Ch.9 Notes PDF

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Notes on Ch.9 in the Chemistry of Life book.
Professor: Dr. Michael...

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Ch.9

Ionic and Covalent Bonding ▸ Ionic Bonds Describing Ionic Bonds ●

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An ionic  bond is a chemical bond formed by the electrostatic attraction between positive and negative ions ○ One or more electrons are transferred from one valence shell to another The cation  is the ion that loses electrons (is positive) The anion  is the ion that gains electrons (is negative)

Lewis Electron-Dot Symbols ● A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element ● Dots are placed one to each side until are 4 sides are occupied Energy Involved in Ionic Bonding ● Coulomb’s law states that the potential energy obtained in bringing two charges, Q1 and Q2, initially far apart, up to a distance r apart is directly proportional to the product of the charges and inversely proportional to the distance between them E= ● ●

kQ1 Q2 r

, k = 8.99 × 109 J • m/C 2 , e = 1.602 × 10−19 C 

The lattice energy is  the change in energy that occurs when an ionic solid is separated into isolated ions in the gas phase 2 elements will bond ionically if the ionization energy of one is sufficiently small and the electron affinity of the other is sufficiently large ○ Exists between a reactive and a nonreactive metal (or in general a metal and a nonmetal)

Lattice Energies from the Born-Haber Cycle ● Follows the reasoning of Hess’s Law ● Steps: 1. Sublimation - transformation of a solid to a gas a. Enthalpy change = enthalpy of sublimation 2. Dislocation - molecules are split into atoms a. Enthalpy change = bond energy 3. Ionization - atoms are ionized into ions a. Enthalpy change = ionization energy of the atom 4. Formation of the other ion - electrons from one atom are transferred to the other atom a. Enthalpy change = negative of the electron affinity for the second atom 5. Formation of the ionic compound - ions join together a. Enthalpy change = negative of the lattice energy (-U)

Ch.9

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Cancel terms that appear on both the right and left sides of the arrows Final equations is simply the formation reaction ○ Enthalpy change = enthalpy of formation

Electron Configurations of Ions Ions of the Main-Group Elements ● Ionization energy increases with each successive ionization ○ First ionization energy = energy required to remove the first valence electron ○ Second ionization energy = energy required to remove the second electron ○ And so on… ● Valence electrons are easily removed, but the energy needed to take an electron from either of the ions that result is extremely high ○ That is why no compounds are found with ions having charges greater than the group number ● The loss of successive electrons from an atom requires increasingly more energy ● Rules: 1. Cations of the Groups 1A to 3A having noble-gas or pseudo-noble-gas configurations. The ion charges equal the group numbers. 2. Cations of the Groups 3A to 5A having ns2 electron configurations. The ion charges equal the group numbers minus two. Examples are TI+, SN2+, Pb2+, and Bi3+. 3. Anions of Group 5A to 7A having noble-gas or pseudo-noble-gas configurations. The ion charge equals the group numbers minus eight. Polyatomic Ions ● These ions are held together by covalent bonds Transition-Metal Ions ● Most transition elements form several cations of different charges ○ Do not have a noble-gas configuration ● In forming ions in compounds, the atoms generally lose to ns electrons first, then they may lose one or more (n − 1)d electrons

Ionic Radii The ionic radius is  a measure of the size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found ● Cations are smaller and anions are larger than the corresponding atom ○ When valence electrons are lost, the ion is smaller (and vice versa) ● Isoelectronic r efers to different species having the same number and configuration of electrons *Across a period the cations decrease in radius. When you reach the anions, there is an abrupt increase in radius, and the the radius again decreases. ●

Ch.9

▸ Covalent Bonds Describing Covalent Bonds ● ● ●

A covalent bond is a chemical bond formed by the sharing of a pair of electrons between atoms Distance between nuclei at the minimum energy is called b  ond length The energy that must be added to break a bond is called the b  ond dislocation energy ○ The larger this is, the stronger the bond

Lewis Formulas ● A bonding  pair is an electron pair shared between two atoms ● A lone, or nonbonding pair is  an electron pair that remains on one atom and is not shared Coordinate Covalent Bonds ● A coordinate covalent bond is  a bond formed when both electrons are donated by one atom ● It still involves the sharing of electrons between two atoms ● Ex. formation of the ammonium ion, where an electron pair on the N atom in NH3 forms a bond with H+ Octet Rule ● The tendency of atoms in molecules to have eight electrons in their valence shells (two for hydrogen atoms) is known as the octet  rule ○ Therefore, they obtain noble-gas configuration Multiple Bonds ● Single bond - a covalent bond in which a single pair of electrons is shared by two atoms ● Double bond - a covalent bond in which two pairs of electrons are shared by two atoms ● Triple bond - a covalent bond in which three pairs of electrons are shared by two atoms ○ Form mostly to C and N atoms

Polar Covalent Bonds; Electronegativity ● ●

When two atoms are of different elements, the bonding electrons need not be shared equally A polar  covalent bond is a covalent bond in which the bonding electrons spend more time near one atom than the other ○ An intermediate between a nonpolar covalent bond and an ionic bond

Ch.9

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○ Results when the bonding pair is drawn more toward on atoms than the other Electronegativity i s a measure of the ability of an atom in a molecule to draw bonding electrons to itself Ionization Energy+Electron Aff inity  E lectronegativity = 2 In general, electronegativity increases from left to right and decreases from top to bottom in the periodic table ○ Metals are the least electronegative elements (they are electropositive) and nonmetals are the most electronegative ○ Fluorine is the most electronegative element Polar molecules have a large electronegativity difference between atoms Electronegativity values also help determine the direction of electron shift (which end of the polar molecule is negative)

Writing Lewis Electron-Dot Formulas Usually consist of a central atom around which are bonded atoms of greater electronegativity ● H cannot be a central atoms because it only forms 1 bond ● Oxoacids a  re substances in which O atoms (and possibly other electronegative atoms) are bonded to a central atom, with one or more H atoms usually bonded to O atoms ● Atomic ions with symmetrical formulas often have symmetrical structures ● Steps: 1. Calculate the total number of valence electrons 2. Write the skeleton structure of the molecule or ion, connecting bonded pairs of atoms 3. Distribute electrons to atoms surrounding the central atom (or atoms) to satisfy the octet rule 4. Distribute the remaining electrons as pairs to the central atom (or atoms) ●

Delocalized Bonding: Resonance ● ●

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Delocalized bonding i s a type of bonding in which a bonding pair of electrons is spread over a number of atoms rather than localized between two According to the resonance description, you describe the electron structure of a molecule having delocalized bonding by writing all possible electron-dot formulas (a.k.a. the resonance formulas) Metals are extreme examples of delocalized bonding → responsible for their electrical conductivity

Exceptions to the Octet Rule ●

A few molecules, such as NO, have an odd number of electrons and so cannot satisfy the octet rule

Ch.9

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Other exceptions to the octet rule fall into two categories: ○ A group of molecules with an atoms having fewer than eight valence electrons around it ○ A group of molecules with an atom having more than eight valence electrons around it ■ The exceptions in which the central atom has >8 electrons are numerous ■ Ex. PF5 From the third period on, the elements have unfilled nd orbitals, which may be used in bonding The other group of exceptions to the octet rule consists mostly of molecules containing Group 2A or 3A elements. ○ Ex. boron trifluoride (BF3)

Formal Charge and Lewis Formulas Multiple bonds like to involve C, N, O, and S atoms The formal charge of  an atom in a Lewis formula is the hypothetical charge you obtain for an atom by assuming that bonding electrons are equally shared between bonded atoms and that electrons of each lone pair belong completely to one atom ● Rules for formal charge: 1. Half of the electrons of a bond are assigned to each atom in the bond (counting each dash as two electrons) 2. Both electrons of a lone pair are assigned to the atom to which the lone pair belongs Formal charge = valence electrons on free atom − 12(# of electrons in the bond) − ( # of lone pair electrons)  ● ●

Rules to decide which resonance formula is best: Whenever you can write several Lewis formulas for a molecule, choose the one having the lowest magnitudes of formal charges 2. When two proposed Lewis formulas for a molecule have the same magnitudes of formal charges, choose the one having the negative formal charge on the more electronegative atom 3. When possible, choose Lewis formulas that do not have like charges on adjacent atoms

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Bond Length and Bond Order ● ● ●

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Bond length (or  bond distance) is the distance between the nuclei in a bond Bond lengths for a given bonding situations can often be predicted within a few picometers from a set of covalent radii The covalent  radius of an atom is the value for that atoms in a set of covalent radii assigned to atoms in such a way that the sum of the covalent radii of atoms A and B predicts the approximate A-B bond length The values for any given set of radii should follow the two major periodic trends for atomic radii:

Ch.9

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Within a period, the covalent radius tends to decrease with increasing atomic number. 2. Within a group, the covalent radius tends to increase with periodic number. ● The bond  order is the number of pairs of electrons in a bond ○ Single bond = 1 ○ Double bond = 2 ○ Triple bond = 3 ● Bond length depends on bond order: as the bond order increases, the bond strength increases and the nuclei are pulled inward, decreasing the bond length

Bond Enthalpy Bond enthalpy functions as a measure of the average strength of a bond in its compounds ● These are often obtained from enthalpies of reaction, ΔH  ● Also known as bond energies ● The A-B bond  enthalpy is the average enthalpy change for the breaking of an A-B bond in a molecule in the gas phase ● Because it takes energy to break a bond, bond enthalpies are always positive numbers ● When a bond is formed, the enthalpy change is equal to the negative of the bond enthalpy (heat is released) *The larger the bond enthalpy, the stronger the bond ● In general, the enthalpy of reaction is (approximately) equal to the sum of the bond enthalpies for bonds broken minus the sum of the bond enthalpies for bonds formed ΔH ≈ Broken − M ade  ●

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A reaction is exothermic (gives off heat) if weak bonds are replaced by strong bonds

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