General Chemistry Study Guide: Periodic Trends, Bonding, and Molecular Structure

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

Periodic Trends and Periodicity

Periodic Trends

Periodic trends are predictable patterns in the properties of elements as you move across periods (rows) or down groups (columns) in the periodic table.

Definition: A specific pattern or direction in a property observed across a period or down a group.
Examples: Atomic radius, ionization energy, and electronegativity.
Key Takeaway: Periodic trends help predict how properties change across the periodic table.

Periodicity

Periodicity refers to the repeating nature of these trends as you move through the periodic table.

Definition: The repetition of trends or patterns at regular intervals in the periodic table.
Example: Atomic radius decreases across a period and increases down a group; elements in the same group (e.g., Li, Na, K) have similar chemical properties.
Key Takeaway: Periodicity explains why trends recur in each period or group.

Atomic Structure and Charges

Positive and Negative Charges

Positive Charge (Cation): Occurs when an atom loses electrons, resulting in more protons than electrons. Example: $\mathrm{Na} \rightarrow \mathrm{Na}^+ + e^-$
Negative Charge (Anion): Occurs when an atom gains electrons, resulting in more electrons than protons. Example: $\mathrm{Cl} + e^- \rightarrow \mathrm{Cl}^-$
Key Takeaway: Positive charge arises from electron loss; negative charge from electron gain.

Outer Electrons (Valence Electrons)

Definition: Electrons in the outermost energy level (shell) of an atom.
Importance: Valence electrons determine chemical bonding and reactivity.
Shielding: Inner electrons shield outer electrons from the full positive charge of the nucleus, reducing the effective nuclear charge felt by valence electrons.
Key Takeaway: Outer electrons dictate how atoms interact and bond.

Effective Nuclear Charge ($Z_{\mathrm{eff}}$)

Definition: The net positive charge experienced by an electron in an atom, accounting for both nuclear attraction and electron shielding.
Formula: $Z_{\mathrm{eff}} = Z - S$ Where $Z$ = atomic number (number of protons), $S$ = shielding constant (number of inner electrons).
Trend: $Z_{\mathrm{eff}}$ increases across a period as protons are added but electrons enter the same shell.
Key Takeaway: $Z_{\mathrm{eff}}$ explains trends in atomic size, ionization energy, and other properties.

Atomic and Ionic Radii

Atomic Radius

Definition: The distance from the nucleus to the outermost electron shell of an atom.
Trend: Decreases across a period (left to right), increases down a group (top to bottom).
Key Takeaway: Atomic radius helps predict bonding and reactivity.

Size of Ions (Ionic Radius)

Definition: The distance from the nucleus to the outermost electron of an ion.
Cations: Smaller than their neutral atoms due to loss of electrons and increased $Z_{\mathrm{eff}}$.
Anions: Larger than their neutral atoms due to electron gain and increased electron-electron repulsion.
Key Takeaway: Ionic size affects packing in solids, interactions in solutions, and chemical reactivity.

Formation of Ions and Electron Configurations

Formation of Ions

Definition: Atoms gain or lose electrons to form ions (charged particles).
Cations: Formed by loss of electrons (e.g., $\mathrm{Li}: 1s^2 2s^1 \rightarrow \mathrm{Li}^+: 1s^2$).
Anions: Formed by gain of electrons (e.g., $\mathrm{O}: 1s^2 2s^2 2p^4 \rightarrow \mathrm{O}^{2-}: 1s^2 2s^2 2p^6$).
Stability: Atoms form ions to achieve noble gas electron configurations.

Electron Configurations

Definition: The arrangement of electrons in an atom's orbitals (e.g., $1s^2 2s^2 2p^4$ for oxygen).
Key Takeaway: Electron configuration determines chemical properties and reactivity.

Table: Electron Configurations of Common Ions

Group	Ion	Electron Configuration
1A	Li+, Na+, K+	[He], [Ne], [Ar]
2A	Be2+, Mg2+, Ca2+	[He], [Ne], [Ar]
6A	O2−, S2−	[Ne], [Ar]
7A	F−, Cl−	[Ne], [Ar]

Ionization Energy

Ionization Energy ($E_i$)

Definition: The energy required to remove the highest-energy electron from a gaseous atom or ion.
Types:
- First Ionization Energy ($E_{i1}$): $\mathrm{M}(g) \rightarrow \mathrm{M}^+(g) + e^-$
- Second Ionization Energy ($E_{i2}$): $\mathrm{M}^+(g) \rightarrow \mathrm{M}^{2+}(g) + e^-$
- Third Ionization Energy ($E_{i3}$): $\mathrm{M}^{2+}(g) \rightarrow \mathrm{M}^{3+}(g) + e^-$
Trend: Increases across a period, decreases down a group.
Successive Ionization Energies: Each is higher than the previous due to increased attraction between electrons and nucleus.
Key Takeaway: Higher ionization energy means electrons are harder to remove; important for predicting reactivity.

Table: First, Second, and Third Ionization Energies (kJ/mol) for Third-Row Elements

Element	$E_{i1}$	$E_{i2}$	$E_{i3}$
Na	496	4562	6912
Mg	738	1451	7733
Al	578	1817	2745
Si	787	1577	3231
P	1012	1903	2912
S	1000	2251	3361
Cl	1251	2297	3822
Ar	1520	2665	3931

Electron Affinity

Electron Affinity ($E_a$)

Definition: The energy change when a neutral atom in the gas phase gains an electron to form an anion.
Equation: $\mathrm{M}(g) + e^- \rightarrow \mathrm{M}^-(g)$
Trend: Generally increases (becomes more negative) across a period; halogens have high electron affinities.
Key Takeaway: Electron affinity helps predict which elements form anions and their reactivity.

Chemical Bonding

Ionic Bonds

Definition: Chemical bonds formed by the transfer of electrons from one atom to another, resulting in oppositely charged ions held together by electrostatic attraction.
Example: $\mathrm{Na} + \mathrm{Cl} \rightarrow \mathrm{Na}^+ + \mathrm{Cl}^- \rightarrow \mathrm{NaCl}$
Properties: High melting/boiling points, conduct electricity when dissolved in water.
Key Takeaway: Ionic bonds create stable compounds by balancing positive and negative charges.

Covalent Bonds

Definition: Bonds formed when two atoms share one or more pairs of electrons.
Example: In $\mathrm{H}_2\mathrm{O}$, each hydrogen shares an electron with oxygen.
Properties: Usually between nonmetals; form molecules.
Key Takeaway: Covalent bonds result in stable molecules with shared electrons.

Bond Length and Bond Energy

Bond Length: The average distance between the nuclei of two bonded atoms.
Bond Energy: The energy required to break a bond between two atoms (measured in $\mathrm{kJ/mol}$).
Relationship: As bond order increases (single < double < triple), bond length decreases and bond strength increases.
Example: $\mathrm{H}-\mathrm{H}$ bond energy in $\mathrm{H}_2$ is about 436 $\mathrm{kJ/mol}$.

Table: Average Bond Dissociation Energies ($\mathrm{kJ/mol}$)

Bond	Energy	Bond	Energy
H–H	436	C–H	410
O–H	460	Cl–Cl	243
C=C	611	C≡C	835
C=O	732	O=O	498

Bond Polarity and Electronegativity

Polar Covalent Bonds

Definition: Covalent bonds where electrons are shared unequally due to differences in electronegativity.
Partial Charges: The more electronegative atom gains a partial negative charge ($\delta^-$), the other a partial positive charge ($\delta^+$).
Example: In $\mathrm{H}_2\mathrm{O}$, O is more electronegative than H, making the O–H bonds polar.

Electronegativity (EN)

Definition: The ability of an atom to attract shared electrons in a bond.
Trend: Increases across a period and up a group; fluorine is the most electronegative element.
Pauling Scale: Assigns numerical values to elements based on their electron-attracting power.

Nonpolar Bonds

Definition: Bonds where electrons are shared equally (identical or similar electronegativities).
Example: $\mathrm{Cl}_2$ molecule has a nonpolar bond.
Key Takeaway: Nonpolar molecules are electrically neutral and often insoluble in water.

Lewis Structures and the Octet Rule

Lewis Symbols and Structures

Lewis Symbols: Dots around an element's symbol represent valence electrons.
Lewis Structures: Show all valence electrons in a molecule, with lines for bonds and dots for lone pairs.
Rules for Drawing Lewis Structures:
1. Count total valence electrons (adjust for charge).
2. Arrange atoms (hydrogen always outside; central atom is least electronegative).
3. Connect atoms with single bonds.
4. Complete octets for outer atoms, then central atom.
5. Form double/triple bonds if central atom lacks an octet (except for H, Be, B).

The Octet Rule

Definition: Main-group atoms tend to be surrounded by eight valence electrons (like noble gases).
Exceptions: Hydrogen (needs 2 electrons), boron (can have 6), and expanded octets for elements in period 3 or beyond.
Key Takeaway: The octet rule guides bonding and molecular structure.

Table: Lewis Symbols, Valence Electrons, and Bonds

Element	Group	Valence Electrons	Typical Bonds	Example
B	3A	3	3	BH3
C	4A	4	4	CH4
N	5A	5	3	NH3
O	6A	6	2	H2O
F	7A	7	1	HF
Ne	8A	8	0	Ne

Bond Order and Resonance

Resonance

Definition: When a molecule can be represented by two or more valid Lewis structures differing only in electron arrangement.
Resonance Hybrid: The actual structure is an average of all resonance forms.
How to Draw: Move only electrons (not atoms); use double-headed arrows ($\leftrightarrow$) between structures.
Stabilization: Resonance delocalizes electrons, stabilizing the molecule.

Bond Order

Definition: The number of chemical bonds between a pair of atoms.
Simple Molecules: Single bond = 1, double bond = 2, triple bond = 3.
Resonance Structures: $\text{Bond Order} = \dfrac{\text{Total number of bonds}}{\text{Number of bond locations}}$
Example: In ozone ($\mathrm{O}_3$), bond order = $\dfrac{3}{2} = 1.5$
Key Takeaway: Higher bond order means stronger, shorter bonds.

Additional info: This guide covers core concepts from periodic trends, atomic structure, ionization energy, electron affinity, chemical bonding, Lewis structures, resonance, and bond order, as relevant to a first-semester general chemistry course.