Chpt. 9, Covalent Compounds Flashcards Preview

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Flashcards in Chpt. 9, Covalent Compounds Deck (25):
1

covalent compound

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

2

binary compound

a compound that contains only two elements

3

naming binary compounds

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.”

4

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

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

5

determining the naming system

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.

6

acids

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.

7

acids without oxygen

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.

8

acids containing oxygen

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.”

9

drawing a Lewis structure

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.

10

Lewis step 1 example

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

11

Lewis step 2 rules (octet electrons)

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

12

Lewis #3, finding the bonding electrons

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.

13

Lewis #4

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

14

guidelines for the bonding of atoms

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.

15

Lewis #6

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.

16

Lewis #7

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.

17

resonance structures

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.

18

expanded octets

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.

19

VSEPR theory

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

20

bond angle

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.

21

polarity

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.

22

polar covalent bond

covalent bonds that exist between two atoms with dissimilar polarities

23

polar covalent molecule

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

24

polarity example

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.

25

covalent compound properties

-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)