Chemistry Chapter 4 and 5

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).

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.)

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

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.

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.

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.

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.

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 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.

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

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-.

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.

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

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, 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

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

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

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

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

Refer to the periodic table to determine the electron configuration of bismuth (Bi).

[Xe] 6s2 4f14 5d10 6p3

Refer to the periodic table to write the electron configuration for iodine (I).

[Kr] 5s2 4d10 5p5

Predict the charges of the monoatomic ions formed by these main- group elements.
A.) N
B.) Rb

A.) N3- B.) Rb+

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

Arrange the elements in order of decreasing radius: S, Ca, F, Rb, Si.

Rb> Ca> Si> S> F

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.

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-

Arrange the following in order of decreasing radius: Ca2+, Ar, Cl-.

Cl-> Ar> Ca2+

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

Arrange the following elements in order of decreasing first ionization energy: S, Ca, F, Rb, Si.

F>S>Si>Ca>Rb

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

Arrange the following elements in order of increasing metallic character: Si, Cl, Na, Rb.

Cl<Si<Na<Rb

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

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

What is the orbital diagram for Vanadium?

[Ar] 4s= 1 up 1 down. Next orbital: 3 up arrows

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

Which element has the smallest atomic radius?
A.) C
B.) Si
C.) Be
D.) F

D.) F

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 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.

What is the electron configuration for Fe2+?

[Ar] 4s0 3d6

Which species is diamagnetic?
A.) Cr2+
B.) Zn
C.) Mn
D.) C

B.) Zn

Arrange these elements in order of increasing radius: Cs+, Xe, I-.

Cs+<Xe<I-

Arrange these elements in order of increasing first ionization energy: Cl, Sn, Si.

Sn<Si<Cl

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

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.

For which element is gaining of an electron most exothermic?
A.) Li
B.) N
C.) F
D.) B

C.) F

What is the charge of the ion most commonly formed by S?

2-

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

___ 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

An ___ for an atom shows which quantum- mechanical orbitals the atom’s electrons occupy.

electron configuration

The order of filling quantum- mechanical orbitals in multi- electron atom is:____

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s.

According to the _____, each orbital can hold a maximum of two electrons (with opposing spins).

Pauli exclusion principal

According to ___, orbitals of the same energy first fill singly with electrons with parallel spins before pairing.

Hund’s rule

An atom’s outermost electrons, also called ___ are most important in determining the atom’s properties.

valance electrons

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

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

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

Elements on the upper right side of the periodic table are __ and tend to gain electrons when they undergo chemical changes.

Nonmetals

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

Many main- group elements form __ with noble gas electron configurations.

ions

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

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

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

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.

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

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

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

Successive ionization energies increase ___ from one valence electron to the next, but the ionization energy increases __ for the first core electrons.

smoothly/dramatically

____- 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

___- 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

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

Determine whether each element is a main-group element?
A.) Tellurium
B.) Potassium
C.) Vanadium
D.) Manganese

A.) Tellurium B.) Potassium

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 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.

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

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

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

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

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

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.

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]

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.

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

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

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-

Arrange this isoelectronic series in order of decreasing radius: F-, Ne, O2-, Mg2+, Na+

O2-, F-, Ne, Na+, Mg2+

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

Arrange these elements in order of increasing first ionization: Si, F, In, N.

In, Si, N, F

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 structural formula for water.

H—O—H

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.

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.

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 –

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.

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 together.
D.) A molecular compound is composed of covalently bonded molecules. An ionic compound is composed of ionically bonded molecules (in the solid phase).

A

The compound NCl3 is nitrogen trichloride, but AlCl3 is simply aluminum chloride. Why?

NCl3 is molecular and AlCl3 is ionic.

Which number if the best estimate for the scaling factor used in these models? In other words, by approximately what number would you have to multiply the radius of an actual oxygen atom to get the radius of the sphere used to represent the oxygen atom in the water molecule shown here?
A.) 10
B.) 10^4
C.) 10^8
D.) 10^16

C

Without doing any calculations, list the elements in C6H6O in order of decreasing mass percent composition.

C>O>H

Which ratio can be correctly derived from the molecular formula for water (H2O)?
A.) 2g H: 1g H2O
B.) 2mL H: 1mL H20
C.) 2 mol H: 1 mol H20

C

Write the empirical formulas for the compounds represented by the molecular formulas.
A.) C4H8
B.) B2H6
C.)CCL4

A.) CH2 B.) BH3 C.) Remains the same

Write the empirical formulas for the compounds represented by the molecular formulas.
A.) C5H12
B.) Hg2Cl2
C.) C2H402

A.) C5H12 B.)HgCl C.) CH20

Use the Lewis model to predict the formula for the compound that forms between calcium and chlorine.

CaCl2

Use the Lewis model to predict the formula for the compound that forms between magnesium and nitrogen.

Mg3N2

Write the formula for the ionic compound that forms between aluminum and oxygen.

Al2O3

Write the formula for the ionic compound that forms between potassium and sulfur.

K2S

Write the formula for the ionic compound that forms between calcium and oxygen.

CaO

Write the formula for the ionic compound that forms between Aluminum and Nitrogen.

AlN

Name the compound CaBr2.

calcium bromide

Name the compound Ag3N.

Silver nitride

Write the formula for rubidium sulfide.

Rb2S

Name the compound PbCl4.

lead(IV) chloride

Name the compound FeS.

Iron(II) Sulfide

Write the formula for ruthenium(IV) oxide.

RuO2

Name the compound Li2Cr2O7.

lithium dichromate.

Name the compound Sn(ClO3)2.

tin(II) chlorate

Write the formula for cobalt(II) phosphate.

Co3(PO4)2

Name each compound.
A.) NI3
B.) PCl5
C.) P4S10

A.) nitrogen triiodide B.) phosphorus pentachloride C.) tetraphosphorus decasulfide

Name the compound N205.

dinitrogen pentoxide

Write the formula for phosphorus tribromide.

PBr3

Calculate the formula mass of glucose, C6H12O6.

180.16amu

Calculate the formula mass of calcium nitrate.

164.10 amu

An asprin tablet contains 225 mg of acetylsalicylic acid (C9H8O4). How many acetylsalicylic acid molecules does it contain?

1.09E21 C9H8O4 molecules

Find the number of ibuprofen molecules in the tablet containing 200.0 mg of ibuprofen (C13H1802).

5.839E20

What is the mass of a sample of water containing 3.55E22 H2O molecules?

1.06 g H2O

Calculate the mass percent of Cl in Freon-112 (C2l4F2), a CFC refrigerant.

69.58%

Acetic acid (C2H4O2) is the active ingredient in vinegar. Calculate the mass percent composition of oxygen is acetic acid.

53.29%

Calculate the mass percent composition of sodium in sodium oxide.

74.19% Na

The FDA recommends that a person consume less than 2.4 g of sodium per day. What mass of sodium chloride (in grams) can you consume and still be within the FDA guidelines? Sodium chloride is 39% sodium by mass.

6.2 NaCl

What mass (in grams) of iron(III) oxide contains 58.7g of iron? Iron (III) oxide is 69.94% iron by mass.

83.9 g Fe2O3

If someone consumes 22g of sodium chloride per day, what mass in grams of sodium does that person consume? Sodium chloride is 39% sodium by mass.

8.6 g Na

What mass of hydrogen (in grams) is contained in 1.00 gallon of water? (The density of water is 1.00 g/mL.)

4.23E2g H

Determine the mass of oxygen in a 7.2-g sample of Al2(SO4)3.

4.0 g O

Butane (C4H10) is the liquid fuel in lighters. How many grams of carbon are present within a lighter containing 7.25 mL of butane? (The density of liquid butane is 0.601 g/mL).

3.60 g C

A compound containing nitrogen and oxygen is decomposed in the laboratoory and produces 24.5 g nitrogen and 70.0 g oxygen. Calculate the empirical formula of the compound.

N2O5.

A sample of a compound is decomposed in the laboratory and produces 165 g carbon, 27.8 g hydrogen, and 220.2 g oxygen. Calculate the empirical formula of the compound.

CH20

A laboratory analysis of asprin determined the following mass percent composition:
C 60.00%
H 4.48%
O 35.52%
Fine the empirical formula.

C9H8O4.

The empirical formula of butanedione is C2H30 and its molar mass is 86.09 g/mol. Find its molecular formula.

C4H6O2

A compound with the percent composition shown here has a molar mass of 60.10 g/mol. Determine its molecular formula.
C, 39.97%
H, 13.41%
N, 46.62%

C2H8N2

A compound has the empirical formula CH and a molar mass of 78.11 g/mol. What is its molecular formula?

C6H6

Upon combustion, a compound containing only carbon and hydrogen produces 1.83 g CO2 and 0.901 g H20. Determine the empirical formula of the compound.

C1H2.4

Upon combustion, a compound containing only carbon and hydrogen produced 1.60 g CO2 and 0.819 g H20. Find the empirical formula of the compound.

C2H5

Upon combustion, a 0.8233-g sample of a compound containing only carbon, hydrogen, and oxygen produces 2.445 g CO2 and 6.003 g H20. Determine the empirical formula of the compound.

C10H12O1

Upon combustion, a 0.8009-g sample of a compound containing only carbon, hydrogen, and oxygen produced 1.6004 g CO2 and 0.6551 g H2O. Find the empirical formula of the compound.

C2H40

What is the empirical formula of a compound with the molecular formula C10H8?
A.) C5H3
B.) C2H4
C.) C5H4
D.) CH

c

Which substance is an ionic compound?
A.) SrI2
B.) N204
C.) He
D.) CCl4

a

What is the correct formula for the compound formed between calcium and sulfur?
A.) CaS
B.) Ca2S
C.) CaS2
D.) CaS3

a

Name the compound SrI2.
A.) Strontium iodide
B.) strontium diiodide
C.) strontium (II) iodide
D.) strontium (II) diiodide

a

What is the formula for manganese (IV) oxide?
A.) Mn4O
B.) MnO4
C.) Mn2O
D.) MnO2

d

Name the compound Pb(C2H3O2)2.
A.) lead (II) carbonate
B.) lead (II) acetate
C.) lead bicarbonate
D.) lead diacetate

b

Name the compound P2I4.
A.) Phosphorus iodide
B.) Phosphorus diiodide
C.) Phosphorus (II) iodide
D.) diphosphorus tetraiodide

d

What is the correct Lewis symbol for S?

How many CH2Cl2 molecules are there in 25.0g of CH2Cl2?
A.) 0.294 molecules
B.) 1.77×10^23 molecules
C.) 1.28×10^27 molecules
D.) 1.51×10^25 molecules

b

List the elements in the compound CF2Cl2 in order of decreasing mass percent composition.
A.) C&gt;F&gt;Cl
B.) F&gt;Cl&gt;C
C.) Cl&gt;C&gt;F
D.) Cl&gt;F&gt;C

d

Determine the mass of potassium in 35.5 g of KBr.
A.) 17.4g
B.) 0.298g
C.)11.7g
D.) 32.9g

c

A compound is 52.14% C, 13.13%H, and 34.73% O by mass. What is the empirical formula of the compound?
A.) C2H8O3
B.) C2H6O
C.) C4HO3
D.) C3HO6

b

A compound has the empirical formula CH2O and a formula mass of 120.10 amu. What is the molecular formula of the compound?
A.) CH2O
B.) C2H4O2
C.)C3H6O3
D.) C4H8O4

d

Combustion of 30.42g of a compound containing only carbon, hydrogen, and oxygen produces 35.21g of CO2 and 14.42g H2O. What is the empirical formula of the compound?
A.) C4H8O6
B.) C2H4O3
C.) C2H2O3
D.) C6HO12

b

chemical bond

The sharing or transfer of electrons to attain stable electron configurations for the bonding atoms.

ionic bond

A chemical bond formed between two oppositely charged ions, generally a metallic cation and a nonmetallic anion, that are attracted to one another by electrostatic forces.

ionic compound

A compound composed of cations and anions bound together by electrostatic attraction.

covalent bond

A chemical bond in which two atoms share electrons that interact with the nuclei of both atoms, lowering the potential energy of each through electrostatic interactions.

molecular compound

A compound composed of two or more covalently bonded nonmetals.

chemical formula

A symbolic representation of a compound that indicates the elements present in the compound and the relative number of atoms of each.

empirical formula

A chemical formula that shows the simplest whole- number ratio of atoms in the compound.

molecular formula

A chemical formula the shows the actual number of atoms of each element in a molecule of a compound.

structural formula

A molecular formula that shows how the atoms in a molecule are connected or bonded to each other.

ball-and-stick molecular model

A representation of the arrangement of atoms in a molecule that shows how the atoms are bonded to each other and the overall shape of the molecule.

space- filling molecular model

A representation of a molecule that shows how the atoms fill the space between them.

Lewis model

A simple model of chemical bonding, which uses diagrams to represent bonds between atoms as lines or pairs of dots. In this model, atoms bond together to abstain stable octets (eight valence electrons).

lewis electron- dot structure

A drawing of a molecule that represents chemical bonds between atoms as shared or transferred electrons; the valence electrons are represented as dots.

Lewis symbol

A symbol of an element in which dots represent valence electrons.

octet

A Lewis symbol with eight dots, signifying a filled outer electron shell for s and p block elements.

duet

A Lewis structure with dots, signifying a filled outer electron shell for the elements H and He.

octet rule

The tendency for most bonded atoms to possess or share eight electrons in their outer shell in order to obtain stable electron configurations and lower their potential energy.

formula unit

The smallest, electrically neutral collection of ions in an ionic compound.

lattice energy

The energy associated with forming a crystalline lattice from gaseous ions.

common name

A traditional name of a compound that gives little or no information about its chemical structure.

systematic name

An official name based on well- established rules for a compound, which can be determined by examining its chemical structure.

binary compound

A compound that contains only two different elements.

polyatomic ion

An ion composed of two or more atoms.

oxyanion

A polyatomic anion containing a nonmetal covalently bonded to one or more oxygen atoms.

hydrate

An ionic compound that contains a specific number of water molecules associated with each formula unit.

bonding pair

A pair of electrons shared between two atoms.

lone pair

pair of electrons associated with only one atom.

double bond

The bond that forms when two electrons are shared between two atoms.

triple bond

The bond that forms when three electron pairs are shared between two atoms.

formula mass

The average mass of a molecule of a compound in amu.

mass percent composition

An element’s percentage of the total mass of a compound containing the element.

empirical formula molar mass

The sum of the masses of all the atoms in an empirical formula.

combustion analysis

A method of obtaining empirical formulas for unknown compounds, especially those containing carbon and hydrogen, by burning a sample of the compound in pure oxygen and analyzing the products of the combustion reaction.

organic compound

A compound containing carbon combined with several other elements including hydrogen, nitrogen, oxygen, or sulfur.

hydrocarbon

An organic compound that contains only carbon and hydrogen.

Formula for formula mass

(# atoms of 1st element in chemical formula x atomic mass of first element)+(# atoms of 2nd element in chemical formula x atomic mass of 2nd element)+…

Formula for Mass Percent Composition

Mass % of element X= mass of X in 1 mol compound/ mass of 1 mol compound x 100%

Formula for Empirical Formula Molar Mass

Molecular formula= n x (empirical formula) n= molar mass/ empirical formula molar mass

Chemical bonds, the forces that hold atoms together in compounds arise from the interactions between __ and __ in atoms.

nuclei/electrons

In an ___ bond, one more electrons are transferred from one atom to another, forming a cation and an anion. The two ions are then drawn together by the attraction between the opposite charges.

ionic

In a ___ bond, one or more electrons are shared between two atoms. The atoms are held together by the attraction between their nuclei and the shared electrons.

covalent

A compound is represented with a ___, which indicates the elements present and the number of atoms of each.

chemical formula

An empirical formula gives only the __ number of atoms, while a molecular formula gives the __ number of atoms present in the molecule.

relative/actual

___ formulas show how the atoms are bonded together while __ models portray the geometry of the molecule.

structural/molecular

In the ___ model, chemical bonds are formed when atoms transfer (ionic bonding) or share (covalent bonding) valence electrons to attain noble gas electron configurations.

Lewis

The ____ represents valence electrons as dots surrounding the symbol for an element. When two or more elements bond together, the dots are transferred or shared so that every atom gets eight dots, an octet.

Lewis model

In a __ Lewis structure involving main group metals, the metal transfers its valence electrons (dots) to the nonmetal.

ionic

The formation of most ionic compounds is ___ because of ___, the energy released when metal cations and nonmetal anions coalesce to form the solid.

exothermic/lattice energy

In a ___ Lewis structure, neighboring atoms share valence electrons to attain octets (or duets).

covalent

A single shared electron pair constitutes a ___ bond while two or three shared pairs constitute ___ or ___ bonds, respectively.

single/double/triple

The ___ of a compound is the sum of the atomic masses of all the atoms in the chemical formula. Like the atomic masses of elements, the ___ characterizes the average mass of the molecule ( or formula unit).

formula mass

The mass of one mole of a compound is its ___ and equals its ___ (in grams).

molar mass/formula mass

The ___ of a compound indicates that each element’s percentage of the total compound’s mass. We can determine this from the compound’s ___ and ___ of its elements.

mass percent composition/chemical formula/molar mass

The ___ of a compound provides the relative number of atoms or moles of each element in a compound, and therefore we can use it to determine numerical relationships between moles of the __ and moles of its ___.

chemical formula/numerical relationships/constituent elements

If we know the mass percent composition and molar mass of a compound, we can determine its ___ and ___ formulas.

empirical/molecular

___ are composed of carbon, hydrogen, and a few other elements such as nitrogen, oxygen, and sulfur.

Organic compounds

The simplest organic compounds are ___, compounds composed of only carbon and hydrogen.

hydrocarbons

Classify each compound as ionic or molecular
A.) CO2
B.)NiCl2
C.) NaI
D.) PCL3

A.) molecular B.) ionic C.) ionic D.) molecular

Determine the number of each type of atom in each formula.
A.) Mg3(PO4)2
B.) BaCl2
C.) Fe(NO2)2
D.) Ca(OH)2

A.) 3 Mg, 2 P, 8 O B.) 1 Ba 2 Cl C.) 1 Fe, 2 N, 4 O D.) 1 Ca, 2 O, 2 H

Write an electron configuration for N. Then write a Lewis symbol for N and show which electrons form the electron configuration are included in the Lewis symbol.

1s2 2s2 2p3

Use Lewis symbols to determine the formula for the compound that forms between each pair of elements.
A.) Sr and Se
B.) Ba and Cl
C.) Na and S
D.) Al and O

a.) SrSe b.)BaCl2 c.)Na2S d.)Al2O3

The lattice energy of CsF is -744 kJ/mol, whereas that of BaO is -3029 kJ/mol. Explain this large difference in lattice energy.

One factor of lattice energy is the product of the charges of the two ions. The product of the ion charges for CsF is +1, while that for BaO is +4. Because this product is four times greater, the lattice energy is also four times greater.

Write a formula for the ionic compound that forms between each pair of elements.
A.) Calcium and Oxygen
B.) Zinc and Sulfur
C.) Rubidium and bromine
D.) Aluminum and Oxygen

a.) CaO b.) ZnS c.)RbBr d.) Al2O3

Name each ionic compound.
A.) Mg3N2
B.) KF
C.) Na2O
D.) Li2S
E.) CsF
F.) KI

a.) magnesium nitride b.) potassium fluoride c.) sodium oxide d.) lithium sulfide e.) cesium fluoride f.) potassium iodide

Use covalent Lewis structures to explain why each element (or family of elements) occurs as diatomic molecules.
A.) hydrogen
B.) the halogens
C.) Oxygen
D.) Nitrogen

a.) H:H, filled duets . . . . b.):Cl:Cl:, filled octets . . . . . . . . c.) O=O, filled octets . . . . d.) :N (triple bond) N:, filled octets

Name each molecular compound.
A.) CO
B.) NI3
C.) SiCl4
D.) N4Se4

a.) carbon monoxide b.) nitrogen triiodide c.) silicon tetrachloride d.) tetranitrogen tetraselenide

Write a formula for each molecular compound.
A.) phosphorus trichloride
B.) chlorine monoxide
C.) disulfur tetrafluoride
D.) phosphorus pentafluoride

a.) PCl3 b.) ClO c.) S2F4 d.) PF5

Name each compound.
a.) SrCl2
b.) SnO2
c.)P2S5

a.) strontium chloride b.) Tin (IV) oxide c.) Diphosphorus pentasulfide

Calculate the formula mass for each compound.
a.) NO2
b.) C4H10
c.) C6H12O6
d.) Cr(NO3)3

a.) 46.01 amu b.) 58.12 amu c.) 180.16 amu d.) 238.03 amu

Calculate the number of moles in each sample.
a.) 72.5 g CCL4
b.) 12.4 g C12H22O11
c.) 25.2 kg C2H2
d.) 12.3 g of dinitrogen monoxide

a.) 0.471 mol b.) 0.0362 mol c.) 968 mol d.) 0.279 mol

Calculate the mass (in g) of each sample.
A.) 5.94E20 SO3 molecules
B.) 2.8E22 H2O molecules
C.) 1 glucose molecule (C6H12O6)

a.) 0.0790 g b.) 0.84 g c.) 2.922E-22 g

Calculate the mass percent composition of carbon in each carbon containing compound.
a.) CH4
b.) C2H6
c.) C2H2
d.) C2H5Cl

a.) 74.87% b.) 79.88% c.) 92.24% d.)37.23%

Copper (II) fluoride contains 37.42% F by mass. Calculate the mass of fluorine in grams contained in 55.5 g of copper (II) fluoride.

20.8 g F

Calculate the empirical formula for each compound.
a.) 1.651 g Ag, 0.1224 g O
b.) 0.672 g Co, 0.569 g As, 0.486 g O
c.) 1.443 g Se, 5.841 g Br

a.) Aag2O b.) Co3As2O8 c.) SeBr4

Classify each compound as organic or inorganic.
a.) CaCO3
b.) C4H8
c.) C4H6O6
d.) LiF

a.) inorganic b.) organic c.) organic d.) inorganic

Determine whether each compound is a hydrocarbon.
a.)H3C—CH2OH
b.)H3C—CH3
O
l l
c.) H3C—C—CH2—CH3
d.)H3C—NH2

a.) functionalized hydrocarbon, alcohol b.) hydrocarbon c.) functionalized hydrocarbon, ketone d.) functionalized hydrocarbon, amine

Rules for writing an organic compound

placing C first in the formula, followed by H, followed by the remaining symbols in alphabetical order.

Rules for writing a molecular inorganic compound

placing elements of groups 13 to 15 first (in that order), followed by the rest of the symbols starting with those furthest to the left in the periodic table. Elements in the same column are listed alphabetically.