Chapter 3: Compounds – How Elements Combine

3.1 Electron Arrangements and the Octet Rule

How Elements Combine
- Elements combine by transferring, accepting, or sharing electrons.
- Why Elements Combine: To achieve greater stability, typically by fulfilling the octet rule.
- The arrangement of electrons fundamentally drives compound formation.
Electron Cloud and Energy Levels
- The exact location of an electron is quantum mechanically uncertain, but electrons reside within a space called the electron cloud.
- Electrons possess energy due to their charge and constant motion.
- Energy Levels: Electrons occupy specific, discrete energy levels (analogous to steps on a staircase).
  - Electrons cannot exist in between these levels.
  - Electrons generally occupy the lowest available energy level first.
  - The lowest energy levels are found closest to the nucleus.
  - Unlike a real staircase, energy levels get closer together as they move farther from the nucleus.
Maximum Electrons per Energy Level
- The maximum number of electrons an energy level can hold is given by the formula $2n^2$ , where $n$ is the number of the energy level.
  - For the first energy level ( $n = 1$ ), the maximum is $2(1)^2 = 2$ electrons.
  - For the second energy level ( $n = 2$ ), the maximum is $2(2)^2 = 8$ electrons.
  - For the third energy level ( $n = 3$ ), the maximum is $2(3)^2 = 18$ electrons.
Determining Energy Levels (n)
- For main group elements, $n$ corresponds to the row/period number.
- For transition metals, $n$ is typically one less than the row/period number ((row/period #) - 1).
Valence Shell and Valence Electrons
- The highest energy level containing electrons is called the valence shell (or valence level).
- Electrons in the valence shell are called valence electrons.
- Valence electrons are farthest from the nucleus and are the primary participants in chemical reactions because of their higher energy and accessibility.
Finding Valence Electron Information
- For main-group elements, the number of valence electrons is indicated by the group number (XA notation).
- Elements in the same group share the same number of valence electrons, leading to similar chemical properties.
  - Example: Boron (B) and aluminum (Al) are in Group 3A, both having three valence electrons.
Chemical Reactions - Electron Style
- A chemical reaction involves substances reacting to form new ones through the gaining, losing, or sharing of electrons.
- Valence electrons participate in these reactions because they are the outermost and highest energy electrons, making them the most reactive.
Periodic Table and Electron Arrangement
- Groups (columns) indicate the number of valence electrons for main-group elements.
- Periods (rows) indicate the outermost energy level ( $n$ ) containing electrons.
Electron Configuration Examples
- Magnesium (Mg): $12$ electrons
  - $n=1$ : $2$ electrons
  - $n=2$ : $8$ electrons
  - $n=3$ : $2$ electrons (valence electrons)
  - $n=4$ : $0$ electrons
- Sulfur (S): $16$ electrons
  - $n=1$ : $2$ electrons
  - $n=2$ : $8$ electrons
  - $n=3$ : $6$ electrons (valence electrons)
  - $n=4$ : $0$ electrons
The Octet Rule
- Noble Gases (Group 8A): Highly unreactive (inert) under normal conditions.
  - They possess eight valence electrons (except helium, which has two, a stable duet).
  - This configuration represents a highly stable atomic state.
- Octet Rule: Most atoms react with other atoms to achieve eight electrons in their valence shell, mimicking the stability of noble gases.
- Isoelectronic: Describes atoms or ions that have the same number of electrons as another element or ion.
- To achieve an octet, atoms can:
  - Give away electrons (forming cations).
  - Accept or take electrons (forming anions).
  - Share electrons (forming covalent bonds).
- These different mechanisms lead to distinct types of compounds: ionic and covalent.

How to Get a Full Valence Shell

Part 1: Electron Transfer (Ionic Bonds)
- Electrons are transferred between elements to form ions.
- Resulting attraction between metal cation and non-metal anion forms an ionic bond.
- Compounds formed are ionic compounds.
Part 2: Electron Sharing (Covalent Bonds)
- Electrons are shared between non-metals to form a covalent bond.
- Compounds formed are covalent compounds.
- Lewis structures are representations of molecules that illustrate electron sharing and arrangement.

3.2 In Search of an Octet, Part 1: Ion Formation

Ions: Charged atoms formed when an atom gains or loses electrons, resulting in an unequal number of protons and electrons.
- Anions: Negatively charged ions formed when atoms gain electrons.
- Cations: Positively charged ions formed when atoms lose electrons.
Trends in Ion Formation (Main Group Elements)
- Metals (Groups 1A, 2A, 3A) tend to lose electrons to form cations.
  - The cation charge is equal to the group number.
  - Example: Group 1A elements form $+1$ ions; Group 2A elements form $+2$ ions.
- Non-metals (Groups 5A, 6A, 7A) tend to gain electrons to form anions.
  - The anion charge is equal to eight minus the group number (8 - Group #).
  - Example: Group 7A elements form $(8-7) = -1$ ions; Group 6A elements form $(8-6) = -2$ ions.
- Elements in Group 4A often achieve stability through different bonding mechanisms (e.g., covalent bonding).
Trends in Ion Formation (Transition Metals)
- Transition metals typically form cations.
- Many transition metals can form more than one type of ion with different charges.
- Their ionic charge cannot be directly determined from their group number.
- Naming: Roman numerals are used in parentheses to state the charge of a transition metal ion.
  - Example: $\text{Fe}^{3+}$ is iron (III) ion.
Polyatomic Ions
- Consist of a group of nonmetals that are covalently bonded but carry a net charge, acting as a single ion.
- The charge applies to the entire group of atoms.
- Many polyatomic ions end in "-ate" (e.g., sulfate, nitrate) or "-ite" (one fewer oxygen atom, e.g., sulfite, nitrite).
- Exceptions: Hydroxide ( $\text{OH}^-$ ), Hydronium ( $\text{H}_3\text{O}^+$ ), Cyanide ( $\text{CN}^-$ ).
Naming Ions
- Metal Ions: Add the word "ion" to the metal's name.
  - Example: $\text{Na}^+$ is sodium ion.
- Transition Metal Ions with Variable Charge: Use a Roman numeral in parentheses after the metal's name to specify the charge.
  - Example: $\text{Fe}^{2+}$ is iron (II) ion.
- Nonmetal Anions (Monatomic): Replace the last few letters of the element name with the suffix "-ide", then add "ion".
  - Example: Fluorine becomes fluoride ion ( $\text{F}^-$ ) from $fluor + -ide$ .

3.3 Ionic Compounds—Electron Give and Take

Formation of Ionic Compounds
- Occurs when a metal (forms cation) and a nonmetal (forms anion) combine.
- Electrons are transferred from the metal atom to the nonmetal atom.
- The electrostatic attraction between the newly formed oppositely charged ions is called an ionic bond.
- The result is an ionic compound.
Formula of Ionic Compounds
- The ratio of cations and anions combines so that the overall net charge of the compound is zero.
- Subscripts are used to indicate the number of each particle in the formula.
- Charges of the individual ions are not shown in the final chemical formula.
Steps to Check a Chemical Formula
1. Determine the charge of the ions (use periodic table for main-group, memorized or given for polyatomic/transition).
2. Combine the ions in a ratio such that the total positive charge cancels the total negative charge, summing to zero.
3. Verify the formula by ensuring charge neutrality with the determined subscripts.
Examples of Ionic Compound Formation
- $\text{Mg}^{2+}$ and $\text{F}^-$ ions: Forms $\text{MgF}_2$ ( $1 \times (+2) + 2 \times (-1) = 0$ ).
- $\text{Ba}^{2+}$ and $\text{CO}3^{2-}$ ions: Forms $\text{BaCO}3$ ( $1 \times (+2) + 1 \times (-2) = 0$ ).
- Lithium ( $\text{Li}^+$ ) and Hydroxide ( $\text{OH}^-$ ) ions: Forms $\text{LiOH}$ ( $1 \times (+1) + 1 \times (-1) = 0$ ).
Naming Ionic Compounds
- Combine the names of the two ions.
- Cation name first (element name).
- Anion name second (element name with appropriate ending, typically "-ide" for monatomic).
- The word "ion" is omitted.
- Crucially: Do not denote the number of cations or anions in the name (e.g., $\text{NaCl}$ is sodium chloride, not sodium monochloride).
Naming Cations
- Fixed Charge Metals (Main Group): Simply use the name of the element.
  - Example: $\text{Cs}^+$ is Cesium.
- Variable Charge Metals (Transition Metals and some main group like Sn, Pb): Use the element name followed by a Roman numeral in parentheses indicating its charge (Stock System).
  - Example: $\text{Fe}^{3+}$ is Iron (III).
- Polyatomic Cations: Use their memorized names.
  - Example: $\text{NH}_4^+$ is Ammonium.
Naming Anions
- Monoatomic Anions: Take the stem of the element name and add the suffix "-ide". The charge is determined by Group \text{#(A)} - 8.
  - Example: $\text{F}^-$ is Fluoride; $\text{S}^{2-}$ is Sulfide; $\text{H}^-$ is Hydride; $\text{N}^{3-}$ is Nitride.
- Polyatomic Anions: Use their memorized names (many end in "-ate" or "-ite").
Examples: Formula to Name & Name to Formula
- $\text{Ca}^{2+}$ + $\text{Cl}^-$ $\rightarrow \text{CaCl}_2$ (calcium chloride)
- $\text{O}^{2-}$ + $\text{Li}^+ \rightarrow \text{Li}_2\text{O}$ (lithium oxide)
- $\text{Al}^{3+}$ + $\text{O}^{2-} \rightarrow \text{Al}2\text{O}3$ (aluminum oxide)
- $\text{F}^-$ + $\text{K}^+ \rightarrow \text{KF}$ (potassium fluoride)
- $\text{Fe}^{3+}$ + $\text{O}^{2-} \rightarrow \text{Fe}2\text{O}3$ (iron (III) oxide)

3.4 In Search of an Octet, Part 2: Covalent Bond Formation

Covalent Bonding Principles
- Nonmetals combine by sharing valence electrons to achieve an octet (or duet for hydrogen).
- This electron sharing forms a covalent bond.
- The shared valence electrons are considered to belong to both bonded atoms.
- The result is a covalent compound, whose smallest fundamental unit is a molecule (e.g., $\text{H}2\text{O}$ , $\text{CO}2$ ).
Number of Covalent Bonds
- An atom typically forms a number of covalent bonds equal to the number of electrons it needs to complete its octet (or duet).
Electron-Dot Symbols (Lewis Structures)
- Developed by G. N. Lewis to visualize electron sharing.
- An electron-dot symbol consists of the elemental symbol plus a dot for each valence electron.
- Unpaired electrons on an electron-dot symbol typically lead to bond formation.
Bonding and Lone Pairs
- Shared pair of electrons (Bonding pair/Bond): Represented by a line or two dots connecting two electron-dot symbols. Both atoms effectively "own" these electrons.
- Lone pairs of electrons (Unshared pair): Electrons located on only one atom in a Lewis structure, not involved in bonding.
Typical Bonding Patterns for Nonmetals
- Hydrogen (H): Forms $1$ bond (to achieve $2$ electrons, a duet).
- Group 7A (Halogens like F, Cl, Br, I): Form $1$ covalent bond.
- Group 6A (Oxygen): Forms $2$ covalent bonds.
- Group 5A (Nitrogen): Forms $3$ covalent bonds.
- Group 4A (Carbon): Forms $4$ covalent bonds.
Types of Covalent Bonds
- Single bonds: Sharing one pair of electrons between two atoms.
- Double bonds: Sharing two pairs of electrons between two atoms.
- Triple bonds: Sharing three pairs of electrons between two atoms.
- In all cases (single, double, triple), carbon still forms four effective bonds to complete its octet.
Formulas and Structures of Covalent Compounds
- Molecular formula: Identifies the types and total number of atoms in a molecule (e.g., glucose is $\text{C}6\text{H}{12}\text{O}_6$ ).
 - This formula does not show the arrangement or structure.
- Lewis structure: A visual representation that shows the arrangement of atoms and electrons (bonding and lone pairs) within a molecule, revealing its connectivity.
 - Example: Methane ( $\text{CH}_4$ ) shows carbon making four single bonds to hydrogen.
Drawing Lewis Structures for Covalent Compounds - Steps
1. Count Total Valence Electrons: Sum the valence electrons from all atoms in the molecule. Adjust the total if the molecule is an ion (add electrons for negative charge, subtract for positive charge).
2. Draw Skeleton Structure: Arrange the atoms and draw single bonds between them.
  - The central atom is usually the least electronegative (never hydrogen).
  - Hydrogens are always terminal (on the outside).
  - Each single bond uses $2$ electrons.
3. Complete Octets (and Duets): Distribute remaining electrons as lone pairs to peripheral (outer) atoms first to satisfy their octets (or duet for H).
4. Place Remaining Electrons on Central Atom: If any electrons are left, place them as lone pairs on the central atom.
5. Check Octets: If the central atom lacks an octet, convert lone pairs from outer atoms into double or triple bonds with the central atom until all atoms (except H) have an octet.
6. Confirm and Check: Ensure all valence shells are complete and the total number of electrons used matches the initial count.
Examples:
- $\text{CH}_4$ (Methane): H—C—H with H's around the C.
- $\text{CO}_2$ (Carbon Dioxide): O=C=O with two lone pairs on each oxygen.
- $\text{CH}_2\text{O}$ (Formaldehyde): H—C=O with H's and a double-bonded O, two lone pairs on O.
- $\text{PBr}_3$ (Phosphorus Tribromide): P bonded to three Br atoms, with a lone pair on P, and three lone pairs on each Br.
Naming Binary Covalent Compounds
- Applies to compounds composed of only two elements (both nonmetals).
1. Name the first element in the formula (typically the one appearing earlier in the sequence: B, Si, C, P, N, H, S, I, Br, Cl, O, F).
2. Name the second element in the formula and change its ending to "-ide".
3. Use Greek prefixes to designate the number of each element present.
 - Prefixes: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7)
 - Exception: The prefix "mono-" is omitted for the first element if there is only one atom of it.
- Traditional Names: Some common compounds retain their traditional names (e.g., water ( $\text{H}2\text{O}$ ), ammonia ( $\text{NH}3$ ), nitrous oxide ( $\text{N}_2\text{O}$ )).
- Examples:
 - $\text{CO}$ : carbon monoxide (mono omitted for carbon, but kept for oxygen since it's the second element)
 - $\text{CO}_2$ : carbon dioxide

3.5 The Mole: Counting Atoms and Compounds

The Need for a Counting Unit
- Atoms and molecules are extremely small, making it impractical to count them individually.
- Similar to how a "dozen" is a unit for counting $12$ items (e.g., eggs, feet in a mile), chemists use a specific unit for counting atoms and molecules.
The Mole
- The mole (mol) is the SI unit for the amount of substance.
- It relates the mass of an element (in grams) to the number of atoms it contains.
- Definition: One mole of carbon-12 (exactly $12$ grams) contains a specific, fixed number of atoms.
Molar Mass
- The molar mass is the mass of one mole of a substance in grams ( $\text{g/mol}$ ).
- Numerically, the molar mass in grams is equal to the atomic mass of an element (in amu/atom) listed on the periodic table.
- Example:
  - Atomic mass of Au atom: $197 \text{ amu}$ .
  - Molar mass of Au: $197 \text{ g/mol}$ .
  - The numbers are the same, only the units differ (amu per atom vs. grams per mole).
Avogadro's Number ( $\text{N}_A$ )
- Named after Amedeo Avogadro.
- The number of atoms or particles present in one mole of any substance is approximately $6.02 \times 10^{23}$ .
  - $6.02 \times 10^{23 \text{ atoms}} = 1 \text{ mole of atoms}$
  - $6.02 \times 10^{23 \text{ molecules}} = 1 \text{ mole of molecules}$
- Avogadro's number is a conversion factor that relates moles to the number of particles (atoms, molecules, ions, etc.).
Relationship between Mass and Moles
- A mole of different elements will have different masses (e.g., a mole of carbon weighs less than a mole of sodium), just as a dozen bowling balls weigh more than a dozen feathers.
Steps to Converting between Units (General Dimensional Analysis)
1. Determine the unit on your final answer.
2. Establish the given information (starting units and values).
3. Decide how to set up the problem using appropriate conversion factors (e.g., molar mass, Avogadro's number).
4. Solve the problem (perform the calculation).
5. Check your answer (units, significant figures, reasonableness).
Conversion Examples
- Molar mass of C: $12.011 \text{ g/mol}$ .
- Equality: $12.011 \text{ g of C} = 1 \text{ mol of C}$ .
- Mass of 1.45 mol C: $(1.45 \text{ mol C}) \times (12.011 \text{ g C / } 1 \text{ mol C}) = 17.4 \text{ g}$ .
- Atoms in 36.0 g C: $(36.0 \text{ g C}) \times (1 \text{ mol C / } 12.011 \text{ g C}) \times (6.02 \times 10^{23} \text{ atoms C / } 1 \text{ mol C}) = 1.80 \times 10^{24} \text{ atoms}$ .
- Molar mass of H $_2$ O: $(2 \times 1.008 \text{ g/mol H}) + (15.999 \text{ g/mol O}) = 18.015 \text{ g/mol}$ .
Moles of Molecules vs. Moles of Atoms
- In $1 \text{ mol of H}_2\text{O}$ , there are $2 \text{ moles of H atoms}$ .

3.6 Getting Covalent Compounds into Shape

3D Molecular Shapes
- Lewis structures are 2D representations, but molecules exist in 3D space.
- For example, methane ( $\text{CH}_4$ ) appears flat with $90^{\circ}$ bond angles in 2D, but in 3D, electrons rearrange to maximize distance, resulting in a tetrahedral shape (approximately $109.5^{\circ}$ bond angles).
Representing 3D on Paper (Wedges and Dashes)
- Wedges: Indicate bonds coming towards the viewer (in front of the plane of the paper).
- Dashes: Indicate bonds going away from the viewer (behind the plane of the paper).
- Normal lines: Indicate bonds lying in the plane of the paper.
Valence-Shell Electron-Pair Repulsion (VSEPR) Model
- Principle: Electron pairs (both bonding and lone pairs) in the valence shell of a central atom repel each other, arranging themselves as far apart as possible to minimize repulsion.
- This repulsion of negative charge dictates the molecule's geometry.
Determining the Shape of a Molecule Using VSEPR
- Step 1: Identify the Central Atom. This atom is typically bonded to two or more other atoms.
- Step 2: Determine Electron Charge Clouds (Electron Domains). Count all regions of electron density around the central atom.
 - A single bond, double bond, and triple bond each count as one bonding cloud (represented by B in VSEPR notation).
 - A lone pair of electrons counts as one nonbonding cloud (represented by N in VSEPR notation).
- VSEPR Form: Notation like $\text{AB}x\text{N}y$ , where A is the central atom, B is a bonding cloud, and N is a nonbonding cloud.
VSEPR Forms, Molecular Shapes, and Bond Angles (Focus for this chapter)
- AB $_4$ (4 bonding clouds, 0 lone pairs):
 - Molecular Shape: Tetrahedral
 - Bond Angle: $109.5^{\circ}$
- AB $_3$ N (3 bonding clouds, 1 lone pair):
 - Molecular Shape: Pyramidal (or Trigonal Pyramidal)
 - Bond Angle: <109.5^{\circ} (lone pair repulsion compresses angles)
- AB $2$ N $2$ (2 bonding clouds, 2 lone pairs):
 - Molecular Shape: Bent (or V-shaped)
 - Bond Angle: <109.5^{\circ} (lone pair repulsion compresses angles further)
- AB $_3$ (3 bonding clouds, 0 lone pairs):
 - Molecular Shape: Trigonal planar
 - Bond Angle: $120^{\circ}$
- AB $_2$ (2 bonding clouds, 0 lone pairs):
 - Molecular Shape: Linear
 - Bond Angle: $180^{\circ}$
Examples: Predicting Molecular Geometries
- $\text{CH}4$ (Methane): Central C, $4$ bonds, $0$ lone pairs ( $\text{AB}4$ ). Shape: Tetrahedral.
- $\text{CO}2$ (Carbon Dioxide): Central C, $2$ double bonds, $0$ lone pairs ( $\text{AB}2$ ). Shape: Linear.
- $\text{CH}2\text{O}$ (Formaldehyde): Central C, $3$ bonds (one double, two single), $0$ lone pairs ( $\text{AB}3$ ). Shape: Trigonal planar.
- $\text{PBr}3$ (Phosphorus Tribromide): Central P, $3$ bonds, $1$ lone pair ( $\text{AB}3\text{N}$ ). Shape: Pyramidal.
VSEPR for Carbon as the Central Atom
- Carbon typically forms no lone pairs, simplifying its VSEPR prediction:
  - If carbon has $4$ atoms bonded to it (4 bonding clouds): Tetrahedral.
  - If carbon has $3$ atoms bonded to it (3 bonding clouds, e.g., with a double bond): Trigonal planar.
  - If carbon has $2$ atoms bonded to it (2 bonding clouds, e.g., with two double bonds or a triple and a single bond): Linear.
Complex Molecules
- For larger molecules with multiple central atoms, the geometry must be considered around each individual central atom.
- Example: In acetic acid ( $\text{CH}_3\text{COOH}$ ), the first carbon is tetrahedral, the second carbon is trigonal planar, and the oxygen is bent.

3.7 Electronegativity and Molecular Polarity

Electronegativity
- Definition: The ability of an atom to attract the bonding electrons of a covalent bond to itself.
- Analogy: Think of it like a magnet's strength for electrons; greater electronegativity means a stronger pull.
- Highest Electronegativity: Fluorine (F) is the most electronegative element.
- Periodic Trends:
  - Across a Period (left to right): Electronegativity increases because the positive charge of the nucleus increases, leading to a greater attraction for electrons.
  - Down a Group (top to bottom): Electronegativity decreases because atomic radius increases, placing valence electrons farther from the nucleus, thus decreasing the attractive force.
Bond Polarity
- Unequal Sharing: When two different atoms form a covalent bond, the electrons are shared unequally due to differing electronegativities. This results in a polar covalent bond.
- Equal Sharing: When two identical atoms form a covalent bond, the electrons are shared equally. This results in a nonpolar covalent bond.
- Indicating Polarity:
  - Partial Charges: Use the lowercase Greek delta symbol ( $\delta$ ) to denote partial charges.
    - $\delta^-$ indicates the atom that is more electronegative and pulls electrons towards itself.
    - $\delta^+$ indicates the atom that is less electronegative and has electrons pulled away from it.
  - Dipole Moment Arrow: An arrow pointing from the less electronegative atom ( $\delta^+$ ) to the more electronegative atom ( $\delta^-$ ). A plus sign on the tail of the arrow indicates the partial positive end.
Electronegativity Difference ( $\Delta\text{EN}$ ) and Bond Type
- Ionic Bond: Elements with an electronegativity difference of $1.8$ or more typically form an ionic bond (electron transfer).
 - Electron density entirely localized on one ion.
- Covalent Bond: Elements with an electronegativity difference of less than $1.8$ form a covalent bond (electron sharing).
 - Nonpolar Covalent: \Delta\text{EN} < 0.4 (electrons shared virtually equally).
 - Polar Covalent: 0.4 \le \Delta\text{EN} < 1.8 (electrons shared unequally, but still shared).
- The greater the difference in electronegativity ( $\Delta\text{EN}$ ), the more polar the bond.
Electron Densities Illustrate Bonding
- Visual representations (e.g., using color gradients like red for high density, blue for low density) show electron distribution.
- Non-polar covalent: Symmetrical electron distribution between atoms.
- Polar covalent: Asymmetrical electron distribution, with higher density near the more electronegative atom.
- Ionic: Complete transfer, resulting in distinct positive and negative ions.
Example: Ordering Bonds by Polarity
- Order of increasing polarity (least polar to most polar) for N-O, C-H, S-F, C-O:
 \text{C-H < N-O < C-O < S-F} (Note: Spectific $\Delta\text{EN}$ values would be provided for such a task).
Molecular Polarity
- Polar Molecule: Has an uneven distribution of electrons over the entire molecule, resulting in a net molecular dipole moment.
- Nonpolar Molecule: Has an even distribution of electrons over the entire molecule, with no net molecular dipole moment.
- For Diatomic Molecules: The polarity of the molecule is the same as the polarity of its single bond.
- For Molecules with Three or More Atoms: Both the electronegativity of the atoms (determining bond polarity) and the overall shape (geometry) of the molecule must be considered.
Determining if a Molecule is Polar - Steps
1. Draw the Lewis structure.
2. Predict the molecular shape (geometry) using VSEPR.
3. Draw on the individual bond dipoles (based on electronegativity differences).
4. Determine if the bond dipoles cancel out (vector sum of dipoles). If they cancel symmetrically, the molecule is nonpolar; if they do not cancel, the molecule is polar.
Examples: CO $2$ vs. H $2$ O
- Carbon Dioxide ( $\text{CO}_2$ ):
 - Lewis structure: O=C=O.
 - Molecular shape: Linear ( $180^{\circ}$ bond angles).
 - Bond dipoles: There are two polar C=O bonds. The oxygen atoms are more electronegative, so each oxygen pulls electrons away from carbon equally, but in opposite directions.
 - Overall polarity: The bond dipoles cancel each other due to the linear shape. Result: Nonpolar molecule.
- Water ( $\text{H}_2\text{O}$ ):
 - Lewis structure: H-O-H with two lone pairs on oxygen.
 - Molecular shape: Bent ( <109.5^{\circ} bond angle).
 - Bond dipoles: There are two polar O-H bonds. Oxygen is more electronegative, pulling electrons away from each hydrogen.
 - Overall polarity: Due to the bent shape, the bond dipoles do not cancel. They add together to create an overall molecular dipole moment. Result: Polar molecule.
Tug-of-War Analogy: Imagine bond dipoles as vectors in a tug-of-war. If the pulls are equal and in opposite directions (like linear CO $2$ ), there's no winner (nonpolar). If the pulls are not balanced (like bent H $2$ O), there's a winner (polar).
Other Examples of Molecular Polarity
- $\text{BF}_3$ (Trigonal planar, symmetrical bond dipoles): Nonpolar
- $\text{NH}_3$ (Pyramidal, asymmetrical bond dipoles due to lone pair and geometry): Polar
- $\text{SF}_4$ (See-saw shape, asymmetrical): Polar
- $\text{CH}_3\text{Cl}$ (Tetrahedral, but C-Cl bond is more polar than C-H, and geometry is not perfectly symmetrical to cancel): Polar
- $\text{CCl}_4$ (Tetrahedral, C-Cl bonds are polar, but the symmetrical tetrahedral geometry causes dipoles to cancel): Nonpolar