Chem 2 AP Homework #8-1: Problems pg. 338-339 #1, 4, 5, 6, 10, 12, 13, 19, 24, 27, 29(f, g, j, k, l), 31, 34,
35, 37, 41, 42, 43, 45, 46, 47
1
Briefly describe the significance of Mendeleev’ periodic table.
Mendeleev’s periodic table grouped elements more accurately, according to their properties, than
previous tables and was able to predict the existence and discovery of elements and their properties.
4
What is the most important relationship among elements in the same group in the periodic
table?
For elements in the same group, the most important relationship is having the same number of valence
electrons, so they have similar chemical properties.
5
Which of the following elements are metals, non-metals, or metalloids? As, Xe, Fe, Li, B, Cl, Ba,
P, I, Si.
Metals: Fe, Li, Ba; Nonmetals: Xe, Cl, P, I; Metalloids: As, B, Si
6
Compare the physical and chemical properties of metals and nonmetals.
Metals: malleable, ductile, generally shiny, electrically and thermally conductive, medium to high melting
points, lose electrons in chemical reactions. Non-metals: dull, brittle, electrically and thermally insulating,
low melting and boiling points, generally gain electrons in chemical reactions.
10
Indicate whether the following elements exist as atomic species, molecular species or extensive
three-dimensional structures in their most stable states at 25°C and 1 atm and write the
molecular or empirical formula for each one.
Atomic: neon, Ne
Molecular: iodine, I2; oxygen, O2
3-D structures: phosphorous, P; magnesium, Mg; carbon, C; sulfur, S8; cesium, Cs
12
What are valence electrons? For representative elements, the number of valence electrons of an
element is equal to its group number. Show that this is true for the following elements: Al, Sr, K,
Br, P, S, C.
Valence electrons are the electrons involved in chemical bonding, generally the s and p electrons from
an element’s highest occupied energy level. Al: 3 v.e., group 3A; Sr: 2 v.e., group 2A; K: 1 v.e., group 1A;
Br: 7 v.e., group 7A; P: 5 v.e., group 5A; S: 6 v.e., group 6A; C: 4 v.e., group 4A.
13
Write the outer electron configurations for the:
(a) alkali metals: ns1; (b) alkaline earth metals: ns2; (c) halogens: ns2np5; (d) noble gases: ns2np6.
19
In the periodic table, the element hydrogen is sometimes grouped with the alkali metals and
sometime with the halogens. Explain why hydrogen can resemble the Group 1A and the Group
7A elements.
Hydrogen forms the H+ ion (resembles the alkali metals) and the H– ion (resembles the halogens).
24
Specify the group of the periodic table in which each of the following elements is found:
(a) [Ne]3s 1: Group 1A
(b) [Ne]3s 23p 3: Group 5A
2
6
(c) [Ne]3s 3p : Group 8A
(d)
[Ar]4s 23d 8: Group 8B
2
27
HOMEWORK #8-1 ANSWER KEY
Write the ground-state electron configurations of the following ions:
Li+: 1s2
(g) Se2–: [Ar]4s23d104p6
(m) Ba2+: [Xe]
(b) H–: 1s2
(h) Br–: [Ar]4s23d104p6
(n) Pb2+: [Xe]6s24f145d10
(i) Rb+: [Kr]
(o) In3+: [Kr]5d10
(j) Sr2+: [Kr]
(p) Tl+: [Xe]6s24f145d10
(q) Tl3+: [Xe]4f145d10
(a)
(c)
N3–: 1s22s22p6
(d) F–: 1s22s22p6
29
(e)
S2–: [Ne]3s23p6
(k) Sn2+: [Kr]5s24d10
(f)
Al3+: [Ne]
(l) Te2–: [Kr]5s24d105p6
Write the ground-state electron configurations of the following transition metal ions:
This exercise simply depends on determining the total number of electrons and using Figure 7.23 and
Table 7.3 of the text.
(f) Fe2+: [Ar]3d6
(g) Fe3+: [Ar]3d5
(j) Zn2+: [Ar]3d10
31
(k) Cu2+: [Ar]3d9
(l) Ag+: [Kr]4d10
Which of the following species are isoelectronic with each other?
Two species are isoelectronic if they have the same number of electrons. Can two neutral atoms of
different elements be isoelectronic? [No]
(a) C and B– are isoelectronic.
(b) Mn2+ and Fe3+ are isoelectronic.
(c) Ar and Cl– are isoelectronic.
(d) Zn and Ge2+ are isoelectronic.
34
How does atomic radius change (a) from left to right across a period and (b) from top to
bottom in a group?
(a) The atomic radius decreases from left to right across a period; (b) the atomic radius increases from top
to bottom in a group.
35
Define ionic radius. How does the size of an atom change when it is converted to (a) and anion
and (b) a cation?
The ionic radius is the radius of a cation or anion. (a) The radius of an anion is larger than the neutral
atom; (b) the radius of a cation is smaller than the neutral atom.
37
On the basis of their positions in the periodic table, select the atom with the larger atomic
radius in each of the following pairs:
(a) Na, Cs: Cs is larger. Cs is below Na in the same group, so its valence electrons are in a more
shielded orbital at a higher energy, thus experience less coulombic attraction to the nucleus..
(b) Be, Ba: Ba is larger, for the same reasoning as in (a).
(c) N, Sb: Sb is larger, for the same reasoning as in (a).
(d) F, Br: Br is larger, for the same reasoning as in (a).
(e) Ne, Xe: Xe is larger, for the same reasoning as in (a).
41
Why is the radius of the lithium atom considerably larger than the radius of the hydrogen atom?
The electron configuration of lithium is 1s22s1. The two 1s electrons shield the 2s electron effectively
from the nucleus. Consequently, the lithium atom is considerably larger than the hydrogen atom.
HOMEWORK #8-1 ANSWER KEY
3
42
Use the second period of the periodic table as an example to show that the size of atoms
decreases as we move from left to right. Explain the trend.
The atomic radius is largely determined by how strongly the outer-shell electrons are held by the
nucleus. The larger the effective nuclear charge, the more strongly the electrons are held and the smaller
the atomic radius. For the second period, the atomic radius of Li is largest because the 2s electron is
well shielded by the filled 1s shell. The effective nuclear charge that the outermost electrons feel
increases across the period as a result of incomplete shielding by electrons in the same shell.
Consequently, the orbital containing the electrons is compressed and the atomic radius decreases.
43
Indicate which one of the two species in each of the following pairs is smaller:
(a) Cl or Cl–: Cl is smaller than Cl–. An atom gets bigger when electrons are added because additional
electron-electron repulsion reduces the coulombic attraction of the electrons to the nucleus.
(b) Na or a Na+: Na+ is smaller than Na. An atom gets smaller when its valence electrons are removed
because the remaining electrons are in less-shielded orbitals with stronger coulombic attraction to
the nucleus.
(c) O2– or S2–: O2 is smaller than S2 . Both elements belong to the same group, and ionic radius
increases going down a group (see #37a).
(d) Mg2+ or Al3+: Al3+ is smaller than Mg2+. The two ions are isoelectronic and in such cases the radius
gets smaller as the nuclear charge becomes more positive and coulombic attraction increases..
(e) Au+ or Au3+: Au3+ is smaller than Au+ for the same reason as part (b).
−
−
In each of the above cases from which atom would it be harder to remove an electron? [The smaller
and/or more positive one]
45
Explain which of the following cations is larger, and why: Cu+ or Cu2+.
The Cu+ ion is larger than Cu2+ because it has one more electron, so increased electron-electron
repulsion reduces the coulombic attraction to the nucleus.
46
Explain which of the following anions is larger, and why: Se2– or Te2–.
Both selenium and tellurium are Group 6A elements. Since valence electrons occupy more shielded
orbitals with lower coulombic attraction going down a column in the periodic table, it follows that Te2
must be larger than Se2 .
−
−
47
Give the physical states (gas, liquid, or solid) of the representative elements in the fourth period
(K, Ca, Ga, Ge, As, Se, Br) at 1 atm and 25°C.
Bromine is liquid; all the others are solids.