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Chemistry Chapter 8
Chemical Bonds, Resonance Structures
| Question | Answer |
|---|---|
| What region of the electromagnetic spectrum is heat related to? (8.1) | Infrared Often has difficulty passing through CO₂ or CH₄. (Carbon Dioxide + Methane) |
| Why are bonded atoms more stable? (8.1) | Principle stems from the concept of electromagnetic potential energy (Eₑₗ) between charged particles. Eₑₗ is directly proportional to the product of the charges (Q₁ and Q₂) on a pair of ions (or subatomic particles). Eₑₗ is inversely proportional to the distance between them (d). (Eₑₗ) ∝ (Q₁ x Q₂) ÷ (distance, d) |
| Ionic Bond (8.1) | Particles of opposite charge (cations + anions) causes Q₁ x Q₂ to be negative, as well as Eₑₗ. The value of Eₑₗ becomes even more negative as d↓. When the two are very close, electrostatic repulsions equalize the forces of attraction- creating a stable compound. |
| Electrostatic Repulsions (8.1) | Occur between the ion's outer-shell electrons and their positive nuclei. |
| Electrostatic Potential Energy (Eₑₗ) Hydrogen Molecule Formation (8.1) | As two H atoms move closer together, the nucleus of one and the electron of another begin to attract. This creates a negative Eₑₗ which will decrease until its min value is met at 74pm (d). The strength of attraction between e₁ and p₂ is EQUAL to the strength of attraction between p₁ and e₂. Creates H₂ molecule, a single covalent bond. If the two are forced together, their Eₑₗ will rise as their e-static repulsion increases. Molecular potential energy increases (forcing two magnets together). |
| Bond Length (8.1) (Introductory) | Eₑₗ's minimum value corresponds to distance (d) In the example hydrogen bond, H―H is 74pm (picometers). The minimum Eₑₗ at this distance, -436 kJ/mol, is the energy released when 2 moles of H atoms bond into 1 mole of H₂. |
| Bond Strength (8.1) (Introductory) | The Eₑₗ value at H₂'s bond length (74pm) is -436 kJ/mol. Breaking one mole of H₂ bonds would require the addition of 436 kilojoules of energy. This energy requirement is a measurement of [term]. |
| Metallic Bonds (8.1) | Formed within a solid piece of metal, nuclei are attracted to the electrons of all surrounding atoms. Unlike covalent bonds, they are not a pair of electrons shared by two atoms. Instead, they form a "sea" of electrons that are highly DELOCALIZED in their attractions. They are mobile and relocate freely between atoms, creating conditions for effective conduction and heat transfer. |
| Bond Characteristics/Comparisons (8.1) | Ionic: metals + nonmetals. Transfer of electrons. Covalent: nonmetals, metalloids. Sharing electrons. Metallic: metals. Delocalized and ∴ conductive. Keep in mind, not all bonds can be classified definitively. A bond may have ionic, covalent, and metallic characteristics. Break structures into polyatomic ions if possible. ex. Ca²⁺ and CO₃²⁻ bond to make calcium carbonate. These two cations/anions form an ionic pair no charge. However, Carbon is a nonmetal and covalently bound to Oxygen. |
| Indicates a Bond is Ionic (8.1) | -ate (contains oxoanions) -ite (contains anions) -ide (binary compound containing a metal and nonmetal) |
| Lewis Symbols (8.2) | Represents the number of bonds usually formed by an element to complete an octet. Creates a simple, 2D representation of bonds and lone pairs. They also indicate BONDING CAPACITY. C = 4 (has four valence electrons) N = 3 (has five valence electrons) [there are often exceptions] Atoms with lone pairs in their [term] typically have the same number in the Lewis Structures of molecules they form. |
| Octet Rule & Exceptions (8.2) | Atoms of most main group elements lose, gain, or share electrons so that each atom may have eight valence electrons. Hydrogen atoms form duets. Exceptions commonly involve [H, Be, B]. Compounds with >8 valence electrons around an atom are called electron deficient compounds. Also may occur when atoms of a compound have odd numbers of electrons in their valence shells [NO and NO₂]. |
| Step One of Drawing Lewis Structures (8.2) | Find the number of valence electrons. For a neutral molecule, simply count them. For polyatomic ions, count and then add (for anions) and subtract (for cations) the number of electrons needed to account for the charge. [CO₃ ²⁻] |
| Step Two of Drawing Lewis Structures (8.2) | Arrange the symbols of the elements to show how the atoms are bonded, and then connect them with single bonds (single pairs of bonding electrons). Put the atom with the greatest bonding capacity in the center. If two atoms tie, use the one that is least electronegative. Place remaining atoms around, and those that create the least bonds on the periphery [H]. |
| Step Three of Drawing Lewis Structures (8.2) | Complete the octets of atoms bonded to the central atom by adding lone pairs of electrons until each one has an octet (including the two electrons used to connect them to the central atom). The same bonding electrons are used to complete the octets of both atoms they connect. Bonding electrons count twice when completing octets, but only once when calculating the total number of valence electrons in the structure. |
| Step Four of Drawing Lewis Structures (8.2) | Compare the number of valence electrons in the structure to the number derived from step one. If valence electrons remain unused, place lone pairs around the central atom (even if it exceeds an octet). |
| Step Five of Drawing Lewis Structures (8.2) | Complete the octet around the central atom. If there is less than an octet, convert one (or more) of the lone pairs of a surrounding atom into a bonding pair. The goal is to find Steric Number and the Electron Group Arrangement. Number of lone pairs around central atom is very important. |
| Lewis Structures with Double and Triple Bonds (8.2) | Double Bonds: 2 atoms share 2 pairs of e⁻ (O═O) O₂ Triple Bonds: 2 atoms share 3 pairs of e⁻ (N≡N) N₂ Required if (after drawing structure) the initial # of valence e⁻ equals the final # of valence e⁻ but the central atom doesn't have a full octet. Convert one or more lone pairs of e- from the other atoms into a bonding pair. For example: H₂CO. C will only have 6 valence electrons while H₂ and O are full. Convert a lone pair from O into another bonding pair w/ C. |
| Bond Angles (8.2) | Electrons repel one another, so branches are drawn as far apart as possible. (4 bonds = 90°) (3 bonds = 120°) (2 bonds = 180°) |
| Carbonyl Group (8.2) (Introductory) | All Aldehydes contain this group: (C═O) Aldehyde general formula: (RCHO) |
| Lewis Symbols of Ions (8.2) | Brackets are placed around anions to indicate that all eight valence e⁻ are associated with them. The ion's charge is written on the outside. Cations with empty valence shells are written with no electrons. ex: [Cl]⁻ (including four lone pairs, can't type them) ex: Na loses an electron and is written as Na⁺ |
| Low Ionization Energies (8.2) | Causes alkali metal elements and alkaline earth elements to lose rather than share their valence electrons when bonding with nonmetals. |
| Polar Covalent Bonds (8.3) | For example: H―Cl is a polar covalent bond, because the shared electron pair is closer to the Cl than to the H. As a result, when this bond breaks apart the bond pair remains with the Chlorine, ionizing it into Cl⁻ and creating H⁺ (occurs when HCl dissolves in water). |
| Bond Polarity (8.3) | The bond functions as a tiny electric dipole- meaning there is a slightly positive charge to the H⁺ side of the molecule, and a slightly negative charge related to the Cl⁻ side. This is analogous to the positive and negative ends of a battery. |
| Direction of Polarity (8.3) | Indicated by: 1. An arrow with a plus sign embedded into its tail. (⇸) The arrow points toward the more negative, electron-rich atom in the bond, and the position of the plus sign indicates the more positive, electron-poor atom. 2. The lowercase Greek delta (δ) followed by a + or − sign. The deltas represent partial electrical charges, as opposed to the full electrical charges that accompany the complete transfer of one or more electrons when atoms become ions. |
| Electronegativity (8.3) | Electronegativity values are assigned based on the idea that no bond is 100% ionic or covalent. Called Pauling Values, and are unitless. The greater an atom's e-negativity, the greater its ability to attract electrons toward itself within a chemical bond. Increasing attraction between an atom's nucleus and outer electrons (+increasing atomic numbers) leads to lower ionization energy and smaller e-negativity. Most: Fluorine, Nitrogen, Oxygen Least: Francium, Cesium, Rubidium |
| Selecting The Central Atom (8.3) | The central atom is often the one with the least electronegativity, and the highest bonding capacity. |
| Difference in Electronegativity (8.3) | The greater the difference in electronegativity (ΔEN) the more uneven the electrons, and ∴ the more polar the covalent bond. For example, a H―F bond is more polar than a H―Cl because the ΔEN between H(2.1) and F(4.0) is (ΔEN=1.9). While H(2.1) and Cl(3.0) is (ΔEN=0.9); a Cl―Cl molecule is nonpolar because (3.0 - 3.0 = 0)ΔEN. Differences > than about (ΔEN=2.0) are associated with the ionic compounds formed between metals and nonmetals. The presence of a double/triple bond does not change polarity. |
| Bond Vibration (8.3) | Chemical bonds are not rigid, and their vibration/oscillation coupled with polarity allows them to absorb infrared radiation. Polar covalent bonds (specifically atmospheric gasses) are able to catch heat that would otherwise dissipate into space. |
| Infrared Active Vibrations (Asymmetric Stretching) (8.3) | Molecules with polar bonds may absorb photons, and enter excited states like e- (More info: Section 7.4). Symmetric Stretching: Two C═O bonds stretch and compress at the same time. They cancel out and no emission/absorption is possible. (Infrared Inactive) Asymmetric Stretching: One gets shorter, one gets longer. When the energy of a polar covalent bond's vibration becomes equal to the energy of a proton (hv), the proton may be absorbed. This is the molecular mechanism behind the greenhouse effect. |
| Ozone (8.4) | A potentially harmful configuration of oxygen (O₃) that is typically found in the upper atmosphere, created when O₂ is struck by lightning. An Allotrope of oxygen (same element, different chemical/physical properties). |
| Resonance (8.4) | Ozone could theoretically have a double bond on the left and a single bond on the right, or vice versa. However, neither is actually correct. Because a double bond between two atoms is always shorter than a single bond between the same two atoms (see Section 8.7), the structure of ozone cannot consist of one single bond and one double bond. More realistically, it is an AVERAGE of the two resonance structures (left vs right double bond), like 1.5 bonds on each side. O₃ structure #1 ↔ O₃ structure #2 |
| Resonance of Polyatomic Ions (8.4) | Before proceeding as normal, account for the charge on each ion- which means adding the appropriate number of valence electrons to a polyatomic anion and subtracting the appropriate number from a polyatomic cation. Don't forget to bracket/label the charge of every structure. Also remember- Lewis theory does not STRICTLY adhere to the octet rule. Nitrogen sometimes makes 4 bonds. |
| Formal Charge (8.5) | To decide which resonance form in a nonequivalent set is the most important and closest to the actual bonding pattern in a molecule, use Formal Charge. (FC) is not a real charge but rather a measure of the number of electrons formally assigned to an atom in a molecular structure, compared with the number of electrons in the free (i.e., nonbonded) atom. To determine a formal charge, we follow a series of steps to calculate the number of electrons formally assigned to each atom in each resonance structure. |
| Steps to Calculating Formal Charge (8.5) | For each atom: 1. Determine valence # of the free atom. 2. Count the number of nonbonding e- of atom inside structure. 3. Count the number of e- in bonds and divide that # by 2. 4. Sum the results of steps 2 & 3, and subtract that # from the number in step one. Basically: compare each (X) atom unbonded with the (X) atom within the Lewis Structure. Subtract nonbonding e- and ½ of bonding e- from free atom's valence number. |
| Determining Preferred Structure (8.5) | 1. Preferred structure has formal charges of zero. 2. If no such structure can be drawn, or if the structure is that of a polyatomic ion, choose the one with the most atoms closest to zero. 3. Any negative formal charges should be on atoms with the highest atomic electronegativities. |
| Assuming The Octet Rule is Obeyed (8.5) | Atoms have formal charges of zero in resonance structures in which the numbers of bonds they form match their bonding capacities. (Nitrogen with a bonding capacity of three, when involved in three bonds, has a formal charge of zero). If an atom forms one more bond than its bonding capacity, its formal charge will be +1, while if it forms one less bond its formal charge will be -1. The more bonds formed, the less electrons are subtracted from the free atom's valence number. |
| Free Radicals (8.6) | Compounds that have odd numbers of valence electrons. They are typically very reactive species because it is often energetically favorable for them to acquire an electron from another molecule or ion. This characteristic makes them excellent oxidizing agents Oppositely: in general, H, Be, B, and Al are the elements most likely to have LESS than an octet of electrons. |
| Hyper valency (8.6) | A characteristic found in third row nonmetals (phosphorus, sulfur, etc.) allowing them to form more than 4 bonds. Consequently, the octet of electrons is exceeded. Occurs with an atomic number greater than 12 (Z = 15). |
| Bond Order (8.7) | The number of bonds between two atoms. As bond order increases, bond length decreases, as we can see by comparing the lengths of C―C, C═C, and C≡C bonds. |
| Exothermic Reaction (8.7) | MORE energy is RELEASED in forming the bonds in molecules of products than is consumed in breaking the bonds in molecules of reactants. |
| Endothermic Reaction (8.7) | LESS energy is RELEASED in forming the bonds in molecules of products than is consumed in breaking the bonds in molecules of reactants. Bond BREAKING is always ENDOTHERMIC. |
| Bond Strength (8.7) | Usually expressed in terms of the enthalpy change (ΔH) that occurs when 1 mole of a bond in the gas phase is broken. The quantity of energy needed to break a bond depends on the identity of the two bonded atoms + their bond order, and is equal in magnitude but opposite in sign to the quantity of energy released when that same bond forms. |
| Ionic Trends (Ex) | Cations are much smaller because of electron loss. Ions are typically solids at room temperature. They consist of a nonmetal and a metal. |
| Period Table Trends: Increasing Atomic Size (Ex) | ←↓ |
| Period Table Trends: Increasing Electronegativity (Ex) | →↑ |
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