Second Test Review - Chemical Formulas and Molecular Structures

Chemical Formulas

• Definition: A chemical formula is a concise representation using chemical symbols and numerical subscripts that specifies the exact elements and relative proportions of atoms present in a compound.

• Empirical Formula: Represents the simplest whole-number ratio of atoms of each element present in a chemical compound.

  • Example: Hydrogen peroxide has a molecular formula of $H_2O_2$ and an empirical formula of $HO$.
  • Example: Glucose has a molecular formula of $C_6H_{12}O_6$ and an empirical formula of $CH_2O$.

• Molecular Formula: Specifies the actual exact number of atoms of each element present in a single discrete molecule of a covalent compound.

  • Example: Water has a molecular formula of $H_2O$.
  • Example: Benzene has a molecular formula of $C_6H_6$ (while its empirical formula is $CH$).

• Structural Formula: Illustrates how atoms are bonded to each other and spatially arranged within a molecule, displaying single, double, or triple covalent bonds.

• Formula Unit: The simplest neutral repeating unit of an ionic lattice structure, representing the lowest whole-number ratio of cations to anions (e.g., $NaCl$, $MgCl_2$).

Lewis Dot Structures of Ions and Molecules

• Valence Electrons: Electrons situated in the outermost energy level (shell) of an atom that actively participate in chemical bond formation.

• Octet Rule: Atoms tend to gain, lose, or share electrons in order to achieve a stable electronic configuration with eight valence electrons, matching the outer electron configuration of noble gases.

  • Duet Rule: Applies specifically to hydrogen ($H$) and helium ($He$), which reach maximum valence stability with two electrons in their outer shell.

• Lewis Representation of Main Group Atoms: Valence electrons are represented as dots placed around the four sides of the atomic symbol.

  • Group 1 (Alkali Metals): 1 valence electron ($Li\text{--}$ or $Na\text{--}$).
  • Group 2 (Alkaline Earth Metals): 2 valence electrons ($\text{--}Mg\text{--}$).
  • Group 13: 3 valence electrons.
  • Group 14: 4 valence electrons.
  • Group 15: 5 valence electrons.
  • Group 16: 6 valence electrons.
  • Group 17 (Halogens): 7 valence electrons.
  • Group 18 (Noble Gases): 8 valence electrons (except $He$, which has 2).

• Lewis Structures for Cations: Formed when a neutral metal atom loses its valence electrons.

  • Represented by placing the atomic symbol in square brackets (or without dots) with a positive superscript charge: $[Na]^+$ or $Na^+$, $[Mg]^{2+}$ or $Mg^{2+}$, $[Al]^{3+}$ or $Al^{3+}$.

• Lewis Structures for Anions: Formed when a nonmetal atom gains electrons to complete its valence octet.

  • Represented by placing eight dots around the symbol inside square brackets with a negative superscript charge: $[:\text{Cl}:]^-$ or $[:\text{O}:]^{2-}$.

• Systematic Rules for Molecular Lewis Structures:

  1. Sum all valence electrons from every atom in the formula.
  2. Adjust electron count for net electrical charge (add electrons for negative charges, subtract for positive charges).
  3. Determine the central atom (the least electronegative element, excluding hydrogen, which is always terminal).
  4. Connect surrounding atoms to the central atom using single covalent bonds (each line represents a shared electron pair of 2 electrons).
  5. Subtract bonding electrons from the total electron count.
  6. Place remaining electrons as lone pairs on terminal atoms to satisfy their octets.
  7. Assign any leftover electrons to the central atom.
  8. If the central atom lacks a complete octet, convert lone pairs from terminal atoms into double or triple bonds.

Predicting the Charge of an Ion

• Main Group Elements Charge Patterns:
  • Group 1 (Alkali Metals): Lose 1 electron to form cations with a charge of $+1$ (e.g., $Li^+$, $Na^+$, $K^+$).
  • Group 2 (Alkaline Earth Metals): Lose 2 electrons to form cations with a charge of $+2$ (e.g., $Mg^{2+}$, $Ca^{2+}$, $Ba^{2+}$).
  • Group 13 Metals: Lose 3 electrons to form cations with a charge of $+3$ (e.g., $Al^{3+}$).
  • Group 15 Nonmetals: Gain 3 electrons to form anions with a charge of $-3$ (e.g., $N^{3-}$, $P^{3-}$).
  • Group 16 Nonmetals: Gain 2 electrons to form anions with a charge of $-2$ (e.g., $O^{2-}$, $S^{2-}$).
  • Group 17 Nonmetals (Halogens): Gain 1 electron to form anions with a charge of $-1$ (e.g., $F^-$, $Cl^-$, $Br^-$, $I^-$).
  • Group 18 (Noble Gases): Do not form ions under normal conditions; charge is 0$.\n\n- **Transition Metals and Variable Oxidation States**:\n - Transition metals can form cations with varying positive charges due to partial loss of d-orbitals electrons.\n - Expressed using Roman numerals in parentheses (e.g., Iron(II) is Fe^{2+}$, Iron(III) is$Fe^{3+}$, Copper(I) is$Cu^+$, Copper(II) is$Cu^{2+}).\n\n- **Common Polyatomic Ions**:\n - Ammonium: NH_4^+\n - Hydronium: H_3O^+\n - Hydroxide: OH^-\n - Nitrate: NO_3^-\n - Nitrite: NO_2^-\n - Sulfate: SO_4^{2-}\n - Sulfite: SO_3^{2-}\n - Carbonate: CO_3^{2-}\n - Hydrogen carbonate (Bicarbonate): HCO_3^-\n - Phosphate: PO_4^{3-}\n - Acetate: C_2H_3O_2^-$or$CH_3COO^-\n - Permanganate: MnO_4^-\n - Cyanide: CN^-\n\n# Predicting the Chemical Formula of an Ionic Compound\n\n- **Principle of Electroneutrality**: Ionic compounds are electrically neutral. The overall total positive charge must equal the overall total negative charge so the net charge is zero.\n\n- **Criss-Cross Strategy**:\n - Write the chemical symbols and charges for the cation first and anion second (A^{m+}$and$B^{n-}).\n - Cross over the numerical charge magnitude of each ion to become the subscript of the opposite ion (A_n B_m).\n - Reduce subscripts to the lowest whole-number ratio (e.g., Ti^{4+}$and$O^{2-}$yields$Ti_2O_4$, which reduces to$TiO_2$;$Ca^{2+}$and$O^{2-}$yields$CaO).\n - Enclose polyatomic ions in parentheses whenever the assigned subscript is greater than 1 (e.g., Ca^{2+}$and$NO_3^-$forms$Ca(NO_3)_2$;$Al^{3+}$and$SO_4^{2-}$forms$Al_2(SO_4)_3).\n\n# Naming Ionic Compounds\n\n- **Binary Ionic Compounds (Fixed-Charge Metal + Nonmetal)**:\n - Name the cation (metal) first using its element name.\n - Name the anion (nonmetal) second, modifying its root name to end with the suffix "-ide".\n - Example: NaCl is sodium chloride.\n - Example: Mg_3N_2 is magnesium nitride.\n\n- **Ionic Compounds with Variable-Charge Metals (Stock System)**:\n - Determine the charge of the metal cation required to balance the negative anion charges.\n - Place a Roman numeral inside parentheses directly following the metal name to indicate its ionic charge.\n - Example: FeCl_2 is iron(II) chloride.\n - Example: FeCl_3 is iron(III) chloride.\n - Example: Cu_2O is copper(I) oxide.\n\n- **Ionic Compounds Containing Polyatomic Ions**:\n - State the name of the cation, followed by the exact un-modified name of the polyatomic anion.\n - Example: Na_2SO_4 is sodium sulfate.\n - Example: NH_4Cl is ammonium chloride.\n - Example: Mg(NO_3)_2 is magnesium nitrate.\n\n# Calculating the Mass of a Compound Based on Its Formula\n\n- **Molar Mass (M$)**: The sum of the atomic masses of all individual atoms present in a chemical formula, expressed in units of$g ,mol^{-1}$or atomic mass units ($u).\n\n- **Calculation Procedure**:\n 1. Count the number of atoms of each element in the chemical formula.\n 2. Multiply the atom count of each element by its standard atomic mass (A_r).\n 3. Sum all calculated masses together to determine total molar mass.\n\n- **Calculation Example: Water (H_2O)**:\n - Hydrogen (H$):$2 \times 1.008 ,g ,mol^{-1} = 2.016 ,g ,mol^{-1}\n - Oxygen (O$):$1 \times 16.00 ,g ,mol^{-1} = 16.00 ,g ,mol^{-1}\n - Total Molar Mass of H_2O$:$2.016 + 16.00 = 18.016 ,g ,mol^{-1}\n\n- **Calculation Example: Calcium Nitrate (Ca(NO_3)_2)**:\n - Calcium (Ca$):$1 \times 40.08 ,g ,mol^{-1} = 40.08 ,g ,mol^{-1}\n - Nitrogen (N$):$2 \times 14.01 ,g ,mol^{-1} = 28.02 ,g ,mol^{-1}\n - Oxygen (O$):$6 \times 16.00 ,g ,mol^{-1} = 96.00 ,g ,mol^{-1}\n - Total Molar Mass of Ca(NO_3)_2$:$40.08 + 28.02 + 96.00 = 164.10 ,g ,mol^{-1}\n\n# Determining if Bonds in a Molecule are Ionic or Covalent\n\n- **Electronegativity Difference (\Delta EN)**:\n - Electronegativity is a measure of an atom's ability to attract shared electrons in a bond.\n - \Delta EN = |EN_A - EN_B|\n\n- **Classification Thresholds**:\n - **Nonpolar Covalent Bond**: \Delta EN \le 0.4 (electrons are shared equally between atoms).\n - **Polar Covalent Bond**: 0.4 < \Delta EN \le 1.7$(electrons are shared unequally, producing partial positive$\delta^+$and partial negative$\delta^- charges).\n - **Ionic Bond**: \Delta EN > 1.7 (complete transfer of valence electrons from a metal to a nonmetal, generating electrostatic attraction between ions).\n\n- **General Elemental Rules**:\n - Metal + Nonmetal = Ionic bonding.\n - Nonmetal + Nonmetal = Covalent bonding.\n - Metalloid + Nonmetal = Covalent bonding.\n\n# Predicting the Number of Bonds an Element Will Form\n\n- **Bonding Capacity (Valence)**: Main group nonmetals form covalent bonds to achieve a full valence octet.\n - Formula for bonding capacity: \text{Number of Bonds} = 8 - \text{Group Number} (under standard 8-column group notation).\n - Group 14 (Carbon, Silicon): 4 valence electrons; forms 4 covalent bonds.\n - Group 15 (Nitrogen, Phosphorus): 5 valence electrons; forms 3 covalent bonds (and retains 1 lone pair).\n - Group 16 (Oxygen, Sulfur): 6 valence electrons; forms 2 covalent bonds (and retains 2 lone pairs).\n - Group 17 (Halogens: Fluorine, Chlorine, Bromine, Iodine): 7 valence electrons; forms 1 covalent bond (and retains 3 lone pairs).\n - Hydrogen (H): 1 valence electron; forms 1 covalent bond.\n\n# Predicting Single, Double, or Triple Bonds\n\n- **Bond Multiplicities**:\n - **Single Bond**: Shares 1 pair of electrons (2 electrons). Has the lowest bond energy and longest bond length.\n - **Double Bond**: Shares 2 pairs of electrons (4 electrons). Has intermediate bond energy and length.\n - **Triple Bond**: Shares 3 pairs of electrons (6 electrons). Has the highest bond energy and shortest bond length.\n\n- **Determination by Electron Counting**:\n - Required octet electrons (N$):$N = 8 \times (\text{non-H atoms}) + 2 \times (\text{H atoms}).\n - Total valence electrons available (V).\n - Shared bonding electrons (S$):$S = N - V$.
  • Total number of bonds: $\text{Bonds} = \frac{S}{2}$.
  • Example ($O_2$): Requires 16 octet electrons, possesses 12 valence electrons; $S = 16 - 12 = 4$ shared electrons $\rightarrow$ forms a double bond ($O=O$).
  • Example ($N_2$): Requires 16 octet electrons, possesses 10 valence electrons; $S = 16 - 10 = 6$ shared electrons $\rightarrow$ forms a triple bond ($N\equiv N$).
  • Example ($CO_2$): Requires 24 octet electrons, possesses 16 valence electrons; $S = 24 - 16 = 8$ shared electrons $\rightarrow$ forms two double bonds ($O=C=O$).

Lewis Dot Structures of Molecules

• Summary of Procedure:

  • Sum total available valence electrons.
  • Connect terminal atoms to the central atom with single bonds.
  • Complete octets of surrounding atoms.
  • Place remaining electrons on the central atom.
  • Form double or triple bonds if the central atom lacks an octet.

• Exceptions to the Octet Rule:

  • Incomplete Octet: Stable compounds with fewer than 8 valence electrons around the central atom.
  • Beryllium ($Be$) is stable with 4 valence electrons (e.g., $BeCl_2$).
  • Boron ($B$) is stable with 6 valence electrons (e.g., $BF_3$).
  • Odd-Electron Molecules: Species with an odd number of total valence electrons (free radicals) cannot fulfill the octet rule for every atom (e.g., $NO$ with 11 valence electrons, $NO_2$ with 17 valence electrons).
  • Expanded Octet: Central elements in Period 3 or higher ($P$, $S$, $Cl$, $Br$, $I$, $Xe$) utilize vacant d-orbitals to hold more than 8 valence electrons.
  • Example: Phosphorus pentachloride ($PCl_5$) has 10 electrons around phosphorus.
  • Example: Sulfur hexafluoride ($SF_6$) has 12 electrons around sulfur.

Naming Molecular Compounds

• Greek Prefix System for Binary Covalent Compounds:

  • Name the first element with its standard element name, using a Greek prefix to indicate the number of atoms (if greater than 1).
  • Name the second element using a Greek prefix indicating the number of atoms, followed by the element root name and the suffix "-ide".

• Greek Prefixes:

  • 1 = mono- (omitted for the first element, retained for the second)
  • 2 = di-
  • 3 = tri-
  • 4 = tetra-
  • 5 = penta-
  • 6 = hexa-
  • 7 = hepta-
  • 8 = octa-
  • 9 = nona-
  • 10 = deca-

• Examples:

  • $CO$: carbon monoxide
  • $CO_2$: carbon dioxide
  • $N_2O_4$: dinitrogen tetroxide
  • $P_4O_{10}$: tetraphosphorus decoxide
  • $SF_6$: sulfur hexafluoride

Determining if a Molecule is Ionic or Covalent

• Structural and Physical Differences:
  • Ionic Compounds: Form extended 3D crystalline lattices held together by strong electrostatic interactions between oppositely charged ions. Typically feature high melting and boiling points, high water solubility, and electrical conductivity when molten $(l)$ or dissolved in water $(aq)$.
  • Covalent (Molecular) Compounds: Consist of discrete, individual molecules held internally by shared electron pairs and externally by weaker intermolecular forces. Typically feature lower melting and boiling points, variable water solubility, and do not conduct electricity in liquid or aqueous states (except ionizable acids).

Identifying Molecules with Polar Bonds

• Bond Polarity Characteristics:

  • Occurs when two bonded atoms have an electronegativity difference ($\Delta EN$) between $0.5$ and $1.7$.
  • Creates a partial positive charge ($\delta^+$) on the less electronegative atom and a partial negative charge ($\delta^-$) on the more electronegative atom.
  • Represented by a dipole arrow directed from $\delta^+$ toward $\delta^-$ with a cross on the positive end.

• Bond Examples:

  • Carbon-Hydrogen bond ($C-H$): $\Delta EN = 2.5 - 2.1 = 0.4 \rightarrow$ Nonpolar bond.
  • Oxygen-Hydrogen bond ($O-H$): $\Delta EN = 3.5 - 2.1 = 1.4 \rightarrow$ Polar bond.
  • Fluorine-Carbon bond ($F-C$): $\Delta EN = 4.0 - 2.5 = 1.5 \rightarrow$ Polar bond.

Predicting the Shape of a Molecule

• Valence Shell Electron Pair Repulsion (VSEPR) Theory:

  • Valence electron pairs (bonding pairs and lone pairs) surrounding a central atom experience mutual electrostatic repulsion and position themselves as far apart as possible in 3D space.

• Steric Number Calculation:

  • $\text{Steric Number} = (\text{Number of atoms bonded to central atom}) + (\text{Number of lone pairs on central atom})$.

• VSEPR Geometries and Angles:

  • Steric Number = 2:
  • 2 bonded atoms, 0 lone pairs: Linear, bond angle = $180^\circ$ (e.g., $CO_2$, $BeCl_2$).
  • Steric Number = 3:
  • 3 bonded atoms, 0 lone pairs: Trigonal Planar, bond angle = $120^\circ$ (e.g., $BF_3$).
  • 2 bonded atoms, 1 lone pair: Bent / V-shaped, bond angle = $< 120^\circ$ (e.g., $SO_2$).
  • Steric Number = 4:
  • 4 bonded atoms, 0 lone pairs: Tetrahedral, bond angle = $109.5^\circ$ (e.g., $CH_4$).
  • 3 bonded atoms, 1 lone pair: Trigonal Pyramidal, bond angle = $107.5^\circ$ (e.g., $NH_3$).
  • 2 bonded atoms, 2 lone pairs: Bent / V-shaped, bond angle = $104.5^\circ$ (e.g., $H_2O$).
  • Steric Number = 5:
  • 5 bonded atoms, 0 lone pairs: Trigonal Bipyramidal, bond angles = $90^\circ$ (axial) and $120^\circ$ (equatorial) (e.g., $PCl_5$).
  • 4 bonded atoms, 1 lone pair: Seesaw (e.g., $SF_4$).
  • 3 bonded atoms, 2 lone pairs: T-shaped (e.g., $ClF_3$).
  • 2 bonded atoms, 3 lone pairs: Linear (e.g., $XeF_2$).
  • Steric Number = 6:
  • 6 bonded atoms, 0 lone pairs: Octahedral, bond angle = $90^\circ$ (e.g., $SF_6$).
  • 5 bonded atoms, 1 lone pair: Square Pyramidal (e.g., $BrF_5$).
  • 4 bonded atoms, 2 lone pairs: Square Planar (e.g., $XeF_4$).

Identifying Molecules that are Polar

• Molecular Polarity vs. Bond Polarity:

  • Molecular polarity requires both the presence of polar bonds and an asymmetrical 3D molecular geometry.
  • Net Dipole Moment ($\mu$): Vector sum of all individual bond dipole moments.

• Nonpolar Molecules:

  • Occur when polar bond dipoles cancel out due to symmetrical geometry, resulting in a net dipole moment of zero ($\mu = 0$).
  • Symmetrical geometries with identical terminal atoms: Linear, Trigonal Planar, Tetrahedral, Trigonal Bipyramidal, Octahedral.
  • Example: $CO_2$ is linear; opposing equal bond dipoles cancel ($\mu = 0$).
  • Example: $CCl_4$ is tetrahedral; four symmetrical bond dipoles cancel ($\mu = 0$).
  • Example: $BF_3$ is trigonal planar; three symmetrical bond dipoles cancel ($\mu = 0$).

• Polar Molecules:

  • Occur when bond dipoles do not cancel out due to asymmetric shape or non-identical terminal atoms, giving a net dipole moment greater than zero ($\mu > 0$).
  • Non-zero lone pairs on central atom typically produce asymmetric geometries (e.g., Bent, Trigonal Pyramidal).
  • Example: $H_2O$ is bent; bond dipoles combine to create a net dipole pointing toward oxygen ($\mu > 0$).
  • Example: $NH_3$ is trigonal pyramidal; net dipole points toward nitrogen ($\mu > 0$).
  • Example: $CH_3Cl$ is tetrahedral but asymmetric due to chlorine ($\mu > 0$).

Identifying and Naming Alkanes, Alkenes, and Alkynes

• Hydrocarbons: Organic molecules composed entirely of carbon and hydrogen atoms.

• Alkanes:

  • Saturated hydrocarbons containing exclusively carbon-carbon single bonds ($C-C$).
  • General chemical formula: $C_n H_{2n+2}$.
  • IUPAC Suffix: "-ane".
  • Straight-Chain Root Prefixes:
  • 1 Carbon ($n=1$): Methane ($CH_4$)
  • 2 Carbons ($n=2$): Ethane ($C_2H_6$)
  • 3 Carbons ($n=3$): Propane ($C_3H_8$)
  • 4 Carbons ($n=4$): Butane ($C_4H_{10}$)
  • 5 Carbons ($n=5$): Pentane ($C_5H_{12}$)
  • 6 Carbons ($n=6$): Hexane ($C_6H_{14}$)
  • 7 Carbons ($n=7$): Heptane ($C_7H_{16}$)
  • 8 Carbons ($n=8$): Octane ($C_8H_{18}$)
  • 9 Carbons ($n=9$): Nonane ($C_9H_{20}$)
  • 10 Carbons ($n=10$): Decane ($C_{10}H_{22}$)

• Alkenes:

  • Unsaturated hydrocarbons containing at least one carbon-carbon double bond ($C=C$).
  • General chemical formula (monounsaturated): $C_n H_{2n}$.
  • IUPAC Suffix: "-ene".
  • Nomenclature: Number the carbon chain starting from the end closest to the double bond; indicate the double bond position using the lower carbon index.
  • Example: $CH_2=CH_2$ is ethene.
  • Example: $CH_2=CH-CH_2-CH_3$ is 1-butene.
  • Example: $CH_3-CH=CH-CH_3$ is 2-butene.

• Alkynes:

  • Unsaturated hydrocarbons containing at least one carbon-carbon triple bond ($C\equiv C$).
  • General chemical formula (monounsaturated): $C_n H_{2n-2}$.
  • IUPAC Suffix: "-yne".
  • Nomenclature: Number the carbon chain starting from the end closest to the triple bond; indicate the triple bond position using the lower carbon index.
  • Example: $HC\equiv CH$ is ethyne (acetylene).
  • Example: $HC\equiv C-CH_3$ is propyne.
  • Example: $HC\equiv C-CH_2-CH_3$ is 1-butyne.
  • Example: $CH_3-C\equiv C-CH_3$ is 2-butyne.

Identifying and Naming Molecules Containing Oxygen

• Alcohols:

  • Functional group: Hydroxyl group ($-OH$) attached to a carbon atom.
  • General formula: $R-OH$.
  • IUPAC Naming: Drop the ending "-e" from the corresponding alkane name and add the suffix "-ol"; use a number to specify the position of the $-OH$ group.
  • Example: $CH_3OH$ is methanol.
  • Example: $CH_3CH_2OH$ is ethanol.
  • Example: $CH_3CH(OH)CH_3$ is 2-propanol.

• Ethers:

  • Functional group: Oxygen atom single-bonded between two carbon/alkyl groups ($-O-$).
  • General formula: $R-O-R'$.
  • IUPAC Naming: Name the two alkyl groups in alphabetical order followed by the word "ether".
  • Example: $CH_3-O-CH_3$ is dimethyl ether.
  • Example: $CH_3CH_2-O-CH_2CH_3$ is diethyl ether.
  • Example: $CH_3-O-CH_2CH_3$ is ethyl methyl ether.

• Aldehydes:

  • Functional group: Carbonyl group ($C=O$) bonded to at least one hydrogen atom at the end of a carbon chain ($-CHO$).
  • General formula: $R-CHO$.
  • IUPAC Naming: Drop the ending "-e" from the alkane name and add the suffix "-al".
  • Example: $HCHO$ is methanal (formaldehyde).
  • Example: $CH_3CHO$ is ethanal (acetaldehyde).

• Ketones:

  • Functional group: Carbonyl group ($C=O$) bonded between two internal carbon atoms.
  • General formula: $R-CO-R'$.
  • IUPAC Naming: Drop the ending "-e" from the alkane name and add the suffix "-one"; use a number to locate the carbonyl position.
  • Example: $CH_3COCH_3$ is propanone (acetone).
  • Example: $CH_3COCH_2CH_3$ is 2-butanone.

• Carboxylic Acids:

  • Functional group: Carboxyl group ($-COOH$), composed of a carbonyl ($C=O$) and hydroxyl ($-OH$) on the same carbon atom.
  • General formula: $R-COOH$.
  • IUPAC Naming: Drop the ending "-e" from the alkane name and add the suffix "-oic acid".
  • Example: $HCOOH$ is methanoic acid (formic acid).
  • Example: $CH_3COOH$ is ethanoic acid (acetic acid).

• Esters:

  • Functional group: Carbonyl attached to an alkoxy group ($-COO-$).
  • General formula: $R-COO-R'$.
  • IUPAC Naming: Name the alkyl group attached to the oxygen atom ($R'$) first, followed by the parent carboxylate root ($R$) ending in "-oate".
  • Example: $HCOOCH_3$ is methyl methanoate.
  • Example: $CH_3COOCH_2CH_3$ is ethyl ethanoate (ethyl acetate).