Study Guide: Electronic Structure, Ionic & Molecular Compounds, and Chemical Quantities

Chapter 5: Electronic Structure of Atoms & Periodic Trends

5.1 Electromagnetic Radiation

Electromagnetic radiation is a form of energy that exhibits wave-like behavior as it travels through space. Understanding its properties is essential for explaining atomic structure and spectra.

• Electromagnetic Radiation: Energy transmitted through space in the form of waves, such as light, radio waves, and X-rays.

• Wavelength (\( \lambda \)): The distance between two consecutive peaks of a wave. Measured in meters (m) or nanometers (nm).

• Frequency (\( \nu \)): The number of wave cycles that pass a given point per second. Measured in hertz (Hz).

• Amplitude: The height of the wave from the center line to the peak or trough; relates to the intensity of the radiation.

• Relationship between Wavelength, Frequency, and Energy:

$c = \lambda \nu$

$E = h \nu$

• Where c is the speed of light (\( 3.00 \times 10^8 \) m/s), h is Planck's constant (\( 6.626 \times 10^{-34} \) J·s).

• Order of the Electromagnetic Spectrum: Radio → Microwave → Infrared → Visible Light → Ultraviolet → X-rays → Gamma Rays.

Example: Visible light has wavelengths from about 400 nm (violet) to 700 nm (red).

5.2 Atomic Spectra & Energy Levels

Atoms absorb and emit energy in discrete amounts, leading to unique atomic spectra. These phenomena are explained by the concept of energy levels.

• Photon: A quantum of electromagnetic energy; a particle of light.

• Atomic Spectra: The set of frequencies of light emitted or absorbed by atoms of an element.

• Photon Absorption: When an electron absorbs energy and moves to a higher energy level (excited state).

• Photon Emission: When an electron falls to a lower energy level, releasing energy as a photon.

• Excited State: A state in which an electron has more energy than its ground state.

Example: The emission spectrum of hydrogen shows distinct lines corresponding to electron transitions between energy levels.

5.3 Sublevels & Orbitals

Electrons in atoms are organized into shells, subshells, and orbitals, which determine the atom's chemical behavior.

• Subatomic Particle Responsible for Chemical Behavior: The electron.

• Electron Organization: Shell (energy level, n) → Subshell (s, p, d, f) → Orbital (specific region within a subshell).

• Number of Electrons per Energy Level: $2n^2$

• Subshell Names and Order of Filling: s, p, d, f (order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, ...)

• Shapes: s = spherical, p = dumbbell-shaped.

• Number of Orbitals and Electrons per Subshell:

Subshell	Number of Orbitals	Max Electrons
s	1	2
p	3	6
d	5	10
f	7	14

• Probability Density: The likelihood of finding an electron in a particular region of space; the exact location is unknown due to the Heisenberg uncertainty principle.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers; each orbital can hold a maximum of two electrons with opposite spins.

5.4 Orbital Diagrams & Electron Configurations

Electron configurations describe the arrangement of electrons in an atom. Orbital diagrams visually represent these arrangements.

• Order of Subshell Filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p (Aufbau principle).

• Hund's Rule: Electrons occupy orbitals singly before pairing up in the same subshell.

• Writing Electron Configurations: List subshells with superscripts indicating the number of electrons (e.g., O: 1s22s22p4).

• Abbreviated Electron Configurations: Use the previous noble gas in brackets (e.g., O: [He]2s22p4).

Example: Sodium (Na): 1s22s22p63s1 or [Ne]3s1

5.5 Electron Configuration & The Periodic Table

The periodic table is organized according to electron configurations, with blocks corresponding to subshells.

• s-block: Groups 1 and 2, plus helium.

• p-block: Groups 13-18.

• d-block: Transition metals (Groups 3-12).

• f-block: Lanthanides and actinides.

• Abnormal Configurations: Chromium (Cr) and copper (Cu) have electron configurations that differ from the expected order due to stability of half-filled and filled d subshells.

Example: Cr: [Ar]4s13d5; Cu: [Ar]4s13d10

5.6 Trends in Periodic Properties

Periodic trends describe how certain properties of elements change across periods and down groups in the periodic table.

• Valence Electrons: Electrons in the outermost shell; determine chemical reactivity.

• Atomic Size: Decreases left to right across a period (due to increasing nuclear charge), increases top to bottom down a group (due to added energy levels).

• Ionization Energy: Increases left to right (harder to remove electrons), decreases top to bottom (easier to remove electrons as size increases).

• Metallic Character: Decreases left to right, increases top to bottom.

Example: Sodium is larger and more metallic than chlorine; chlorine has a higher ionization energy.

Chapter 6: Ionic & Molecular Compounds

6.1 Ionic Transfer of Electrons

Ionic and covalent bonds are two major types of chemical bonds. Ionic bonds involve the transfer of electrons, while covalent bonds involve sharing.

• Ionic Bond: Electrostatic attraction between oppositely charged ions.

• Covalent Bond: Sharing of electrons between nonmetal atoms.

• Ion: An atom or group of atoms with a net charge due to loss or gain of electrons.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

• Ionic Charge: $\text{Ionic charge} = \text{number of protons} - \text{number of electrons}$

• Metals: Lose electrons to form cations (+).

• Nonmetals: Gain electrons to form anions (−).

Example: Mg loses 2 electrons to form Mg2+; O gains 2 electrons to form O2−.

6.2 Ionic Compounds

Ionic compounds are composed of cations and anions held together by ionic bonds. Their formulas reflect the ratio of ions needed for charge neutrality.

• Ionic Compound: A compound composed of positive and negative ions.

• Ionic Formula: The simplest ratio of ions that results in a neutral compound (e.g., NaCl).

6.3 Naming & Writing Ionic Formulas

Naming ionic compounds follows specific rules based on the ions present. The periodic table helps deduce ionic charges for representative elements.

• Deducing Ionic Charges: Group number indicates charge for s- and p-block elements (e.g., Group 1: +1, Group 17: −1).

• Naming Rules:

  • Cation name (element name) + anion name (element root + "-ide").

  • No prefixes for ionic compounds.

  • Roman numerals for variable charge metals (e.g., FeCl2: iron(II) chloride).

• Variable Charge Metals: The charge is indicated in the name (e.g., Copper(II) = Cu2+).

• Representative Elements: s- and p-block elements form only one type of ion; their charges are based on group number.

Example: MgCl2: magnesium chloride; CuO: copper(II) oxide (since O is 2−, Cu must be 2+).

6.4 Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge. They form ionic compounds with other ions.

• Polyatomic Ion: A charged group of covalently bonded atoms (e.g., NO3−, SO42−).

• Naming Compounds with Polyatomic Ions: Name the cation first, then the polyatomic ion (e.g., NaNO3: sodium nitrate).

Example: CaCO3: calcium carbonate.

Additional info: The 20 most common polyatomic ions should be memorized (see course slides or Quizlet set).

6.5 Molecular Compounds Sharing Electrons

Molecular compounds are formed when nonmetals share electrons, resulting in discrete molecules.

• Molecular Compound: A compound composed of molecules formed by covalently bonded nonmetals.

• Molecule: A group of atoms bonded together, representing the smallest unit of a molecular compound.

Example: H2O (water), CO2 (carbon dioxide).

Chapter 7: Chemical Quantities

7.1 The Mole

The mole is a fundamental unit in chemistry for counting particles. Avogadro's number relates moles to the number of particles.

• Avogadro's Number: $6.022 \times 10^{23}$ particles/mol.

• Conversions:

  • Particles to moles: $\text{moles} = \frac{\text{number of particles}}{6.022 \times 10^{23}}$

  • Moles to particles: $\text{number of particles} = \text{moles} \times 6.022 \times 10^{23}$

• Mole Ratios: 1 mol CH4 contains 1 mol C atoms and 4 mol H atoms.

7.2 Molar Mass

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol). It is calculated using the atomic masses from the periodic table.

• Molar Mass: The mass of 1 mole of a substance (g/mol).

• Finding Molar Mass: Sum the atomic masses of all atoms in the chemical formula.

Example: H2O: (2 × 1.01) + 16.00 = 18.02 g/mol.

7.3 Calculations Using Molar Mass

Molar mass allows conversion between the mass of a substance and the number of moles.

• Conversions:

  • Mass to moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

  • Moles to mass: $\text{mass (g)} = \text{moles} \times \text{molar mass (g/mol)}$

Example: 36.04 g H2O × (1 mol / 18.02 g) = 2.00 mol H2O.

7.4 Mass Percent Composition

Mass percent composition expresses the percentage by mass of each element in a compound.

• Mass Percent: $\text{Mass percent of X} = \frac{\text{mass of X in 1 mol compound}}{\text{molar mass of compound}} \times 100\%$

Example: In H2O, %H = (2.02/18.02) × 100% = 11.2%; %O = (16.00/18.02) × 100% = 88.8%.

7.5 Empirical Formulas

The empirical formula is the simplest whole-number ratio of elements in a compound. It can be determined from the molecular formula or from mass percent composition.

• Empirical Formula: The simplest ratio of atoms in a compound.

• Finding Empirical Formula from Molecular Formula: Divide subscripts by their greatest common factor.

• From Mass Percent Composition:

  1. Assume 100 g sample; convert % to grams.

  2. Convert grams to moles for each element.

  3. Divide by the smallest number of moles to get the simplest ratio.

Example: A compound with 40% C, 6.7% H, 53.3% O yields empirical formula CH2O.

7.6 Molecular Formulas

The molecular formula shows the actual number of atoms of each element in a molecule. It is a whole-number multiple of the empirical formula.

• Difference: Empirical formula = simplest ratio; molecular formula = actual number of atoms.

• Determining Molecular Formula:

  1. Find empirical formula and its molar mass.

  2. Divide the compound's molar mass by the empirical formula mass to find the multiple (n).

  3. Molecular formula = (empirical formula) × n.

• From Mass Percent and Molar Mass: Use mass percent to find empirical formula, then use molar mass to find molecular formula.

Example: Empirical formula CH2O, molar mass 180 g/mol; empirical mass = 30 g/mol; n = 180/30 = 6; molecular formula = C6H12O6.