Chemistry 11 Test 2 Review Solutions
1.
What is the difference when forming an ionic and a covalent bond?
An ionic bond is formed from transferring electrons from the metal atom to the nonmetal atom. A covalent bond is formed by electrons sharing.
2.
Use electron dot symbols (Lewis structure) to determine the formula unit and then name the ionic compounds formed when the following elements combine. a) potassium and chlorine b) calcium and chlorine c) aluminum and oxygen d) beryllium and phosphorus
3.
Write the electron dot structures for each of the following elements. a.
Na b. Be c. Ge d. Be
4.
What are cations and anions? How and why are cations and anions produced? e. F
Cations are positively charged ions while anions are negatively charged ions. Cations are produced when an atom gives away its valence electrons making it have more of a positive charge. Anions are produced when an atom accepts electrons making have more of a negative charge.
5.
Write the correct chemical formula for the compound formed by the following pairs of ions: a. Na + and F NaF b. K + and S 2K
2
S c. Ca 2+ and N 3Ca
3
N
2
d. Al 3+ and O 2- Al
2
O
3
6.
Which of the following pairs of elements are most likely to form ionic compounds? a. chlorine and bromine b. potassium and helium c. lithium and fluorine d. iodine and sodium

7.
Write the formula for the ions in the following compounds. a. LiF Li + and F b. BaO Ba 2+ and O 2d. Ga
2
S
3
Ga 3+ and S 2e. Ca
3
N
2
Ca 2+ and N 3-
8.
Draw the electron dot structure for the following covalent molecules: a) H
2
S b) PH
3 c) BrF c. Na
2
S Na + and S 2-
9.
What is a coordinate covalent bond? Provide an example.
A coordinate covalent bond is when an atom donates two electrons to the shared pair of electrons.
An example is in the compound NH
4
+
10.
What is a resonance structure? Provide an example.
A resonance structure is a diagram that shows different placements of the electrons around the central atom. An example is SO
3
11.
Name or write the formula for the given compounds. a) MgS Magnesium sulfide r) magnesium oxide MgO b) KBr Potassium bromide s) potassium phosphate
K
3
PO
4 c) As
4
O
10 d) Al
2
O
3
Tetrarsenic decoxide
Aluminum oxide t) calcium nitride u) tin(IV) sulfate
Ca
3
N
2
Sn(SO
4
)
2 e) NH
4
NO
3 f) SrF
2 g) BrO
3 h) Ni(CH i) PbBr
4
3
COO)
2
Ammonium nitrate
Strontium fluoride
Bromine trioxide nickel (II) acetate lead (IV) bromide v) potassium iodide KI w) phosphoric acid H
3
PO
4 x) sodium sulfide Na
2
S y) tin(II) carbonate SnCO
3 z) hydrobromic acid HBr(aq)

j) H
2
CO
3(aq) carbonic acid aa) lead(IV) iodide PbI
4 k) Li2S l) BN m) Ba n) H
2
3
N
2
SO
3(aq) lithium sulfide boron mononitride barium nitride sulfurous acid bb) boron trifluoride cc) chlorine monoxide dd) oxygen difluoride ee) dinitrogen monoxide
BF
3
ClO
OF
N
2
2
O o) H
3 p) Pb
3
PO
N
2
3(aq)
Yieks! This polyatomic ion is not on your table! lead (II) nitride ff) hydrochloric acid gg) copper(I) sulfide
HCl (aq)
Cu
2
S q) HNO
3(aq) nitric acid hh) nitrous acid HNO
2(aq)
12.
Which of the bonds listed below have a high melting and boiling point? Explain why for each. a) Ionic Bond b) Covalent Bond c) Metallic Bond d) Network Solid
The substances with high melting points and high boiling points are network solids, ionic bonds and metallic bonds. Network solids are large covalently bonded structures that are intermingled. It will require a lot of energy to break the different molecules apart for the network solids.
Ionic bonds are very strong attractive forces. The ions involved are strongly attracted to one another and therefore will require a lot of energy to separate them from a solid state to a liquid state.
Metals have high melting points and boiling points because they are not easily separated from one another.
13.
Which of the bonds listed below are brittle? Explain why. a) Ionic Bond b) Covalent Bond c) Metallic Bond d) Network Solid
Ionic substances are brittle because they are made of oppositely charged particles. When the like charges are forced to come close to each other they repel and break apart the substance.
14.
Which of the bonds listed below are malleable? Explain why. a) Ionic Bond b) Covalent Bond c) Metallic Bond d) Network Solid
Covalent bonds and metallic bonds are malleable. Covalent molecules can move easily among each other and they do not repel each other so they are considered malleable. When metals bond they

create a sea of electrons. When the atoms are pushed around they also do not repel each other and will be malleable.
15.
What does it mean to conduct electricity?
When substance conducts electricity it can carry a charge. This means that electrons will easily be moved around.
16.
Which types of compounds are able to conduct electricity? Explain the reasons for your choices.
The elements that can conduct electricity are metallic substances, ionic substances when they are dissolved in water or melted. Graphite can also conduct electricity because there are electrons that are free to move.
17.
What is an allotrope? Give an example.
An allotrope when elements can exist in two or more different forms. Carbon can bond to other carbons to create diamond and also graphite.
18.
How does the structure of Buckminster Fullerene differ from that of graphite?
Buckminster Fullerene has a pentagonal and hexagonal arrangement of carbon atoms with only single bonds whereas graphite has carbon that has two single bonds and one double bond.
19.
What is an alloy?
An alloy is a mixture of different metals atoms bonded.
20.
How is the lattice of network solids different from the ionic lattice?
Network solids are made of large covalent molecules together and an ionic lattice is made of oppositely charged ions mixed together.
21. Draw the electron dot structure for the following covalent molecules or ions. Use VSEPR theory to state their shape. a) H
2
S b) PH
3 c) BrF d) NH
4
+ e) CO
3
2- f)HCN g) I
2
h)H
2
S i) OF
2
j) NI
3
(a) bent (b) pyramidal (c) linear (d) tetrahedral (e) trigonal planar
(g) linear (h)bent (i)bent (i) pyramidal (f) linear
22. Explain how the VSEPR theory can be used to predict the shapes of molecules.
VSEPR Theory - Valence Shell Electron Pair Repulsion
VSEPR theory takes into consideration that electrons are of the same charge and will therefore repel each other. The shapes that are predicted from these repulsions are dependent on the number of electron groups surrounding the central atom, and the number of lone pairs and bonded pairs on the central atom.
23. Explain the difference between a nonpolar covalent bond and a polar covalent bond .
Polar covalent bond – atoms that share electrons to create a covalent bond, but the electrons are not shared equally. One atom has a stronger attraction for the shared pair, therefore making that end of the bond slightly more negative δ- and the other end of the bond slightly more positive δ+.
Non-polar covalent bond (pure covalent) – atoms that share electrons creating a covalent bond but the atoms have the same desire for these electrons (similar electronegativity), so the they are shared equally creating a pure covalent bond.

24. Use electronegative differences and types of atoms to identify the types of bonds between atoms in the following pairs of elements (ionic, polar covalent or pure covalent). a) SF
2 b) O
2 c) CaO d) SiO
2 e) BrCl f ) LiBr g) Ca
3
N
2 h) Br
2 i) H
2
Se
(a) S-F polar covalent
(i) H – Se polar covalent
(e) Br – Cl polar covalent
(b) O –O pure covalent
(f) LiBr ionic
(c) CaO ionic (d) Si – O polar covalent
(g) CaN ionic (h) Br-Br pure covalent
25. Which covalent bond is the most polar and which is the least polar? Explain what this means. a) H – Cl b) H – Br c) H – S d) H – C e) F – F
The least polar bond is the F – F bond. They will have the same electronegativity value, meaning that both atoms want the shared pair of electrons the same. There will be no overall net dipole (no slight charges)
The most polar pair is H – Cl. This can be identified from the largest electonegativity difference. This means that the chlorine atom wants the shared pair of electrons more than the hydrogen, they will be pulled slightly towards the chlorine creating a slight negative charge. The slight negative charge in the
H – Cl will be larger than other slight charges in the other bonds.
26. Based on molecular shapes and polarity of bonds, which of the following molecules are polar molecules? a) SO
2 b) H
2
S c) SBr
2 d) HCl
(a) bent, polar bonds, polar molecule (b) bent, polar bonds, polar molecule
(c) bent, polar bonds, polar molecule (d) linear, polar bonds, polar molecule
27. Using the elements Cl, Mg, F, H, and K, construct as many examples (give formulas) as possible of substances with pure covalent bonding, polar covalent bonding, and ionic bonding.
Pure Covalent Polar Covalent Ionic
H
2
, Cl
2,
F
2
HF, HCl, FCl Mg(Cl)
2
, MgF
2
, KF, KCl
28. Which intermolecular force of attraction is the strongest? The weakest? What does this mean?
The strongest intermolecular force of attraction is the attraction of one very polar molecule to another
(hydrogen bonding)).
The weakest intermolecular force of attraction is the London dispersion forces of attraction.
This means that the particles that undergo the London dispersion forces are not very attracted to one another; they only create temporary slight charges. It will not require very much energy to break apart these non-polar molecules. Ionic bonding is formed from the transfer of electrons; therefore these are very strong bonds and require a lot of energy to separate the particles.
29. Assuming all things are equal, arrange the following from the lowest predicted boiling point to the highest predicted boiling point a) NCl
2
H
(a) Hydrogen bonding (b) ionic
b) NaCl
(c) dipole dipole
Lowest BP  Highest BP Br
2
, Br
2
S, NCl
2
H, NaCl c) Br
2
S d)Br
2
(d) London dispersion forces

30. Using PCl
3
as an example, draw and label diagrams to show the difference between intramolecular bonds and intermolecular bonds.
.
Intramolecular force  -------
Intermolecular force
31. Which bonds are harder to break, intramolecular bonds or intermolecular bonds? Explain.
The bonds that are more difficult to break are the bonds created by the transfer of electrons or electron sharing.
These are known as intramolecular bonds and are stronger than intermolecular bonds.. Intermolecular bonds are formed from the slight attraction from one molecule to another. These are not as strong; they are easier to separate the particles and therefore will require less energy to break the intermolecular bond.
32. Draw the Lewis structure for F- I and indicate which end of the molecule is
 
and which end is
 
.
Explain what this means.
δ- F – I δ+
When I and F share electrons to complete their octet the fluorine has a much stronger attraction for those shared electrons (larger electronegativity), so it will pull them toward its nucleus, making the fluorine atom slightly negative. The iodine atom will have partially lost its lone electron, so it will become slightly positive.
33. Why do ionic compounds conduct an electrical current in the liquid state?
Ionic compounds will conduct an electric current when in the liquid state because the ions will be free to move around (no longer in the crystal lattice structure) and carry the electrons and complete the circuit.
34. Why do covalent compounds usually have low melting points?
Covalent compounds usually have low melting points because their intermolecular forces of attraction are weak in comparison to ionic compounds. The molecules do not have a strong attraction to one another, so they will not require a lot of energy to break them apart (low MP )
35. Covalent compounds do not conduct an electrical current in the liquid state, why?
To complete an electrical current in a liquid state a mobile charge must be present. Covalent compounds may only have slight charges from polar molecules but they will not be able to transfer the electrons to complete the circuit.
36. Why are metals excellent conductors of electricity and heat?
Metals are excellent conductors of electricity because they have electrons that are delocalized and free to move around the metallic atoms to complete the circuit.
Metals are also excellent conductors of heat because the particles can gain the energy and the vibrations of the particles are easily passed along from one atom to another.
37. Why don't nonmetals form metallic bonds?
Non-metals do not form metallic bonds because to form metallic bonds the atoms must give up their electrons for sharing. Non-metals have a tendency to gain electrons, not give them up easily .
38. A substance was analysed and had the following properties: a) conducted electricity in the solid state

b) solid at room temperature c) ductile
Was the substance an ionic compound, a molecular compound, or a metal? Explain how you can tell.
The unknown substance was a metal. You can tell because it conducted electricity in a solid state. This indicates the charge to complete the circuit was free to pass along – this is only a trait of metallic bon ding.