatomic? Iodine, Argon, Copper → Argon is monoatomic (Ar); Iodine is diatomic (I2) in its pure form; Copper is monoatomic in its standard state (Cu).

Connections to Foundational Principles and Real-World Relevance

  • The periodic table reflects electron configurations and periodic trends, underpinning chemical reactivity and bonding.
  • Understanding ions and oxidation states is essential for predicting compound formation, electrochemistry, and environmental chemistry.
  • Isotopes and atomic mass underpin dating methods (e.g., radiocarbon dating) and medical/industrial isotopic labeling.
  • The concept of diatomic molecules helps explain natural diatomic species in the atmosphere and their roles in chemistry and biology.

Practical and Ethical Considerations

  • Isotopic dating methods influence archaeology, paleontology, and anthropology; accuracy and limitations must be considered.
  • Ion formation and redox chemistry are central to battery technology, corrosion, and environmental processes; proper handling of reactive species is essential for safety.
  • The use of radiocarbon dating and isotopes in medicine and industry requires careful management of radiation exposure and ethical considerations around access and interpretation of data.

Summary of Key Formulas and Notation

  • Atomic composition: A = Z + N
  • Isotope notation: ^{A}_{Z}\mathrm{X}
  • Ion formation (loss of electrons): ext{M}
    ightarrow ext{M}^{+} + e^{-}
  • Ion formation (gain of electrons): ext{N} + e^{-}
    ightarrow ext{N}^{-}
  • Neutral atom net charge: 0
  • Isotopic average mass (weighted): ar{m} = \sumi fi m_i
  • Diatomic molecule representation: \mathrm{X{2}} (e.g., \mathrm{H{2}}, \mathrm{N{2}}, \mathrm{O{2}})
  • Isotope notation example: ^{12}_{6}\mathrm{C} for Carbon-12
  • Nuclear model statements (summary):
    • Nucleus contains protons and neutrons; most of the atom’s mass and positive charge reside in the nucleus.
    • Most of the atom’s volume is empty space with dispersed negative charges (electrons).
    • The number of protons equals the number of electrons in a neutral atom; overall net charge is zero.