Study Guide: Periodic Properties and Molecules & Compounds (Chapters 3 & 4)

Chapter 3: Periodic Properties of the Elements

Periodic Trends

Periodic trends describe how certain properties of elements change across periods and groups in the periodic table. Understanding these trends is essential for predicting element behavior.

• Atomic Radius: Generally decreases across a period (left to right) and increases down a group (top to bottom).

• Ionization Energy: The energy required to remove an electron. Increases across a period and decreases down a group. Successive ionization energies (1st, 2nd, etc.) are higher for each subsequent electron removed.

• Electron Affinity: The energy change when an atom gains an electron. Generally becomes more negative across a period.

• Metallic Character: Increases down a group and decreases across a period.

• Exceptions: Some elements deviate from these trends due to electron configurations.

Ion Size and Charge

Ion size depends on the number of electrons and the charge of the ion.

• Cations (positive ions) are smaller than their parent atoms.

• Anions (negative ions) are larger than their parent atoms.

• Charge Prediction: Main-group elements form ions based on their group number.

Effective Nuclear Charge, Shielding, and Penetration

These concepts explain how electrons experience the nucleus's pull.

• Effective Nuclear Charge (Zeff): The net positive charge felt by an electron. Calculated as $Z_{\text{eff}} = Z - S$ where Z is the atomic number and S is the shielding constant.

• Shielding: Inner electrons block outer electrons from the nucleus.

• Penetration: The extent to which an electron can approach the nucleus.

Electron Configurations and Orbital Diagrams

Electron configurations show the arrangement of electrons in an atom or ion.

• Aufbau Principle: Electrons fill the lowest energy orbitals first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund’s Rule: Electrons fill degenerate orbitals singly before pairing.

• Orbital Diagrams: Visual representations using arrows for electrons.

Magnetism: Paramagnetism and Diamagnetism

• Paramagnetic: Atoms/ions with unpaired electrons; attracted to magnetic fields.

• Diamagnetic: Atoms/ions with all electrons paired; repelled by magnetic fields.

Classification of Elements

• Main-group elements: s and p blocks.

• Transition elements: d block.

• Inner transition elements: f block.

Core vs. Valence Electrons

• Core electrons: Inner electrons not involved in bonding.

• Valence electrons: Outermost electrons involved in chemical reactions.

Types of Elements

• Metals: Conductive, malleable, and typically form cations.

• Nonmetals: Insulative, brittle, and typically form anions.

• Metalloids: Properties intermediate between metals and nonmetals.

Special Groups

• Noble gases: Group 18, inert.

• Halogens: Group 17, highly reactive.

• Alkaline earth metals: Group 2.

• Alkali metals: Group 1.

Chapter 4: Molecules and Compounds

Ionic vs. Covalent Bonds

Chemical bonds form between atoms to create compounds.

• Ionic bonds: Formed between metals and nonmetals; involve transfer of electrons.

• Covalent bonds: Formed between nonmetals; involve sharing of electrons.

• Bond Types: Single, double, and triple covalent bonds indicate the number of shared electron pairs.

Lewis Electron-Dot Structures

Lewis structures represent valence electrons as dots around element symbols.

• Purpose: Visualize bonding and lone pairs.

• Application: Used to predict molecular shape and reactivity.

Writing Formulas for Compounds and Hydrates

• Ionic Compounds: Combine cations and anions to achieve charge neutrality.

• Molecular Compounds: Use prefixes to indicate the number of atoms.

• Hydrates: Compounds with water molecules attached; indicated by a dot (e.g., CuSO4·5H2O).

Naming Compounds

• Ionic Compounds: Name cation first, then anion. Use Roman numerals for metals with variable charge.

• Molecular Compounds: Use prefixes (mono-, di-, tri-, etc.) and end the second element with "-ide".

• Hydrates: Add "hydrate" with appropriate prefix (e.g., pentahydrate).

Formula Mass, Molar Mass, and Mass Percent Composition

• Formula Mass: Sum of atomic masses in a formula unit.

• Molar Mass: Mass of one mole of a substance (g/mol).

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

Calculations: Grams, Moles, and Molecules

• Avogadro’s Number: $6.022 \times 10^{23}$ things per mole.

• Conversions: Use molar mass and Avogadro’s number to convert between grams, moles, and molecules.

• Example: To find molecules in 10 g of H2O: $\text{moles} = \frac{10\ \text{g}}{18.02\ \text{g/mol}}$ $\text{molecules} = \text{moles} \times 6.022 \times 10^{23}$

Conversion Factors and Mole Ratios

• Mole Ratios: Derived from chemical formulas; used in stoichiometric calculations.

• Example: In H2O, the mole ratio of H to O is 2:1.

Empirical Formula Determination

• Empirical Formula: Simplest whole-number ratio of atoms in a compound.

• Combustion Analysis: Used to determine empirical formula from experimental data.

Common Ions and Metals

Memorize monoatomic anions (Table 4.2), polyatomic ions (Table 4.4), and metals that only make one cation (Fig. 4.6).

Monoatomic Anions	Polyatomic Ions
Cl- (chloride)	NO3- (nitrate)
O2- (oxide)	SO42- (sulfate)
N3- (nitride)	CO32- (carbonate)
Additional info: See textbook for full lists.	Additional info: See textbook for full lists.

Metals with Only One Cation
Na+ (sodium)
K+ (potassium)
Ca2+ (calcium)
Additional info: See Fig. 4.6 for more examples.

Hydrocarbons and Organic vs. Inorganic Compounds

• Hydrocarbons: Compounds containing only carbon and hydrogen.

• Organic Compounds: Contain carbon, often with hydrogen, oxygen, nitrogen, etc.

• Inorganic Compounds: All other compounds, including salts, metals, and minerals.

Additional info: For detailed lists of ions and naming conventions, refer to Tables 4.2, 4.4, and Fig. 4.6 in your textbook.