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Chem.200-9.Bonds
General Chemistry Ch. 9 - Models of Chemical Bonding
| Question | Answer |
|---|---|
| Review: the three major types of bonding | Ionic, covalent, and metallic |
| Why do atoms bond? | Bonding lowers the potential energy between positive and negative particles. |
| Type of bonding: metal with nonmetal | Typically results in ionic bonding. The metal atom (high IE) loses valence electrons, whereas the nonmetal atom (highly negative EA) gains the electron. Electron transfer occurs and each atom forms an ion with a noble gas electron config. |
| Type of bonding: nonmetal with nonmetal | Typically results in covalent bonding. Each nonmetal atom holds onto its own electrons tightly (high IE) while attracting other electrons (highly negative EA). |
| Type of bonding: metal with metal | Results in metallic bonding. All the metal atoms in a sample ‘pool’ their valence electrons into an evenly distributed “sea” of electrons that “flows” between and around the metal-ion cores. |
| In covalent bonding, share electrons are said to be… | A shared electron pair is considered to be “localized” between the two atoms. |
| In metallic bonding, electrons are said to be… | Unlike the localized electrons in covalent bonding, electrons in metallic bonding are “delocalized”, moving freely throughout the piece of metal. |
| Why are metallic bonds structured the way they are? | Metal atoms are relatively large, and their few outer electrons are well shielded by filled inner levels. Thus, they lose outer electrons comparatively easily (low IE) but do not gain them very readily (small EA). |
| Are there exceptions to the general bonding trends? | Yes, e.g. beryllium and chlorine form a covalent bond rather than an ionic bond. |
| Lewis electron-dot symbol | The element symbol represents the nucleus and inner electrons. The surrounding dots represent the valence electrons. The order in which you place the dots does not matter. |
| Note, the _____ number gives the number of valence electrons (for the Lewis symbol) | A-group |
| Octet rule | When atoms bond, they lose, gain, or share electrons to attain a filled outer level of eight (or two, i.e. helium) electrons. |
| The total number of electrons lost by the metal atoms in an ionic bond equals… | The total number of electrons gained by the nonmetal atoms. |
| The electron transfer process absorbs or expels energy? | Absorbs. The ionizing energy IE exceeds the electron affinity EA. (IE > EA) |
| If the electron transfer process absorbs energy, why does it occur at all? | The reason ionic substances exist at all is because of the enormous release of energy that occurs when the ions come together as a solid. |
| Why is there a release of energy when the ions come together as a solid? | This occurs from the strong attraction among many oppositely charged ions. E.g. when 1 mol of Li+(g) and 1 mol of F-(g) form 1 mol of gaseous LiF, a large quantity of heat is released. Much more energy if LiF is solid. |
| Lattice energy | The enthalpy change that occurs when 1 mol of ionic solid separates into gaseous ions. It indicates the strength of ionic interactions, which influence melting point, hardness, solubility, and other properties. |
| How is lattice energy calculated? | They are calculated by means of a Born-Haber cycle, a series of chosen steps from elements to ionic solid for which all the enthalpies are known except lattice energy. (Practice problems to master this technique) |
| BE | Bond energy; energy required to break a bond, e.g. F2 to F |
| Recap: ionic solids exist only because… | The lattice energy exceeds the energetically unfavorable electron transfer. |
| How can lattice energy be related to Coulomb’s law? | Since the strength of the ionic solid determines how much lattice energy it contains, that means the lattice energy is directly proportional to the electrostatic energy which is directly proportional to Coulomb’s law equation. |
| How can Coulomb’s law (constant * charge A*chargeB / distance^2) be restated in terms of ionic bonds | (constant) * (cation charge * anion charge)/((cation radius + anion radius)^2) |
| Effect of ionic size on lattice energy | As radius increases, electrostatic force decreases, thus lattice energies decrease. |
| Effect of ionic charge on lattice energy | Increases electrostatic force, thus increases lattice energy |
| The first and most important job of any model is to… | Explain the facts |
| Why do ionic solids (e.g. NaCl) crack when applied force rather than bend? | Once enough force is supplied to break the electrostatic attractions between the individual atoms, the ions of like charge will be brought next to each other and repulsive forces will crack the sample suddenly. |
| Ionic solids are typically brittle, rigid, and hard. Define those terms | Brittle: cracks without deforming, rigid: does not bend, hard: does not dent. |
| Why doesn’t cracked NaCl re-form their bonds? | Obviously you can’t put cracked NaCl back together because the endothermic electron transfer can only be overcome if enough energy is supplied at the atomic level, e.g. lattice energy. |
| Do ionic compounds conduct electricity? | No, the ions are all bonded and thus static. But when dissolved in water they are conductive because the aqueous ions can now move freely and respond to electric fields. |
| Ion pairs | Interionic attractions are so strong that the vapor consists of ion pairs, gaseous ionic molecules rather than individual ions. Note: in solid form there are no separate molecules; ionic compounds consist of arrays of alternating ions. |
| Which is stronger: ionic bonds or covalent bonds? | Ionic. As a result they have much higher melting points. |
| Which is larger: the number of known ionic compounds or the number of known covalent compounds | The number of known covalent compounds. Covalent molecules range from diatomic hydrogen to biological and synthetic macromolecules consisting of many hundreds or thousands of atoms. |
| What is the principal way that atoms interact chemically? | By sharing electrons (covalent bonding). |
| Covalent bonds arise when | Nucleus electron attractions and electron-electron repulsions balance each other out. |
| Formation of a covalent bond always results in | Greater electron density between the nuclei. |
| What is the notation to indicate a “shared pair”, or “bonding pair”, e.g. for a diatomic hydrogen molecule? | H:H or H--H |
| What is a lone pair, or unshared pair? | An outer-level electron pair that is not involved in bonding. E.g. the bonding pair in HF fills the outer level of the H atom and fills up one of F’s outer level together with three other lone pairs. |
| Bond order | The number of electron pairs being shared by any pair of bonded atoms. E/g F2 or HF are single bonds because they each consist of a single pair of bonding electrons. Ethylene (C2H4) has a double bond between the carbon atoms. |
| We know single bonds and double bonds, what about triple bond? | A triple bond consists of three bonding pairs; two atoms share six electrons so the bond order is 3. E.g. N2. |
| Bond energy (BE) | AKA bond enthalpy or bond strength, is the energy required to overcome this attraction and is defined as the standard enthalpy change for breaking the bond in 1 mol of gaseous molecules. |
| Is bond breakage endothermic or exothermic? Ergo, is the bond energy positive or negative? | Endothermic, so the bond energy is always positive |
| The energy absorbed to break the bond is released when… | The bond forms. |
| Important: Stronger bonds are ____ in energy; weaker bonds are ____ in energy. | Lower; higher. |
| Bond length | A covalent bond has a bond length: the distance between the nuclei of two bonded atoms. |
| How does bond order relate to bond length and bond energy? | For a given pair of atoms, a higher bond order results in a shorter bond length and a higher bond energy. For a given pair of atoms, a shorter bond is a stronger bond. |
| Recap: how does bond strength relate to bond energy? | The weaker the bond (ergo longer bond length) = higher the energy. |
| If covalent bonds are so strong, why do covalent substances melt and boil at such low temperatures? | You must distinguish between strong covalent bonding forces holding the atoms together within the molecule and the weak intermolecular forces holding the separate molecules near each other in the macroscopic sample. |
| How does pentane (C5H12) boil? | The weak interactions between the pentane molecules (intermolecular forces) are affected, not the strong C-C and C-H covalent bonds within each molecule. |
| Network covalent solids | Do NOT consist of separate molecules rather they are held together by covalent bonds that extend in three dimensions throughout the sample. Intermolecular forces aren’t an issue. |
| Examples of network covalent solids: Quartz | Quartz (SiO2) has silicon-oxygen covalent bonds that extend throughout the sample; no separate SiO2 molecules exist. Quartz in very hard and melts at 1550deg C. |
| Examples of network covalent solids: Diamond | Diamond has covalent bonds connecting each of its carbon atoms to four others throughout the sample. It is the hardest natural substance known and melts at around 3550deg C. |
| Briefly, why is there such a large difference in strength of diamond and graphite? | The carbon in each substance is arranged differently. |
| Unlike ionic compounds, most covalent substances are… | Poor electrical conductors, even when dissolved. An electric current is carried by either mobile electrons or mobile ions. In covalent substances, the electrons are localized as either shared or unshared pairs; no ions are present. |
| When 1 mol of H2 and 1 mol of F2 react at 298K, 2 mol of HF forms and 546 kJ of heat is ____. Write an equation that shows this | Released. H2 + F2 -> 2HF + 546 kJ. Exothermic. deltaH = -546kJ |
| From where is the (released or absorbed) heat derived in a reaction? | The two forces that affect heat energy (ergo heat) is kinetic and potential energy. Kinetic doesn’t change (unless temperature changes); the movement of the atoms stays constant. The potential energy changes. |
| How does the potential energy (and thus heat) change in a reaction? | The only significant change in potential energy is in the strength of attraction of the nuclei for a shared electron pair (bond energy). In other words: it’s due to differences between the reactant and product bond energies. |
| What about the potential energy change of all the other electrons in the atom | Only the bonding electrons are affected, nothing else in the atom (i.e. the electrons other than the valence electrons) will change. |
| What two step process can we think of any reaction as performing (even though it technically doesn’t occur this way, it’s a good logical model)? | A quantity of heat is absorbed (deltaH is positive) to break the reactant bonds and form separate atoms and a different quantity is released (deltaH is negative) when the atoms rearrange to form product bonds. |
| Equation to determine deltaH_rxn using bond energies (add note regarding signs) | SUM(deltaH_ReactantBondsBroken) + SUM(deltaH_ProductBondsFormed) note: the sign for reactant bonds broken are positive, the product bonds formed are negative |
| When calculating deltaH using the bond energy equation and comparing it to the same reaction as observed in a calorimeter, why is there a small discrepancy? | Because the bond energies are average values obtained from many different compounds. The energy of the bond is a particular substance is usually close, but not equal, to this average. |
| Fuel | A material that reacts with atmospheric oxygen to release energy and is available at a reasonable cost. |
| The most common fuels for machines | Hydrocarbons and coal. |
| The most common fuels for organisms | Fats and carbohydrates |
| Similarities between machine and organism fuels | Both types of fuels are composed of large organic molecules with mostly C-C and C-H bonds (the foods also contain come C-O and O-H bonds). |
| As with any reaction, the energy released from the combustion of a fuel arises from… | Differences in bond energies between the reactants (fuel plus O2) and the products (CO2 and H2O). |
| What do we know about the total strength of the resulting bonds in the combustion of fuel? | Because the reactions are exothermic, we know that the total strength of the bonds in the products is greater than that of the bonds in the reactants (i.e. the bonds of H2O and CO2 are stronger than that of gas/carbs/fat/coal. |
| Stronger bonds means… | Lower energy/more stable. |
| Fuels with more weak bonds yield ____. As a result… | More energy than fuels with fewer weak bonds. As a result we can say that the fewer bonds to O in a fuel, the more energy it releases when burned because C-O and O-H have higher bond energies than C-C and C-H. |
| How is the energy stored in carbs and fats? | Fats consist of chains of carbon atoms: C-C bonds attached to hydrogen atoms C-H bonds, with few C-O and O-H bonds. Carbs have fewer C-C and C-H bonds and many more C-O and O-H bonds. |
| How does the molecular composition of carbs and fats affect their abilities to yield energy? | Fats contain more Calories per gram than carbohydrates because fats have more weaker bonds and fewer stronger bonds (fewer bonds to O). Thus fats release more energy than carbs. |
| While our models (which are nothing more than idealized descriptions of reality) describe bonds as either covalent or ionic, in reality… | Bonding lies somewhere between these extremes: the great majority of bonds are more accurately thought of as “polar covalent”, that is partially covalent and partially ionic. |
| Electronegativity (EN) | The relative ability of a bonded atom to attract the shared electrons. |
| Who developed the most common scale of relative EN values for elements? | Linus Pauling |
| How did Linus determine the magnitude of electronegativity in bonds? | We might expect the bond energy of the HF bond to be the average of the energies of an H-H bond and F-F bond, but the H-F is much higher than those bonds. Why? Because electrons are drawn very closely to F due to electrostatic forces. |
| Trends in electronegativity | Electronegativity is inversely related to atomic size. It increases up a group and across a period. Nonmetals are more electronegative than metals. |
| The most electronegative element is… | Fluorine, with oxygen at a close second. |
| The least electronegative is | Francium, which is radioactive and extremely rare. So for all intents and purposes, the least electronegative is cesium. |
| How does electronegativity relate to assigning oxidation numbers to elements in compounds? | The more electronegative atom in a bond is assigned all shared electrons. Each atom in a bond is assigned all of its unshared electrons. |
| Polar covalent bonds | Whenever atoms of different electronegativities form a bond, the bonding pair is shared unequally. This unequal distribution of electron density gives the bond partially negative and positive poles, forming a polar covalent bond. |
| Nonpolar covalent bonds | E.g. H-H and F-F bonds. |
| By knowing the EN values of the atoms in a bond, we can find the direction of… | The bond polarity. |
| Partial ionic character | An estimate of the actual charge separation in a bond (caused by the electronegativity difference of the bonded atoms) relative to complete separation. |
| Electronegativity difference (deltaEN) | The differences in electronegativities between the atoms in a bond. |
| A greater deltaEN results in | Larger partial charges and a higher partial ionic character. |
| The deltaEN of LiCl is 3 - 1 = 2. The deltaEN of HCl is 3 - 2.1 = 0.9. What does this tell us? | It tells us that the bond in LiCl has more ionic character than the H-Cl bond. |
| Percent ionic character | Determined by comparing the actual behavior of a polar molecule in an electric field with the behavior it would show if the electron were transferred completely (a pure ionic bond). |
| Threshold percent between what we would call covalent and ionic? | 50%. Anything above that = ionic. |
| The percent ionic character generally increases with | deltaEN |
| The continuum of bonding across a period | When bonding Cl with all elements across period 3 the trend is visible on a macroscopic level. As deltaEN decreases (as you move across period) the compounds move from solid -> liquid -> gas and the bonds become more covalent |
| As deltaEN becomes smaller, bonds become more ____ | Covalent. |
| What model explains metallic bonding? | Electron-sea model |
| Electron-sea model | A qualitative description of metallic bonding proposing that metal atoms pool their valence electrons into a delocalized sea of electrons in which the metal cores (metal ions) are submerged in an orderly array. |
| Although there are metallic compounds, two or more metals typically form ____ | Alloys. |
| Alloy | A mixture with metallic properties that consists of solid phases of two or more pure elements, a solid-solid solution or distinct intermediate phases. |
| What’s the difference between metals and ionic substances with regard to conductivity? | Metals: conduct in both liquid and solid form. Ionic: only in liquid form. |
| Melting points of metal | Related to the energy of metallic bonding. Melting points are only moderately high because the attractions between moveable cations and electrons need not be broken during melting. |
| Boiling points of metal | Boiling a metal requires each cation and its electrons to break away from the others, so boiling points are very high. |
| Metals periodic trends with regard to melting points | Increases as you move across period: valence electrons increase which increase attraction between cations and more valence electrons means stronger metallic bonding. |
| How does the electron-sea model account for the malleable, ductile, and conductive characteristics of metal? | When a piece of metal is deformed by a hammer, the metal ions slide past each other through the electron sea and end up in new positions. The metal-ion cores don’t repel each other (like ionic compounds). Mobile electrons = conduct. |
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