BackGeneral Chemistry Exam Master Study Guide: Atomic Structure, Bonding, and Molecular Geometry
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Atomic Structure
Subatomic Particles
The atom is the smallest unit of an element that retains its chemical properties. Atoms are composed of three main subatomic particles:
Proton: Positively charged (+1), located in the nucleus.
Neutron: Neutral (0 charge), located in the nucleus.
Electron: Negatively charged (−1), found in the electron cloud surrounding the nucleus.
Atomic number is the number of protons in an atom. Mass number is the sum of protons and neutrons.
Example: Carbon (C): Atomic number = 6, Protons = 6, Neutrons = 6, Electrons = 6, Mass number = 12.
Isotopes are atoms of the same element with different numbers of neutrons (e.g., C-12 vs. C-14).
Periodic Table Basics
Organization and Groups
The periodic table arranges elements by increasing atomic number. Rows are called periods, and columns are called groups or families. Elements in the same group have similar chemical properties due to the same number of valence electrons.
Metals: Left side of the table.
Nonmetals: Right side of the table.
Metalloids: Located along the staircase line.
Important Groups:
Group 1: Alkali metals (+1 charge)
Group 2: Alkaline earth metals (+2 charge)
Group 17: Halogens (−1 charge)
Group 18: Noble gases (very stable, inert)
Valence Electrons
Determining Valence Electrons
Valence electrons are the outermost electrons involved in bonding. For main group elements, the group number indicates the number of valence electrons.
Group | Valence Electrons |
|---|---|
1 | 1 |
2 | 2 |
13 | 3 |
14 | 4 |
15 | 5 |
16 | 6 |
17 | 7 |
18 | 8 |
Example: Nitrogen (Group 15) has 5 valence electrons; Oxygen (Group 16) has 6 valence electrons.
Ions
Cations and Anions
An ion is an atom or molecule with a net electric charge due to the loss or gain of electrons.
Cation: Positively charged ion (lost electrons).
Anion: Negatively charged ion (gained electrons).
Examples: Na+ (sodium ion), Mg2+ (magnesium ion), Cl− (chloride ion), O2− (oxide ion).
Common Ion Charges by Group
Group | Elements | Common Charge |
|---|---|---|
1 | Li, Na, K, Rb, Cs | +1 |
2 | Be, Mg, Ca, Sr, Ba | +2 |
13 | Al | +3 |
17 | F, Cl, Br, I | −1 |
16 | O, S | −2 |
15 | N, P | −3 |
Polyatomic Ions
Ion | Name |
|---|---|
NH4+ | Ammonium |
OH− | Hydroxide |
NO3− | Nitrate |
NO2− | Nitrite |
SO42− | Sulfate |
SO32− | Sulfite |
CO32− | Carbonate |
HCO3− | Bicarbonate |
PO43− | Phosphate |
C2H3O2− | Acetate |
Ionic Compounds
Naming and Writing Formulas
Ionic compounds are formed between metals and nonmetals. The naming convention is:
Name the metal (cation) first.
Name the nonmetal (anion) second, changing the ending to -ide.
Examples: NaCl (sodium chloride), MgO (magnesium oxide), CaF2 (calcium fluoride), KBr (potassium bromide).
Writing Ionic Formulas
Write the symbols for the ions.
Determine their charges.
Criss-cross the charges to balance the formula.
Reduce subscripts if necessary.
Example: Magnesium chloride: Mg2+ and Cl− → MgCl2
Example: Aluminum oxide: Al3+ and O2− → Al2O3
Covalent Compounds
Naming Covalent Compounds
Covalent compounds are formed between nonmetals. Prefixes are used to indicate the number of each atom:
Number | Prefix |
|---|---|
1 | mono- |
2 | di- |
3 | tri- |
4 | tetra- |
5 | penta- |
6 | hexa- |
7 | hepta- |
8 | octa- |
9 | nona- |
10 | deca- |
Examples: CO (carbon monoxide), CO2 (carbon dioxide), SO2 (sulfur dioxide), P2O4 (diphosphorus tetroxide), N2O5 (dinitrogen pentoxide).
Lewis Structures
Drawing Lewis Structures
Lewis structures represent the arrangement of valence electrons in a molecule. Steps:
Count total valence electrons.
Draw the skeleton structure of the molecule.
Connect atoms with single bonds.
Fill octets with lone pairs.
Create double or triple bonds if needed to satisfy the octet rule.
Example: CO2: C = 4 electrons, O = 6 each, total = 16 electrons. Structure: O=C=O.
VSEPR Shapes and Molecular Geometry
VSEPR Theory
Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules based on electron group repulsion.
Shape | Electron Groups | Example | Bond Angle |
|---|---|---|---|
Linear | 2 | CO2 | 180° |
Trigonal planar | 3 | BF3 | 120° |
Tetrahedral | 4 | CH4 | 109.5° |
Trigonal pyramidal | 4 (1 lone pair) | NH3 | ~107° |
Bent | 4 (2 lone pairs) | H2O | ~104.5° |
Polarity
Polar vs. Nonpolar Molecules
Polar molecules: Have unequal sharing of electrons, resulting in a dipole moment (e.g., H2O, NH3).
Nonpolar molecules: Have symmetrical shapes and dipoles cancel out (e.g., CO2, CH4, O2).
Moles and Molar Mass
Avogadro's Number and Calculations
One mole contains particles (Avogadro's number).
Formula:
Example: CO2 molar mass = 44 g/mol; 44 g CO2 = 1 mole.
Dilution
Dilution Equation
To dilute a solution, use the equation:
Example: 2.0 M solution diluted to 1.0 M, original volume = 50 mL: mL.
Bond Types
Ionic vs. Covalent Bonds
Ionic bond: Involves electron transfer (metal + nonmetal).
Covalent bond: Involves electron sharing (nonmetal + nonmetal).
Single bond: 1 shared pair of electrons.
Double bond: 2 shared pairs.
Triple bond: 3 shared pairs.
Quick Chemistry Rules
Metal + nonmetal → ionic compound.
Nonmetal + nonmetal → covalent compound.
Second element in ionic compounds ends in -ide.
Covalent compounds use prefixes.
Use the periodic table to determine charges quickly.
Practice Problems
Naming Ionic Compounds
NaCl: sodium chloride
KBr: potassium bromide
MgO: magnesium oxide
CaF2: calcium fluoride
Li3P: lithium phosphide
K3N: potassium nitride
Rb2O: rubidium oxide
CsCl: cesium chloride
Naming Covalent Compounds
CO2: carbon dioxide
NO: nitrogen monoxide
SO2: sulfur dioxide
N2O5: dinitrogen pentoxide
P2O4: diphosphorus tetroxide
Writing Formulas from Names
Sodium chloride: NaCl
Potassium bromide: KBr
Magnesium oxide: MgO
Calcium phosphide: Ca3P2
Potassium nitride: K3N
Barium sulfide: BaS
Lewis Structure Practice
CO2: O=C=O
H2O: H–O–H (with two lone pairs on O)
NH3: N with three H atoms and one lone pair
CH4: Tetrahedral structure
SO2: S with two O atoms, one double bond, one single bond, and lone pairs
Polarity Practice
Molecule | Polarity |
|---|---|
CO2 | Nonpolar |
H2O | Polar |
NH3 | Polar |
CH4 | Nonpolar |
BF3 | Nonpolar |
Additional info: This guide covers foundational concepts in atomic structure, periodic trends, bonding, molecular geometry, and basic calculations, suitable for first-semester general chemistry students.
