Ionic and Covalent Compounds Flashcards

Introduction to Chemical Bonding

• Definition of Chemical Bonding: Chemical bonding refers to the joining of two or more atoms in a stable arrangement to form chemical species with lower potential energy than the isolated atoms.

• Noble Gas Stability:

  • Noble gases located in Group 8A (Group 18) of the periodic table do not readily react to form bonds under normal conditions.

  • This lack of reactivity is due to their exceptionally stable electronic configurations, characterized by a completely filled outer valence shell.

• Driving Force of Bonding:

  • Elements other than noble gases gain, lose, or share electrons during chemical reactions to achieve an electron configuration identical to the noble gas closest to them on the periodic table.

• Primary Categories of Chemical Bonds:

  • Ionic Bonds: Formed when valence electrons are transferred completely from one atom to another. One atom loses valence electrons to become a positively charged cation, while another atom gains those valence electrons to become a negatively charged anion.

  • Covalent Bonds: Formed when two atoms share one or more pairs of valence electrons between them.

• Predicting Bond Types via Periodic Table Position:

  • Ionic Bonding: Occurs when a metal (found on the far left side of the periodic table) combines with a nonmetal (found on the far right side of the periodic table).

  • Covalent Bonding: Occurs when two nonmetals combine, or when a metalloid combines with a nonmetal.

Ions and the Octet Rule

• The Octet Rule:

  • During bond formation, main group atoms adjust their valence electrons by gaining, losing, or sharing to equal the noble gas electronic arrangement.

  • A main group element achieves maximal stability when it possesses an octet of eight valence electrons ($8\text{ e}^-$) in its outermost shell.

  • Exception: Helium ($\text{He}$) achieves maximal stability with a duet of two valence electrons ($2\text{ e}^-$) in its outer energy level ($1s^2$).

• General Definition of Ions:

  • An ion is a charged species formed when an atom or group of atoms carries an unequal number of protons and electrons.

  • Ionic compounds consist of oppositely charged ions held together by powerful electrostatic forces of attraction.

• Cations (Positively Charged Ions):

  • Formed when a neutral atom loses one or more valence electrons.

  • A cation possesses fewer electrons than protons, yielding a net positive charge.

  • Metals characteristically lose valence electrons during chemical reactions to form cations.

• Anions (Negatively Charged Ions):

  • Formed when a neutral atom gains one or more valence electrons.

  • An anion possesses more electrons than protons, yielding a net negative charge.

  • Nonmetals characteristically gain valence electrons during chemical reactions to form anions.

• Relating Group Numbers to Ionic Charges for Main Group Elements:

  • Elements belonging to the same group in the periodic table form ions carrying identical charges because they possess the same number of valence electrons.

  • Group 1A (Alkali Metals): Lose $1\text{ e}^-$ to form cations with a $+1$ charge (e.g., $\text{Li}^+$, $\text{Na}^+$, $\text{K}^+$, $\text{Rb}^+$, $\text{Cs}^+$).

  • Group 2A (Alkaline Earth Metals): Lose $2\text{ e}^-$ to form cations with a $+2$ charge (e.g., $\text{Be}^{2+}$, $\text{Mg}^{2+}$, $\text{Ca}^{2+}$, $\text{Sr}^{2+}$, $\text{Ba}^{2+}$).

  • Group 13A Metals: Aluminum ($\text{Al}$) loses $3\text{ e}^-$ to form a cation with a $+3$ charge ($\text{Al}^{3+}$).

  • Group 15A Nonmetals: Gain $3\text{ e}^-$ to form anions with a $-3$ charge (e.g., $\text{N}^{3-}$, $\text{P}^{3-}$).

  • Group 16A Nonmetals: Gain $2\text{ e}^-$ to form anions with a $-2$ charge (e.g., $\text{O}^{2-}$, $\text{S}^{2-}$, $\text{Se}^{2-}$).

  • Group 17A Nonmetals (Halogens): Gain $1\text{ e}^-$ to form anions with a $-1$ charge (e.g., $\text{F}^-$, $\text{Cl}^-$, $\text{Br}^-$, $\text{I}^-$).

Ionic Compounds: Formation, Properties, and Lattice Structures

• Formation of Ionic Compounds:

  • Ionic compounds are composed of oppositely charged cations and anions organized into crystalline lattices.

  • Formation involves the direct transfer of electrons from a metal atom to a nonmetal atom.

  • Example reaction of elemental sodium and chlorine gas yielding crystalline sodium chloride:         $\text{Na}(s) + \frac{1}{2}\text{Cl}_2(g) \rightarrow \text{NaCl}(s)$

  • Individual atoms alter their chemical and physical properties entirely upon forming ionic compounds. Metallic sodium (a highly reactive, soft metal) and chlorine gas (a toxic, yellow-green gas) combine to form sodium chloride (a benign table salt).

• Detailed Electronic Representation of Formation:

  • Sodium ($\text{Na}$) atom: $11\text{ protons}$, $12\text{ neutrons}$, $11\text{ electrons}$; electron configuration $1s^2 2s^2 2p^6 3s^1$. Losing $1\text{ e}^-$ yields a sodium cation ($\text{Na}^+$) with configuration $1s^2 2s^2 2p^6$ (isoelectronic with Neon).

  • Chlorine ($\text{Cl}$) atom: $17\text{ protons}$, $18\text{ neutrons}$, $17\text{ electrons}$; electron configuration $1s^2 2s^2 2p^6 3s^2 3p^5$. Gaining $1\text{ e}^-$ yields a chloride anion ($\text{Cl}^-$) with configuration $1s^2 2s^2 2p^6 3s^2 3p^6$ (isoelectronic with Argon).

• Principles of Writing Formulas for Ionic Compounds:

  • Ionic compounds are electrically neutral overall; therefore, the sum of positive charges from cations must equal the sum of negative charges from anions.

  • Step 1: Identify the cation and anion from the elements and assign their correct charges based on periodic position.

  • Step 2: Determine the lowest whole-number ratio of cations and anions necessary to make the net compound charge sum to zero.

    • If charges are equal in magnitude (e.g., $\text{Ca}^{2+}$ and $\text{O}^{2-}$), combine in a 1:1 ratio ($\text{CaO}$).

    • If charges differ in magnitude (e.g., $\text{Ca}^{2+}$ and $\text{Cl}^-$), combine two $\text{Cl}^-$ anions for every one $\text{Ca}^{2+}$ cation ($\text{CaCl}_2$).

  • Step 3: Write the formula placing the cation symbol first, followed by the anion symbol. Omit explicit charge signs and use subscripts to denote atom counts. Subscripts equal to 1 are omitted.

  • The Crisscross Method: The numerical value of the cation charge becomes the subscript of the anion, and the numerical value of the anion charge becomes the subscript of the cation. Subscripts must always be simplified to the lowest whole-number ratio.

• Physical Properties of Ionic Compounds:

  • Physical State: They form dense, highly ordered crystalline solids composed of alternating positive and negative ions arranged in a three-dimensional lattice.

  • Melting and Boiling Points: They possess exceptionally high melting points and boiling points due to the strong electrostatic forces holding the lattice together. For example, sodium chloride ($\text{NaCl}$) has a melting point of $801^\text{o}\text{C}$ and a boiling point of $1413^\text{o}\text{C}$.

  • Solubility and Electrical Conductivity:

    • When dissolved in water, the crystalline lattice breaks apart, releasing free cations and anions hydrated by water molecules.

    • Because free ions act as mobile charge carriers, aqueous solutions of ionic compounds conduct electricity efficiently.

Nomenclature of Ionic Compounds and Polyatomic Ions

• Naming Fixed-Charge Cations:

  • Main group metal cations that exhibit only one charge state are named simply by using the name of the parent metal element followed by the word "ion".

  • Examples: $\text{Na}^+$ = sodium, $\text{K}^+$ = potassium, $\text{Ca}^{2+}$ = calcium, $\text{Mg}^{2+}$ = magnesium.

• Naming Variable-Charge Cations (Transition Metals and Heavy Main Group Metals):

  • Transition metals and high atomic number metals can lose different numbers of outer electrons to form multiple distinct cations. Most transition metals form at least two different cations, with a $+2$ charge state being widespread.

  • Method 1 (Systematic / Stock System): Write the English name of the metal followed immediately by a Roman numeral in parentheses corresponding to the charge magnitude.

  • Method 2 (Common / Classical System): Use the root Latin/English element name with specific suffixes:

    • Suffix -ous designates the ion with the lower positive charge.

    • Suffix -ic designates the ion with the higher positive charge.

• Detailed Table of Variable-Charge Metals:

Note on Mercury(I): Mercury(I) exists natively as a diatomic complex $\text{Hg}_2^{2+}$, containing two bonded mercury atoms that together carry a net $+2$ charge ($+1$ charge contribution per atom).

• Exceptions to Roman Numeral Usage:

  • Roman numerals are never used for main group metals in Group 1A or Group 2A.

  • Roman numerals are omitted for five specific metal cations that possess fixed charges: Silver ($\text{Ag}^+$), Zinc ($\text{Zn}^{2+}$), Cadmium ($\text{Cd}^{2+}$), Aluminum ($\text{Al}^{3+}$), and the ammonium cation ($\text{NH}_4^+$).

• Rule for Heavy Main Group Metals:

  • Metals of high atomic number in Groups 3A–5A (e.g., $\text{Sn}$, $\text{Pb}$) typically form two cations:

    1. Charge equal to the group number minus 2 ($\text{group } - 2$).

    2. Charge equal to the group number ($\text{group }$).

• Naming Monatomic Anions:

  • Monatomic anions are named by taking the root of the nonmetal element name and adding the suffix -ide.

• Comprehensive Monatomic Anion Reference:

  • Bromine ($\text{Br}$) $\rightarrow$ Bromide ($\text{Br}^-$)

  • Chlorine ($\text{Cl}$) $\rightarrow$ Chloride ($\text{Cl}^-$)

  • Fluorine ($\text{F}$) $\rightarrow$ Fluoride ($\text{F}^-$)

  • Iodine ($\text{I}$) $\rightarrow$ Iodide ($\text{I}^-$)

  • Nitrogen ($\text{N}$) $\rightarrow$ Nitride ($\text{N}^{3-}$)

  • Oxygen ($\text{O}$) $\rightarrow$ Oxide ($\text{O}^{2-}$)

  • Phosphorus ($\text{P}$) $\rightarrow$ Phosphide ($\text{P}^{3-}$)

  • Sulfur ($\text{S}$) $\rightarrow$ Sulfide ($\text{S}^{2-}$)

• Polyatomic Ions:

  • A polyatomic ion is a cation or anion composed of two or more covalently bonded atoms carrying a net electrical charge.

• Polyatomic Anion Reference List:

  • Carbon-based:

    • Carbonate: $\text{CO}_3^{2-}$

    • Hydrogen carbonate (or bicarbonate): $\text{HCO}_3^-$

    • Acetate: $\text{CH}_3\text{CO}_2^-$ (or $\text{C}_2\text{H}_3\text{O}_2^-$)

    • Cyanide: $\text{CN}^-$

  • Nitrogen-based:

    • Nitrate: $\text{NO}_3^-$

    • Nitrite: $\text{NO}_2^-$

  • Oxygen-based:

    • Hydroxide: $\text{OH}^-$

  • Phosphorus-based:

    • Phosphate: $\text{PO}_4^{3-}$

    • Hydrogen phosphate: $\text{HPO}_4^{2-}$

    • Dihydrogen phosphate: $\text{H}_2\text{PO}_4^-$

  • Sulfur-based:

    • Sulfate: $\text{SO}_4^{2-}$

    • Hydrogen sulfate (or bisulfate): $\text{HSO}_4^-$

    • Sulfite: $\text{SO}_3^{2-}$

    • Hydrogen sulfite (or bisulfite): $\text{HSO}_3^-$

• Writing Formulas with Polyatomic Ions:

  • Treat polyatomic ions as intact single units.

  • When multiple polyatomic ions are required to balance charge, enclose the entire polyatomic formula inside parentheses and place the balancing subscript outside.

  • Example: Combining $\text{Mg}^{2+}$ and $\text{OH}^-$ requires two hydroxide ions to balance one magnesium cation, giving $\text{Mg(OH)}_2$

• Health and Medical Applications:

  • Calcium carbonate ($\text{CaCO}_3$): Active antacid compound in Tums.

  • Magnesium hydroxide ($\text{Mg(OH)}_2$): Active antacid and laxative component in Maalox.

  • Iron(II) sulfate ($\text{FeSO}_4$): Essential iron supplement used medically to treat anemia.

Periodic Trends in Atomic Properties

• Classification of Elements:

  • Metals: Shiny/lustrous, malleable, ductile, good conductors of heat and electricity; tend to lose electrons to form cations.

  • Nonmetals: Non-lustrous, brittle in solid form, poor conductors of heat and electricity; tend to gain electrons to form anions or share electrons in covalent bonds.

  • Metalloids: Display properties intermediate between metals and nonmetals.

  • Hydrogen Exception: Hydrogen is located in Group 1A but exhibits chemical properties characteristic of nonmetals.

• Atomic Radius Trends:

  • Down a Group: Increases from top to bottom due to the addition of principal energy shells ($n$ levels).

  • Across a Period: Decreases from left to right due to increasing nuclear charge (more protons in the nucleus) pulling the electron cloud closer.

  • Cation vs. Parent Atom: Cations are smaller than their neutral parent atoms because the loss of valence electrons often vacates an outer shell, and reduced electron-electron repulsion allows the nucleus to pull remaining electrons tighter (e.g., $\text{Ca} > \text{Ca}^{2+}$).

  • Anion vs. Parent Atom: Anions are larger than their neutral parent atoms because adding electrons increases electron-electron repulsion within orbitals, expanding the electron cloud (e.g., $\text{Br} < \text{Br}^-$).

• Ionization Energy (IE):

  • Definition: The amount of energy required to remove an electron from a neutral atom in its gaseous state.

  • Down a Group: Decreases down a group because atomic size increases, placing valence electrons farther from the nucleus where they experience weaker electrostatic pull.

  • Across a Period: Increases from left to right across a period because atomic size decreases, nuclear charge increases, and elements approach a filled octet.

Lewis Structures of Atoms and Molecules

• Lewis Structures for Atoms:

  • Represent valence electrons as dots drawn around the chemical symbol of an element.

  • For main group elements, the number of valence electrons equals the group number (A-system numbering).

  • Dots are placed singly on the four sides of the element symbol before pairing them.

  • Example: Boron ($\text{B}$) has atomic configuration $[\text{He}]2s^2 2p^1$; belonging to Group 3A, it possesses 3 valence electrons represented by 3 dots.

• Step-by-Step Guidelines for Drawing Molecular Lewis Structures:

  • Step 1: Calculate Total Valence Electrons:

    • Sum the group numbers of all constituent main group atoms in the chemical formula.

    • For polyatomic anions, add electrons equal to the negative charge magnitude.

    • For polyatomic cations, subtract electrons equal to the positive charge magnitude.

  • Step 2: Position Atoms and Form Single Bonds:

    • Select the central atom. The central atom is usually the least electronegative element and is never hydrogen. It is frequently the single unique atom in the formula.

    • Connect surrounding terminal atoms to the central atom using single bonds. Each single bond represents a shared pair of electrons ($2\text{ e}^-$).

  • Step 3: Distribute Remaining Electrons as Lone Pairs:

    • Place remaining electrons in pairs onto terminal atoms first until each achieves an octet (hydrogen receives only $2\text{ e}^-$).

    • Place any left-over electrons on the central atom as lone pairs.

    • Always place electrons around the most electronegative atoms first.

    • Crucial Rule: Do not exceed the total electron count calculated in Step 1.

  • Step 4: Form Multiple Bonds if Octets are Incomplete:

    • If a central atom lacks an octet after distributing all available electrons, convert one or more lone pairs from adjacent terminal atoms into shared double or triple bonds.

• Common Bonding Patterns and Lone Pairs:

  • Carbon ($\text{C}$): 4 bonds, 0 lone pairs. (Forms single, double, or triple bonds).

  • Nitrogen ($\text{N}$): 3 bonds, 1 lone pair.

  • Oxygen ($\text{O}$): 2 bonds, 2 lone pairs. (Forms single or double bonds).

  • Halogens ($\text{F}, \text{Cl}, \text{Br}, \text{I}$): 1 bond, 3 lone pairs. (Form single bonds only as terminal atoms).

  • Sulfur ($\text{S}$): Typically 2 bonds, 2 lone pairs (forms single or double bonds).

Covalent Bonding and Molecular Compounds

• Fundamentals of Covalent Bonding:

  • A covalent bond is a two-electron bond formed when two nonmetal atoms (or a nonmetal and a metalloid) share a pair of valence electrons.

  • A molecule is a discrete group of atoms joined together by covalent bonds.

  • Unshared electron pairs localized on a single atom are termed nonbonded electron pairs or lone pairs.

• Covalent Bond Order:

  • Single Bond: 1 shared pair of electrons ($2\text{ e}^-$).

  • Double Bond: 2 shared pairs of electrons ($4\text{ e}^-$).

  • Triple Bond: 3 shared pairs of electrons ($6\text{ e}^-$).

• Orbital Overlap Model:

  • Covalent bonds form when atomic orbitals containing unpaired electrons overlap in space.

  • In the $\text{H}_2$ molecule, two spherical $1s$ atomic orbitals overlap to form a localized bonding region.

  • In the $\text{Cl}_2$ molecule, two dumbbell-shaped $3p$ atomic orbitals overlap end-to-end to form the covalent bond.

Nomenclature of Covalent Compounds

• Rules for Naming Binary Molecular Compounds:

  • Step 1: Name the first element using its full English name. Name the second element using the root of its element name plus the suffix -ide.

  • Step 2: Add Greek numerical prefixes to indicate the exact number of atoms of each element present.

  • Prefix Selection Rule: Omit the prefix mono- for the first named element. For example, $\text{CO}$ is named carbon monoxide, not monocarbon monoxide.

  • Vowel Elision Rule: When a prefix ending in a or o precedes the word oxide, drop the final vowel of the prefix to prevent awkward double vowels (e.g., monoxide instead of monooxide, tetroxide instead of tetraoxide, pentoxide instead of pentaoxide).

• Greek Prefixes Reference:

  • 1 = Mono-

  • 2 = Di-

  • 3 = Tri-

  • 4 = Tetra-

  • 5 = Penta-

  • 6 = Hexa-

  • 7 = Hepta-

  • 8 = Octa-

  • 9 = Nona-

  • 10 = Deca-

• Nomenclature Examples:

  • $\text{NO}_2$: Nitrogen dioxide

  • $\text{N}_2\text{O}_4$: Dinitrogen tetroxide

  • $\text{SO}_2$: Sulfur dioxide

  • $\text{P}_3\text{Cl}_8$: Triphosphorus octachloride

Valence Shell Electron Pair Repulsion (VSEPR) Theory and Molecular Geometry

• VSEPR Theory Principles:

  • Electron group geometry around a central atom is determined by minimizing electrostatic repulsion between negative electron clouds.

  • Electron groups (bonding electron pairs and nonbonding lone pairs) space themselves as far apart as possible in three-dimensional space.

• Molecular Shapes Summary Table:

Electronegativity, Bond Polarity, and Molecular Polarity

• Electronegativity Definition:

  • Electronegativity ($\text{EN}$) is a measure of an atom's ability to attract shared electrons toward itself within a chemical bond.

  • Fluorine ($\text{F}$) is the most electronegative element and is assigned a maximum value of $4.0$ on the Pauling Scale.

  • Periodic Trends: Electronegativity increases from left to right across a period and decreases down a group.

• Classifying Bond Types via Electronegativity Differences (\begin{equation*}\Delta\text{EN}\end{equation*}):

• Nonpolar Covalent Bond (\begin{equation*}\Delta\text{EN} < 0.5\end{equation*}):

  • Electrons are shared equally between atoms.

  • Occurs between identical elements (e.g., $\text{Cl}_2$, $\text{N}_2$) or elements with similar electronegativities (e.g., $\text{C}-\text{H}$ bond where $\text{EN}_\text{C} = 2.5$ and $\text{EN}_\text{H} = 2.1$, yielding \begin{equation*}\Delta\text{EN} = 0.4\end{equation*}).

• Polar Covalent Bond (\begin{equation*}0.5 \le \Delta\text{EN} \le 1.9\end{equation*}):

  • Electrons are shared unequally, pulled toward the more electronegative element.

  • Produces a permanent bond dipole, characterized by a partial negative charge (\begin{equation*}\delta^-\end{equation*}) on the more electronegative atom and a partial positive charge (\begin{equation*}\delta^+\end{equation*}) on the less electronegative atom.

  • Example: $\text{C}-\text{O}$ bond where $\text{EN}_\text{O} = 3.5$ and $\text{EN}_\text{C} = 2.5$, yielding \begin{equation*}\Delta\text{EN} = 1.0\end{equation*} (\begin{equation*}\delta^+\text{C}\rightarrow\text{O}\delta^-\end{equation*}).

• Ionic Bond (\begin{equation*}\Delta\text{EN} > 1.9\end{equation*}):

  • Electrons are fully transferred from the less electronegative element to the more electronegative element.

• Determining Overall Molecular Polarity:

  • Molecular polarity depends on both the presence of polar bonds and the 3D molecular geometry.

  • Nonpolar Molecules:

    • Possess no polar bonds, OR

    • Possess individual bond dipoles that cancel out due to symmetric geometry (e.g., linear $\text{CO}_2$, trigonal planar $\text{BF}_3$, tetrahedral $\text{CCl}_4$).

  • Polar Molecules:

    • Possess a single polar bond, OR

    • Possess multiple polar bonds whose dipoles reinforce each other without cancelling due to asymmetric geometry (e.g., bent $\text{H}_2\text{O}$, trigonal pyramidal $\text{NH}_3$).

Comprehensive Practice Problems and Applications

• Combining Elements into Ionic Formulas:

  • $\text{Mg}$ ($+2$) and $\text{F}$ ($-1$) $\rightarrow \text{MgF}_2$ (Magnesium fluoride)

  • $\text{Li}$ ($+1$) and $\text{P}$ ($-3$) $\rightarrow \text{Li}_3\text{P}$ (Lithium phosphide)

  • $\text{Al}$ ($+3$) and $\text{O}$ ($-2$) $\rightarrow \text{Al}_2\text{O}_3$ (Aluminum oxide)

  • $\text{Be}$ ($+2$) and $\text{S}$ ($-2$) $\rightarrow \text{BeS}$ (Beryllium sulfide)

• Naming Transition Metal Ionic Compounds:

  • $\text{CuO}$: Copper(II) oxide or cupric oxide

  • $\text{Cu}_2\text{O}$: Copper(I) oxide or cuprous oxide

  • $\text{CrS}$: Chromium(II) sulfide or chromous sulfide

  • $\text{PbO}_2$: Lead(IV) oxide or plumbic oxide

  • $\text{AuP}$: Gold(III) phosphide or auric phosphide

  • $\text{Cu}_2\text{S}$: Copper(I) sulfide or cuprous sulfide

  • $\text{AuBr}_3$: Gold(III) bromide or auric bromide

  • $\text{ZnCl}_2$: Zinc chloride

• Deriving Ionic Formulas from Names:

  • Lead(IV) bromide $\rightarrow \text{Pb}^{4+}$ and $\text{Br}^- \rightarrow \text{PbBr}_4$

  • Silver iodide $\rightarrow \text{Ag}^+$ and $\text{I}^- \rightarrow \text{AgI}$

  • Magnesium nitride $\rightarrow \text{Mg}^{2+}$ and $\text{N}^{3-} \rightarrow \text{Mg}_3\text{N}_2$

  • Potassium selenide $\rightarrow \text{K}^+$ and $\text{Se}^{2-} \rightarrow \text{K}_2\text{Se}$

  • Barium phosphide $\rightarrow \text{Ba}^{2+}$ and $\text{P}^{3-} \rightarrow \text{Ba}_3\text{P}_2$

  • Lead(IV) carbonate $\rightarrow \text{Pb}^{4+}$ and $\text{CO}_3^{2-} \rightarrow \text{Pb(CO}_3)_2$

• Multiple-Choice Practice Solutions:

  • Question: Symbol and name of ion with $16\text{ protons}$ and $18\text{ electrons}$.

    • Answer: $\text{S}^{2-}$ (Sulfide ion)

  • Question: Correct formula for ionic compound formed between potassium and oxygen.

    • Answer: $\text{K}_2\text{O}$

  • Question: Correct name for $\text{Cu}_2\text{S}$.

    • Answer: Copper(I) sulfide

  • Question: Correct formula for baking soda (sodium bicarbonate).

    • Answer: $\text{NaHCO}_3$

  • Question: Classify $\text{NO}_2$ and state its name.

    • Answer: Molecular, nitrogen dioxide

  • Question: Identify the nonpolar covalent bond among $\text{N}-\text{P}$, $\text{Br}-\text{Br}$, $\text{Br}-\text{F}$, $\text{Si}-\text{Br}$.

    • Answer: $\text{Br}-\text{Br}$

  • Question: Shape of a molecule with four bonded atoms and no lone pairs.

    • Answer: Tetrahedral

  • Question: Correct chemical formula for compound made from aluminum and sulfur.

    • Answer: $\text{Al}_2\text{S}_3$

  • Question: Correct chemical name for $\text{PbO}_2$.

    • Answer: Lead(IV) oxide

  • Question: Correct chemical name for \text{Fe}_3(\text{PO}_4)_2$.\n * *Answer:* Iron(II) phosphate\n * *Question:* Correct chemical name for ext{P}_2 ext{O}_5$.

    • Answer: Diphosphorus pentoxide

  • Question: Correct chemical name for \text{KNO}_3$.\n * *Answer:* Potassium nitrate\n * *Question:* Which molecule is nonpolar: ext{HBr}$,$ ext{NF}_3$,$ ext{CBr}_4$,$ ext{HCl}?\n * *Answer:* ext{CBr}_4\n * *Question:* Is the molecule ext{SCl}_2 ionic, polar, or nonpolar?\n * *Answer:* Polar (bent geometry with 2 lone pairs on ext{S})\n* **Structural and Polarity Analysis Exercises:**\n * *Phosphate Ion (egin{equation}\text{PO}_4^{3-}\end{equation}):*\n * Total valence electrons = 5 ( ext{P}) + 4 imes 6 ( ext{O}) + 3 ( ext{charge}) = 32 ext{ e}^-.\n * Shape: Tetrahedral around central phosphorus atom.\n * Polarity: Nonpolar ion/species overall due to symmetric cancellation of polar ext{P}- ext{O} bond dipoles.\n * *Silicon Tetrafluoride (egin{equation}\text{SiF}_4\end{equation}):*\n * Total valence electrons = 4 ( ext{Si}) + 4 imes 7 ( ext{F}) = 32 ext{ e}^-.\n * Shape: Tetrahedral around central silicon atom.\n * Polarity: Nonpolar molecule due to symmetric cancellation of polar ext{Si}- ext{F} bond dipoles.\n * *Methylamine Structure Analysis:*\n\n\n\n * Number of polar bonds = 3 (1 ext{C}- ext{N}$bond and 2$ ext{N}- ext{H} bonds).\n * Number of nonpolar bonds = 3 (3 ext{C}- ext{H}$bonds, ignoring the small$egin{equation}\Delta\text{EN} = 0.4\end{equation}$$ difference).