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Unit 6: Ionic Compounds

We discussed compounds in an earlier unit, but to revisit the topic:

A chemical compound is an example of matter that is made of two or more different elements that are chemically bonded together.

Something else we need to know about compounds:

They are put together with elements in a fixed proportion – that means that water, for example, always has two hydrogen atoms and one oxygen atom; table sugar (sucrose) always has twelve carbon atoms, twenty-two hydrogen atoms, and 11 oxygen atoms, potassium chloride (found in lite salt products) always has one potassium atom and one chlorine atom.

The “glue” that holds compounds together is called a bond. There are two different kinds of bonds, ionic (also known as electrovalent) and covalent. For this unit, we will be working with ionic (or electrovalent) bonds. Those are the bonds that hold ions together to form compounds. We will get to covalent bonds when we discuss a little biochemistry in a later unit. Covalent bonds deal with the sharing of electrons rather than ionic charges to form compounds.

Ions are “un-neutral” atoms; atoms with either more protons than electrons or more electrons than protons. They carry a charge, either positive or negative.

To figure out the charges of ions, we have to understand the Rule of Octets. This is not the rule of octaves that our friend John Newlands came up with, this is a different concept. It says that elements, except for hydrogen and helium, want 8 electrons in their outer shells. Hydrogen and helium are exceptions because all of their electrons can be contained in the K shell (1st energy level) and it only holds 2 electrons, right?

Here’s how it works:

Let’s say we are looking at sodium. Take a look on your Periodic Table. What is sodium’s atomic number? 11 – so as a neutral atom, sodium has 11 protons and 11 electrons. Visualize that balance this way:

Protons: + + + + + + + + + + +

There would be 2 electrons in the K shell, 8 electrons in the L shell, and 1 electron in the M shell

According to the Rule of Octets, sodium has two choices to get 8 electrons in its outer shell.

  1. Grab on to 7 more electrons so that it would have 8 in the M shell, or
  2. Lose the 1 electron in the M shell so that the L shell becomes its outer shell with 8 electrons

What do you think would be easier for sodium? If you said lose 1 you are correct; it’s easier for sodium to lose one electron than to gain 7 more electrons.

Protons: + + + + + + + + + + +

After it loses one electron, it has one more proton than electron. It’s charge as an ion is +1.

(Check that out on the ion chart)

Common Ions and their Charges

Position Ions (Cations)

Aluminum: Al

Ammonium: NH+


Barium: Ba

Cadmium: Cd

Calcium: Ca

Chromium (II): Cr
2+ a.k.a Chromous

Chromium (III): Cr
3+ a.k.a. Chromic

Cobalt: Co

Copper (I): Cu
1+ a.k.a. Cuprous

Copper (II): Cu
2+ a.k.a. Cupric

Hydrogen: H

Iron (II): Fe
2+ a.k.a. Ferrous

Iron (III): Fe
3+ a.k.a. Ferric

Lead: Pb

Lithium: Li

Magnesium: Mg

Manganese: Mn

Nickel: Ni

Potassium: K

Silver: Ag

Sodium: Na

Zinc: Zn

Negative Ions (Anions)

Acetate: CH
or C


Bromide: Br

Carbonate: CO2-


Bicarbonate: HCO-

3 a.k.a. Hydrogen carbonate

Chlorate: ClO-


Chloride: Cl

Chlorite: ClO-


Chromate: CrO2-


Cyanide: CN

Dichromate: Cr


Fluoride: F

Hydroxide: OH

Hypochlorite: ClO

Iodide: I

Nitrate: NO-


Nitrite: NO-


Oxalate: C


Oxide O

Perchlorate: ClO-


Permanganate: MnO-


Phosphate: PO3-


Sulfate: SO2-


Sulfide: S

Sulfite: SO2-


Now let’s try chlorine. Look on the Periodic Table. What’s its atomic number? 17. So as a neutral atom chlorine has 17 protons and 17 electrons. 2 in the K shell, 8 in the L shell, and 7 in the M shell

What’s it probably going to do: pick up 1 electron in the M shell or drop 7 so that the L shell becomes it outer shell?

If you say pick up 1 electron in the M shell, you are correct; it’s easier for chlorine to pick up one electron than to drop 7 electrons.

So if it has 17 protons and 17 electrons as a neutral atom and it picks up an extra electron what will be its charge as an ion? – 1.

By the way, the ions of the halogens will all have an -ide ending (chloride, bromide, fluoride, iodide)

(check it out on the ion chart to be sure of the ending and the charge)

The fun begins! We’re going to make compounds from ions! First, there are a few things we have to understand about compounds:

  1. Compounds have to be neutral
    • They have to have the same number of positive charges as negative charges
    • Their net charge has to equal zero
  2. The positive ion always goes first in the compound and the negative ion always goes last
  3. The “glue” that holds these ions together is called an ionic bond or electrovalent bond. The ions are held together in the compound by a strong electrostatic attraction between the positive and the negative ions.

So here we go:

Example 1:

Let’s take the two ions we figured out earlier, Na+ and Cl, to figure out the formula for sodium chloride.

Whenever the charges of the ions automatically adds up to zero (net charge of zero), all you have to do is put those symbols together and you’re done!

That’s the case with the sodium ion and the chloride ion. +1 and -1 add up to 0

Put those symbols together, my friend! What do you have: NaCl

Example 2:

How about the formula for barium sulfide? Ba2+ and S2-

+2 and -2 adds up to 0, put ’em together!

BaS is the correct formula

OK, that’s easy enough, but what about when the charges don’t add up to zero right off the bat?

In that case you can approach the formula with a couple of different strategies:

Example 3:

What is the compound formula for aluminum chloride?

Take a look on the chart. aluminum’s charge is +3

We’ve already seen that chloride’s charge is -1

+3 and -1 = +2. The net charge is +2 not 0. It wouldn’t work to simply put the two symbols together.

One way to look at this is to say: How many chlorides (-1) is it going to take to counter the one aluminum (+3)? It would take 3 chlorides (3 x -1 = -3) in the compound formula to add up to the one aluminum’s charge of +3. So, +3 – 3 = 0 (net charge = 0). Since you only need 1 aluminum atom in the compound and 3 chlorides, you would write the formula for aluminum chloride:


Please note: in a compound formula, if there is not a subscript besides an element’s symbol, that means there’s just one of them (ex: for water the formula is H2O, not H2O1)

Another way to figure out the compound formula when the ions’ charges don’t add up to zero at first is to do the old cris-cross method.

Aluminum’s charge is +3 and chloride’s is -1. That could be written:

Al+3 Cl-1

Take the 3 superscript (just the number, not the charge, and cris-cross it so that it becomes the subscript of the other guy, Cl in this case). In other words, you need 3 chlorides

Take the 1 superscript and cris-cross it so that it goes with the Al (remember, since it’s just 1 you don’t need to write it as a subscipt). In other words, you just need one aluminum

The formula for aluminum chloride:


The cris-cross method works every time the charges don’t add up to 0 initially!

Polyatomic ions

Polyatomic ions are ions that have more than one element in them. There are a bunch on the negative (anion) side of the ion chart and a few on the positive (cation) side. Take a look: acetate, carbonate, bicarbonate, chlorate, chlorite, chromate, and on and on. If it has more than one element in it and it had a charge, it’s a polyatomic ion. In a very important way, polyatomic ions work just like regular ions. If the charge of the cation and the charge of the anion = 0 immediately, just put those symbols together!

Example 1:

Sodium acetate: Na+, CH3COO
+1 and -1 = 0 (net charge is 0)
Put ’em together!

The compound formula for sodium acetate:


Example 2:

Ammonium chloride: NH+
4 and Cl
+1 and -1 = 0
Put’em together!

The compound formula for ammonium chloride:


Here’s where it gets tricky, so pay close attention to this part!

Let’s say your charges don’t add to zero right off the bat and it turns out you are going to need more than one of a polyatomic ions unit in the final compound formula. There’s a particular way to show that!

Example 3:

Barium hydroxide: Ba2+ and OH
Note that hydroxide is polyatomic because it is made of both oxygen and hydrogen!

Since there is not a net charge of zero initially, we have to do the cris-cross. That means that we will need 2 hydroxides (we’re cris-crossing the 2 from the barium to the hydroxide) and we will need 1 barium (we’re cris-crossing the 1 from the hydroxide). But we can’t just write it: BaOH2 that’s wrong! That says that there is one barium, one oxygen, and two hydrogens and that’s not what we want to say. We have to show that we need two hydroxide units and here’s how we do that:

We have to enclose the OH in parenthesis and then put the 2 subscript on the outside:

2 that’s right!

Example 4:

Ammonium sulfide: NH+
4 and S2-. Ammonium is polyatomic because it is made of nitrogen and hydrogen. There is not a net charge of zero initially so we have to do the cris>-cross. We are going to need 2 ammonium units (cris-cross the 2 from sulfide’s charge) and 1 sulfide (cris-cross the 1 from the ammonium charge).

Since ammonium is polyatomic and we need more than one ammonium unit in the compound formula, we have to use parenthesis around the ammonium and place the 2 subscript on the outside of the parenthesis.


Something else to deal with: some of the transition metals!

Take a look at the ion chart again. Look at Chromium, Copper, and Iron. Those rascals have more than one charge as ions! It all has to do with their being transition metals – that’s the way some transition metals can act.

But it’s not really a big deal.

Chromium can either have a +2 charge and it can be called Chromium (II) or Chromous. The Chromium (II) name is the more recent form, but the older Chromous is also still being used. The (II) indicates the +2 charge and the -ous ending indicates that it is the lower of the two possible charges.

Chromium (III) or Chromic is the other form and it has a +3 charge. (III) for +3 and -ic for the higher of the two charges.

Copper (I) or Cuprous has a +1 charge

Copper (II) or Cupric has a +2 charge

Iron (II) or Ferrous has a +2 charge

Iron (III) or Ferric has a +3 charge

Naming compounds

There are a lot of official rules to follow when you are naming compounds, but we’re going to keep it simple. Since you will always be able to use your ion chart any old time (tests included), I just want you to be able to work backwards from the chart and put the names of each ion together – then you will have the whole name for the compound. Most of the time it’s pretty darn straight forward – those transition metals that have more than one charge are the only ones you really have to watch out for because you do have to be specific about their particular form when you name a compound involving them.

Example 1: KOH

The K is potassium and the OH is hydroxide, put ’em together and what do you have:

Potassium hydroxide

Example 2: Al2(SO4)3

The Al is aluminum and the SO4 is sulfate, put ’em together and you have:

Aluminum Sulfate

So it’s pretty straight forward when you are dealing with just putting the ions together. When you are dealing with some of the transition metals (chromium, copper, and iron, for example) that can have more than one charge, you have to be specific about which form is in the compound.

Example 3: FeCl3

Sure, the Fe is iron and the Cl is chloride, but we can’t just call it iron chloride because we don’t know which form of iron it is. So in this example we’re going to do the cris-cross in reverse. That 3 subscript by the Cl can be cris-crossed back to the Fe to become a +3 charge. If it’s iron with a +3 charge it can either be called:

Iron(III) or Ferric – both naming systems are in play currently – so our compound can be called Iron(III) chloride or Ferric chloride.

Example 4: Cu2O

Wait a minute, there’s not a subscript by the O. If there’s not a subscript, what’s the number? The number is 1. So drag that 1 back up to the Cu. The Cu has a +1 charge. Copper with a +1 charge can either be called:

Copper(I) or Cuprous

So our compound can be called:

Copper(I) oxide or Cuprous oxide

Example 5: FeSO4

What the heck is going on with this one? SO4 is sulfate. Looks like the Fe and the SO4 have just been put together. When do you do that? When the charges add up to zero right off the bat. So if SO4 has a -2 charge (that’s what the chart tells us), then the Fe would have to have a +2 charge. And Fe with a +2 charge is called:

Iron(II) or Ferrous so our compound can be called:

Iron(II) sulfate or Ferrous sulfate

Here is a link to a pretty good series of YouTube videos concerning ionic compounds you may want to check out. This link will take you to the first video (it’s only about 10 minutes) and from there you can go to the others in the series if you choose.


Order the answer to view it

Assignment Solutions

Assignment Solutions