Exam 2 Study Guide: Ions, Compounds, Electronegativity, and Molecular Structure

Exam 2 Study Guide: Ions, Compounds, Electronegativity, and Molecular Structure

Overview

This study guide summarizes the key topics and skills required for Exam 2 in a GOB Chemistry course. The focus is on ions, chemical formulas, polyatomic ions, molecular compounds, valence electrons, electronegativity, bond polarity, and molecular shapes.

Ions and Their Symbols

Identifying Ions from Protons and Electrons

Understanding how to write the symbol for an ion requires knowing the number of protons (atomic number) and electrons (charge).

Ion Symbol: The chemical symbol is followed by the charge in superscript (e.g., Na+, Cl-).

Determining Charge: $\text{Charge} = \text{Number of protons} - \text{Number of electrons}$

Example: An atom with 11 protons and 10 electrons is Na+ (sodium ion).

Writing Chemical Formulas

Formulas for Ionic Compounds

Compounds formed between pairs of ions must be electrically neutral.

• Steps:

Write the symbol and charge for each ion.

Balance the charges so the total positive and negative charges are equal.

Write the formula using subscripts to indicate the number of each ion.

Example: Magnesium chloride: Mg2+ and Cl- combine to form $\text{MgCl}_2$.

Formulas for Compounds with Metal Cations and Nonmetal Anions

• Metal cations (e.g., Na+, Ca2+) combine with nonmetal anions (e.g., Cl-, O2-).

• Example: Calcium oxide: Ca2+ and O2- combine to form $\text{CaO}$.

Polyatomic Ions

Naming Polyatomic Anions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge.

• Common Polyatomic Anions:

  • Nitrate: NO3-

  • Sulfate: SO42-

  • Phosphate: PO43-

  • Hydroxide: OH-

• Example: Sodium sulfate: Na+ and SO42- combine to form $\text{Na}_2\text{SO}_4$.

Naming Nonmetallic Molecular Compounds

Rules for Naming Molecular Compounds

Molecular compounds are formed between nonmetals. Prefixes indicate the number of atoms.

• Prefixes: mono-, di-, tri-, tetra-, penta-, etc.

• First element: Use the full element name.

• Second element: Use the root plus "-ide" ending.

• Example: CO2 is carbon dioxide; N2O is dinitrogen monoxide.

Valence Electrons in Simple Molecules

Determining Total Number of Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom and are important for bonding.

• Steps:

  1. Identify the group number for each atom in the molecule.

  2. Multiply the number of atoms by their group number (for main group elements).

  3. Add up the total for all atoms in the molecule.

• Example: For H2O: H (1 valence electron each) × 2 + O (6 valence electrons) = 8 valence electrons.

Electronegativity

Definition and Trends

Electronegativity is the ability of an atom to attract electrons in a chemical bond.

• Definition: Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.

• Periodic Trend: Electronegativity increases across a period (left to right) and decreases down a group (top to bottom).

• Example: Fluorine is the most electronegative element.

Bond Polarity and Electronegativity Differences

Predicting Bond Type

The difference in electronegativity between two atoms determines the type of bond formed.

• Bond Types:

  • Nonpolar Covalent: Electronegativity difference < 0.5

  • Polar Covalent: Electronegativity difference between 0.5 and 1.7

  • Ionic: Electronegativity difference > 1.7

• Example: H–Cl bond: H (2.1), Cl (3.0); difference = 0.9 → polar covalent.

Bond Dipoles

Indicating Positive and Negative Ends

A bond dipole shows the direction of electron density in a polar bond.

• Arrow Notation: Draw an arrow pointing toward the more electronegative atom; the tail indicates the positive end.

• Example: In H–Cl, the arrow points from H to Cl.

VSEPR Theory and Molecular Shapes

Predicting Shapes of Simple Molecules

Valence Shell Electron Pair Repulsion (VSEPR) Theory is used to predict the geometry of molecules based on electron pair repulsion.

• Basic Shapes:

  • Linear: 2 electron groups (e.g., CO2)

  • Trigonal planar: 3 electron groups (e.g., BF3)

  • Tetrahedral: 4 electron groups (e.g., CH4)

  • Bent: 2 bonds + 2 lone pairs (e.g., H2O)

• Example: Water (H2O) is bent due to two lone pairs on oxygen.

Table: Molecular Shapes Predicted by VSEPR Theory

Electron Groups	Bonding Pairs	Lone Pairs	Shape	Example
2	2	0	Linear	CO2
3	3	0	Trigonal planar	BF3
4	4	0	Tetrahedral	CH4
4	2	2	Bent	H2O

Additional info: Some details, such as specific electronegativity values and attachments referenced in the original notes, were inferred or expanded for completeness.