Lecture 7 Bond Properties and Inorganic Compounds PDF

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Bond Properties and Inorganic Compounds ●

Formal Charge Patterns ○ The best Lewis structure or resonance contributing structure has the least number of atoms with formal charge ○ Equivalent atoms have the same formal charge. For example, all the hydrogen atoms of methane (CH4 ) are equivalent and therefore have the same formal charge. All six hydrogens of ethane (H3C-CH3 ) have the same formal charge, as do the two carbon atoms ○ Any hydrogen bearing one covalent bond always has a formal charge of zero ○ Formal charges other than +1, 0 or -1 are uncommon except for metals ○ The vast majority of organic structures are made up of a small set of atoms with a limited number of bonding possibilities. Recognizing these cases will allow you to avoid FC calculations most of the time, and speed your understanding of how charge influences reactions and properties of molecules. These patterns are summarized in the table below ○ Some elements fail to obey octet rule when bonding ■ For the following elements, “less is more” ● Hydrogen (1 bond always) ● Beryllium (2 bonds yield FC = 0) ● Boron & Aluminum (3 bonds yield FC = 0)

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Organic compounds ○ Based on a framework of carbon atoms and contain mainly H, N, O, halogens ○ Carbon makes 4 bonds ○ Nitrogen makes 3 bonds ○ Oxygen makes 2 bonds ○ Halogens make 1 bond ○ Hydrogen makes 1 bond Question: What is the relationship between having full valence shells and formal charges? Do full valence shells always result in a formal charge of zero? ○ Answer: Valence shell occupancy alone does not determine formal charge. The

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element involved also matters. For example, the full valence shell for carbon in methane (CH4) results in a formal charge of zero for carbon, whereas the full valence shell for nitrogen in the ammonium cation (NH4+) results in a +1 formal charge for nitrogen. Bond characteristics – amount of e- sharing ○ Bond polarity (lecture 5) ○ Bond order (BO) ■ Number of bonding pairs of electrons between atoms ■ 1 for single bond, 2 for double bond, 3 for triple bond ○ Bond length (l) ■ Equilibrium distance between pair of nuclei joined by covalent bond ○ Bond dissociation energy (D) ■ Energy that must be added to break one mole of bonds in the vapor phase Relationships among characteristics ○ All 3 characteristics are related ■ Bond length decreases as bond order increases ● E.g., Bond length for triple bond < bond length for double bond ■ Bond dissociation energy increases as bond order increases ● E.g., D for triple bond > D for double bond ○ Bond dissociation energies were studied by Linus Pauling and it was from them that he set up the relative scale of electronegativities Comparison of bond lengths ○ Picometers = 10-12 m ○ Ångstrom = 10-10 m Formation of a Covalent Bond ○ Diagram shows formation of H—H bond 1. Initially, atoms behave independently 2. Nucleus of one atom attracts electron cloud of other atom 3. Bonding achieved at point of minimum energy 4. Should nuclei grow closer together, nuclei repel, causing energy increase ○ Down arrow is energy released upon bond formation (-432 kJ) ○ Bottom of well is average bond length (74 pm) Average Bond Length ○ Consider the bond between atoms as a flexible spring, not as a rigid bar ○ Thermal energy increases vibrational amplitude about equilibrium position ○ Function varies depending on molecule under consideration ○ Minimum energy represents average, internuclear distance ****

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Lengths of Commonly Encountered Bonds

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**** Only some molecules absorb infrared (IR) light – Why? ○ When pairs of partial charges in a polar covalent bond move asymmetrically, their motion distorts electric fields associated with the bond ○ Only light at the frequency of vibration of the bond interacts with these electric fields ○ All bonds vibrate with frequencies corresponding to IR light ○ Only molecules with polar covalent bonds and that can vibrate asymmetrically absorb IR light. ○ CO2 - 2 bends, one asymmetrical stretch, 1 symmetrical stretch ○ N2 and O2 - non-polar covalent bonds; do not absorb IR radiation Types of compounds ○ Nomenclature - a formal system to match names and formulas for different compounds ○ Compounds classified into organic and inorganic compounds ■ Not a precise definition ■ Organic compounds are those that contain carbon and are typically found in living systems ■ Inorganic compounds are those that are not organic as well as all ionic compounds Types of inorganic compounds ○ Binary inorganic compounds – three types ■ Binary ionic compounds formed from metal cation and non-metal anion ■ Binary compounds formed from two non-metals ■ Binary acids formed from hydrogen and halogen (Group 7A element) ○ Inorganic compounds with 3 or more elements ■ Frequently occur with polyatomic ions (many of them oxo-acids) Rules for naming binary ionic compounds ○ First name is neutral element from which the cation is derived ■ For metals not in Group 1A or 2A, enclose within parentheses metal’s positive charge in Roman numerals after the name ○ The second part of the name is the base name of the non-metal with the suffix “ide.” ■ Examples: MgCl2 magnesium chloride; ZnO zinc (II) oxide

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Common monatomic non-metal anions

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Rules for naming binary inorganics with two non-metals ○ Name of the first non-metallic element is the first name ○ Second part of the name is the base name of the second non-metal with the suffix “-ide.” ○ Append Greek prefixes for number of atoms of element to each name* ■ 1 is mono■ 2 is di■ 3 is tri■ 4 is tetra■ 5 is penta■ 6 is hexa■ *Exception – Prefix is omitted if there is only one atom for the first element ○ Examples: ■ P2S3 diphosphorus trisulfide ■ N2O4 dinitrogen tetroxide ■ SF6 sulfur hexafluoride ■ CO carbon monoxide Naming Ionic Compounds ○ Metal nonmetal-ide ○ -ine → -ide (halogens) ■ oxygen → oxide ■ sulfur → sulfide ■ nitrogen → nitride ■ phosphorus → phosphide ■ E.g., ● NaCl = sodium chloride ● Na2O = sodium oxide ● Na2S = sodium sulfide ● Na3P = sodium phosphide

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**** Identifying and Naming Oxoacids ○ H + polyatomic anion ○ Hydrogen (can have more than one) is listed first in formula ○ Naming: ■ -ite → -ous ■ polyatomic ion-ous acid ■ -ate → -ic ■ polyatomic ion-ic acid ○ For example ■ H2SO3 = sulfurous acid ■ H2SO4 = sulfuric acid ■ HNO2 = nitrous acid ■ HNO3 = nitric acid

Inorganic compounds with 3 or more elements ○ Frequently occur with polyatomic ions ○ Cation name is name of the metal ○ This table only shows some of the many polyatomic ions **** Naming Ionic Compounds Containing Polyatomic Ions ○ Metal polyatomic ion ○ Metal(charge) polyatomic ion ○ Names of polyatomic ions must be memorized ○ For example ■ NaC2H3O2 = sodium acetate ■ Na2CO3 = sodium carbonate ■ Na3PO4 = sodium phosphate ■ CuC2H3O2 = copper(I) acetate ■ FeCO3 = iron(II) carbonate ■ FePO4 = iron(III) phosphate

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