<p>Name : ______Class : ______Date : ______</p><p>Bonding Review</p><p>1. Distinguish between an ionic bond and a covalent bond</p><p>2. Why do ionic compounds ordinarily have such high melting and boiling points?</p><p>3. List properties of metals, ionic compounds, and covalent compounds</p><p>4. Metallic bonds are different from ionic and covalent bonds because the electrons in metallic bonds are delocalized. The electrons are not assigned to a specific atom. This is important because it gives metals the ability to conduct heat and electricity, makes them malleable (easily shaped) and ductile (very thin). When two atoms are smushed together as they would be if the metal was being hammered into a new shape, the electrons repel but are able to move apart easily. In an ionic bond, because electrons are localized (assigned a specific atom) if two like charges come into contact, they repel and cause the material to crack. This explains the brittleness of ionic compounds. The localized electron model also explains why ionic compounds can’t conduct electricity – the ions can’t move past each other unless they are either liquefied or dissolved in water. </p><p>Although metals and salts have similar lattice structures, metals make good materials for electrical wiring. Why aren’t salts used instead? (more than one reason)</p><p>1 5. Why is the strontium nitrate found in roadside emergency flairs represented at Sr(NO3)2 rather than SrN2O6?</p><p>6. What is wrong with the following formulas a. RbCl2 c. NaCs</p><p> b. Ge12S24 d. NaNe</p><p>7. Propose a reason why magnesium ions form Mg2+ ions and not Mg6- ions? </p><p>8. Which has greater potential energy, a noble gas or a metal? Explain.</p><p>9. A classmate insists that sodium gains a negative charge when it becomes an ion because it loses an electron. Explain this student’s error. </p><p>2 10.Explain why when Na loses and electron to be like neon, it is not actually becoming a neon atom. </p><p>Remember that electronegativity gives a measure of the ability of one atom to pull electrons to itself in a bond. Electronegativity values are not measures of energy, nor do they correlate, necessarily to bond strength. Linus Pauling made up the scale to be able to predict the ionic strength vs. covalent strength of a particular bond in a molecule. He did this using quantum mechanics to theoretically calculate the strength of a bond between two atoms as if the electrons were evenly shared. He then compared the theoretical bond strength to actual bond strength and was able to calculate the difference. In this way, he was able to map out the relative ionic character of various bonds. From there, he could come up with a scale that ranked the electron pulling ability of each element. </p><p>Chemists use the electronegativity scale to distinguish between ionic, polar, and non-polar compounds. Generally speaking, the larger the difference in electronegativity between two atoms, the more ionic the bond. The closer the difference in electronegativity, the more covalent the bond. Not all chemists agree on the exact ranges for ionic, polar and non-polar bonds, but the following are useful guides:</p><p>Ionic bonds EN 4.0 – 2.0 Polar Bonds EN 2.0-0.5 Non-Polar EN 0.5-0</p><p>11.Using Table 7 in the data booklet, calculate the difference in electronegativity between the atoms for the following bonds and state whether they are ionic, polar or non-polar. Circle the atom that is the most electronegative.</p><p> a. Na-F e. C-O b. Br-Br f. C-H c. K-I g. C-Cl d. S-O</p><p>3 12.Define: a. Dispersion Forces</p><p> b. Dipole/Dipole Moment</p><p> c. Temporary/Induced Dipole</p><p> d. Permanent Dipole</p><p> e. Dipole-Dipole Forces</p><p> f. Hydrogen Bonding</p><p> g. Van der Waals Forces</p><p>13.Rank the following in order of increasing strength a. Dispersion forces b. Hydrogen Bonding c. Dipole-Dipole Forces</p><p>4</p>