Introduction to Chemistry: Exam 1 Study Guide

Exam Instructions and Academic Integrity

Exam Protocols

This section outlines the rules and expectations for students during the chemistry exam. Adhering to these protocols ensures fairness and academic honesty.

• Allowed Materials: Pencil/eraser and a non-programmable calculator. No cell phones, laptops, translators, notes, or books are permitted.

• Answer Sheet: Only answers on the bubble sheet will be graded. Completely darken the bubbles for your answers.

• Academic Misconduct: Any form of cheating or unauthorized collaboration is considered academic dishonesty and will be penalized according to university policy.

• Reference Materials: A periodic table and other reference sheets are provided within the exam.

Atomic Structure and Electron Configuration

Atomic Structure

Atoms consist of a nucleus (protons and neutrons) surrounded by electrons in defined energy levels or shells.

• Valence Electrons: Electrons in the outermost shell, important for chemical bonding and reactivity.

• Electron Shells: Labeled as n = 1, 2, 3, etc., with each shell holding a specific maximum number of electrons.

• Electron Configuration: The arrangement of electrons in an atom's orbitals, following the Aufbau principle, Pauli exclusion principle, and Hund's rule.

Example: The electron configuration for Sulfur (S, atomic number 16) is:

$1s^2 2s^2 2p^6 3s^2 3p^4$

Aufbau Model: Electrons fill orbitals in order of increasing energy, as shown in the provided diagram.

Electron Configuration Notation

• Noble Gas Notation: Use the previous noble gas in brackets to simplify configuration. For example, for Sulfur: $[Ne] 3s^2 3p^4$

• Valence Shell: The highest principal quantum number (n) in the configuration indicates the valence shell.

Chemical Bonding and Molecular Structure

Types of Chemical Bonds

Chemical bonds form when atoms share or transfer electrons to achieve stable electron configurations.

• Covalent Bonds: Electrons are shared between atoms. Can be nonpolar (equal sharing) or polar (unequal sharing).

• Ionic Bonds: Electrons are transferred from one atom to another, resulting in oppositely charged ions that attract each other.

• Bond Polarity: Determined by the difference in electronegativity ($\Delta EN$) between bonded atoms.

Example Table: Electronegativity and Bond Type

Bond	Electronegativity Difference ($\Delta EN$)	Type of Bond
H–H	0.0	Nonpolar Covalent
C–H	0.4	Considered Nonpolar
Cl–H	0.9	Polar Covalent
F–H	1.9	Polar Covalent (close to ionic)

Key Points:

• A perfectly covalent bond is perfectly nonpolar.

• $\Delta EN$ greater than around 1.9 is considered an ionic bond.

Lewis Structures and Resonance

Lewis structures represent the arrangement of valence electrons in molecules. Resonance occurs when more than one valid Lewis structure can be drawn for a molecule.

• Lone Pairs: Non-bonding pairs of electrons on an atom.

• Bond Pairs: Shared pairs of electrons between atoms.

• Resonance: The actual structure is a hybrid of all possible Lewis structures.

Example: The carboxylate group in aspirin exhibits resonance, with the negative charge delocalized over two oxygen atoms.

Octet Rule and Exceptions

• Atoms tend to form bonds to achieve eight electrons in their valence shell (octet rule).

• Hydrogen is an exception, achieving a stable configuration with two electrons (helium configuration).

Periodic Table and Periodic Trends

Periodic Table Organization

The periodic table arranges elements by increasing atomic number and groups elements with similar properties into columns.

• Groups: Vertical columns; elements in the same group have similar chemical properties.

• Periods: Horizontal rows; properties change progressively across a period.

Periodic Trends

• Electronegativity: Increases across a period (left to right) and decreases down a group. Fluorine (F) is the most electronegative element.

• Valence Electrons: The number of valence electrons increases across a period.

Molecular Geometry and Bond Angles

Bond Angles

The geometry of a molecule determines the angles between bonds. For example, the bond angle in a linear molecule is 180°, while in a tetrahedral molecule it is approximately 109.5°.

• Example: The C–O–C bond angle in the aspirin molecule is closest to 109.5°, indicating a tetrahedral geometry around the central carbon atom.

Practice Questions and Key Concepts

Sample Multiple Choice Questions

• Polarity: Which molecule is least polar? (e.g., HCl, H2O, CH4, NH3)

• Electron Configuration: Which configuration matches a given element?

• Valence Electrons: How many valence electrons are in a given atom or ion?

• Lewis Structures: Which is the correct Lewis structure for a molecule?

• Bond Types: Identify whether a bond is covalent, ionic, or binary.

Reference Materials Provided

• Electronegativity Table: Values for common elements and guidelines for bond classification.

• Aufbau Diagram: Order of orbital filling for electron configurations.

• Periodic Table: Full table for reference during the exam.

Summary Table: Key Chemistry Concepts

Concept	Definition/Key Point	Example
Valence Electrons	Electrons in the outermost shell	Carbon has 4 valence electrons
Octet Rule	Atoms tend to have 8 electrons in their valence shell	Oxygen forms two bonds to complete its octet
Electronegativity	Ability of an atom to attract electrons	F (4.0) > O (3.5) > N (3.0)
Covalent Bond	Shared pair of electrons	H2O
Ionic Bond	Transfer of electrons	NaCl
Resonance	Delocalization of electrons across multiple structures	Carboxylate ion

Additional info:

• Some context and explanations have been expanded for clarity and completeness, as the original exam questions are brief.

• All key terms and concepts are standard for an Introduction to Chemistry college course.