Language of Chemistry
NOTE: Many of the terms used in this module are defined elsewhere.
alkali metals family of elements characterized by their vigorous reaction with water. The elements in this family are lithium, sodium, potassium, rubidium, cesium and francium.
anion negatively charged ion.
atomic radius one half the distance between nuclei of two adjacent atoms of the same element.
cation positively charged ion.
charge density charge on an ion divided by its surface area.
Coulomb's Law relationship between electrical forces, charges and distance: the electrical force between two charged objects varies directly as the product of the charges and inversely as the square of the distance between them [F = k x (q+ x q ) / r2].
critical pressure minimum pressure that must be applied to bring about liquefaction at the critical temperature.
critical temperature temperature above which a gas will not liquefy.
crystal lattice energy energy required to completely separate one mole of a solid ionic compound into gaseous ions.
electrolysis process, involving either the molten state or an electrolytic solution, by which compounds are decomposed electrically.
electron affinity energy associated with the gain of an electron by a neutral gaseous atom.
electronegativity measure of electron attracting power of an atom; metals have low electronegativities, nonmetals have high electronegativities.
hydration energy energy associated with dissolving gaseous ions in water, usually expressed per mole of ions ion charged atom or group of atoms, formed when the atom or group of atoms loses or gains electrons.
ionic radius radius of a spherical ion; it is the radius associated with an element in its ionic compounds.
ionization energy amount of energy needed to remove a single electron from a neutral, isolated (gaseous) atom.
lattice geometric model showing the regular arrangement of atoms or ions in a crystalline solid.
metal one of a group of substances characterized by luster, malleability, ductility, and good electrical and heat conductivity; metals tend to form positive ions in ionic compounds; elements that are metals are located on the left side of the Periodic Table.
ore mineral deposit containing sufficiently high concentration to allow economical recovery of a desired metal.
oxidation number a number assigned to an atom in a neutral molecule or ion to reflect its state of oxidation.
photoelectric cell a chemical cell which requires for its operation the ejection of electrons from specific metal atoms when exposed to light.
second ionization energy amount of energy needed to remove a second electron after a single electron has already been removed from a neutral, isolated ion (gaseous) atom.
unit cell smallest unit of a crystal that, if repeated indefinitely, could generate the whole crystal.
Pattern Recognition
1. Figure 2 gives crystal ionic radii in picometers, as measured by X-ray crystallography. (1 pm =1 x 1012 m)
Figure 2. Crystal ionic radii (pm).
a. Considering chemical periodicity, what trend would you predict for the interionic distances for the fluorides of the alkali metals? For the cesium salts of the four halide ions? (The interionic distance is the distance between the centers of the two ions.) [Interionic distances will increase as you go down the alkali metal family of fluoride compounds.]
b. Calculate the interionic distances between the centers of all combinations of halide and alkali metal ions. [Calculated interionic distances of all combinations of alkali metal and halogen ionic compounds:
2. Use the results of your calculations from Problem 1b above and Pauling's values for interionic distances in Figure 4 to compare calculated values of ionic radii of alkali and halide ions to the actual values of interionic distances between centers of alkali metal and halide ions measured from X-ray crystallography data. [Calculated values are generally lower than Pauling values.]
Figure 4. Calculated interionic distances between center of halide and alkali metal ions (pm).
3. Consider the forces of attraction between alkali metal ions and halide ions in terms of their size and charge (Figure 2). Which pair of ions would form an ionic crystal with the greatest crystal lattice energy? Least crystal lattice energy? [Lithium fluoride, with the two smallest ions, would have the greatest crystal lattice energy because they will be held together most tightly. Cesium iodide, with the largest ion sizes, should have the least crystal lattice energy. See Figure 5.]
Figure 5. Crystal lattice energies for alkali halides (kJ/mol).
4. Using Figures 4 and 5 and Coulomb's Law, explain whether trends in crystal lattice energies in Figure 5 appear consistent with the interionic distances in Figure 4. [The force of attraction between ions increases with increasing charge on the ions and with decreasing size of ions. Since all alkali metal ion pairings with halide ions are identical with respect to charge, size is the determining factor. Therefore the combination involving smallest ion sizes (Li +and F ) should have the greatest attraction and hence the greatest crystal lattice energy (1034 kJ/mol). On this basis Cs+ and I, the largest ions, should have the smallest lattice energy (585 kJ/mol). Trends in crystal lattice energy seem consistent with interionic distances in general; e.g., ion combinations with similar interionic distances have similar crystal lattice energies. (For examples, see RbF and LiBr, 282 and 275 pm interionic distances and 780 and 781 kJ/mol lattice energies.)]
5. The Table of Properties of Alkali Metals in the Appendix summarizes an extensive set of properties for the alkali metal elements. A useful procedure is to take the properties one by one and ask students to predict the trend in each property going down the group. Additionally, one can query students as to the relative values for a given property as they relate to properties already covered. The overall objective is to underscore patterns of behavior for the properties of a family of elements.
A possible sequence of discussion questions follows. While displaying on the overhead projector the data needed for each question, you can uncover the table of properties sequentially.
a. On the basis of atomic number of the alkali metal elements, write the electron configurations. (Note similarities
[Similarity: All have the same outer energy level electron configuration (ns 1).
Difference: Each has its own inner core of filled energy levels (its own noble gas configuration).]
b. What valence (oxidation) state do you predict for the alkali metals? Why? [M+ oxidation state because each has only one electron in its outer energy level to lose.]
c. How do the sizes of the atoms change with increase in atomic number? Explain. [As atomic number increases, size of the alkali metal atom will increase also due to increasing number of filled energy levels of electrons within each atom.]
d. Which alkali metal has the lowest (highest) ionization energy and loses its electrons most (least) readily? Explain. How does the first ionization energy value change with increasing atomic number? Why? [Francium has the lowest ionization energy and will lose electrons most easily. This is due to its large atomic size and its relative inability to hold on to its outer electron. Lithium has the highest ionization energy and loses electrons least easily due to its small atomic size and its relative ability to hold on to its outer electron. First ionization energy decreases with increasing atomic number due to the increasing atomic size.]
e. What can you predict about the second ionization energies of the alkali metals? Would you expect this family of elements to form compounds with 2+ ions? Explain in terms of electronic structures. [Second ionization energies of alkali metals will be large because, in order to remove a second electron, one must attack the next innermost energy level, which is filled in each case. Since this is so difficult a task, compounds with 2+ ions are not likely to be stable.]
f. How do the ionic radii and atomic radii of alkali metal elements compare? Explain. [Ionic radii of alkali metals are all smaller than their respective atomic radii, since alkali metal ions form by removal of the outermost electron which in effect removes the outside energy level and decreases the size of the remaining ion.]
g. The properties of an ion depend on its charge, radius, and inner electronic structure. How does the attractive coulombic force of an ion change with decreasing size? Increasing charge? [As the ion gets smaller, the force of attraction becomes larger (as the square of the distance); as the ionic charge increases, the force of attraction becomes larger also.]
h. List the alkali metal ions in order of increasing electron affinity. [Electron affinities: Fr < Cs < Rb < K < Na < Li]
i. How does electronegativity of the alkali metals change going from Cs to Li? [Electronegativity increases from Cs to Li.]
j. Inview of the decrease in attractive forces in the metallic lattices going from Li to Cs, predict how the following properties change in going down the column of alkali metals: (1) melting point, (2) boiling point, (3) heat of fusion, (4) heat of vaporization, (5) heat of atomization, (6) hardness, (7) critical pressure.[(1) melting points should decrease, (2) boiling points decrease, (3) heats of fusion decrease, (4) heats of vaporization decrease, (5) heats of atomization decrease,(6) hardnessdecreases, (7) criticalpressuredecreases.]
Optional Questions
k. Discuss the trend in density with respect to trends in metallic radius and atomicmass. [Densitygenerallyincreases,meaning thattheatomicmass must increase at a greater rate than does the volume.]
l. Discuss the relationship between trends in heat of fusion, metallic radius, and meltingpoint.[Asone goesdownthePeriodic Table,themetallicradius ofalkalimetalsincreases. Theincreaseinradius resultsinmuchdecreased attractive forces between atoms within the lattice structure, resulting in a decrease in heat of fusion and melting points because less heat is needed to break apart the solid lattice among the larger alkali metals.]
m. Discuss the relationship between trends in heat of vaporization, metallic radius, andboilingpoint.[As onegoesdownthe PeriodicTable,themetallic radius of alkali metals increases. The increase in radius results in much decreasedattractive forcesbetweenatomswithi n theliquidmetals,resulting in a decrease in heat of vaporization and boiling points because less heat is needed to separate atoms from the liquid state within the larger alkalimetals.]
n. The "ionic potential" combines the properties of both charge and ionic radius into one numerical value. For Li+
What can you conclude from the numerical values of the ionic potentials, goingfromLi+to Cs+? [Ionic potential decrease going from Li+ to Cs+. Since the ionic charge is +1 for all alkali metals, this decrease in ionic potential must be due to the increasing ionic radius (in the denominator of the term).]
o. Consider the surface charge on the alkali metal cations and the resulting difference in attraction for anions.
Compare the charge density of Li and Cs. Which ion will exert more attractive force on nearby anions? Why?[Lithium ion will exert more attractive force onnearbyanionsbecause ithasacharge densityof22nm 2 vs.2.8nm2 for cesium.The greater charge density for lithium indicates a greater charge per surface area ratio that will result in greater attractive forces (for anions).]
p. Remembering the charge density of Li+ and Cs+, predict their relative abilities to attract water molecules in solution. What effect would this have on the trend in the size of hydrated alkali metal ions? (See the table ofdatatocheck yourpredictions.)[Lithiumshould haveagreaterattractive force for water molecules than cesium, resulting in the hydrated lithium ionÕs being larger than the hydrated cesium ion. LithiumÕs hydrated ion is the largest of all the alkali metals.]Ionic potential = +1 0.060nm = 16.7nm 1 Charge density = Chargeoncation Surfaceareaofcation( +1 4pr ) 2
In the horizontal direction the entries correspond to the principal quantum number, n, and to the period (horizontal row) in the Periodic Table. The vertical direction relates to the azimuthal quantum number, l, and lists the number of electrons in the sublevels (s, p, d, and f) of each main energy level. As one goes down a family (group) of elements, electrons are found in one additional main energy level. Except for the first element in each family, the electron arrangement in the outer two levels remains the same. If the atomic number increases by 8, a new column of 2,6 is added; if it increases by 18 in addition to the new column of 2,6, ten (10) electrons are added to the third from the outermost column; whereas, with an increase of 32 in the atomic number, the addition adds a group of 14 to the fourth outermost column. Going down the alkali metal column, additions are made internally, but the outermost energy level stays the same; e.g.,
Going across a period; e.g., Na to Ar, no additional vertical columns are added, but the outermost energy level increases one electron at a time; e.g., l
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