Chapter 3: Structure and Properties of Ionic and Covalent Compounds

3.1 Chemical Bonding

• Chemical Bond:

  • The force of attraction between any two atoms in a compound.

  • This attractive force counteracts the repulsion of the positively charged nuclei of the two atoms.

  • Interactions involving valence electrons are responsible for chemical bonding.

• Lewis Symbols:

  • A method to represent atoms using the element symbol and valence electrons as dots.

  • Only valence electrons participate in bonding, simplifying the application of the octet rule.

  • The number of dots directly corresponds to the number of valence electrons in the outermost shell.

• Writing Lewis Symbols:

  • The four sides around the atomic symbol can each hold two dots, for a maximum of eight (an octet of electrons).

  • Procedure:

    1. Place one dot on each side until four dots surround the symbol.

    2. Then, add a second dot to each side in rotation.

    3. The actual number of valence electrons limits the total number of dots placed.

    4. Each unpaired dot (representing an unpaired valence electron) is available to form a chemical bond.

• Principal Types of Chemical Bonds:

  • Ionic bond: An attractive force resulting from the complete transfer of one or more electrons from one atom to another.

    • This attraction arises from the opposite charges of the resulting ions.

  • Covalent bond: An attractive force resulting from the sharing of electrons between atoms.

  • Some bonds exhibit characteristics of both types, making their classification ambiguous.

• Ionic Bonding:

  • Representative elements form ions that typically obey the octet rule.

  • Electrons are lost by a metal atom and gained by a nonmetal atom.

  • Each atom achieves a stable “noble gas” electron configuration.

  • Two types of ions are formed: a positively charged cation and a negatively charged anion.

  • These oppositely charged ions attract each other, forming the ionic bond.

• Ionic Bonding: Example of NaCl Formation:

  • $Na + Cl \rightarrow NaCl$

  • Sodium (Na): Has a low ionization energy, meaning it readily loses an electron.

    • When sodium loses its valence electron, it achieves the electron configuration of Neon (Ne).

  • Chlorine (Cl): Has a high electron affinity.

    • When chlorine gains an electron, it achieves the electron configuration of Argon (Ar).

• Essential Features of Ionic Bonding:

  • Metals: Tend to form cations because they possess low ionization energies and low electron affinities.

  • Nonmetals: Tend to form anions because they possess high ionization energies and high electron affinities.

  • Ions are inherently formed by the transfer of electrons.

  • The oppositely charged ions are held together by a strong electrostatic force.

  • Reactions between metals and nonmetals typically lead to the formation of ionic compounds.

• Ion Arrangement in a Crystal:

  • When a sodium atom loses an electron, it becomes a smaller sodium ion $(Na^+)$.

  • When a chlorine atom gains that electron, it becomes a larger chloride ion $(Cl^-)$.

  • The attraction between $(Na^+)$ cations and $(Cl^-)$ anions leads to the formation of NaCl ion pairs, which then aggregate into a regular, repeating three-dimensional structure called a crystal lattice.

• Covalent Bonding: Example of H2 Formation:

  • $H + H \rightarrow H_2$

  • Each hydrogen atom has one electron in its valence shell.

  • If bond formation were purely ionic, it would involve electron transfer, which is generally unfavorable under normal conditions since both hydrogen atoms have an equal tendency to gain or lose electrons.

• Covalent Bond:

  • Instead of transferring electrons, each atom attains a stable noble gas configuration (like Helium, with two electrons) by sharing electrons.

  • The shared pair of electrons constitutes a covalent bond.

• Covalent Bonding in Hydrogen (Visualized):

  • Hydrogen atoms approach each other at high velocity.

  • The nuclei of the hydrogen atoms begin to attract each other's electrons.

  • The hydrogen atoms form a hydrogen molecule, held together by the shared electrons, which form the covalent bond.

• Features of Covalent Bonds:

  • Covalent bonds form between atoms with similar tendencies to gain or lose electrons.

  • Compounds containing covalent bonds are known as covalent compounds or molecules.

  • Diatomic Elements: Have completely covalent bonds, meaning electrons are shared absolutely equally. Examples include $(H<em>2, N</em>2, O<em>2, F</em>2, Cl<em>2, Br</em>2, I_2)$.

  • In $(F_2)$, each fluorine atom is surrounded by eight electrons, achieving the electron configuration of Neon (Ne).

• Examples of Covalent Bonding:

  • Water $(H_2O)$: $2H \cdot + \cdot \ddot{O} \cdot \rightarrow H:\ddot{O}:H$

    • 2 electrons from 2 H atoms, 6 electrons from O atom.

    • Each H has 2 electrons; O has 8 electrons.

  • Methane $(CH_4)$: $4H \cdot + \cdot \dot{C} \cdot \rightarrow H:\ddot{C}:H$

    • 4 electrons from 4 H atoms, 4 electrons from C atom.

    • Each H has 2 electrons; C has 8 electrons.

• Polar Covalent Bonding and Electronegativity:

  • Polar Covalent Bond: Bonds formed by unequally shared electron pairs.

  • Contrast:

    • Ionic bonding: involves the transfer of electrons.

    • Covalent bonding: involves the sharing of electrons.

• Polar Covalent Bond (Example: HF):

  • In the H-F bond, hydrogen becomes somewhat positively charged, and fluorine becomes somewhat negatively charged.

  • The two electrons in the bond are not shared equally; they spend more time associated with the fluorine atom.

  • This unequal sharing establishes a polar covalent bond.

  • A truly nonpolar covalent bond only occurs when the two bonded atoms are identical (e.g., $(H<em>2, O</em>2)$).

• Polar Covalent Bonding in HF:

  • Fluorine is considered electron-rich.

  • Hydrogen is considered electron-deficient.

  • This discrepancy results in the unequal sharing of electrons within the bond, leading to a polar covalent bond.

• Electronegativity:

  • $Electronegativity$ is a measure of an atom's ability to attract electrons within a chemical bond.

  • Elements with high electronegativity attract electrons more strongly than elements with low electronegativity.

  • A covalent bond can be viewed as a competition for electrons between two positively charged nuclei.

  • The difference in electronegativity between two bonded atoms quantifies the extent of the bond's polarity.

• Electronegativities of Selected Elements (Periodic Table Trend):

  • The most electronegative elements are located in the upper right portion of the periodic table (e.g., F, O, N).

  • The least electronegative elements are found in the lower left portion of the periodic table (e.g., Fr, Cs).

• Electronegativity Calculations:

  • A larger difference in electronegativity between two atoms corresponds to a greater polarity of their bond.

  • Question: Which bond is more polar, H-F or H-Cl?

    • For H-F: $Electronegativity ext{ }difference = 4.0 (F) - 2.2 (H) = 1.8$

    • For H-Cl: $Electronegativity ext{ }difference = 3.2 (Cl) - 2.2 (H) = 1.0$

  • Conclusion: The H-F bond is more polar than the H-Cl bond because its electronegativity difference is greater.

3.2 Naming Compounds and Writing Formulas of Compounds

• Nomenclature: The systematic assignment of a correct and unambiguous name to every chemical compound.

  • There are two primary naming systems: one for ionic compounds and one for covalent compounds.

• Formulas of Compounds:

  • A chemical formula represents a fundamental compound unit using chemical symbols and numerical subscripts.

  • The formula identifies the number and type of various atoms that constitute the compound unit.

  • Subscripts denote the number of like atoms in the unit.

  • If no subscript is present for an atom, its presence of one atom is understood.

• Ionic Compounds:

  • Typically formed when metals and nonmetals react.

  • Metals form cations (positive ions), and nonmetals form anions (negative ions).

  • Cations and anions arrange themselves in a regular, repeating three-dimensional array called a crystal lattice.

  • The formula of an ionic compound represents the smallest whole-number ratio of ions present in the substance.

• Writing Formulas of Ionic Compounds from Ion Identities:

  • Steps:

    1. Determine the charge of each ion:

       • For main group metals, the charge is usually equal to their group number.

       • For main group nonmetals, the charge is typically their group number minus eight.

    2. Cations and anions must combine in a ratio that results in a net charge of zero for the compound.

       • The total number of positive charges must equal the total number of negative charges.

• Predict Formulas (Examples):

  1. Sodium (Na) and Oxygen (O): $Na^+$ (Group 1), $O^{2-}$(Group 16, 16-8=8-2=6 or 8-6=2)

     • Formula: $Na_2O$

  2. Lithium (Li) and Bromine (Br): $Li^+$ (Group 1), $Br^-$ (Group 17, 17-8=9-1=8 or 8-7=1)

     • Formula: $LiBr$

  3. Aluminum (Al) and Oxygen (O): $Al^{3+}$ (Group 13), $O^{2-}$

     • Formula: $Al<em>2O</em>3$

  4. Barium (Ba) and Fluorine (F): $Ba^{2+}$ (Group 2), $F^-$ (Group 17)

     • Formula: $BaF_2$

• Writing Names of Ionic Compounds from the Formula (Rule 1):

  • Name the cation first, followed by the anion.

  • The positive ion (cation) retains the name of the element.

  • The negative ion (anion) uses the stem of the element name with the suffix "-ide".

  • Examples:

    • NaCl: sodium chloride

    • $Na_2O$: sodium oxide

    • $Li_2S$: lithium sulfide

    • $AlBr_3$: aluminum bromide

    • CaO: calcium oxide

• Writing Names of Ionic Compounds from the Formula (Rule 2 - for Transition Metals):

  • If the cation is an element that can form several ions with different charges (common for transition metals), a Roman numeral is used after the metal's name to indicate its charge.

  • Examples:

    • $FeCl_3$: iron(III) chloride (Iron has a $+3$ charge)

    • $FeCl_2$: iron(II) chloride (Iron has a $+2$ charge)

    • CuO: copper(II) oxide (Copper has a $+2$ charge)

• Common Nomenclature System (for Transition Metals):

  • Uses suffixes to indicate charge, primarily for metals with two common oxidation states.

  • The suffix "-ic" indicates the higher of the two common charges for that ion.

  • The suffix "-ous" indicates the lower of the two common charges for that ion.

  • Examples:

    • $FeCl_2$: ferrous chloride ($Fe^{2+}$ is the lower charge)

    • $FeCl_3$: ferric chloride ($Fe^{3+}$ is the higher charge)

• Monatomic Ions: Ions consisting of a single charged atom.

• Polyatomic Ions:

  • Ions composed of two or more atoms bonded together that collectively carry an overall positive or negative charge.

  • Within the polyatomic ion itself, the atoms are held together by covalent bonds.

  • These polyatomic ions then form ionic bonds with other oppositely charged ions.

  • Examples: ammonium ion ($NH<em>4^+$), sulfate ion ($SO</em>4^{2-}$).

• Common Polyatomic Cations and Anions (Table 3.3):

  • Cations: Hydronium ($H<em>3O^+$), Ammonium ($NH</em>4^+$)

  • Anions: Carbonate ($CO<em>3^{2-}$), Bicarbonate (HCO3^-$), Nitrite (NO2^-$), Nitrate (NO3^-$), Sulfite (SO3^{2-}$), Sulfate ($SO4^{2-}$), Hydrogen sulfate ($HSO4^-$), Hydroxide (OH^-$), Cyanide (CN^-$), Phosphate ($PO</em>4^{3-}$), Hydrogen phosphate ($HPO<em>4^{2-}$), Dihydrogen phosphate (H2PO4^-$), Hypochlorite (ClO^-$), Chlorite (ClO2^-$), Chlorate (ClO3^-$), Perchlorate (ClO4^-$), Acetate (CH3COO^-$or$C2H3O2^-$), Permanganate (MnO4^-$), Dichromate (Cr2O7^{2-}$), Chromate ($CrO4^{2-}$), Peroxide ($O_2^{2-}$).</p></li></ul></li><li><p><strong>Name These Compounds (Examples):</strong></p><ol><li><p>$NH_4Cl$: Ammonium chloride</p></li><li><p>$BaSO_4$: Barium sulfate</p></li><li><p>$Fe(NO3)3$: Iron(III) nitrate (or Ferric nitrate)</p></li><li><p>$CuHCO3$: Copper(I) bicarbonate (if$Cu^+$) or Copper(II) bicarbonate (if$Cu^{2+}$- usually$Cu(HCO3)_2$)</p></li><li><p>$Ca(OH)_2$: Calcium hydroxide</p></li></ol></li><li><p><strong>Writing Formulas of Ionic Compounds from the Name:</strong></p><ul><li><p><strong>Steps:</strong></p><ol><li><p>Determine the charge of each ion specified in the name.</p></li><li><p>Write the formula so that the overall compound is electrically neutral (net charge of zero).</p></li></ol></li><li><p><strong>Example:</strong> Barium chloride</p><ul><li><p>Barium is$(Ba^{2+})$.</p></li><li><p>Chloride is$(Cl^-)$.</p></li><li><p>Formula:$BaCl_2$</p></li></ul></li></ul></li><li><p><strong>Determine the Formulas from Names (Examples):</strong></p><ol><li><p>Sodium sulfate:$Na2SO4$</p></li><li><p>Ammonium sulfide:$(NH4)2S$</p></li><li><p>Magnesium phosphate:$Mg3(PO4)_2$</p></li><li><p>Chromium(II) sulfate:$CrSO_4$</p></li></ol></li><li><p><strong>Covalent Compounds:</strong></p><ul><li><p>Typically formed between two nonmetal atoms.</p></li><li><p>Are characterized by covalent bonding.</p></li><li><p>Exist as discrete molecules in solid, liquid, and gas states, unlike ionic compounds which form mass crystal structures.</p></li></ul></li><li><p><strong>Naming Covalent Compounds (Rule 1):</strong></p><ol><li><p>The names of the elements are written in the order they appear in the formula.</p></li><li><p>Prefixes are used to indicate the number of each kind of atom present.</p><ul><li><p><strong>Prefixes Used to Denote Numbers of Atoms (Table 3.4):</strong></p><ul><li><p>Mono- (1), Di- (2), Tri- (3), Tetra- (4), Penta- (5), Hexa- (6), Hepta- (7), Octa- (8), Nona- (9), Deca- (10)</p></li></ul></li></ul></li></ol></li><li><p><strong>Naming Covalent Compounds (Rule 2):</strong></p><ol start="3"><li><p>If there is only one atom of the <strong>first</strong> element in the molecule, the prefix "mono-" is usually omitted.</p></li><li><p><strong>Example:</strong> CO is named carbon monoxide (not monocarbon monoxide).</p></li><li><p>The stem of the name of the <strong>last</strong> element is used, followed by the suffix "-ide".</p></li><li><p>When a prefix ends in a vowel (like "tetra" or "penta") and the element name begins with a vowel (like "oxide"), the final vowel in the prefix is often dropped (e.g., "tetroxide" instead of "tetraoxide").</p></li></ol></li><li><p><strong>Name These Covalent Compounds (Examples):</strong></p><ol><li><p>$SiO_2$: Silicon dioxide</p></li><li><p>$N2O5$: Dinitrogen pentoxide</p></li><li><p>$CCl_4$: Carbon tetrachloride</p></li><li><p>$IF_7$: Iodine heptafluoride</p></li></ol></li><li><p><strong>Writing Formulas of Covalent Compounds:</strong></p><ul><li><p>Use the prefixes in the compound's name to determine the subscripts for each element in the formula.</p></li><li><p><strong>Examples:</strong></p><ul><li><p>Nitrogen trichloride:$NCl_3$</p></li><li><p>Diphosphorus pentoxide:$P2O5$</p></li></ul></li><li><p><strong>Some Common Names (not systematic):</strong></p><ul><li><p>$H_2O$: water</p></li><li><p>$NH_3$: ammonia</p></li><li><p>$C2H5OH$: ethanol or ethyl alcohol</p></li><li><p>$C6H12O_6$: glucose</p></li></ul></li></ul></li><li><p><strong>Provide Formulas for These Covalent Compounds (Examples):</strong></p><ol><li><p>Nitrogen monoxide: NO</p></li><li><p>Dinitrogen tetroxide:$N2O4$</p></li><li><p>Diphosphorus pentoxide:$P2O5$</p></li><li><p>Nitrogen trifluoride:$NF_3

3.3 Properties of Ionic and Covalent Compounds

• 
  

• Physical State (at room temperature):

  • Ionic compounds: Primarily solids.

  • Covalent compounds: Can exist as solids, liquids, or gases.

• 
  

• Introduction to Melting and Boiling Points:

  • Melting point: The specific temperature at which a solid transforms into a liquid.

  • Boiling point: The specific temperature at which a liquid transforms into a gas.

• 
  

• Physical Properties: Melting and Boiling Points Comparison:

  • Ionic compounds: Possess significantly higher melting points and boiling points compared to covalent compounds.

    • This is because a substantial amount of energy is required to overcome the strong electrostatic attractions between the ions in the crystal lattice.

    • Ionic compounds typically melt at several hundred degrees Celsius.

• 
  

• Structure of Compounds in the Solid State:

  • Ionic compounds: Are always crystalline (characterized by a regular, repeating arrangement of atoms or ions).

  • Covalent compounds: Can be either crystalline or amorphous (lacking a regular, ordered structure).

• 
  

• Electrolytes and Nonelectrolytes:

  • Solutions of Ionic and Covalent Compounds:

    • Ionic compounds: Often dissolve in water and dissociate, meaning they break apart into their constituent positive and negative ions in solution.

    • Electrolytes: Substances (like dissolved ionic compounds) that produce ions in solution, thereby allowing the solution to conduct electricity.

    • Covalent compounds: Most covalent solids do not dissociate into ions when dissolved and consequently do not conduct electricity. They are termed nonelectrolytes.

• 
  

• Comparison of Ionic vs. Covalent Compounds (Summary Table):
  

  Property	Ionic	Covalent
  Often composed of	Metal + nonmetal	2 nonmetals
  Electrons are	Transferred	Shared
  Physical state is	Solid and crystalline	Any; crystal or amorphous
  Dissociation	Yes, they are electrolytes	No, they are nonelectrolytes
  Boiling and Melting point	High	Low
  3.4 Drawing Lewis Structures of Molecules and Polyatomic Ions

  • Lewis Structure Guidelines (Step 1: Skeletal Structure):

    1. Use chemical symbols to write the skeletal arrangement of the compound.

    2. The least electronegative atom is typically placed in the central position.

    3. Hydrogen atoms always occupy terminal (outer) positions.

    4. Halogen atoms usually occupy terminal positions, unless more electronegative elements are present.

    5. Carbon atoms frequently form chains of carbon-carbon covalent bonds.

  • Lewis Structure Guidelines (Step 2: Total Valence Electrons):

    1. Determine the number of valence electrons contributed by each atom in the compound.

    2. Combine these to find the total number of valence electrons for the entire compound.

    3. For polyatomic cations, subtract one electron for every positive charge.

    4. For polyatomic anions, add one electron for every negative charge.

  • Lewis Structure Guidelines (Step 3 & 4: Bonds and Lone Pairs):

    3. Connect the central atom to each of the surrounding (terminal) atoms with a single bond (representing two shared electrons).

    4. Place remaining electrons as lone pairs around the terminal atoms first, to complete an octet for each (except hydrogen, which only needs two electrons).

       • Electrons not involved in bonding are represented as lone pairs.

       • After all terminal atoms have an octet, if valence electrons are still available, provide the central atom with an octet.

  • Lewis Structure Guidelines (Step 5 & 6: Multiple Bonds and Recheck):

    5. If the octet rule is not satisfied for all atoms (especially the central atom, if it's not a known exception), move one or more lone pairs from a terminal atom to form an additional bond (double or triple bond) with the central atom.

       • Continue shifting electrons until all atoms satisfy the octet rule.

    6. After a satisfactory Lewis structure is constructed, perform a final electron count to ensure the total number of valence electrons used matches the initial calculation.

  • Drawing Lewis Structures of Covalent Compounds (Example: CO_2$):</strong></p><ol><li><p><strong>Skeletal Structure:</strong> Carbon (2.5) is less electronegative than Oxygen (3.5), so Carbon is central: O-C-O.</p></li><li><p><strong>Valence Electrons:</strong> Carbon (Group 14) has 4 valence electrons. Each Oxygen (Group 16) has 6 valence electrons. Total =$4 + (2 \times 6) = 16$valence electrons.</p></li><li><p><strong>Single Bonds and Initial Lone Pairs:</strong> Connect C to each O with a single bond. ($16 - 4 = 12$remaining electrons). Place the 12 electrons as lone pairs on the O atoms to give them octets (6 each). Each O now has 8 electrons, but C only has 4.</p></li><li><p><strong>Form Multiple Bonds (Redistribute Electrons):</strong> Move two lone pair electrons from each Oxygen atom to form double bonds with the central Carbon atom. This satisfies the octet rule for all atoms:$O=C=O$. Each O has 2 lone pairs + 2 shared pairs = 8 electrons. C has 2 shared pairs + 2 shared pairs = 8 electrons.</p></li><li><p><strong>Recheck Electron Distribution:</strong> 4 bonds$(\times 2e^- ext{ each}) = 8e^-$. 4 lone pairs$(\times 2e^- ext{ each}) = 8e^-$. Total =$16e^-$. This matches the initial count, and all atoms have octets.</p></li></ol></li><li><p><strong>Lewis Structure of Polyatomic Anions (Example: Carbonate ion,$CO_3^{2-}$):</strong></p><ol><li><p><strong>Skeletal Structure:</strong> Carbon is less electronegative than oxygen, so carbon is the central atom: surrounded by 3 oxygens. Include notation for the charge ($C_3O^{2-}$).</p></li><li><p><strong>Total Valence Electrons:</strong> Carbon (4), 3 Oxygen (3 x 6 = 18). Add 2 electrons for the$2-$charge. Total =$4 + 18 + 2 = 24$valence electrons.</p></li><li><p><strong>Single Bonds and Initial Lone Pairs:</strong> Connect C to each O with a single bond ($24 - 6 = 18 remaining electrons). Distribute the 18 electrons as lone pairs around the terminal O atoms, giving each an octet ($ ext{6 each}$). Each O has 8 electrons, but C only has 6 electrons (from 3 single bonds).

  • Form Multiple Bonds: To satisfy the octet for carbon, move one lone pair from one of the oxygen atoms to form a double bond with the carbon atom.

• Lewis Structure, Stability, Multiple Bonds, and Bond Energies:

  • Single bond: One pair of electrons is shared between two atoms.

  • Double bond: Two pairs of electrons are shared between two atoms.

  • Triple bond: Three pairs of electrons are shared between two atoms.

• Bond Energy and Bond Length:

  • Bond energy: The amount of energy required to break a bond holding two atoms together.

    • Trend: Triple bond > Double bond > Single bond (stronger bonds require more energy to break).

  • Bond length: The distance separating the nuclei of two adjacent atoms.

    • Trend: Single bond > Double bond > Triple bond (stronger bonds are typically shorter).

• Lewis Structures and Resonance (Example: Carbonate ion, CO_3^{2-}$):</strong></p><ul><li><p>In some cases, it's possible to draw more than one valid Lewis structure that satisfies the octet rule for a particular compound (e.g., placing the double bond on different oxygen atoms in carbonate).</p></li><li><p>Experimental evidence (e.g., all C-O bonds in carbonate having the same length) shows that none of these individual Lewis structures fully represent the actual structure.</p></li><li><p>The actual structure is an average or hybrid of these contributing Lewis structures.</p></li><li><p><strong>Resonance:</strong> This phenomenon describes two or more Lewis structures that collectively contribute to represent the true (hybrid) structure of a molecule or ion.</p></li></ul></li><li><p><strong>Lewis Structures and Exceptions to the Octet Rule:</strong></p><ol><li><p><strong>Incomplete Octet:</strong> Atoms with less than eight electrons around them (excluding hydrogen).</p><ul><li><p><strong>Example:</strong>$BeH_2$. Only 4 electrons surround the beryllium atom.</p></li><li><p>The Lewis structure will show Be with only 4 electrons.</p></li></ul></li><li><p><strong>Odd Electron Species:</strong> If there is an odd number of total valence electrons, it is impossible for every atom to achieve a full octet.</p><ul><li><p><strong>Example:</strong> NO (nitric oxide). Total valence electrons = 5 (N) + 6 (O) = 11 electrons. One atom will necessarily have an odd number of electrons.</p></li></ul></li><li><p><strong>Expanded Octet:</strong> Elements in the 3rd period or below (having available d-orbitals) can accommodate more than eight electrons (10 or 12 electrons) around them.</p><ul><li><p>This is the most common exception to the octet rule.</p></li><li><p><strong>Example:</strong>$PF_5$. Phosphorus is a third-period element. Its Lewis structure shows 10 electrons around the central phosphorus atom (5 bonding pairs).</p></li></ul></li></ol></li></ul><h4 id="eff31f08-935f-48fd-83c9-98b356a8a565" data-toc-id="eff31f08-935f-48fd-83c9-98b356a8a565" collapsed="false" seolevelmigrated="true">3.5 Lewis Structures and Molecular Geometry; VSEPR Theory</h4><ul><li><p><strong>Molecular Shape and Properties:</strong> The geometric shape of molecules significantly influences their physical and chemical properties.</p></li><li><p><strong>VSEPR Theory (Valence Shell Electron Pair Repulsion Theory):</strong></p><ul><li><p><strong>Purpose:</strong> Used to predict the three-dimensional shape of molecules.</p></li><li><p><strong>Underlying Principle:</strong> All electron pairs (both bonding pairs and lone pairs) around a central atom will arrange themselves in space to be as far away from each other as possible.</p><ul><li><p>This arrangement minimizes electron-electron repulsion, leading to the most stable geometry.</p></li></ul></li></ul></li><li><p><strong>VSEPR Theory (Cont.):</strong></p><ul><li><p>In a covalent bond, bonding electrons are localized between specific atomic nuclei.</p></li><li><p>Covalent bonds are <strong>directional</strong>, meaning they have a specific orientation in space between the bonded atoms.</p></li><li><p>Ionic bonds are based on electrostatic forces, which are non-directional and extend uniformly in all spatial orientations around the ions.</p></li></ul></li><li><p><strong>A Stable Exception to the Octet Rule (Example:$BeH_2$):</strong></p><ul><li><p>The central beryllium (Be) atom only has 4 electrons (two bonding pairs) around it.</p></li><li><p>These electron pairs minimize repulsion by positioning themselves on opposite sides of the beryllium atom.</p></li><li><p>This arrangement results in a <strong>linear structure</strong> with a bond angle of$180^ ext{o}$. (Electron Groups: 2, Bonded: 2, Lone Pairs: 0)</p></li></ul></li><li><p><strong>Another Stable Exception to the Octet Rule (Example:$BF_3$):</strong></p><ul><li><p>The central boron (B) atom has 3 bonded atoms around it (three bonding pairs), totaling 6 electrons.</p></li><li><p>These bonded atoms achieve minimal repulsion when arranged in a plane, forming a triangle.</p></li><li><p>This results in a <strong>trigonal planar structure</strong> with bond angles of$120^ ext{o}$. (Electron Groups: 3, Bonded: 3, Lone Pairs: 0)</p></li></ul></li><li><p><strong>Basic Electron Pair Repulsion of a Full Octet (Example:$CH_4$):</strong></p><ul><li><p>The central carbon (C) atom has 4 bonded atoms around it (four bonding pairs), totaling 8 electrons.</p></li><li><p>Minimal electron repulsion is achieved when these electron pairs are directed towards the four corners of a tetrahedron.</p></li><li><p>This results in a <strong>tetrahedral structure</strong>, where each H-C-H bond angle is$109.5^ ext{o}$. (Electron Groups: 4, Bonded: 4, Lone Pairs: 0)</p></li></ul></li><li><p><strong>Basic Electron Pair Repulsion of a Full Octet with One Lone Pair (Example:$NH_3$):</strong></p><ul><li><p>The central nitrogen (N) atom has three bonded atoms and one lone pair, making a total of four electron groups.</p></li><li><p>Lone pairs are more electronegative and exert greater electron repulsion than bonding pairs, effectively occupying more space.</p></li><li><p>The lone pair takes one of the corners of the tetrahedral electron geometry, distorting the arrangement of the bonded atoms.</p></li><li><p>Ammonia thus has a <strong>trigonal pyramidal structure</strong> with bond angles of approximately$107^ ext{o}$. (Electron Groups: 4, Bonded: 3, Lone Pairs: 1)</p></li></ul></li><li><p><strong>Basic Electron Pair Repulsion of a Full Octet with Two Lone Pairs (Example:$H_2O$):</strong></p><ul><li><p>The central oxygen (O) atom has two bonded atoms and two lone pairs, totaling four electron groups.</p></li><li><p>All four electron pairs are approximately tetrahedral in their spatial arrangement around the oxygen.</p></li><li><p>The two lone pairs occupy two corners of the tetrahedron, effectively distorting the arrangement of the hydrogen atoms.</p></li><li><p>Water consequently has a <strong>bent</strong> or <strong>angular structure</strong> with bond angles of approximately$104.5^ ext{o}$.</p></li></ul></li><li><p><strong>Predicting Geometric Shape Using Electron Pairs (Summary Table):</strong><br></p><table style="min-width: 150px;"><colgroup><col style="min-width: 25px;"><col style="min-width: 25px;"><col style="min-width: 25px;"><col style="min-width: 25px;"><col style="min-width: 25px;"><col style="min-width: 25px;"></colgroup><tbody><tr><th colspan="1" rowspan="1" style="text-align: left;"><p>Bonded Atoms</p></th><th colspan="1" rowspan="1" style="text-align: left;"><p>Nonbonding Lone Electron Pairs</p></th><th colspan="1" rowspan="1" style="text-align: left;"><p>Bond Angle</p></th><th colspan="1" rowspan="1" style="text-align: left;"><p>Molecular Geometry</p></th><th colspan="1" rowspan="1" style="text-align: left;"><p>Example</p></th><th colspan="1" rowspan="1"><p><br></p></th></tr><tr><td colspan="1" rowspan="1" style="text-align: left;"><p>2</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>0</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>$180^ ext{o}$</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>Linear</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>$CO2$| | 3 | 0 |$120^ ext{o}$| Trigonal planar |$SO3$</p></td><td colspan="1" rowspan="1"><p><br></p></td></tr><tr><td colspan="1" rowspan="1" style="text-align: left;"><p>2</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>1</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>$<120^ ext{o}$</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>Bent</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>$SO2$| | 4 | 0 |$109.5^ ext{o}$| Tetrahedral |$CH4$</p></td><td colspan="1" rowspan="1"><p><br></p></td></tr><tr><td colspan="1" rowspan="1" style="text-align: left;"><p>3</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>1</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>$~107^ ext{o}$</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>Trigonal pyramidal</p></td><td colspan="1" rowspan="1" style="text-align: left;"><p>$NH3$| | 2 | 2 |$~104.5^ ext{o}$| Bent |$H2O$</p></td><td colspan="1" rowspan="1"><p></p></td></tr><tr><td colspan="1" rowspan="1" style="text-align: left;"><ul><li><p><strong>Basic Procedure to Determine Molecular Shape:</strong></p></li></ul></td><td colspan="1" rowspan="1" style="text-align: left;"><p></p></td><td colspan="1" rowspan="1" style="text-align: left;"><p></p></td><td colspan="1" rowspan="1" style="text-align: left;"><p></p></td><td colspan="1" rowspan="1" style="text-align: left;"><p></p></td><td colspan="1" rowspan="1"><p></p></td></tr></tbody></table><ol><li><p>Draw the Lewis structure for the molecule or ion.</p></li><li><p>Count the total number of bonded atoms and lone pairs around the central atom.</p></li><li><p>If no lone pairs are present on the central atom, the geometry is straightforward:</p><ul><li><p>2 bonded atoms: Linear.</p></li><li><p>3 bonded atoms: Trigonal planar.</p></li><li><p>4 bonded atoms: Tetrahedral.</p></li></ul></li><li><p>If one or more lone pairs are present on the central atom, consider their influence on the bonded atom arrangement and name the geometry accordingly (e.g., Bent, Trigonal pyramidal).</p></li></ol></li><li><p><strong>More Complex Molecules (Example: Dimethyl ether,$CH3OCH3$):</strong></p><ul><li><p>This molecule has two different central atoms (one oxygen and two carbons). VSEPR theory must be applied to each central atom independently.</p></li><li><p>Each$CH3$(methyl group) will have a tetrahedral geometry around its carbon atom, similar to methane ($CH4$).</p></li><li><p>The portion of the molecule linking the two methyl groups (the oxygen atom) will have bond angles similar to those in water ($H_2O$) due to two lone pairs and two bonded atoms around the oxygen, resulting in a bent geometry at the oxygen atom.</p></li></ul></li><li><p><strong>Determine the Molecular Geometry (Examples):</strong></p><ul><li><p>$PCl_3$: Trigonal pyramidal (4 electron groups, 3 bonded, 1 lone pair)</p></li><li><p>$SO_2$: Bent (3 electron groups, 2 bonded, 1 lone pair)</p></li><li><p>$PH_3$: Trigonal pyramidal (4 electron groups, 3 bonded, 1 lone pair)</p></li><li><p>$SiH_4$$: Tetrahedral (4 electron groups, 4 bonded, 0 lone pairs)