Chapter 2 Notes: Atoms, Ions, and Molecules

The Nuclear Atom

• Nucleus: positively charged center of an atom, containing nearly all of the atom’s mass
• Consists of two types of particles:
  • Protons – positively charged subatomic particles
  • Neutrons – electrically neutral subatomic particles
• The nucleus contains most of the atom’s mass; electrons orbit around the nucleus

Atomic Mass Units

• Atomic mass units (u) are used to express the relative masses of atoms and subatomic particles
• 1 u = 1 Dalton (Da)
• Named for the English chemist John Dalton
• Formula/idea: masses of atoms are typically given in atomic mass units when discussing isotopes

The Nuclear Atom: Summary

• (Key takeaway) The atom has a tiny, dense, positively charged nucleus containing protons and neutrons; electrons occupy surrounding space with mostly empty space between nucleus and electrons

Isotope Symbol or Nuclide Symbol

• Isotope Symbol (Nuclide Symbol): ^{A}_{Z}X
  • A = mass number = total number of nucleons (protons + neutrons) in the nucleus
  • Z = atomic number = number of protons in the nucleus; determines the identity of the element
  • X = the element’s symbol (one- or two-letter)
• Isotopes are atoms of the same element (same Z) with different mass numbers (A) due to different numbers of neutrons
• Example concepts:
  • 1p 0n for protium (hydrogen-1, the most common hydrogen isotope)
  • 1p 1n for deuterium (hydrogen-2, D)
  • 1p 2n for tritium (hydrogen-3, T)
• Nuclear notation emphasizes the relation A = Z + N (N is the number of neutrons)

Practice: Writing Nuclide Symbols

• Format: ^{A}_{Z}X to write nuclides
• Example: write nuclide with 11 protons and 12 neutrons
  • Z = 11 → X = Na (sodium)
  • A = Z + N = 11 + 12 = 23
  • Nuclide symbol: ^{23}_{11} ext{Na}
• Another: atom with 19 protons and 20 neutrons
  • Z = 19 → X = K (potassium)
  • A = 39; Symbol: ^{39}_{19} ext{K}

Isotopes

• Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A)
• Consequence: different nuclei masses, slightly different physical properties

Isotopes: Carbon-14 Dating

• The method is used to determine the age of organic material
• Process:
  • Living organisms take in carbon-14 (C-14) from the atmosphere while alive
  • After death, C-14 decays at a known rate
  • Measuring remaining C-14 in bone or plant material estimates time since death
• Limitation: C-14 dating is effective up to about ext{50,000 years}; older samples require other dating methods
• Notation: isotopes of carbon include ^{14}_{6} ext{C} (C-14)

Average Atomic Mass

• Average atomic mass = weighted average of masses of all isotopes of an element
• Formula:ar{M} = igg( ext{fractional abundance}1igg) imes m1 + igg( ext{fractional abundance}2igg) imes m2 + \, ext{…}
• Masses are given in atomic mass units (u)
• Natural abundance: the proportion of a particular isotope, usually expressed as a percentage relative to all isotopes for that element in a natural sample
• Example concept: for copper with two isotopes, the average atomic mass is the sum of each isotope’s mass times its natural abundance

Mass Spectrometer (Example Calculation)

• Naturally occurring copper has two isotopes:
  • ^{63}_{29} ext{Cu} with abundance 63 ext{Cu}: 69.09\, ext{%} and mass 62.93 ext{ amu}
  • ^{65}_{29} ext{Cu} with abundance 65 ext{Cu}: 30.91\, ext{%} and mass 64.9278 ext{ amu}
• Average atomic mass:
  • ar{M}_{ ext{Cu}} = (0.6909)(62.93 ext{ amu}) + (0.3091)(64.9278 ext{ amu})
  • Result: ar{M}_{ ext{Cu}} \,=\, 63.55 ext{ amu}

Average Atomic Mass of Li

• Natural lithium composition:

  • ^{6} ext{Li}: 7.42\% ext{ with mass } 6.015 ext{ amu}
  • ^{7} ext{Li}: 92.58\% ext{ with mass } 7.016 ext{ amu}

• Calculation:

  • \bar{M}_{ ext{Li}} = (0.0742)(6.015 ext{ amu}) + (0.9258)(7.016 ext{ amu}) = 6.941 ext{ amu}

• Note: Anecdote about lithium in early soda beverages

  • In the early 1920s, a lime soda named 7up advertised that it included lithium citrate and promoted benefits of “lithiated soda”
  • Lithium salts are used as antipsychotic drugs; 7-Up no longer contains Li
  • The note is included here as a historical aside from the slides

Periodic Table of the Elements

• A tabular arrangement showing regular repetition of properties
• Key terms:
  • Periods: horizontal rows
  • Groups: vertical columns
  • Elements in the same group tend to have similar physical and chemical properties
• Broad categories of elements:
  1) Metals (left side and bottom of the table)
  2) Nonmetals (right side and top of the table)
  3) Metalloids (between metals and nonmetals)
• Broad groupings and representative examples:
  • Group 1: Alkali metals
  • Group 2: Alkaline earth metals
  • Group 17: Halogens
  • Group 18: Noble gases
• Periodic table also contains metals, metalloids, and nonmetals in characteristic regions
• (Visual students should refer to the large block showing H, He, Li, Be, B, C, N, O, F, Ne, etc.)

Groups of Elements: Cont. (Overview of layout)

• Main elements (representative elements) in Groups 1, 2, 13–18; Transition metals occupy central block
• Notation highlights periodic locations of elements like H, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, As, Se, Br, Kr, etc.

Ionic Compounds

• Ionic compounds consist of charged particles (ions) formed by transfer of electrons between atoms
• Ions are held together by electrostatic forces
• Cations: ions with positive charge
• Anions: ions with negative charge
• Example questions:
  • How many protons and electrons are in ^{27}_{13} ext{Al}^{3+}?
  • Protons: 13; Electrons: 13 − 3 = 10
  • How many protons and electrons are in ^{78}_{34} ext{Se}^{2-}?
  • Protons: 34; Electrons: 34 + 2 = 36

Ionic Compounds: Composition and Nomenclature

• Ions are typically formed from a metal and a nonmetal
• Metals form cations; nonmetals form anions
• The charges on ions depend on their location in the periodic table
• Formula unit: the smallest electrically neutral unit of an ionic compound
• Naming ions:
  • 1) Cation is named first using the element’s name: Mg → magnesium
  • 2) Anion is named by adding the -ide suffix to the element’s name: Cl → chloride
  • Example: MgCl2

Oxoanions and Polyatomic Ions

• Oxoanions: polyatomic anions that include oxygen with another element
• Naming oxoanions: the name of the element that appears first in the formula with its ending changed to -ite or -ate
• Common rules:
  • SO4^{2-} → Sulfate
  • SO3^{2-} → Sulfite
  • ClO4^{-} → Perchlorate
  • ClO3^{-} → Chlorate
  • ClO2^{-} → Chlorite
  • ClO^{-} → Hypochlorite
• Some elements form more than two oxoanions; the prefix -per- means one more oxygen than -ate, and hypo- means one fewer oxygen than -ite

Polyatomic Ions

• Polyatomic ions are ions composed of two or more atoms acting as a unit
• Examples include: sulfate, sulfite, nitrate, nitrite, carbonate, ammonium, etc.

Molecular Compounds

• Molecular compounds are formed by covalent bonds between nonmetals
• Covalent bonds are created by sharing one or more pairs of electrons
• Molecular compounds are composed of nonmetals

Molecular Formulas and Empirical Formulas

• Molecular formula shows the number and type of atoms in one molecule (e.g., glucose)
  • Molecular formula for glucose: ext{C}{6} ext{H}{12} ext{O}_{6}
• Empirical formula shows the smallest whole-number ratio of elements in a compound (e.g., CH2O for glucose)
• Example:
  • If a compound has the same elemental makeup as glucose but in simplest ratio, empirical formula would be CH2O

Naming Molecular Compounds

• Rules:
  • The first word is the name of the first element as it appears in the formula
  • The second element’s name ends with -ide
  • Use prefixes to indicate the number of atoms of each element in the molecule (Table 2.2): example SO3 = sulfur trioxide
• Examples:
  • NF3 → nitrogen trifluoride
  • SO2 → sulfur dioxide
  • N2Cl4 → dinitrogen tetrachloride
  • NO2 → nitrogen dioxide
  • N2O → dinitrogen monoxide (dinitrogen oxide is less common; common is monoxide)
  • CO → carbon monoxide
  • N2O3 → dinitrogen trioxide
  • (These illustrate how prefixes are applied to second element as needed)

Prefixes and Naming Rules

• Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-
• Rules: 1) Do not use mono- for the first element in a molecule
  • Example: SO3 is sulfur trioxide, not sulfur monioxide
    2) Prefixes ending with -o or -a are modified when the following element begins with a vowel
  • Example: P4O10 → tetraphosphorus decaoxide
• Additional examples:
  • NF3 → nitrogen trifluoride
  • NO2 → nitrogen dioxide
  • N2O → dinitrogen monoxide
  • CO → carbon monoxide
  • N2Cl4 → dinitrogen tetrachloride

What Are the Names of the Following Compounds?

• OF2 → oxygen difluoride
• S4N4 → tetrasulfur tetranitride
• BCl3 → boron trichloride

Naming Ionic Compounds (Recap)

• Ionic compounds consist of cations (usually metals) and anions (usually nonmetals)
• 1) Cation is named first using the element’s name (e.g., Mg → magnesium)
• 2) Anion is named by adding the -ide suffix to the element’s name (e.g., Cl → chloride)
• Formulas for ionic compounds must be neutral: e.g., Mg^{2+} + 2 Cl^{-} → MgCl_{2}

Formula Examples for Ionic Compounds

• Al2O3 is formed by Al^{3+} and O^{2-} ions since: 2(Al^{3+}) + 3(O^{2-}) = +6 + (-6) = 0
  • So the formula unit is Al{2}O{3}
• CaBr2 forms from Ca^{2+} and Br^{-}: 1 Ca^{2+} and 2 Br^{-} → CaBr_{2}
• Na2CO3 forms from Na^{+} and CO3^{2-}: 2 Na^{+} and 1 CO3^{2-} → Na{2}CO{3}

Binary Ionic Compounds: Classical vs Stock System

• Classical (Latin-based) system uses Latin names modified with suffixes:
  • Fe^{2+} → ferrous; Fe^{3+} → ferric
  • Cu^{+} → cuprous; Cu^{2+} → cupric
  • Examples: CuCl → cuprous chloride; CuCl2 → cupric chloride
• Stock system uses Roman numerals in parentheses to indicate the charge on the cation:
  • Fe^{2+} → iron(II); Fe^{3+} → iron(III)
  • Cu^{+} → copper(I); Cu^{2+} → copper(II)
  • Examples: CuCl → copper(I) chloride; CuCl2 → copper(II) chloride

Practice: Naming Ionic Compounds

• FeCl2 → iron(II) chloride
• FeCl3 → iron(III) chloride
• MnO → manganese(II) oxide
• Mn2O3 → manganese(III) oxide
• MnO2 → manganese(IV) oxide

Practice: Writing Formulas with Polyatomic Ions

• Sodium sulfite → Na2SO3

Formulas for Ionic Compounds: Additional Examples

• Potassium carbonate → K2CO3
• Manganese(II) sulfate → MnSO4

Practice: Naming Ionic Compounds (More Examples)

• BaO → barium oxide
• Cr2(SO4)3 → chromium(III) sulfate

Naming Acids

• Binary (hydrogen + a halide) acids:
  • Naming rule: Hydro- + the base name of the halogen + -ic + acid
  • Example: HBr → hydrobromic acid

Oxoanions and Related Acids

• If the oxoanion ends in -ate, the corresponding acid ends in -ic; if ending is -ite, the corresponding acid ends in -ous
• Examples:
  • NO3^{-} (nitrate) → nitric acid (HNO3)
  • NO2^{-} (nitrite) → nitrous acid (HNO2)
  • SO4^{2-} (sulfate) → sulfuric acid (H2SO4)

Summary of Key Points from Chapter 2

• The nucleus contains protons and neutrons; electrons are arranged around it
• Atomic mass units (u) express masses; 1 u = 1 Da
• Nuclide symbol encodes A (mass number) and Z (atomic number): ^{A}_{Z}X
• Isotopes differ by neutron number; isotopes have same Z but different A
• Carbon-14 dating uses the known decay rate of C-14; effective up to ~5.0\times10^{4} years
• Average atomic mass is a weighted average of isotope masses by their natural abundances
• The mass spectrum example highlights how isotope masses and abundances yield the standard atomic mass
• Periodic table: periods, groups; metals, nonmetals, metalloids; major groupings and representative elements
• Ionic compounds: formed by transfer of electrons; charge balance yields neutral formula units
• Cations are named first; anions end with -ide unless they are polyatomic ions
• Oxoanions use -ate vs -ite naming; per- and hypo- prefixes indicate extra or fewer oxygens
• Polyatomic ions exist as units (e.g., sulfate, nitrate, carbonate)
• Molecular compounds form covalent bonds; use prefixes to indicate atom counts; empirical vs molecular formulas differ
• Prefix usage rules include omitting mono- for the first element and vowel effects on prefixes
• Binary acids use hydro- prefix and -ic suffix; oxoacids use -ic or -ous depending on the oxoanions
• Stock vs Classical naming conventions help indicate oxidation states in ionic compounds
• Practice problems reinforce naming, formula writing, and identification of ions and their charges

ext{Key formulas to remember:}

• Average atomic mass: ar{M} = \,\sumi fi m_i\,
  
• Nuclide symbol: ^{A}_{Z}X
• Oxoanions and acids: \text{SO}4^{2-} \rightarrow \text{Sulfate}; \text{H}2\text{SO}_4 \rightarrow \text{sulfuric acid}
• Ionic compound neutrality: \text{cation}^{m+} + \text{anion}^{n-} \to \text{neutral formula unit}
• Example balances: \text{Al}^{3+}\text{ + }\text{O}^{2-} \to \text{Al}2\text{O}3
• Example charges:
  • \text{Mg}^{2+} + 2\text{Cl}^{-} \to \text{MgCl}_2
  • \text{Ca}^{2+} + 2\text{Br}^{-} \to \text{CaBr}_2
• Binary acid naming: \text{HBr} \rightarrow \text{hydrobromic acid}
• Per-/hypo- prefixes for oxyanions: \text{ClO}_4^{-} \rightarrow \text{perchlorate}, \text{ClO}^{-} \rightarrow \text{hypochlorite}