Chemistry Chapter 4 and 5 - CHE 111
The order of filling quantum- mechanical orbitals in multi- electron atom is:____
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s.
Write an electron configuration for phosphorus. Identify the valence electrons and core electrons.
1s2 2s2 2p6 3s2 3p3 or [Ne] 3s2 3p3. Valence electrons: 5 Core electrons: 10
Write the electron configuration for Ge. Identify the valence electrons and the core electrons.
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2. 4 valence electrons and 28 core electrons.
Write the electron configuration for Se.
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 or [Ar] 4s2 3d10 4p4
Write the orbital diagram for Ar and determine its number of unpaired electrons.
1s= 1 up 1 down 2s= 1 up 1 down 2p= 1 up 1 down 1 up 1 down 1 up 1 down 3s= 1 up 1 down 3p= 1 up 1 down 1 up 1 down 1 up 1 down. No unpaired electrons
What is the charge of the ion most commonly formed by S?
2-
Use the periodic table to determine each quantity. A.) the number of 2s electrons in Li B.) the number of 3d electrons in Cu C.) the number of 4p electrons in Br D.) the number of 4d electrons in Zr
A.) 1 B.) 10 C.) 5 D.) 2
List the number of valence electrons for each element and classify each element as an alkali metal, alkaline earth metal, halogen, or noble gas. A.) Sodium B.) Iodine C.) Calcium D.) Barium E.) Krypton
A.) 1 valence electron, alkali metal B.) 7 valence electrons, halogen C.) 2 valence electrons, alkaline Earth metal D.) 2 valence electrons, alkaline Earth metal E.) 8 valence electrons, noble gas
Write the full electron configuration for each element. A.) Si B.) O C.) K D.) Ne
A.) 1s2 2s2 2p6 3s2 3p2 B.) 1s2 2s2 2p4 C.) 1s2 2s2 2p6 3s2 3p6 4s1 D.) 1s2 2s2 2p6
Write an electron configuration for each element. A.) Mg B.) P C.) Br D.) Al
A.) 1s2 2s2 2p6 3s2 or [Ne] 3s2 B.) 1s2 2s2 2p6 3s2 3p3 or [Ne] 3s2 3p3 C.) 1s2 2s2 2p6 3s23p6 4s2 3d10 4p5 or [Ar] 4s2 3d10 4p5 D.) 1s2 2s2 2p6 3s2 3p1 or [Ne] 3s2 3p1
Write the full orbital diagram for each element. A.) N B.) F C.) Mg D.) Al
A.) 1s= 1 up 1 down. 2s= 1 up 1 down. 2p= 3 up B.) 1s= 1 up 1 down. 2s= 1 up 1 down. 2p= 1 up 1 down 1 up 1 down 1 up. C.) 1s= 1 up 1 down. 2s= 1 up 1 down. 2p= 1 up 1 down 1 up 1 down 1 up 1 down. 3s= 1 up 1 down. D.) 1s= 1 up 1 down. 2s= 1 up 1 down. 2p= 1 up 1 down 1 up 1 down 1 up 1 down. 3s= 1 up 1 down. 3p= 1 up.
Determine the number of valence electrons in each element. A.) Ba B.) Cs C.) Ni D.) S
A.) 2 B.) 1 C.) 10 D.) 6
Predict the charge of the ion formed by each element and write the electron configuration of the ion. A.) O B.) K C.) Al D.) Rb
A.) 2- [Ne] B.) 1+ [Ar] C.) 3+ [Ne] D.) 1+ [Kr]
Choose the element with the higher first ionization energy in each pair. A.) Br or Bi B.) Na or Rb C.) As or At D.) P or Sn
A.) Br B.) Na C.) Not possible D.) P
On the basis of periodic trends, determine the element in each pair with the higher first ionization energy (if possible). A.) Sn or I B.) Ca or Sr C.) C or P D.) F or S
A.) I B.) Ca C.) not possible D.) F
Choose the larger atom in each pair. A.) Al or In B.) Si or N C.) P or Pb D.) C or F
A.) In B.) Si C.) Pb D.) C
Choose the larger atom or ion from each pair. A.) K or K+ B.) F or F- C.) Ca2+ or Cl-
A.) K B.) F- C.) Cl-
Which is the larger species in each pair? A.) Li or Li+ B.) I- or Cs+ C.) Cr or Cr3+ D.) O or O2-
A.) Li B.) I- C.) Cr D.) O2-
On the basis of periodic trends, choose the larger atom in each pair (if possible). Explain your choices. A.) N or F B.) C or Ge C.) N or Al D.) Al or Ge
A.) N atoms are larger than F atoms because, as you trace the path between N and F on the periodic table, you move to the right within the same period. As you move right across a period, the effective nuclear charge experienced by the outermost electrons increases, resulting in smaller radius. B.) Ge atoms are larger than C toms because, as you trace the path between C and Ge on the periodic table, you move down a column. Atomic size increases as you move down a column because the outermost electrons occupy orbitals with a higher principal quantum number that are therefore larger, resulting in a larger atom. C.) Al atoms are larger than N atoms because as you trace the path between N and Al on the periodic table, you move down a column (atomic size increases) and then to the left across a period (atomic size increases). These effects add together for an overall increase. D.) Based on the periodic trends alone, you cannot tell which atom is larger, because as you trace the path between Al and Ge you go to the right across a period (atomic size decreases) and then down to a column (atomic size increases.) These effects tend to counter each other, and it is not easy to tell which will predominate.
Predict the charges of the monoatomic ions formed by these main- group elements. A.) N B.) Rb
A.) N3- B.) Rb+
Identify the polyatomic ion including its charge in each compound. A.) KNO2 B.) CaS04 C.) Mg(NO3)2
A.) NO2 - B.) S04 2- C.) NO3 -
Choose the element with the more negative (more exothermic) electron affinity in each pair. A.) Na or Rb B.) B or S C.) C or N D.) Li or F
A.) Na B.) S C.) C D.) F
Write the name of each element and classify it as a metal, nonmetal, or metalloid. A.) K B.) Ba C.) I D.) O E.) Sb
A.) Potassium, metal B.) Barium, metal C.) Iodine, nonmetal D.) Oxygen, nonmetal E.) Antimony, metalloid
On the basis of periodic trends, determine which element in each pair has the higher first ionization energy (if possible) A.) Al or S B.) As or Sb C.) N or Si D.) O or Cl
A.) S B.) As C.) N D.) not possible
Choose the larger atom or ion from each pair. A.) S or S2- B.) Ca or Ca2+ C.) Br- or Kr
A.) S2- B.) Ca C.) Br-
On the basis of periodic trends, chose the more metallic element from each pair (if possible) A.) Sn or Te B.) P or Sb C.) Ge or In D.) S or Br
A.) Sn B.) Sb C.) In D.) Not possible
On the basis or periodic trends, choose the more metallic element for each pair (if possible). A.) Ge or Sn B.) Ga or Sn C.) P or Bi D.) B or N
A.) Sn B.) not possible C.) Bi D.) B
On the basis of periodic trends, choose the larger atom in each pair (if possible) A.) Sn or I B.) Ge or Po C.) Cr or W D.) F or Se
A.) Sn B.) not possible C.) W D.) Se
Determine whether each element is a main-group element? A.) Tellurium B.) Potassium C.) Vanadium D.) Manganese
A.) Tellurium B.) Potassium
Write the electron configuration and orbital diagram for each ion and predict whether each will be paramagnetic or diamagnetic. A.) Co2+ B.) N3- C.) Ca2+
A.) [Ar] 4s0 3d7. Paramagnetic. Orbital diagram: [Ar] 4s= none. 3d= 1 up 1 down 1 up 1 down 1 up 1 up 1 up. B.) [He] 2s2 2p6. Diamagnetic. Orbital diagram: [He] 2s= 1 up 1 down. 2p= 1 up 1 down 1 up 1 down 1 up 1 down. C.) [Ne] 3s2 3p6. Diamagnetic. Orbital diagram= [Ne] 3s= 1 up 1 down. 3p= 1 up 1 down 1 up 1 down 1 up 1 down.
Write the electron configuration for each ion. A.) O2- B.) Br- C.) Sr2+ D.) C03+ E.) Cu2+
A.) [Ne] B.) [Kr] C.) [Kr] D.) [Ar] 3d6 E.) [Ar] 3d9
Use the periodic table to write an electron configuration for each element. Represent core electrons with the symbol of the previous noble gas in brackets. A.) P B.) Ge C.) Zr D.) I
A.) [Ne] 3s2 3p3 B.) [Ar] 4s2 3d10 4p2 C.) [Kr] 5s2 4d2 D.) [Kr] 5s2 4d10 5p5
Which set of four quantum numbers corresponds to an electron in a 4p orbital? A.) n=4, l=1, ml=0, ms=1/2 B.) n=4, l=3, ml=3, ms=-1/2 C.) n=4, l=2, ml=0, ms=1/2 D.) n=4, l=4, ml=3, ms=-1/2
A.) n=4, l=1, ml=0, ms=1/2
How can you determine the electron configuration of an ion?
Add or subtract the corresponding number of electrons to the electron configuration of the neutral atom.
Write the electron configuration and orbital diagram for each ion and determine whether each is diamagnetic or paramagnetic. A.) Al3+ B.) S2- C.) Fe3+
Al: [Ne] 3s2 3p1 Al 3+: [Ne] or [He] 2s2 2p6 Orbital diagram: [He] 2s= 1 up 1 down. 2p= 1 up 1 down 1 up 1 down 1 up 1 down. Diamagnetic S: [Ne] 3s2 3p4 S2-: [Ne] 3s2 3p6 Orbital Diagram: [Ne] 3s= 1 up 1 down. 3p= 1 up 1 down 1 up 1 down 1 up 1 down. Diamagnetic Fe: [Ar] 4s2 3d6 Fe3+: [Ar] 4s0 3d5 Orbital Diagram: [Ar] 4s= none. 3d= 5 ups. Paramagnetic
Predict the charges of the monoatomic (single atom) ions formed by these main- group elements. A.) Al B.) S
Aluminum is a main- group metal and tends to lose electrons to form a cation with the same electron configuration as the nearest noble gas. The electron configuration of aluminum is 1s2 2s2 2p6 3s2 3p1. The nearest noble gas is neon, which has an electron configuration of 1s2 2s2 2p6. Therefore, aluminu loses three electrons to form the cation Al3+. Sulfur is a nonmetal and tends to gain electrons to form an anion with the same electron configuration as the nearest noble gas. The electron configuration of sulfur is 1s2 2s2 2p6 3s2 3p4. The nearest noble gas is argon, which has an electron configuration of 1s2 2s2 2p6 3s2 3p6. Therefore, sulfur gains two electrons to form the anion S2-.
Choose the correct electron configuration for Se. A.) 1s2 2s2 2p6 3s2 3p4 B.) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 C.) 1s2 2s2 2p6 3s2 3p6 4s2 4p4 D.) 1s2 2s2 2p6 3s2 3p6 4s2 3d4
B.) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4
Which statement is true about nuclear charge? A.) Effective nuclear charge decreases as you move to the right across a row in the periodic table. B.) Effective nuclear charge increases as you move to the right across a row in the periodic table. C.) Effective nuclear charge remains relatively constant as you move to the right across a row in the periodic table. D.) Effective nuclear charge increases then decreases at regular intervals as you move to the right across a row in the periodic table.
B.) Effective nuclear charge increases as you move to the right across a row in the periodic table.
Which species is diamagnetic? A.) Cr2+ B.) Zn C.) Mn D.) C
B.) Zn
According to Coulumb's law, if the separation between two particles of the same charge is doubled, the potential energy of the two particles__. A.) become twice as high as it was before the distance separation. B.) becomes one- half as high it was before the separation. C.) does not change D.) becomes one- fourth as high as it was before the separation.
B.) becomes one- half as high it was before the separation.
Which electron in S is most shielded from nuclear charge? A.) An electron in the 1s orbital B.) An electron in a 2p orbital C.) An electron in a 3p orbital D.) None of the above
C.) An electron in the 3p orbital
Which statement is true about electron shielding of nuclear charge? A.) Outermost electrons efficiently shield one another from nuclear charge. B.) Core electrons efficiently shield one another from nuclear charge. C.) Outermost electrons efficiently shield core electrons from nuclear charge. D.) Core electrons efficiently shield outermost electrons from nuclear charge.
D.) Core electrons efficiently shield outermost electrons from nuclear charge.
Which element has the smallest atomic radius? A.) C B.) Si C.) Be D.) F
D.) F
The periodic table was primarily developed by___ in the 19th century. This person arranged the elements in a table so the ___ increased from left to right in a row and elements with ___ fell in the same columns.
Dmitri Mendeleev/atomic mass/similar properties
____- the energy associated with an element in its gaseous state gaining an electron- does not show a general trend as we move down a column in the periodic table, but it generally becomes more __ (more exothermic) to the right across row.
Electron affinity/negative
Write the orbital diagram for sulfur and determine its number of unpaired electrons.
Electron configuration: 1s2 2s2 2p6 3s2 3p4. Orbital diagram: 1s= 1 up 1 down. 2s= 1 up 1 down. 2p= 1 up 1 down 1 up 1 down 1 up 1 down. 3s= 1 up 1 down. 3p= 1 up 1 down 1 up 1 up. Two unpaired electrons.
___- the tendency to lose electrons in a chemical reaction- generally __ down a column in the periodic table and __ to the right across a row.
Metallic character/increases/decreases
Elements on the left side and in the center of the periodic table are __ and tend to lose electrons when they undergo chemical changes.
Metals
The compound NCl3 is nitrogen trichloride, but AlCl3 is simply aluminum chloride. Why?
NCl3 is molecular and AlCl3 is ionic.
Which electrons experience a greater effective nuclear charge: the valence electrons in beryllium or the valence electrons in nitrogen? Why?
Nitrogen because the valence electrons of both atoms are screened by two core electrons, but Nitrogen has a greater number of protons and therefore a greater net nuclear charge.
Elements on the upper right side of the periodic table are __ and tend to gain electrons when they undergo chemical changes.
Nonmetals
Arrange this isoelectronic series in order of decreasing radius: F-, Ne, O2-, Mg2+, Na+
O2-, F-, Ne, Na+, Mg2+
In the previous sections, we have seen how the number of electrons and the number of protons affect the size of an atom or an ion. However we have not considered how the number of neutrons affects the size of the atom. Why not? Would you expect isotopes- for example, C-12 and C-13- to have different atomic radii?
The isotopes of an element all have the same radii for two reasons: 1.) neutrons are highly negligibly small compared to the size of an atom and therefore extra neutrons do not increase atomic size. 2.) Neutrons have no charge and therefore do not attract electrons in the way that protons do.
What is wrong with the following statement? "Atoms form bonds in order to satisfy the octet rule."
The reasons that atoms form bonds are complex. One contributing factor is the lowering of their potential energy. The octet rule is just a handy way to predict the combinations of atoms that will have a lower potential energy when they bond together.
Based on what you learned in Chapter 1 about atoms, what part of the atom do you think the spheres in the molecular models shown here represent? If you were to superimpose a nucleus on one of these spheres, how big would you draw it?
The spheres represent the electron cloud of the atom. It would be nearly impossible to draw a nucleus to scale on any of the space- filling molecular models in this book- on this scale, the nucleus would be too small to see.
According to Coulomb's law, what happens to the potential energy of two oppositely charges particles as they get closer together? A.) Their potential energy decreases. B.) Their potential energy increases. C.) Their potential energy does not change.
Their potential energy increases. Since the charges are opposite, the potential energy of the interaction is negative. As the charges get closer together, r becomes smaller and the potential energy decreases (it becomes more negative).
Elements with one or two valence electrons are among the most ___, readily losing their valence electrons to attain noble gas configurations.
active metals
Elements with six or seven valence electrons are among the most ___, readily gaining enough electrons to attain a noble gas configuration.
active nonmetals
Because quantum- mechanical orbitals fill sequentially with increasing ___, we can infer the electron configuration of an element from its position in the periodic table.
atomic number
As we move across a row in the periodic table, atomic radii __ because the effective nuclear charge- the net or average charge experienced by the atom's outermost electrons-___.
decrease/increases
An ___ for an atom shows which quantum- mechanical orbitals the atom's electrons occupy.
electron configuration
The ___- the energy required to remove an electron from an atom in the gaseous state- generally __ as we move down a column in the periodic table and ___ when we move to the right across a row.
ionization energy/decreases/increases
Many main- group elements form __ with noble gas electron configurations.
ions
What are the four quantum numbers for each of the two electrons in a 4s orbital?
n=4, l=0, ml=0, ms=1/2 n=4, l=0, ml=0, ms=-1/2
The most stable (or chemically unreactive) elements in the periodic table are the ___. These elements have completely full principal energy levels, which have particularly low potential energy compared to other possible electron configurations.
noble gases
The size of an atom is largely determined by its ___. As we move down a column in the periodic table, the principal quantum number (n) of the outermost electrons ___, resulting in successively larger orbitals and therefore larger atomic radii.
outermost electrons/increases
Which outer electron configuration would you expect to correspond to reactive metal? To a reactive nonmetal. A.) ns^2 B.)ns^2 np^6 C.)ns^2 np^5 D.)ns^2 np^2
reactive metal: A reactive nonmetal: C
The radius of a cation is much __ than that of the corresponding atom, and the radius of an anion is much __ than that of the corresponding atom.
smaller/larger
Successive ionization energies increase ___ from one valence electron to the next, but the ionization energy increases __ for the first core electrons.
smoothly/dramatically
The atomic radii of the transition elements ___ as we move across each row because electrons are added to the n highest -1 orbitals while the number of highest n electrons stay ___.
stay roughly constant/roughly constant
For main group ions, the order of removing electrons is ___ as the order in which they are added in building up the electron configuration.
the same
For ___, ns electrons are removed before (n-1) d electrons.
transition metal atoms
An atom's outermost electrons, also called ___ are most important in determining the atom's properties.
valance electrons
Which statement best summarizes the difference between ionic and molecular compounds? A.) Molecular compounds contain highly directional covalent bonds which results in the formation of molecules- discrete particles that do not covalently bond to each other. Ionic compounds contain nondirectional ionic bonds, which results (in solid phase) in the formation of ionic lattices- extended networks of alternating cations and anion. B.) Molecular compounds contain covalent bonds in which one of the atoms shares an electron with the other one, resulting in a new force that holds the atoms together in a covalent molecule. Ionic compounds contain ionic bonds in which one atom donates an electron to the other, resulting in a new force that holds the ions together in pairs (in the solid phase). C.) The key difference between ionic and covalent compounds is the types of elements that compose them, not the way that the atoms bond
A
Based on what you just learned about ionization energies, explain why valence electrons are more important than core elements in determining the reactivity and bonding in atoms.
As you can see from the successive ionization energies of any element, valence electrons are held most loosely and can therefore be transferred or shared most easily. Core electrons on the other hand, are held tightly and are not easily transferred or shared. Consequently, valence electrons play a central role in chemical bonding.
Which statement is true? A.) An orbital that penetrates into the region occupied by core electrons is more shielded from nuclear charge than an orbital that does not penetrate and will therefore have a higher energy. B.) An orbital that penetrates into the region occupied by core electrons is less shielded from nuclear charge than an orbital that does not penetrate and will therefore have a higher energy. C.) An orbital that penetrates into the region occupied by core electrons is less shielded from nuclear charge than an orbital that does not penetrate and will therefore have a lower energy. D.) An orbital that penetrates into the region occupied by core electrons is more shielded from nuclear charge than an orbital that does not penetrate and will therefore have lower energy.
C.) An orbital that penetrates into the region occupied by core electrons is less shielded from nuclear charge than an orbital that does not penetrate and will therefore have a lower energy (Penetration results in less shielding from nuclear charge and therefore lower energy.)
For which element is gaining of an electron most exothermic? A.) Li B.) N C.) F D.) B
C.) F
The ionization energies of an unknown third period element are shown here. Identify the element. IE1= 786 kJ/ mol; IE2= 1580 kJ/mol; IE3= 3230 kJ/mol; IE4= 4360 kJ/mol; IE5= 16100 kJ/mol A.) Mg B.) Al C.) Si D.) P
C.) Si
Which electrons experience the greatest effective nuclear charge? A.) The valence electrons in Mg B.) The valence electrons in Al C.) The valence electrons in S
C.) The valence electrons in S. Since Zeff increases from left to right across a row in a periodic table, the valence electrons in S experience a greater effective nuclear charge than the valence electrons in Al or Mg.
Which pair of element do you expect to be the most similar? Why? A.) N and Ni B.) Mo and Sn C.) Na and Mg D.) Cl and F E.) Si and P
Chlorine and Fluorine because they are in the same group or family. Elements in the same group or family have similar chemical properties.
Arrange the following in order of decreasing radius: Ca2+, Ar, Cl-.
Cl-> Ar> Ca2+
Write electron configurations for each element. A.) Cl B.) Si C.) Sr D.) O
Cl: 1s2 2s2 2p6 3s2 3p5 or [Ne] 3s2 3p5 Si: 1s2 2s2 2p6 3s2 3p2 or [Ne] 3s2 3p2 Sr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 or [Kr] 5s2 O: 1s2 2s2 2p4 or [He] 2s2 2p4
Arrange the following elements in order of increasing metallic character: Si, Cl, Na, Rb.
Cl<Si<Na<Rb
Arrange these elements in order of increasing radius: Cs+, Xe, I-.
Cs+<Xe<I-
Arrange the following elements in order of decreasing first ionization energy: S, Ca, F, Rb, Si.
F>S>Si>Ca>Rb
Write the structural formula for water.
H---O---H
According to ___, orbitals of the same energy first fill singly with electrons with parallel spins before pairing.
Hund's rule
Arrange these elements in order of increasing first ionization: Si, F, In, N.
In, Si, N, F
What is the metallic trend?
Metallic character decreases as we move to the right across a row in the periodic table an increases as we move down a column.
According to the _____, each orbital can hold a maximum of two electrons (with opposing spins).
Pauli exclusion principal
___ are predictable based on an element's position within the periodic table. This includes atomic radius, ionization energy, electron affinity, density, and metallic character.
Periodic properties
____ explains the periodic table by showing how electrons fill the quantum- mechanical orbitals within the atoms that compose the elements.
Quantum mechanics
Arrange the elements in order of decreasing radius: S, Ca, F, Rb, Si.
Rb> Ca> Si> S> F
Arrange these elements in order of increasing first ionization energy: Cl, Sn, Si.
Sn<Si<Cl
Use the trends in ionization energy and electron affinity to explain why sodium chloride has the formula NaCl and not Na2Cl2.
The 3s electron in sodium has a relatively low ionization energy (496 kJ/mol) because it is a valence electron. The energetic cost for sodium to lose a second electron is extraordinarily high (4560 kJ/mol) because the next electron to be lost is a core electron (2p). Similarly, the electron affinity of chlorine to gain one electron (-349 kJ/mol) is highly exothermic since the added electron completes chlorine's valence shell. The gain of a second electron by the negatively charged chlorine anion is not so favorable. Therefore, we expect sodium and chlorine to combine in a 1:1 ratio.
Use the ionic bonding model developed in this section to determine which has the higher melting point, NaCl or MgO. Explain your answer.
You would expect MgO to have the higher melting point because, in our bonding model, the magnesium and oxygen ions are held together in a crystalline lattice by charges of 2+ for magnesium and 2- for oxygen. In contrast, the NaCl lattice is held together by charges of 1+ for sodium and 1- for chlorine. According to Coloumb's law, as long as the spacing between the nation and the anion in the two compounds is not much different, the higher charges in MgO should result in lower potential energy (more stability), and therefore a higher melting point. The experimentally measured melting points of these compounds are 801 degrees Celsius for NaCl and 2852 degrees Celsius for MgO, in accordance with the model.
What is the electron configuration for Fe2+?
[Ar] 4s0 3d6
What is the orbital diagram for Vanadium?
[Ar] 4s= 1 up 1 down. Next orbital: 3 up arrows
Refer to the periodic table to write the electron configuration for iodine (I).
[Kr] 5s2 4d10 5p5
Refer to the periodic table to determine the electron configuration of bismuth (Bi).
[Xe] 6s2 4f14 5d10 6p3
