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Chemistry Notes: Atomic Structure, Isotopes, Ions, Electron Configuration, and Periodic Trends

Atomic Structure and Matter

  • Chemistry is the study of matter. Matter is composed of atoms.
  • Atoms are made up of three fundamental particles: protons (positive charge), neutrons (no charge), and electrons (negative charge).
  • In a neutral atom, number of protons equals number of electrons.
  • Ions are atoms that have gained or lost electrons, resulting in a net charge.
  • An ion with a charge is called an ion; a neutral species is not an ion.

Major and Minor Points about Protons, Neutrons, and Electrons

  • Protons: positively charged.
  • Neutrons: neutral.
  • Electrons: negatively charged.
  • The nucleus contains protons and neutrons; electrons reside in surrounding electron shells/orbitals.
  • The force between positively charged protons and negatively charged electrons is what holds the atom together.
  • There is mostly empty space between the nucleus and the outer electrons; electrons are not located in fixed orbits but in orbitals (conceptually).

Atomic Number, Mass Number, and Isotopes

  • Atomic number (Z): number of protons in the nucleus; defines the identity of the element.
  • Mass number (A): total number of protons and neutrons in the nucleus, A = Z + N.
  • Number of neutrons (N) = A − Z.
  • Isotopes: atoms with the same Z (same element) but different A (different number of neutrons).
  • Changing protons -> a different element (identity changes).
  • Changing neutrons -> isotope of the same element.
  • Changing electrons -> ion (net charge changes).
  • Neutral atoms have equal numbers of protons and electrons.

Isotope Notation and How to Read It

  • Isotopes are often written as ^A_Z X, where X is the element symbol, A is the mass number (top), and Z is the atomic number (bottom).
  • Example: ^{12}_{6}C represents carbon-12 with 6 protons and 6 neutrons (A = 12, Z = 6).
  • In some contexts, the mass number (A) can be written as the superscript, and the atomic number (Z) as the subscript next to X.
  • The mass number is not the number of neutrons; it is the sum of protons and neutrons (A = Z + N).
  • Your periodic-table reference can tell you Z for a given element (e.g., C has Z = 6).
  • To figure out neutrons from a given isotope: N = A − Z. For carbon-12: N = 12 − 6 = 6 neutrons.

Examples and Practice with Isotopes and Ions

  • Hydrogen isotopes: protium (protons=1, neutrons=0), deuterium (protons=1, neutrons=1), tritium (protons=1, neutrons=2).
  • If you change neutrons but keep Z constant, you get isotopes of hydrogen.
  • If you change protons (even by 1), you get a different element entirely.
  • If you change electrons while keeping Z the same, you get an ion.
  • Example: An ion with 7 protons and 7 + x electrons where the total mass is 14 and x is computed as above leads to a charge. If there are 3 extra electrons, the ion is -3 charge (a monatomic ion).
  • Monatomic ion: an ion made from a single atom (e.g., Na^+, Cl^−, N^3−).
  • Polyatomic ion: an ion made from two or more atoms (e.g., SO4^{2−}, PO4^{3−}).

Monatomic vs Polyatomic Ions and Nomenclature

  • Monatomic ion: ion formed from one atom. Prefix: mono- (ionic charge definitions use cation for positive and anion for negative).
  • Polyatomic ion: ion formed from more than one atom.
  • Rules of thumb from the lecture:
    • A cation is a positive ion.
    • An anion is a negative ion.
    • If a species has a charge, it is an ion (superscript indicating the charge is a hint).
  • In chemistry, the term monatomic ion is used for ions from a single atom; polyatomic ions are composed of multiple atoms.
  • Common monatomic ions include many alkali and alkaline earth metals (e.g., Na^+, Mg^{2+}), halides (e.g., Cl^−), etc.
  • Common aqueous polyatomic ions are listed in standard reference tables; you will encounter them on homework and exams.

Prefixes and Ionic Nomenclature

  • Prefix for one atom is mono- (e.g., monatomic ion).
  • In practice, vocabulary like cation (positive) and anion (negative) is essential; mnemonic: Cats are positive (cations); ants are negative (anions).

Periodic Table Structure and Trends

  • Periods: horizontal rows on the periodic table.
  • Groups: vertical columns on the periodic table.
  • Metals vs nonmetals: metals are generally left and in the lower portion; nonmetals are on the right and upper portion.
  • The most metallic elements are at the bottom-left; the least metallic elements are at the top-right.
  • Common trends include atomic radius, ionization energy, and metallic character.

Atomic Radius (Size) Trends

  • Atomic radius increases as you move down a group (more electron shells).
  • Atomic radius decreases as you move across a period from left to right (increasing nuclear charge with similar shielding pulls electrons closer).
  • Periods vs Groups:
    • Groups: atoms get bigger going down (more shells).
    • Periods: atoms get smaller across (increasing effective nuclear charge with similar shielding).
  • Important definitions:
    • Groups: columns; Periods: rows.

Ionization Energy (Energy to Remove an Electron)

  • Ionization energy is the energy required to remove the outermost electron from an atom.
  • Higher ionization energy: harder to remove an electron.
  • Trend: ionization energy is larger in the top-right region (noble gases and near them) and smaller in the bottom-left region (larger atoms with more shielding).
  • Example intuition: outer electrons in small, highly charged atoms are held tightly; larger atoms with shielding have lower ionization energies.
  • Conceptual pairing: top-right atoms have high ionization energy; bottom-left atoms have low ionization energy.

Electron Configuration and Subshells

  • Electrons fill shells and subshells in a patterned order to minimize energy.
  • Subshell capacities:
    • s subshell holds up to 2 electrons.
    • p subshell holds up to 6 electrons.
    • d subshell holds up to 10 electrons.
    • f subshell holds up to 14 electrons (not needed for this course's examples).
  • Shells and subshells terminology:
    • Each shell is labeled by a principal quantum number n (1, 2, 3, …).
    • Each shell contains subshells (s, p, d, f) with characteristic shapes and capacities.
  • Notation for electron configuration: write the shells from low to high energy, filling subshells in order of increasing energy.
  • Example: Bromine (Br, Z = 35) electron configuration:
    • ext{Br}: 1s^2\, 2s^2\, 2p^6\, 3s^2\, 3p^6\, 3d^{10}\, 4s^2\, 4p^5
    • This reflects the arrangement across shells 1 through 4 with the appropriate subshells filled.
  • Conceptual model introduced:
    • The first architect fills the first shell (1s).
    • The second architect fills the second shell (2s, 2p).
    • The third architect fills the third shell (3s, 3p, 3d), then the fourth shell (4s, 4p).
  • Tendency: electrons fill to achieve a noble gas core, i.e., the electron configuration often resembles that of the nearest noble gas with additional valence electrons.
  • Example of the noble gas rule:
    • Fluorine (Z = 9) has a [Ne] core with 2 more electrons in the 2p orbital: ext{F}: 1s^2\, 2s^2\, 2p^5
    • Neon (Z = 10) is a noble gas with closed shells: 1s^2\, 2s^2\, 2p^6
    • Gaining or losing electrons tends to move the electron configuration toward that of a noble gas.
  • Important caution:
    • Do not rely solely on visual direction on the periodic table to predict which noble gas is closest; compare actual numeric atomic numbers to determine the closest noble gas in terms of electronic configuration.

Practical Application: Electron Configuration Practice

  • Given a configuration, you should identify the element or verify whether a proposed configuration matches the element by counting electrons in each subshell and shell.
  • For ex., Br's configuration above corresponds to Z = 35.
  • If asked to determine whether a given configuration corresponds to a particular element, compare total electrons to Z and compare to the nearest noble gas core.
  • Common teaching heuristic: elements strive to achieve the electron configuration of the nearest noble gas; this is the driving force behind chemical reactivity and bonding.

Common Concepts for Homework and Exams

  • Distinguish between:
    • Isotopes: same Z, different A (neutron count changes).
    • Ions: different electron count, charged species.
    • Monatomic ions vs Polyatomic ions: one-atom ions vs multi-atom ions.
  • Identify ion types from charge notation and from the presence or absence of superscripts in formulas.
  • Determine the identity of an element from its atomic number Z; count protons, neutrons, and electrons in a specified ion or isotope.
  • Recognize common ion charges by group classification (e.g., +1 for alkali metals, +2 for alkaline earth metals, +3 for group 13; -1 for halogens; -2 for chalcogens; -3 for pnictogens; noble gases are neutral or non-ionizing under normal conditions).
  • Understand exceptions (e.g., mercury) to standard ionic charges; these are usually discussed as notable exceptions in advanced contexts.

Common Mixtures and Properties

  • Mixtures can be classified as:
    • Homogeneous (solutions): uniform composition, cannot see different parts.
    • Heterogeneous: visibly different parts.
  • In early notes, a simple model described particles and shells; later, the distinction between physical and chemical properties was introduced via examples like shiny rings (physical property) and mixtures (heterogeneous vs homogeneous).
  • Physical properties: can be observed without changing the composition (e.g., color, density, hardness, phase).
  • Chemical properties: describe how substances react to form new substances (not deeply covered in this portion, but part of the test scope).

Density, Units, and Conversions (Worked Examples)

  • Density formula:
    ho = \frac{m}{V} where m is mass and V is volume.
  • Given mass and volume, convert to desired units:
    • Example: mass = 45 g, volume = 30.7 cm^3 = 30.7 mL = 0.0307 L
    • Convert to kg/L if needed: 45 g = 0.045 kg; 0.0307 L is the same as 0.0307 L
    • Density: \rho = \frac{0.045\ \text{kg}}{0.0307\ \text{L}} \approx 1.46\ \text{kg/L}
  • Key concept: 1 cm^3 = 1 mL; 1 L = 1000 mL; 1 kg = 1000 g; 1 g = 1000 mg.

Unit Conversions and LD50 (Practical Data Handling)

  • LD50 (lethal dose 50): amount of a substance that kills 50% of the test population; typically reported as mg per kg of body mass (mg/kg).
  • To convert mg/kg to g per pound (g/lb):
    • First convert mg to g: 1 g = 1000 mg → multiply by 1/1000
    • Then convert kg to pounds: 1 kg ≈ 2.20462 lb → multiply by (1 kg / 2.20462 lb)
    • Overall factor: x\ \text{mg/kg} \times \frac{1}{1000} \times \frac{1}{2.20462} = x\times 4.53592\times 10^{-4} \frac{\text{g}}{\text{lb}}
  • Example: 10 mg/kg → 0.00454 g/lb approximately.
  • Important practical note: Such problems are usually more complex on homework; exact numbers depend on provided conversion factors. The key skill is set up and proper unit cancellation.
  • For unit cancellation problems, the order of multiplication/division does not matter as long as units cancel properly.

Sig Figs and Measurement Precision

  • Determine the number of significant figures in a given number:
    • Nonzero digits are always significant.
    • Leading zeros are not significant.
    • Captive zeros (between nonzero digits) are significant.
    • trailing zeros are significant if and only if there is a decimal point somewhere in the number.
  • Examples (conceptual):
    • 0.00560 has 3 sig figs (5, 6, and 0 is captive after the decimal point).
      1. has 3 sig figs (decimal point makes trailing zeros significant).
    • 100 has ambiguous sig figs without a decimal point (could be 1, 2, or 3).
  • The goal in the course is to apply the sig figs rules consistently in calculations.

Quick Reference: Common Notation and Terminology

  • Atom identity is determined by atomic number Z.
  • Isotopes differ in neutron number; ions differ in electron count and charge.
  • Monatomic vs polyatomic ions differentiation based on the number of atoms involved.
  • Noble gases are generally nonreactive due to full electron shells.
  • The column and row terminology:
    • Groups (columns): main block categories (alkali metals, halogens, noble gases, etc.).
    • Periods (rows): indicate energy level occupancy; across a period, radius decreases; down a group, radius increases.
  • Electron configuration serves as a bridge between the periodic table and chemical properties; it explains reactivity, ionization energy, and periodic trends.

Final Tips for the Exam

  • Be able to distinguish isotopes, ions, monatomic ions, and polyatomic ions from notation and descriptions.
  • Memorize the general trends: metallic character (bottom-left), atomic radius (increases down a group, decreases across a period), ionization energy (increases up and to the right).
  • Practice writing electron configurations, especially for Br and nearby elements, using the shell/subshell framework and the filling order.
  • Use the noble gas core shortcut to simplify electron configurations: write [Noble gas] followed by valence electron configuration.
  • Be comfortable with unit conversions and dimensional analysis; practice multi-step conversions with proper cancellation.
  • For sig figs, practice identifying significant digits in a variety of numbers and applying the rules consistently.

Practice Prompt Reflections (from Lecture)

  • If given a configuration, determine the element or verify the plausibility with the nearest noble gas core.
  • Identify whether a given species is an ion and whether it is monatomic or polyatomic.
  • Be prepared to classify materials as metallic or nonmetal and to discuss relative trends across periods and down groups.
  • Be prepared to perform and explain density calculations, including unit conversions between mass/volume units.