Chpt. 9, Covalent Compounds

Question 1

covalent compound

Answer 1

a compound in which two negatively charged atoms share their electrons in order to be more like the nearest noble gas

Question 2

binary compound

Answer 2

a compound that contains only two elements

Question 3

naming binary compounds

Answer 3

Basically to name these compounds, all you have to do is name the first element, name the second element with “-ide” at the end, and then tell people how many of each atom you have.

For example, let’s consider the case of P2O3. The name of this compound is “diphosphorous trioxide”, or literally “two phosphorus atoms and three oxygen atoms.”

Question 4

binary compound prefixes:
1
2
3
4
5
6
7
8
9
10

Answer 4

mono (only for O, only when it is the 2nd element)
di
tri
tetra
penta
hexa
hepta
octa
nona
deca

Question 5

determining the naming system

Answer 5

If the compound starts either with a metal ion or NH4, name it like an ionic compound.

If the compound contains only nonmetals and doesn’t start with H, name it like a covalent compound.

If the compound starts with H, name it like an acid.

Question 6

acids

Answer 6

These are compounds that give off H+ ions in water. Though there are some exceptions to this rule, the formulas of acids usually start with “H”, which makes them easy to identify.

Question 7

acids without oxygen

Answer 7

If it doesn’t contain oxygen, the name of the acid is “hydro[something]ic acid.” For example, HF is “hydrofluoric acid”, HCN is “hydrocyanic acid”, H2Se is “hydroselenic acid” and so forth.

Question 8

acids containing oxygen

Answer 8

If the acid contains oxygen, it’s named “[something][suffix] acid.” Because polyatomic ions are what give these acids their oxygen, we need to know the names of these ions. If the name of the ion ends with “-ate”, the suffix is “-ic” and if the name of the ion ends with “-ite”, the suffix is “- ous.” Thus, H2SO3, which contains the sulfite ion, is called “sulfurous acid.” H3PO4, which contains the phosphate ion, is named “phosphoric acid.”

Question 9

drawing a Lewis structure

Answer 9

1. Determine the number of valence electrons that are present in the compound.
2. Determine the number of “octet electrons” in the molecule.
3. Determine the number of bonding electrons in the compound.
4. Determine the number of bonds in the compound.
5. Draw the compound.
6. Add dots to represent “lone pair” or “unbonded pair” electrons.
7. For polyatomic ions, you need to figure out where the charged atoms are.

Question 10

Lewis step 1 example

Answer 10

In NH3, we have a total of eight valence electrons – five for nitrogen and one for each of the three hydrogen atoms.

Question 11

Lewis step 2 rules (octet electrons)

Answer 11

H = 2 oe-
Be = 2 oe-
B = 6 oe- (or, if it’s in a polyatomic ion, 8)
all other elements = 8 oe-

Question 12

Lewis #3, finding the bonding electrons

Answer 12

This is done by subtracting the number of valence electrons from the number of octet electrons. In our example, this is 14-8 = 6 bonding electrons.

Question 13

Lewis #4

Answer 13

This is done by dividing the number above by two. In our example, 6/2 = 3 bonds.

Question 14

guidelines for the bonding of atoms

Answer 14

a. Hydrogen and the halogens always bond once.
b. Oxygen’s family and beryllium bond twice (unless it’s a polyatomic ion, in which case they may bond once or twice).
c. Nitrogen’s family and boron bond three times (unless it’s a polyatomic ion, in which case they may bond three or four times.
d. Carbon’s family bonds four times.

Question 15

Lewis #6

Answer 15

These are electrons which aren’t involved in bonding, but are present on an atom so that it might have a full octet of valence electrons. Basically, this is done by just adding pairs of dots representing electrons until each element has the number of octet electrons we determined they needed earlier.

Question 16

Lewis #7

Answer 16

To do this, determine how many electrons each of them has around it (two per lone pair, one for each bond) and compare that to the number of valence electrons each wants.

If the number of electrons around the atom is greater than the number of valence electrons, the atom is negatively charged. If it’s less, then it’s positively charged.

Question 17

resonance structures

Answer 17

Resonance structures are a set of different Lewis structures, each of which show one of the ways that a given molecule can be arranged. When all these version are combined, the result is the true structure of the molecule.

Question 18

expanded octets

Answer 18

Every so often, we run into a chemical compound with a Lewis structure that can’t be solved using the method above. That’s most likely because the central atom has an “expanded octet”, meaning that it has more than 8 valence electrons on it. This is caused because d-orbitals get involved in the bonding, making following the octet rule more difficult. This can be dealt with by doing the following:
c. Surround each of the outer atoms with pairs of electrons until they each have eight electrons.

a. Determine the number of valence electrons in the compound.
b. Draw the molecule so that the central atom (usually the first thing in the formula) is bonded once to each of the other atoms.
c. Surround each of the outer atoms with pairs of electrons until they each have eight electrons.
d. Add the remaining electrons to the central atom in pairs until you run out.

Question 19

VSEPR theory

Answer 19

valence shell electron pair repulsion theory:

the shapes of molecules are determined by how the outer electrons repel one another and try to move away from each other as much as possible

this is important to take into consideration when drawing Lewis diagrams; electrons should always be positioned as far apart from one another as possible

Question 20

bond angle

Answer 20

1 or 2 = linear or straight (if it’s straight)

2 = bent (if it’s bent)

3 = bent

4 = tetrahedral

For every e- on the side, subtract 2 from the bond angle.

Question 21

polarity

Answer 21

According to the octet rule, elements on the left side of the periodic table (metals) have low electronegativities while those on the right (nonmetals) have high electronegativities.

Because nonmetals and metals have very dissimilar electronegativities, there is a transfer of electrons to form an ionic compound. Because two nonmetals will have similar electronegativities, there’s a sharing of electrons to form a covalent compound.

But what happens when two nonmetals have similar, but not identical electronegativities? The more electronegative atom will pull harder on the electrons than the less electronegative atom, causing an imbalance in the distribution of electrons.

This, in turn, leads to a partial negative charge on the more electronegative atom because it has more than its share of electrons, and a partial negative charge on the less electronegative atom because some of its share of electrons has been pulled away from it.

Question 22

polar covalent bond

Answer 22

covalent bonds that exist between two atoms with dissimilar polarities

Question 23

polar covalent molecule

Answer 23

any molecule in which the electrons are unevenly distributed; this occurs in any asymmetric molecule

Question 24

polarity example

Answer 24

A covalent molecule can be nonpolar even while having polar covalent bonds. Use the example of carbon dioxide: though the O=C=O bonds are polar (oxygen is more electronegative than carbon), the symmetry of the molecule makes it nonpolar overall.

Question 25

covalent compound properties

Answer 25

• lower MP and BP than ionic (they are not bonded in a rigid position, so the move around more easily)
• less hard and brittle than ionic (they have more latitude to move around when struck with a hard force)
• more likely to burn than ionic compounds (C and H are required in order for fire to occur, and they are much more prevalent in covalent compounds)
• generally don’t conduct electricity; almost uniform insulators (electricity can be conducted by either electrons or moving ions, and in a covalent compound there are no ions present, and the electrons are not free to move around)