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Archive for the ‘Chemistry’ Category

What sections should I know before attempting to learn this section?

—> Covalent, Ionic, and Metallic Bonds

—> Forming Ionic Compounds

—> Breaking Apart Ionic Compounds
 

How do you use an ionic name to create the ionic compound?

Going in the other direction (taking the name and making the ionic compound), however, is a littler harder. First, you have to find the charges of the elements, and then you have to get them to total up to zero. How many of each element you need depends on how many each take to get to zero. In math, they call this concept the least common multiple.

 

Examples: Give the ionic compound from the name using a regular periodic table. Don’t forget to use the ion periodic table if you need it.  VIDEO Ionic Naming Examples 2.

Lithium Phosphide	Li3P
Radium Iodide	RaI2

 

VIDEO Ionic Naming Demonstrated Example 2: Give the ionic compound from the name using a regular periodic table. Don’t forget to use the ion periodic table if you need it.

 

Sodium Chloride

 

Step 1:

What is the symbol for Chloride?

Answer: Cl

		Cl
1
1
Total =

 

Step 2:

What is the ion of Chloride?

Answer: -1

		Cl–
1
1
Total =

 

Step 3:

What is the symbol for Sodium?

Answer: Na

	Na	Cl–
1
1
Total =

 

Step 4:

What is the ion of Sodium?

Answer: +1

	Na+	Cl–
1
1
Total =

 

Step 5:

What is the least common denominator between 1 and 1?

Answer: 1

	Na+	Cl–
1
1
Total =	+1	-1

 

Step 6:

What is the complete chemical formula?

COMPLETE ANSWER: NaCl

 

VIDEO Ionic Naming Demonstrated Example 3: Give the ionic compound from the name using a regular periodic table. Don’t forget to use the ion periodic table if you need it.

 

Calcium Phosphide

 

Step 1:

What is the symbol for phosphorous?

Answer: P

		P
1
1
Total =

 

Step 2:

What is the charge for phosphorous?

Answer: -3

		P3-
1
1
Total =

 

Step 3:

What is the symbol for calcium?

Answer: Ca

	Ca	P3-
1
1
Total =

 

Step 4:

What is the charge for calcium?

Answer: +2

	Ca2+	P3-
1
1
Total =

 

Step 5:

What is the least common denominator between 2 and 3?

Answer: 6

	Ca2+	P3-
1
1
Total =	+6	-6

 

Step 6:

How many phosphorous do you need to add up to a total of -6?

Answer: You need 2 phosphorus

	Ca2+	P3-
1		P3-
1
Total =	+6	-6

 

Step 7:

How many calcium do you need to add up to a total of +6?

Answer: You need 3 Ca

	Ca2+	P3-
	Ca2+	P3-
	Ca2+
Total =	+6	-6

 

Step 8:

What is the complete chemical formula?

COMPLETE ANSWER: Ca₃P₂

 

PRACTICE PROBLEMS: Give the ionic compound from the name using a regular periodic table. Don’t forget to use the ion periodic table if you need it.

Lithium Chloride	LiCl
Calcium Iodide	CaI2
Beryllium Oxide	BeO
Strontium Phosphide	St3P2
Potassium Silicide	K4Si
Sodium Sulfide	Na2S

 

What sections should I know before attempting to learn this section?

—> Covalent, Ionic, and Metallic Bonds

—> Forming Ionic Compounds

—> Breaking Apart Ionic Compounds

 

How do you name ionic compounds?

For the ionic system of naming, it does not depend on how many of each element are in the compound. The key for the ionic system of naming is what the ion of each element is. The ions can be found on this periodic table. Any ionic compound should add up to a charge of zero unless otherwise stated.

 

Examples: Name the ionic compound using a regular periodic table. Don’t forget to use the ion periodic table if you need it.
VIDEO Ionic Naming Examples 1.

LiF	Lithium Fluoride
MgBr2	Magnesium Bromide
Al2O3	Aluminum Oxide
K2S	Potassium Sulfide

 

VIDEO Ionic Naming Demonstrated Example 1: Name the ionic compound using a regular periodic table. Don’t forget to use the ion periodic table if you need it.

 

BaF₂

 

Step 1:

What is the name for the symbol Ba?

Answer: Barium

 

Step 2:

What is the name for the symbol F?

Answer: Fluorine

 

Step 3:

How do you modify the ending of Fluorine?

Answer: Becomes fluoride

 

Step 4:

What is the complete chemical name?

COMPLETE ANSWER: Barium Fluoride

 

PRACTICE PROBLEMS: Name the ionic compound using a regular periodic table. Don’t forget to use the ion periodic table if you need it.

NaI	Sodium Iodide
BeF2	Beryllium Fluoride
Ba3P2	Barium Phosphide
Rb3N	Rubidium Nitride
K2O	Potassium Oxide
CaCl2	Calcium Chloride

 

 

What sections should I know before attempting to learn this section?

—> Covalent, Ionic, and Metallic Bonds

 

How do you use a covalent (molecular) name to create the covalent compound?

Now that we have taken the covalent chemical compound and produced the chemical name (as we did in the previous section) we should try the other direction. That is taking the chemical name and producing the chemical compound. Remember to use the covalent prefixes from the picture below.

Examples: Give the chemical formula from the name using a regular periodic table and the covalent prefixes. VIDEO Covalent Naming Examples 2.

 

Triboron Tetracarbide	B3C4
Nitrogen Trifluoride	NF3
Heptasulfur Diselenide	S7Se2
Bromine Monochloride	BrCl

 

VIDEO Covalent (Molecular) Naming Demonstrated Example 2: Give the chemical formula from the name using a regular periodic table and the covalent prefixes.

 

Pentaphosphorus Dibromide

 

Step 1:

What does the prefix “penta” mean?

Answer: 5

 

Step 2:

What is the symbol for phosphorus?

Answer: P

 

Step 3:

How much phosphorus do you have?

Answer: 5

 

Step 4:

What does “di” mean?

Answer: 2

 

Step 5:

What element does bromide refer to?

Answer: Bromine

 

Step 6:

What is the symbol for bromine?

Answer: Br

 

Step 7:

How much bromine do you have?

Answer: 2

 

Step 8:

What is the chemical formula?

COMPLETE ANSWER: P₅Br₂

 

PRACTICE PROBLEMS: Give the chemical formula from the name using a regular periodic table and the covalent prefixes.

Silicone Tetrahydride	SiH4
Diarsenic Hexasulfide	As2S6
Trinitrogen Heptaoxide	N3O7
Octaphosphorus Nonaxenide	P8Xe9

 

What sections should I know before attempting to learn this section?

—> Covalent, Ionic, and Metallic Bonds

 

How do you name covalent (molecular) compounds?

For the covalent system of naming, it depends on how many of each element there is in a compound. You describe the amount of each element by using a prefix word that indicates a number. The prefixes names for the system are as follows:

 

 

These prefixes are used in front of each elemental name in a compound. Sometimes the “a” on the end of the prefixes, like penta or octa, will be omitted. Don’t worry about it too much, right now.

 

Examples: Give the name of each compound using the covalent prefixes and a regular periodic table. VIDEO Covalent Naming Examples 1.

C2O4	Dicarbon Tetroxide
P3Br10	Triphosphorus Decabromide

 

One trick in this whole system has to do with the mono- prefix. Mono- is a special prefix that you do not always use. You have heard this exception before but may not have realized it. Take the compound CO₂. Think… what is the name of it? Right, carbon dioxide. So why does carbon not have a mono- prefix on it? The exception I was talking about earlier is just that. If the first element in the compound is only one of that element, you do not use the mono- prefix. Here’s an additional example to drive the point home. Another common compound you have heard of is carbon monoxide. So what does it look like? That is right: CO. It is a perfect example of how the first element, carbon, has no prefix, but the second has a prefix. This rule is only true for the mono- prefix, so do not use it for any other purposes in covalent naming.

 

VIDEO Covalent (Molecular) Naming Demonstrated Example 1: Give the name of each compound using a regular periodic table and the covalent prefixes.

 

ClF₂

 

Step 1:

What does Cl stand for?

Answer: Chlorine

 

Step 2:

How much chlorine do you have?

Answer: 1

 

Step 3:

What is the prefix for the chlorine?

Answer: It should be mono but since it is the first element you don’t write it.

 

Step 4:

What does F stand for?

Answer: Fluorine

 

Step 5:

How much fluorine do you have?

Answer: 2

 

Step 6:

What is the prefix for the fluorine?

Answer: Di

 

Step 7:

How do you modify the ending of fluorine?

Answer: Becomes fluoride

 

Step 8:

What is the complete chemical name?

COMPLETE ANSWER: Chlorine Difluoride

 

PRACTICE PROBLEMS: Give the name of each compound using a regular periodic table and the covalent prefixes.

O2F2	Dioxygen Difluoride
PCl3	Phosphorous Trichloride
N4S5	Tetranitrogen Pentasulfide
Se7I10	Heptaselenium Deciodide
H2O	Dihydrogen Monoxide
BBr3	Boron Tribromide

 

 

How do you change the ending of a chemical name?

The second part of the naming system that both covalent and ionic compounds have in common is that any binary compound (compound made up of two elements) has the ending -IDE in the name. This -IDE usually replaces the -INE suffix in most elemental names. If you take the element fluorine, it becomes fluoride when it is part of a compound. Oxygen can be strange. It becomes oxide.

 

Examples: I will supply you with the beginning of the name and you give me the ending of the name. The periodic table link may be useful.  VIDEO Changing Chemical Name Ending Examples 1.

KI	Potassium Iodide
CaF2	Calcium Flouride
Ba3N2	Barium Nitride
SO2	Sulfur Dioxide
Si2Br5	Disilicon Pentabromide

 

PRACTICE PROBLEMS: I will supply you with the beginning of the name and you give me the ending of the name. The periodic table link may be useful.

LiCl	Lithium Chloride
SrBr2	Strontium Bromide
Fr2Se	Francium Selenide
PO	Phosphorus Monoxide
KI	Potassium Iodide

Which element comes first in the compound or name?

Naming a chemical compound involves first identifying it as either a covalent or ionic compound. However, both covalent and ionic chemicals have a couple rules in common.

ONE is that both the chemical formulas and the chemical names (nomenclature) tend to start with the element furthest to the left on the periodic table.

TWO is that both can end their names with IDE. The IDE endings will be discussed and demonstrated in the next section changing ending of name to IDE.

Here we will show you how to predict the first element in any compound. For example, CsF is Cesium Fluoride. Another example is SiO₂ which is Silicon Dioxide. That means if we pick out two random elements and put them together we want them to be in the correct order.

 

Examples: Give the element name that would appear first when you go to name the compound regular periodic table.  VIDEO First Element in a Chemical Name Examples 1.

C and O	Carbon
Na and P	Sodium
Se and Ca	Calcium
F and Sn	Tin
Fe and I	Iron
As and V	Vanadium

 

PRACTICE PROBLEMS: Give the elemental name that would appear first when you go to name the compound using the regular periodic table.

Mg and Br	Magnesium
Si and Ni	Nickle
Li and Br	Lithium
Te and Au	Gold
F and Sr	Strontium

What is the lesson about?

Nomenclature means naming. So this lesson is all about naming different chemical compounds. There are different ways to name different combinations of elements so the rules can get confusing.

 

Why is it critical to understand?

That is because naming of compounds tends to be added to questions from other lessons. Nomenclature can also come in handy when we are buying food products or medicines. Most chemicals in food and drugs are labeled according to the nomenclature system here.

 

What you should know before attempting this lesson?

If you have trouble in this lesson go back to sections on Ions, Covalent Ionic and Metallic Bonds, Introduction to Polyatomic Ions, Identifying Polyatomic Ions in Compounds, Forming Ionic Compounds, and Breaking Apart Ionic Compounds.

 

New Learning Sections:

—> First Element in Chemical Name

—> Changing Ending of name to IDE

—> Naming Covalent Compounds Part 1

—> Naming Covalent Compounds Part 2

—> Naming Ionic Compounds Part 1

—> Naming Ionic Compounds Part 2

—> Ionic Compounds with Transition Metals Part 1

—> Ionic Compounds with Transition Metals Part 2

—> Ionic Compounds with Polyatomic Ions Part 1

—> Ionic Compounds with Polyatomic Ions Part 2

 

Other Nomenclature Areas:

I STRONGLY BELIEVE THAT ALL TEACHERS SHOULD NOT REQUIRE THESE OTHER AREAS TO BE COVERED ALONG SIDE THE ABOVE SECTIONS. However, some teachers insist you to learn them early in chemistry so I put the links below in case you needed them all in one place.  They are integrated sections in other lessons on this website.

LINKS NOT IN / PAGES NOT COMPLETE YET

—> Naming Acids

—> Naming Organic Compounds

 

References Pages:

—> Standard Periodic Table

—> Metal / Non-Metal Periodic Table

—> Covalent Prefixes

—> Ion Periodic Table

—> Polyatomic Ions List

 

Worksheets:

—> Nomenclature Worksheet 1

—> Nomenclature Worksheet 1 WITH ANSWERS

 

 

What sections should I know before attempting to learn this section?

—> Representation of Compounds and Molecules with Subscripts

—> Covalent, Ionic, and Metallic Bonds

—> Introduction to Polyatomic Ions Part 1

—> Forming Ionic Compounds

 

How do you break apart ionic compounds?

As we go forward into the future of learning chemistry it will become essential for us to be able to take ionic compounds and break them up into their component ions. It is a mindset that you have to develop as you go but right here we want to give you a boost or head start on it. All ionic compounds break up into positive ions (cations) and negative ions (anions). The positive ions (usually metals) are on the left side of the chemical formula and the negative ions (usually non-metals) are on the right side of the chemical formula.

 

Examples: Separate the CHEMICAL FORMULAS into their ionic components. Give the POSITIVE ion and the NEGATIVE ion and tell how many of each are in the CHEMICAL FORMULA. Use the ionic periodic table or the polyatomic ions list if needed. VIDEO Breaking Apart Ionic Compounds Examples 1.

Formula	Positive	Negative
LiF	1 Li1+	1 F1-
Al2O3	2 Al3+	3 O2-
HCN	1 H1+	CN1-
Mg(NO3)2	1 Mg2+	2 NO31-
Rb2Se	2 Rb1+	1 Se2-
NH4OH	1 NH41+	1 OH1-
Mn3N4	3 Mn4+	4 N3-

ADD ONE DEMONSTRATED EXAMPLE

 

PRACTICE PROBLEMS: Pick out the POLYATOMIC IONS from the chemical formula. There may be ONE, TWO, or NONE. Use the ionic periodic table or the polyatomic ions list if needed.

Formula	Positive	Negative
CsBr	1 Cs1+	1 Br1-
Be3N2	3 Be2+	2 N3-
KNO3	1 K1+	1 NO31-
Ca(OH)2	1 Ca2+	2 OH1-
Ba3(PO4)2	3 Ba2+	2 PO43-
Ti2S	2 Ti1+	1 S2-
CrSO4	1 Cr2+	1 SO42-

What sections should I know before attempting to learn this section?

—> Representation of Compounds and Molecules with Subscripts

—> Introduction to Polyatomic Ions Part 1

—> Covalent, Ionic, and Metallic Bonds

 

How do you form ionic compounds?

The formation of ionic compounds comes from the ion rules of the periodic table. The first and most important rule of forming compounds is that they have to add up to a charge of zero unless otherwise stated. If we put two ions together, one positive and one negative, then we have to add as many elements to each side until we get a sum of zero. In math they call this the smallest common multiple. You can usually do this in chemistry by multiplying together the two numbers you are dealing with.

 

Examples: Put these two elements together to form an ionic compound.

K and Br	KBr
Na and O	Na2O
Mg and N	Mg3N2

 

VIDEO Forming Ionic Compounds Demonstrated Example 1: Put these two elements together to form an ionic compound: Na and S

 

Step 1:

What is the charge on Na?

Answer: +1

	Na+
1
1
Total =

 

Step 2:

What is the charge on S?

Answer: -2

	Na+	S2-
1
1
Total =

 

Step 3:

What is the least common multiple between 1 and 2?

Answer: 2

	Na+	S2-
1
1
Total =	+2	-2

 

Step 4:

How much Na do you need to add up to +2?

Answer: 2 Na

	Na+	S2-
	Na+
1
Total =	+2	-2

 

Step 5:

How much S do you need to add up to -2?

Answer: 1 S

	Na+	S2-
	Na+
1
Total =	+2	-2

 

Step 6:

If you have 2 Na and 1 S what does the compound look like?

COMPLETE ANSWER: Na₂S

 

VIDEO Forming Ionic Compounds Demonstrated Example 2: Put these two elements together to form an ionic compound: Ba and P

 

Step 1:

What is the charge on Ba?

Answer: +2

	Ba2+
1
1
Total =

 

Step 2:

What is the charge on P?

Answer: -3

	Ba2+	P3-
1
1
Total =

 

Step 3:

What is the least common multiple between 2 and 3?

Answer: 6

	Ba2+	P3-
1
1
Total =	+6	-6

 

Step 4:

How much Ba do you need to add up to +6?

Answer: 3 Ba

	Ba2+	P3-
	Ba2+
	Ba2+
Total =	+6	-6

 

Step 5:

How much P do you need to add up to -6?

Answer: 2 P

	Ba2+	P3-
	Ba2+	P3-
	Ba2+
Total =	+6	-6

 

Step 6:

If you need 3 Ba and 2 P what does the compound look like?

COMPLETE ANSWER: Ba₃P₂

 

VIDEO Forming Ionic Compounds Demonstrated Example 3: Put these two ions together to form an ionic compound: Mg²⁺ and OH^–

 

Step 1:

What is the charge on Mg?

Answer: +2

	Mg2+
1
1
Total =

 

Step 2:

What is the charge on OH^–?

Answer: -3

	Mg2+	OH–
1
1
Total =

 

Step 3:

What is the least common multiple between 2 and 1?

Answer: 2

	Mg2+	OH–
1
1
Total =	+2	-2

 

Step 4:

How much Mg do you need to add up to +2?

Answer: 1

	Mg2+	OH–
1
1
Total =	+2	-2

 

Step 5:

How much OH do you need to add up to -2?

Answer: 2

	Mg2+	OH–
		OH–
1
Total =	+2	-2

 

Step 6:

If you need 1 Mg and 2 OH what does the compound look like?

COMPLETE ANSWER: Mg(OH)₂

 

PRACTICE PROBLEMS: Put these two elements or ions together to form an ionic compound. Don’t Forget to use the ion rules of the periodic table, if you need them.

Li and O	Li2O
Al and P	AlP
Cs and N	Cs3N
Ca and I	CaI2
Ra and As	Ra3As2
Mg and C	Mg2C
NH4+ and Te2-	(NH4)2Te
Ba2+ and CN–	Ba(CN)2
Be2+ and PO43-	Be3(PO4)2

 

How do you learn to find polyatomic ions in compounds?

What you want to start doing in this section is getting good at spotting polyatomic ions whenever and wherever you see them. Many of them will be hidden in the chemical formulas you will see in the future of chemistry. The better you become at identifying them, the better you will be at solving future chemistry problems.

 

Examples: Pick out the POLYATOMIC IONS from the chemical formula. There may be ONE, TWO, or NONE. Use the polyatomic ions list if needed.VIDEO Identifying Polyatomic Ions in Compounds Example 1.

Formula	Polyatomic Ions
Mg(NO3)2	NO31-
HCN	CN1-
CaBr2	NONE
NH4OH	NH41+   OH1-
(H3O)2CO3	H3O1+   CO32-

 

PRACTICE PROBLEMS: Pick out the POLYATOMIC IONS from the chemical formula. There may be ONE, TWO, or NONE. Use the polyatomic ions list if needed.

Formula	Polyatomic Ions
Fe3(PO4)2	PO43-
NaC2H3O2	C2H3O21-
NH4CN	NH41+   CN1-
K2S	NONE
BeSO4	SO42-

How do you Better and More Quickly Memorize Polyatomic Ion Charges?

VIDEO Explanation of Tips to Memorize Polyatomic Ions

For many of the polyatomic ions on the list even though their names change like ATE, ITE, PER, HYPO, the charge of their ion does not change.

 

Examples:

Sulfate	SO42-
Sulfite	SO32-
1
Nitrate	NO31-
Nitrite	NO21-

 

PRACTICE PROBLEMS: Without looking at the polyatomic ions list try to write in the missing charge. The first problem is done for you as an example.

Name	Charge
1
Hypochlorite	-1
Chlorate	-1
1
Carbonate	-2
Carbonite	-2
1
Hypoiodite	-1
Periodate	-1
1
Phosphate	-3
Phosphite	-3

 

 

 

When you finish the lesson, try out the worksheet on polyatomic ions.

What does PER Versus HYPO mean in the Beginning of Polyatomic Ion Names?

Video Explanation of PER Versus HYPO in the Beginning of Polyatomic Ion Names

Group 2 of the polyatomic ion list starts introducing us to the addition of PER and HYPO. Like ATE and ITE, PER and HYPO are a continuation of the counting of oxygens. In most cases (but not all) PER adds one more oxygen from ATE. In most cases (but not all) HYPO is minus one oxygen from ITE.

 

Examples: Adding Per or Hypo can changed the count of the oxygens.

IO4–	Periodate
IO3–	Iodate
1
IO2–	Iodite
IO–	Hypoiodite

 

PRACTICE PROBLEMS: Without looking at the polyatomic ions list try to write in how many oxygen each ion has by comparing it to its neighbor above or below. The first problem is done for you as an example.

 

Name	Amount of Oxygen
1
Perchlorate	4 oxygen
Chlorate	3 oxygen
1
Chlorite	2 oxygen
Hypochlorite	1 oxygen
1
Bromite	2 oxygen
Hypobromite	1 oxygen
1
Perbromate	4 oxygen
Bromate	3 oxygen

 

What does ATE Versus ITE mean on the end of Polyatomic Ion Names?

 VIDEO Explanation of ATE Versus ITE on Polyatomic Ion Names

In Group 1 and 2 of the polyatomic ions list we can notice that many of the polyatomic ions have a name ending in -ATE or -ITE. The -ATE or -ITE is telling the reader each ion has certain a number of oxygens. To be clear, it does not exactly tell you how many oxygens, but it gives you an idea about them. Its purpose is to give you a consistent comparison between -ATE and -ITE.

 

Examples: A difference of one oxygen between them.

Carbonate has three oxygens

Carbonite has two oxygens

 

The hard rule here is -ate always has one more oxygen than -ite. Look at how each -ate and -ite are organized so that you could compare easily. However, -ate DOES NOT always mean it has three oxygens and the ending -ite does not always mean it has two oxygens.

 

PRACTICE PROBLEMS: Without looking at the polyatomic ions list try to write in how many oxygens each ion has by comparing it to its neighbor above or below. The first problem is done for you as an example.

Name	Amount of Oxygen
1
Sulfate	4 oxygens
Sulfite	3 oxygens
1
Chlorate	3 oxygens
Chlorite	2 oxygens
1
Nitrite	2 oxygens
Nitrate	3 oxygens
1
Phosphite	3 oxygens
Phosphate	4 oxygens
1
Carbonite	2 oxygens
Carbonate	3 oxygens
1
Bromate	3 oxygens
Bromite	2 oxygens

 

What is the Polyatomic Ion Two Part Naming System?

VIDEO Explanation of the Two Part Polyatomic Ion Naming System

First, open up the polyatomic ion list in this link and have it ready to look at as you are reading or viewing this section. This way you can look at the examples in the different groups as they are talked about. The first and most obvious thing I would like to point out is that, in most of the polyatomic ions, the first element in the polyatomic ion is what the name of the polyatomic ion starts with.

 

Examples:

Carbonate (CO₃^2-) has carbon as its first element.

Nitrite (NO₂^1-) has nitrogen as its first element.

 

The name is also split into two parts.

 

Examples:

Part 1	Part 2
Carbon	ate
Nitr	ite

 

PRACTICE PROBLEMS: Without looking at the polyatomic ions list try to fill in the first half of the name of the ions below. The first problem is done for you as an example.

Ion	First Half of Name
NO31-	Nitr
PO43-	Phosph
SO32-	Sulf

 

What are polyatomic ions?

In the previous section of ions we looked at how ions could be created from a single atom of a single element. There are, however, ions that are made up of many different atoms bonded together. These are what we call the polyatomic ions. If you break down the word, poly means “many” and atomic means “atoms”.  I designed several of the sections after this one to help you understand the polyatomic ions better.  The polyatomic ion sections also help you find tricks to more easy memorize the polyatomic ions if needed.

Why are polyatomic ions important?

We interact with polyatomic ions all the time without realizing it.  They are in the soil and streams and they are critical components to all life forms.  Some of them are the nutrients that make up the components of all living things body.   For example, phosphate is critical for farmers to spread in the field as part of fertilizer.  At the same time phosphate makes up part of our blood, our bones, and our DNA.

Do I need to memorize the polyatomic ions?

YOU CAN LOOK AT THE POLYATOMIC ION LIST IN THE PICTURE BELOW OR IN THIS LINK. Whether or not you need to memorize the polyatomic ions will depend mostly on what kind of chemistry class you are in.  In most general chemistry high school classes, you do not need to memorize them.  However, honors high school classes may make you memorize a small number of these polyatomic ions but AP and college classes will make you memorize most of them.

What do I need to know or memorize about the polyatomic ions?

What you have to know is the name, chemical formula, and charge of everything in the polyatomic ion list below.  I ranked the following polyatomic ions into groups. Group 1 you absolutely have to know how to spot or memorize and you will see them very often. Group 2 would be useful to know, but not critical. Some of the Group 3 polyatomic ions will be used in AP and college students. While you are reading through and learning this chapter you should always have the polyatomic ions link open so you can look back at it. All other sections in this lesson are just to help you understand the names of the polyatomic ions and to help you memorize them. If you know how to memorize them on your own then you don’t need to look at any more sections. However, I strongly suggest that you do continue with the rest of the sections.

 

 

What is the difference between covalent, ionic, and metallic bonds?

Two or more elements are held together by their electrons, but the electrons can do this in different ways. Electrons are like the skin of the atom because they are on the outer parts of the atom. Atoms interact with each other mostly by contact or trading of electrons, just like people interact with each other by contacting skin in the form of a handshake. As you read through the explanations of the different types of bonds have the periodic table of metal versus non-metal elements open. If you are confused by this section then it is suggested you go back to the metals and non-metals section.

Covalent: One way to create a bond between elements is if those elements share electrons. The covalent bonds form between two or more NON-METAL elements. These bonds are typically strong and flexible. One good example of covalent bonds is the bonds in your skin. Your skin can be twisted to show its flexibility and it takes a very sharp edge, like a knife, to cut it, which shows its strength.

Ionic: Another way to create bond is if those elements steal electrons from each other. The ionic bonds form between at least one METAL and at least one NON-METAL element. These bonds are typically strong and brittle. One good example is salt. You can crush it with a hammer, which shows its strength. When that happens, it will crumble into smaller pieces showing its’ brittle nature.

Metallic: The third and least common type of bond forms if the elements create a sea of electrons. This means the electrons become delocalized from their original atom and flow easily from one atom to another. It is like if you were swimming in a lake or ocean. Your sweat delocalizes from your body and mixes in with all the other water around you. The metallic bonds form between two or more METALS. These bonds are typically strong, have moderate flexibility and conduct electricity well.

These 3 types of bonds (covalent, ionic, and metallic) make up what are called intramolecular bonds (Notice the first 5 letters: INTRAmolecular bonds). They are the bonds formed within one compound or molecule.

With this information we want to be able to identify different compounds that we come across. Any time we see a chemical in front of us we want to be able to label it as covalent, ionic, or metallic. The first stage of this is to be able to take two elements together and say what kind of bond would form between them? Examples are below.

 

Examples: Label each pair of elements as covalent, ionic, or metallic in terms of the bonds using the metals versus non-metals periodic table. VIDEO Explanation of bonding examples 1.

Na to Br	Ionic
N to O	Covalent
Ni to Zn	Metallic
Ca to S	Ionic
Si to Cl	Covalent
F to F	Covalent
Sr to Cr	Metallic

 

PRACTICE PROBLEMS: Label each pair of elements as covalent, ionic, or metallic in terms of the bonds bonds using the metals versus non-metals periodic table.

Ba to P	Ionic
Li to C	Ionic
Cu to Pb	Metallic
Al to H	Ionic
I to Kr	Covalent
Fe to O	Ionic
Pt to Ba	Metallic
B to H	Covalent

You also want to be able to look at complete compounds and be able to tell if they are covalent, ionic, or metallic.

Examples: Label each compound as covalent, ionic, or metallic in terms of the bonds using the metals versus non-metals periodic table. VIDEO Explanation of bonding examples 2.

BeO	Ionic
MnAg2	Metallic
SCl2	Covalent
W2Se	Ionic
SiF4	Covalent
CaSn	Metallic
AsBr2	Covalent

PRACTICE PROBLEMS: Label each compound as covalent, ionic, or metallic in terms of the bonds using the metals versus non-metals periodic table.

BH3	Covalent
CrF2	Ionic
TeI2	Covalent
Hg2Co	Metallic
Fr2S	Ionic
KAg	Metallic
N2O3	Covalent

How different compounds and molecules form the way they do depends on whether they are covalent, ionic, or metallic. Covalent will be discussed when you get to the valence electron dot structures lesson (Lewis structures lesson) later in the website. Ionic will be discussed in another section of this lesson called how to form ionic compounds. Metallic will more than likely not be discussed on this website.

 

 

How do you calculate the molar mass?

After we have learned how to form these ionic compounds, we can move on to calculating the mass of a compound. This is what teachers will call a molar mass, or atomic mass, or molar weight, or an atomic weight. Most chemistry teachers will use all of those phrases in the previous sentence interchangeably. As far as you need to know, the molar mass, atomic mass, molar weight, or atomic weight are all talking about the same thing and you use the same resource to answer these problems. That resource is the periodic table. At the bottom of each individual element box on the periodic table will be the molar mass of each element. The way to solve the problems is to total up all the masses of the individual atoms. In order to keep the numbers simple and easy to read, I will usually round to the nearest whole number. If you are confused as to what the parenthesis mean then look back to a previous section representing compounds and molecules with subscripts.

 

Examples: Give the molar mass of the compounds or molecules.

H2O	18 g/mol
CO2	44 g/mol
N4S3	152 g/mol
(H3O)3As	130 g/mol
Ra(BrO3)2	482 g/mol

 

VIDEO Calculating Molar Mass Demonstrated Example 1: Give the molar mass of the compound or molecule: Al₂S₃

 

Step 1:

What is the mass of a single Aluminum?

Answer: 27 g/mol

 

Step 2:

How much Aluminum do we have?

Answer: 2

 

Step 3:

What is the total mass of the Aluminum?

Answer: 2 * 27 g/mol = 54 g/mol

 

Step 4:

What is the mass of a single Sulfur?

Answer: 32 g/mol

 

Step 5:

How much Sulfur do you have?

Answer: 3

 

Step 6:

What is the total mass of the Sulfur?

Answer: 3 * 32 g/mol = 96 g/mol

 

Step 7:

What is the total mass of the compound?

96 g/mol + 54 g/mol = 150 g/mol

COMPLETE ANSWER: 150 g/mol

 

VIDEO Calculating Molar Mass Demonstrated Example 2: Give the molar mass of the compound or molecule: Be(NO₃)₂

 

Step 1:

What is the mass of a single Beryllium?

Answer: 9 g/mol

 

Step 2:

How much Beryllium do we have?

Answer: 1

 

Step 3:

What is the total mass of the Beryllium?

Answer: 1 * 9 g/mol = 9 g/mol

 

Step 4:

What is the mass of a single Nitrogen?

Answer: 14 g/mol

 

Step 5:

How much Nitrogen do you have?

Answer: 2

 

Step 6:

What is the total mass of the Nitrogen?

Answer: 2 * 14 g/mol = 28 g/mol

 

Step 7:

What is the mass of a single Oxygen?

Answer: 16 g/mol

 

Step 8:

How much Oxygen do we have?

Answer: 6

 

Step 9:

What is the total mass of the Oxygen?

Answer: 6 * 16 g/mol = 96 g/mol

 

Step 10:

What is the total mass of the compound?

96 g/mol + 28 g/mol + 9 g/mol = 133 g/mol

COMPLETE ANSWER: 133 g/mol

 

PRACTICE PROBLEMS: Give the molar mass of the compounds or molecules. Use the periodic table.

BaSe	216 g/mol
Si5As8	740 g/mol
K3P	131 g/mol
CH4	16 g/mol
C6H12O6	180 g/mol
BClF2	84 g/mol
(NH4)3P	85 g/mol
Pb(Cr2O7)2	639 g/mol
Ba(C2H3O2)2	255 g/mol

 

We will be reviewing the molar mass continuously throughout the first half of chemistry, so make sure you know how to solve the above problems and keep the techniques in mind.

 

What is the difference between atoms, molecules, and compounds?

Definitions are important to address in this lesson. At this point in chemistry, many people get confused by the words atom, molecule, and compound. An atom is the smallest indivisible part of matter. When you look at a chemical like Rb₂S, 1 Rb is 1 atom. An atom is simply any single element written. So if a single element is written alone like P then it is only an atom. However, that P is also a molecule because it is separate from other chemicals. A molecule is simply any arrangement of atoms that are bound to each other but separate from other chemicals. I like to think of a molecule as a team. Most of the time a team consists of 2 or more people but in theory a team could be only one person. So most of the time a molecule consists of many different atoms, but it is possible for a molecule to consist of only 1 atom. The last definition we need to know is a compound. A compound is when two or more different elements are bound together. Our original example Rb₂S is a compound. Compounds will be by far the most common chemical arrangement you will see.

Examples: Label whether each is a compound, molecule, or atom. You can give more than one answer if needed. VIDEO Atoms, Molecules, and Compounds Examples Video 1.

NH3	Compound and Molecule
He	Molecule and Atom
Br2	Molecule
Xe	Molecule and Atom
Si4	Molecule
Li2S	Compound and Molecule

 

PRACTICE PROBLEMS: Label whether each is a compound, molecule, or atom. You can give more than one answer if needed.

NaCl	Compound and Molecule
CO2	Compound and Molecule
F2	Molecule
S8	Molecule
Ar	Molecule and Atom
C2H6O2	Compound and Molecule

 

 

Before we discuss the formation of any compounds, we have to make sure we have an understanding of how those compounds are represented in chemistry. In any type of chemistry class, the main focus of how many elements you have in a compound is what they call the subscripts. Subscripts are numbers that are placed to the right of and down towards the bottom of an element. The subscripts are a number that multiplies the element they are next to. If there is no subscript present then that means it is counted as a 1.

 

Examples: Identify what the subscripts of each element in the compounds below.

K2O	K is 2 and O is 1
SrF2	Sr is 1 and F is 2
Mg3P2	Mg is 3 and P is 2

 

Some chemical representations also have parenthesis in them. These parenthesis are much like the ones you use in math class. That is parenthesis are used to multiply everything inside them by a number that is put outside of them.

 

Examples: How many atoms of each element are present in the following compounds.

(NH4)2S	2 N and 8 H and 1 S
Be(NO3)2	1 Be and 2 N and 6 O
Fe2(S2O3)3	2 Fe and 6 S and 9 O

 

 

VIDEO Counting Compound Subscripts Demonstrated Example 1: How many atoms of each element are present in the compound below.

Ca₃(PO₄)₂

 

Step 1:

What is the first element?

Answer: Ca

 

Step 2:

What is the subscript of the first element?

Answer: 3

 

Step 3:

Is Ca in parenthesis?

Answer: No

 

Step 4:

How many Ca atoms do we have?

Answer: 3 Ca

 

Step 5:

What is the next element?

Answer: P

 

Step 6:

What is the subscript of P?

Answer: 1

 

Step 7:

Is P in parenthesis?

Answer: Yes

 

Step 8:

What is the subscript of the parenthesis?

Answer: 2

 

Step 9:

How many P atoms do we have?

Answer: 1 * 2 = 2 P

 

Step 10:

What is the last element?

Answer: O

 

Step 11:

What is the subscript of O?

Answer: 4

 

Step 12:

Is O in parenthesis?

Answer: Yes

 

Step 13:

What is the subscript of the parenthesis?

Answer: 2

 

Step 14:

How many O atoms do we have?

Answer: 4 * 2 = 8 O

 

Step 15:

COMPLETE ANSWER: 3 Ca and 2 P and 8 O

 

PRACTICE PROBLEMS: How much of each element are present in the following compounds.

LiCl	1 Li and 1 Cl
Rb2S	2 Rb and 1 S
Al2Se3	2 Al and 3 Se
CH4	1 C and 4 H
Cs2O	2 Cs and 1 O
Cr3P2	3 Cr and 2 P
Sr(OH)2	1 Sr and 2 O and 2 H
Cr2(SO3)3	2 Cr and 3 S and 9 O
Pb3(PO4)4	3 Pb and 4 P and 16 O

 

What is this lesson about?

This lesson is mostly about how different elements interact to form different compounds. You will also learn how to represent different elements together in one compound and some of the measurements you can take with those compounds.

 

Why is it critical to understand?

If you are not able to understand how different elements come together and how those elements and compounds are recognized, then you will not be able to continue with lessons like nomenclature, valene electron dot structures (Lewis structures), and chemical equations. Beyond that, knowing how elements are put together also allows you to know how to take them apart. Whether you can dissolve them in water or whether you need to burn them in a fire. The type of the bonding that occurs also give you clues on how strong that bond is and therefore, what that material you have created can be used for.

 

What you should know before attempting this lesson?

If you have trouble in this lesson go back to the section on Ions.

 

New Learning Sections:

—> Representation of Compounds and Molecules with Subscripts

—> Atoms, Molecules, and Compounds

—> Calculating the Molar Mass of Compounds

—> Covalent, Ionic, and Metallic Bonds

—> Introduction to Polyatomic Ions Part 1

—> Two Part Name for Polyatomic Ions Part 2

—> ATE versus ITE in Polyatomic Ion Names Part 3

—> PER versus HYPO in Polyatomic Ion Names Part 4

—> Memorizing Polyatomic ion Charges Part 5

—> Identifying Polyatomic Ions in Compounds Part 6

—> Forming Ionic Compounds

—> Breaking Apart Ionic Compounds

 

Reference Pages:

—> Polyatomic Ions List

—> Ion Rules of the Periodic Table

 

Worksheets:

—> Compounds and Bonding Worksheet 1

—> Compounds and Bonding Worksheet 1 WITH ANSWERS

—> Polyatomic Ions Worksheet 1

—> Polyatomic Ions Worksheet 1 WITH ANSWERS

—> College:   Polyatomic Ions Worksheet 2

—> College:   Polyatomic Ions Worksheet 2 WITH ANSWERS

 

What sections should I know before attempting to learn this section?

—> Solving for an Unknown

—> Calculating the Wavelength and Frequency of Light

 

How do you calculate the energy of light?

We can also link the frequency calculations to the energy of a wave of light. When the frequency of a wave is HIGH the energy is LARGE. If the frequency is LOW the energy is SMALL. This relationship is also in the equation below.

 

Energy = Plank’s constant * (frequency)

E = h (f)

 

In the equation E represents energy and has the units of Joules (J), f stands for frequency and has the units of Hz, and h represents what they call Plank’s constant. Plank’s constant is 6.626 *10^-34 m²kg/s.

 

Notice you can link together the equation for the energy of a wave and the equation for the speed of a wave because they both contain frequency.

 

VIDEO Energy of Light Calculation Demonstrated Example 1: If the frequency of light is 7.5 * 10²¹ Hz, how much energy does one this wave contain?

 

Step 1:

What information are we given?

Answer:

frequency = f = 7.5 * 10²¹ Hz

Plank’s constant = h = 6.626 * 10^-34 m²kg/s (this is a constant that you should always have even if it does not state it in the problem)

 

Step 2:

What is the problem asking for?

Answer: Energy = E

 

Step 3:

What is the formula the question involves?

Answer: E = h (f)

 

Step 4:

How do we fill in the numbers for the formula?

Answer: E = 6.626 * 10^-34 m²kg/s (7.5 * 10²¹ Hz)

 

Step 5:

COMPLETE ANSWER: about 5.0 * 10^-12 J

 

VIDEO Energy of Light Calculation Demonstrated Example 2: If the wavelength of light is 2.8 * 10⁴ m, how much energy does one of these waves contain?

 

Step 1:

What information are we given?

Answer:

wavelength = λ = 2.8 * 10⁴ m

speed of light = c = 3.0 * 10⁸ m/s (this is a constant that you should always have even if it does not state it in the problem)

Plank’s constant = h = 6.626 * 10^-34 m²kg/s (this is a constant that you should always have even if it does not state it in the problem)

 

Step 2:

What is the problem asking for?

Answer: Energy = E

 

Step 3:

What is the first formula the question involves?

Answer: λ (f) = c

 

Step 4:

How do we fill in the numbers for the formula?

Answer: 2.8 * 10⁴ m(f) = 3.0 * 10⁸ m/s

 

Step 5:

How do we rearrange the equation to solve for frequency (f)?

Answer: Divide both sides by 2.8 * 10⁴ m

2.8 * 104 m(f) =	3.0 * 108 m/s
2.8 * 104 m	2.8 * 104 m

 

Step 6:

Cross out like terms

2.8 * 104 m(f) =	3.0 * 108 m/s
2.8 * 104 m	2.8 * 104 m

 

Step 7:

Simplify

f =	3.0 * 108 m/s
	2.8 * 104 m

 

Step 8:

What is the answer for the frequency (f)?

f = 10714 Hz or about 1.07 * 10⁴ Hz

 

Step 9:

What is the second formula the question involves?

E = h (f)

 

Step 10:

How do we fill in the numbers for the formula?

E = 6.626 * 10^-34 m²kg/s (1.07 * 10⁴ Hz)

 

Step 11:

COMPLETE ANSWER: about 7.1 * 10^-30 J

 

PRACTICE PROBLEMS: Solve for the unknown wavelength, frequency, or energy. Don’t forget you that you have your constants of speed of light = c = 3.0 * 10⁸ m/s and Plank’s constant = h = 6.626 * 10^-34 m²kg/s.

 

If the frequency of light is 3.5 * 106 Hz then what is the energy of the wave?

Answer: 2.3 * 10^-27 J

 

If the energy of a light wave is 2.9 * 10^-9 J, what is the frequency?

Answer: 4.4 * 10²⁴ Hz

 

If the wavelength of a light wave is 1.4 * 10^-8 m, what is the energy?

Answer: 1.4 * 10^-17 J

 

If the energy of light is 8.7 * 10^-11 J, what is the wavelength of the light?

Answer: 2.3 * 10^-15 m

What sections should I know before attempting to learn this section?

—> Solving for an Unknown

 

How do you calculate the wavelength or frequency of light?

In addition to the wavelength and frequency being related to the color of light, they are also related to each other. If you have a LONG wavelength then the frequency is LOW and if you have a SHORT wavelength then the frequency is HIGH. This relationship is also demonstrated in the equation below.

 

Wavelength * (Frequency) = speed

λ (f) = speed

 

The equation above is the general way to calculate the speed of a wave. However, the equation below is usually the one you see in a chemistry class. It is more specific to calculating the frequency and wavelength of the speed of light but all the concepts behind it are the same. Wavelength (λ) is measured in the units of meters (m). Frequency (f) is measured in the units of Hertz (Hz). The speed of light (c) is a constant number and its value is 3.0 * 10⁸ m/s.

 

Wavelength * (Frequency) = speed of light

λ (f) = c

 

VIDEO Speed of light, Frequency, and Wavelength Calculation Demonstrated Example 1: If the frequency of the light is 240 Hz what is the wavelength?

 

Step 1:

What information are we given?

Answer:

frequency = f = 240 Hz

speed of light = c = 3.0 * 10⁸ m/s (this is a constant that you should always have even if it does not state it in the problem)

 

Step 2:

What is the problem asking for?

Answer: wavelength = λ

 

Step 3:

What formula does this question involve?

Answer: λ (f) = c

 

Step 4:

How do we fill in the numbers for the formula?

Answer: λ (240 Hz) = 3.0 * 10⁸ m/s

 

Step 5:

How do we rearrange the equation to solve for wavelength (λ)?

Answer: Divide both sides by 240 Hz

λ   (240 Hz) =	3.0 * 108 m/s
240 Hz	240 Hz

 

Step 6:

Cross out like terms

λ (240 Hz) =	3.0 * 108 m/s
240 Hz	240 Hz

 

Step 7:

Simplify

λ =	3.0 * 108 m/s
	240 Hz

 

Step 8:

COMPLETE ANSWER: 1.25 * 10⁶ m

 

VIDEO Speed of Light, Frequency, and Wavelength Calculation Demonstrated Example 2: If the wavelength of light is 3.7 * 10^-7 m what is the frequency?

 

Step 1:

What information are we given?

Answer:

wavelength = λ = 3.7 * 10^-7 m

speed of light = c = 3.0 * 10⁸ m/s (this is a constant that you should always have even if it does not state it in the problem)

 

Step 2:

What is the problem asking for?

Answer: wavelength = f

 

Step 3:

What formula does this question involve?

Answer: λ (f) = c

 

Step 4:

How do we fill in the numbers for the formula?

Answer:

λ (3.7 * 10-7 m) =	3.0 * 108 m/s
1

 

Step 5:

How do we rearrange the equation to solve for frequency (f)?

Answer: Divide both sides by 3.7 * 10^-7 m

(3.7 * 10-7 m)  f =	3.0 * 108 m/s
3.7 * 10-7 m	3.7 * 10-7 m

 

Step 6:

Cross out like terms

(3.7 * 10-7 m) f =	3.0 * 108 m/s
3.7 * 10-7 m	3.7 * 10-7 m

 

Step 7:

Simplify

f =	3.0 * 108 m/s
	3.7 * 10-7 m

 

Step 8:

COMPLETE ANSWER: 8.1 * 10¹⁴ Hz

 

PRACTICE PROBLEMS: Solve for the unknown wavelength or frequency. Don’t forget that you have the speed of light constant = c = 3.0 * 10⁸ m/s.

 

If the wavelength of a light wave is 14.6m, what is the frequency?

Answer: 2.0 * 10⁷ Hz

 

If the frequency of a light wave is 0.89 Hz, what is the wavelength?

Answer: 3.4* 10⁸ m

 

If the wavelength of a light wave is 6.7 * 10^-5 m, what is the frequency?

Answer: 4.5 * 10¹² Hz

 

If the frequency of a light wave is 5.3 * 10¹² Hz, what is the wavelength?

Answer: 5.7 * 10^-5 m

 

 

So what does the interaction of light and electrons look like?

Like light, electrons are also tiny particles. In addition, they move like light both very quickly and tend to follow wave patterns. What this means is that the particle of light, the photon, and the electron particle can interact. When these two particles (electrons and photons) interact, it turns out to be something like a pool (billiards) game. The photons can knock around the electrons and cause the electrons to move to different positions within an atom.

 

VIDEO demonstration of the interactions between light and electrons that is explained in text and pictures below.

If we look at an atom as a Bohr model like the picture below, it allows us to gain a pretty easy understanding of the interactions between electrons and light. We will use the element hydrogen as our model atom. The first picture is what happens before the light particle (photon) hits the electrons. Notice how the electron starts on the closest electron shell (energy level) to the nucleus. The closest available shell (energy level) to the nucleus is where you will find the electron assuming no energy is put into it.

 

 

The second picture is what happens after the light particle (photon) hits the electron. See how the electron is moving from the first electron shell to the second electron shell. That was because the photon hit the electron and knocked or pushed it up one electron shell (energy level).

 

 

In the third picture, the electron has moved to the second shell.

 

 

An instant (very very small) amount of time after the third, picture happens the fourth picture happens. The fourth picture is an electron falling back down to the closest shell to the nucleus while releasing a photon. It basically had this photon in storage and when it falls it releases it. Therefore, the light will travel away from the electron at this point.

 

 

In the fifth and last picture the electron in the hydrogen atom has returned to its original shell (energy level), closest to the nucleus.

 

 

This concept of electrons moving when they are struck by something else is actually the same basic concept that a soccer player uses to juggle a ball. Just like the soccer ball is being pulled to the Earth by gravity, the electron is being pull close to the nucleus because the electron (negatively charged) has an attraction to the protons (positively charged) in the nucleus. If you want to force the soccer ball away from the Earth, you have to knock (kick) your foot into the soccer ball. If you want to force the electron away from the nucleus, you have to knock the electron with a photon. When you knock the soccer ball, it flies away from the Earth and comes to some level for a split second and then falls back down. When you knock the electron, it flies away from the nucleus comes to a certain level for split second and then falls back down. You can even describe the energy that the electron stores as it rises and falls because light is a form of energy. The electron absorbs or releases that light energy as it changes position. Likewise, the soccer ball stores the energy in terms of sound. When you kick the soccer ball the collision of your foot and the ball make a sound. When the soccer ball falls back down it releases that sound once it hits something.

 

In fact, there is only one small difference between what happens with a soccer ball versus what happens with an electron. The electron follows a quantum (or is quantized) where as the soccer ball is not. Quantized means that the electron can only be at the first shell or the second shell or the third shell and so on. The electron cannot ever occupy the space in between the shells. So, if we run a video of the pictures in the example I gave you this is what it would look like. You see the electron at the first shell. You see the electron disappear. You see the electron reappear at the second shell. You see the electron disappear. You see the electron reappear at the first shell. It is kind of strange to think about but that is how it works. On the other hand, the soccer ball can move and occupy any space in between its starting point and its’ highest point after it is kicked by a person.

 

VIDEO Here is a video demonstration of the electron versus the football (soccer ball).

 

During the electron’s travel between different energy levels (shells), you can see different wavelengths of light coming off of the atom. The WAVELENGTH of light that the electron gives off when it falls back from a higher energy level to a lower energy level depends on the DISTANCE traveled by that electron. The GREATER the distance the HIGHER the energy, the HIGHER the frequency, and therefore the SHORTER the wavelength. For example, in most atoms, if an electron jumps only one energy level (shell) you will see a long wavelength (red side of the spectrum). However, if an electron jumps 4 or more energy levels then you will see a very short wavelength (purple side of the spectrum).

 

 

 

What is light?

VIDEO Explanation of light.

The technical name for light is electromagnetic radiation. Light is made up of tiny particles (spheres) called photons and is a form of energy. The sensation you get of feeling warmth on your skin when you walk outside on a sunny day is actually billions and billions of photon particles smashing into your skin and the reaction of your nerves sensing them. So, whenever there is light around, you are being bombarded by countless numbers of tiny particles. It is as if a thousand or more people all threw a ball at you at once and all those balls hit you. This is the easy part about light to understand. However, light also behaves in a second way. Unlike a ball or bullet, the photon of light does not travel in a straight line. The photon (particle) of light travels in a wave. That means it travels up and down slightly as it travels in the forward direction. This is kind of like a stone skipping across water except a stone will eventually sink, whereas the photon never stops the up and down motion. These two components, or behaviors of light,are what scientists call the wave particle duality. That is light acts both like a wave and a particle. A picture of the light wave is shown below. In science, anything that is called a wave or has wave properties travels in this way.

 

 

What is the difference between different colors of light?

The most important feature of the wave is what they call the wavelength. A wavelength is the measurement of how long the wave is from crest to crest (highest point to the next highest point). You can also make the wavelength measurement at any repeating set of points like trough to trough (lowest point to lowest point). The difference in color of light that we see is the difference in how long or short those wavelengths are. If we look at a representation of the different colors that we see (the visible spectrum), we observe that the colors run from red to purple (violet) like in the electromagnetic spectrum picture below. Thereare also types of electromagnetic radiation that we, as humans, cannot see. Past the red side of the spectrum this includes things like radio waves and infrared. Past the purple side of the spectrum this includes things like ultraviolet and x-rays.

 

 

What is the relationship between the wavelength and the color of light?

Notice in the picture above that the red side of the spectrum has longer wavelengths than the purple side. Therefore, we can use color of light that we can see as a relative indicator of wavelength. Toward the RED side of the electromagnetic spectrum the WAVELENGTH IS LONGER. Toward the PURPLE side of the electromagnetic spectrum the WAVELENGTH IS SHORTER.

 

What is the frequency?

VIDEO Explanation of frequency.

Frequency is the other way to measure waves. In common terms it is how many times the wave goes up and down. In scientific terms, frequency is how many waves can you observe passing in a particular time period. The units of frequency are 1 / seconds or Hertz (Hz). A good way to think about frequency is to imagine you are standing at the edge of a body of water like a lake or ocean. Water waves behave the same laws as light waves so we can use water waves to test measurements and theories. Let’s bring a stopwatch and count how many waves reach our feet in one minute, then we can say something about the frequency of those waves. If a lot of waves hit your feet in one minute, like 10 waves, then the frequency of that particular set of waves is high. If only a handful of waves hit your feet in one minute, like 3 waves, then the frequency of that particular set of waves is low.

 

What is the relationship between frequency and the color of light?

We can look back to the electromagnetic spectrum and also see the relationship between frequency and color of light. Toward the RED side of the electromagnetic spectrum the FREQUENCY IS LOW. Toward the PURPLE side of the electromagnetic spectrum the FREQUENCY IS HIGH.

 

PRACTICE PROBLEMS: Given information about the wavelength or frequency, state whether the color of the light is more toward the red side of the spectrum or the purple side of the spectrum.

 

Wave has a short wavelength	Purple side
Wave has a low frequency	Red side

 

 

How do we calculate the energy required to move an electron between energy levels?

The answer is the Rydberg equation. Note this Rydberg equation can only be used for hydrogen energy levels or shells. It is shown below. The equations are written two different ways. Both of the equations are the same but the second one is organized better for future demonstrated examples.

 

Δ E = R_H((1/n_I²) – (1/n_f²))

 

Δ E =	1           –	1
RH	nI2	nf2

 

The Δ E stands for the change in energy with the units of Joules (J). The R_H is the Rydberg constant. It is simply the number 2.178 * 10^-18 J that only applies to hydrogen and allows you to put the equation together. n_I stands for the electron shell (or energy level) initial. It will be an integer like 1,2,3… The n_f stands for the electron shell (or energy level) final and will also be an integer like 1,2,3…

 

The Rydberg equation also connects to other equations in this lesson. The most obvious one is E = h (f) in the section calculating the energy of light. Just treat Δ E and E as the same thing. The only difference between them is that Δ E can be a positive or negative number where as E has to be a positive number.

 

VIDEO Rydberg Equation Demonstrated Example 1: If the electron of a hydrogen atom goes from the second electron shell to the first electron shell, what is the maximum amount of energy that can be released as light? The Rydberg constant is 2.178 * 10^-18 J.

 

Step 1:

What information are we given?

Answer:

n_I = 2

n_f = 1

R_H = 2.178 * 10^-18 J

 

Step 2:

What is the problem asking for?

Answer: Δ E = ?

 

Step 3:

What formula does the question involve?

Answer:

Δ E =	1                      –	1
RH	nI2	nf2

 

Step 4:

How do we fill in the numbers for the equation?

Answer:

Δ E =	1                     –	1
2.178 * 10-18 J	22	12

 

Step 5:

How do we solve for the Δ E?

Answer: First simplify the exponents

Δ E =	1                   –	1
2.178 * 10-18 J	4	1

 

Step 6:

Then put the fractions in decimals

Δ E =	0.25               –	1
2.178 * 10-18 J

 

Step 7:

Now subtract the right side

Δ E =	– 0.75	1
2.178 * 10-18 J

 

Step 8:

Finally multiply both sides by 2.178 * 10^-18 J

2.178 * 10-18 J * Δ E =	– 0.75 * 2.178 * 10-18 J	1
2.178 * 10-18 J

 

Step 9:

Cross out like terms on left side

2.178 * 10-18 J * Δ E =	– 0.75 * 2.178 * 10-18 J	1
2.178 * 10-18 J

 

Step 10:  simplify

Δ E =	– 0.75 * 2.178 * 10-18 J	1
1

 

Step 11:  calculations

Answer: -0.75 * (2.178 * 10^-18 J) = 1.6 * 10^-18

1.6 * 10-18 =	– 0.75 * 2.178 * 10-18 J	1
1

 

Step 12:

What is the final answer?

COMPLETE ANSWER: – 1.6 * 10^-18 J

 

VIDEO Rydberg Equation Demonstrated Example 2: When a single hydrogen atom absorbs 2.09 * 10^-18 J, and its electron starts from the first energy level, what energy level does the electron end at? The Rydberg constant is 2.178 * 10^-18 J.

 

Step 1:

What information are we given?

Answer:

Δ E = 2.09 * 10^-18 J

n_I = 1

R_H = 2.178 * 10^-18 J

 

Step 2:

What is the problem asking for?

Answer: n_f = ?

 

Step 3:

What formula does the question involve?

Answer:

Δ E =	1          –	1
RH	nI2	nf2

 

Step 4:

How do we fill in the numbers for the equation?

Answer:

2.09 * 10-18 J =	1          –	1
2.178 * 10-18 J	12	nf2

 

Step 5:

How do we solve for n_f?

Answer: First divide the left side.

0.96 =	1          –	1
	12	nf2

 

Step 6:

Then solve the exponent on the right side.

0.96 =	1          –	1
	1	nf2

 

Step 7:

Now solve the fraction on the right side.

0.96 =	1          –	1
		nf2

 

Step 8:

Minus 1 to both sides

0.96 – 1 =	1 – 1      –	1
		nf2

 

Step 9:

Cross out right side

0.96 – 1 =	1 – 1      –	1
		nf2

 

Step 10:

Simplify left side

– 0.04 =	–	1
		nf2

 

Step 11:

Eliminated the negatives from both sides

0.04 =		1
		nf2

 

Step 12:

Multiply both sides by n_f²

nf2 * 0.04 =	nf2	1
		nf2

 

Step 13:

Cross out like terms

nf2 * 0.04 =	nf2	1
		nf2

 

Step 14:

Simplify

nf2 * 0.04 =		1
1

 

Step 15:

Divide both sides by 0.04

nf2 * 0.04 =		1
0.04	0.04

 

Step 16:

Cross out like terms

nf2 * 0.04 =		1
0.04	0.04

 

Step 17:

Simplify

nf2 =		1
	0.04

 

Step 18:

Divide the right side ( 1 / 0.04 = 25 )

nf2 =		25
		1

 

Step 19:

Take the square root of both sides

Sqrt nf2 =		Sqrt 25
		1

 

Step 20:

Cross out the square root and the square on the left side

Sqrt nf2 =		Sqrt 25
		1

 

Step 21:

Simply

nf =		Sqrt 25
		1

 

Step 22:

Take the square root of 25

nf =		5
		1

 

Step 23:

What is the final answer?

COMPLETE ANSWER: The 5^th energy level

 

PRACTICE PROBLEMS: Solve the energy level (electron shell) or energy required to move electrons around in a hydrogen atom. Don’t forget to use the Rydberg constant R_H of 2.178 * 10^-18 J when needed.

 

If the electron of a hydrogen atom goes from the third electron shell to the first electron shell, what is the maximum amount of energy that can be released as light?

Answer: -1.94 * 10^-18 J

 

If the electron of a hydrogen atom goes from the second electron shell to the fifth electron shell, what is the maximum amount of energy that can be absorbed as light?

Answer: 4.5 * 10^-19 J

 

When a single hydrogen atom absorbs 2.04 * 10^-18 J, and its electron starts from the first energy level, what energy level does the electron end at?

Answer: Fourth

 

When a single hydrogen atom releases -3.03 * 10^-19 J, and its electron ends in the second energy level, what energy level does the electron start at?

Answer: Third

 

 

How do electrons move between shells or energy levels?

VIDEO Explanation of electrons moving between shells.

With all this information about the position of the electrons compared to the nucleus in sections like electron shells and energy levels, we should also discuss how electrons can move from one shell or energy level to anther. In these next few sections I will use the words electron shells and energy levels to mean that same thing. Teachers can use either of those words to describe where electrons can move to. Make sure you understand both the sections on electron shells and energy levels before you go any further in this section.

 

So how do electrons move from one shell to the next? The explanation is actually quite simple. Let us discuss how an electron moves from a lower energy level to a higher energy level first. Electrons are negatively charged and move around the other edges of the atom, and protons are positively charged and are at the center of the atom (nucleus). That means that the negatively charged electron and the positively charged proton attract each other. So how do I move the electron away from the nucleus to a shell that is further away or to a higher energy level? Simple I expend energy or I put energy into the electron to do it. Pulling the electron away from the nucleus (center of an atom) requires the same kind of action just like moving an object off the floor against the force of gravity.

 

We can create a scenario that is easy to think about. I am going to use three objects or areas to describe it to you. Lets put a soccer ball on the ground. The ground represents the nucleus and the soccer ball represents the electron and the air above the ball represents higher energy levels. If I want to move the ball away from the ground and higher what do I do? I pick the ball up with my hands and move it away from the ground, but what is required for me to do that. I have to expend energy in my muscles and that energy gets stored in the ball. So moving the ball upward requires energy. The exact same requirements of energy are present when you move an electron away from the nucleus. In fact, the formula that allows us to calculate the force or energy between the electron and the proton (electromagnetic) is extremely similar to the formula that allows you to calculate the force or energy between the ball and the ground (gravity). Therefore, my analogy of the ball to the electron not only works well in thought, but also has mathematical proof behind it.

 

What if we now want to determine what happens when the electron moves from an outer shell or energy level to a lower shell or energy level? Again simple. We run the soccer ball analogy in reverse. If you are holding the soccer ball up high in your hands and you want it to move toward the floor what do you do? You drop it. What happens when you drop it? The ball moves toward the floor until something stops it and when the ball stops it lets out a sound. Sound is just a form or energy so all the ball is doing is letting go of some of its energy that was put into it when it became raised up. What does that mean the electron does? When the electron moves from a higher energy level to a lower energy level it must therefore give away some of its energy. However, the electron does not give away energy in the form of sound but in the form of light.

 

What sections should I know before attempting to learn this section?

—> Orbitals Part 1

—> Energy Level Part 2

—> Complete Electron Configuration Part 3

—> Electron Configuration Diagram Part 5

—> Quantum Numbers Part 1 (n)

—> Quantum Numbers Part 2 (L)

—> Quantum Numbers Part 3 (m_L )

—> Quantum Numbers Part 4 (m_s ) 

—> Quantum Numbers Part 5 (All Together)

 

What are the possible quantum numbers?

The other common type of problem with quantum numbers is, what are the possible quantum numbers given in each of the quantum categories? For these questions I usually turn to a formula because it is easier to understand that way. You can also think about it conceptually if you want and I will explain this way as I go. In these problems they usually propose or give you one quantum category, like the energy level (n), and then ask you for the possibilities of another quantum category like the orbital types (L). The formulas I use to help me out are below.

n = number

L = n – 1 (and all integers to zero)

m_L = L + the negative integers of L

m_s = +1/2 or -1/2

 

How do you use these formulas? Pick a number for your energy level. For an example I am going to pick 3. (You can take a look at this periodic table if you don’t remember where your energy levels are) Take your energy level and minus 1 from it to get your possible L.

3 – 1 = 2….So L can equal 2, 1, 0. That means you can have an electron in the d, p, or s orbitals.

Now take your L possibilities add the negative integers to get m_l.

2, 1, 0, (-1, -2)….So your m_l can equal 2, 1, 0, -1, -2. That means in your d type orbitals you can have up to 5 orbitals labeled 2, 1, 0, -1, or -2. (Take a look at this periodic table if you don’t remember where your orbitals are.)

The last one, as always, is the easiest. No matter what the other quantum numbers are m_s can only be either +1/2 or -1/2.

 

VIDEO Quantum Number Demonstrated Example 3: What are the possible quantum numbers if your n = 2?

 

Step 1:

If your n = 2 then what are your possible L?

n – 1 = L ….. 2 – 1 = 1

Answer: L = 1 or 0

 

Step 2:

If your L = 1 or 0 what are your possible m_L?

Answer: m_L = 1 or 0 or -1

 

Step 3:

What are your possible m_s?

Answer: +1/2 or -1/2

COMPLETE ANSWER: L = 1,0, m_L = 1,0,-1, m_s = +1/2, -1/2

 

VIDEO Quantum Number Demonstrated Example 4: What are the possible quantum numbers if your m_L = -3

 

Step 1:

If your m_l = -3 what are your possible L?

We are dealing with the f orbitals here so the L for the f orbitals is 3

Answer: L = 3 or higher

 

Step 2:

If your L = 3 what are your possible n?

The first energy level that contains f orbitals is the 4^th but energy levels beyond the 4^th also have f orbitals.

Answer: n = 4 or higher.

 

Step 3:

What are your possible m_s?

Answer: +1/2 or -1/2

COMPLETE ANSWER: n = 4 or higher, L = 3 or higher, m_s = +1/2, -1/2

Notice, depending on which quantum number was picked for the question it can severely limit or open up the possibilities of the other quantum numbers.

 

PRACTICE PROBLEMS: What are the possible quantum numbers for the following problems?  Remember to have your regular periodic table handy.  Also if you need it have the quantum periodic table and the energy level periodic table ready.

L = 2	n = 3,4,5… mL = 2,1,0,-1,-2 ms = +1/2, -1/2
n = 4	L = 3,2,1,0 mL = 3,2,1,0,-1,-2,-3 ms = +1/2, -1/2
mL= 1	n = 2,3,4… mL = 1,0,-1 ms = +1/2, -1/2

 

What sections should I know before attempting to learn this section?

—> Orbitals Part 1

—> Energy Level Part 2

—> Complete Electron Configuration Part 3

—> Electron Configuration Diagram Part 5

—> Quantum Numbers Part 1 (n)

—> Quantum Numbers Part 2 (L)

—> Quantum Numbers Part 3 (m_L )

—> Quantum Numbers Part 4 (m_s ) 

 

How do we use the 4 different types of quantum numbers together in one answer?

Now that we know what each quantum number means we also want to be able to put all the different quantum numbers together to answer one question.  You may have noticed that I have used the exact same elements as my example problems for all the different quantum categories. That is because for most questions asked about quantum numbers they usually go one of two ways. Tests tend to ask either what are the all the quantum numbers for the last electron in a certain element? Or they tend to ask what are the possible quantum numbers given in each of the quantum categories? We will go through each type of question below.

The first question, what are all the quantum numbers for the last electron in a certain element is answered by analyzing the energy level (n), orbital type (L), specific orbital (m_L), and spin of the electron (m_s). This is very similar to the practice problems we have had so far it is just combining them all into one.

Remember to use the periodic table links like your regular periodic table handy.  Also if you need it have the quantum periodic table and the energy level periodic table ready.

 

VIDEO Quantum Number Demonstrated Example 1: What are the quantum numbers for the last electron of phosphorus? Remember to use these periodic table links if you need them.

 

Step 1:

What energy level is phosphorus in?

Answer: The third energy level…..so n = 3

 

Step 2:

What orbital type is phosphorus in?

Answer: The p orbitals…..so L = 1

 

Step 3:

What specific orbital is phosphorus in?

Answer: The second p orbital…..so m_L = 1

 

Step 4:

What is the spin of the last electron in phosphorus?

Answer: m_s = -1/2

 

Step 5:

What are the quantum numbers for the last electron of phosphorus?

COMPLETE ANSWER: n = 3, L = 1, m_L = 1, m_s = -1/2

 

VIDEO Quantum Number Demonstrated Example 2: What are the quantum numbers for the last electron of yttrium?  Remember to use these periodic table links if you need them.

 

Step 1:

What energy level is yttrium in?

Answer: The fourth energy level…..so n = 4

 

Step 2:

What orbital type is yttrium in?

Answer: The d orbitals…..so L = 2

 

Step 3:

What specific orbital is yttrium in?

Answer: The first d orbital…..so m_L = -2

 

Step 4:

What is the spin of the last electron in phosphorus?

Answer: m_s = -1/2

 

Step 5:

What are the quantum numbers for the last electron of phosphorus?

COMPLETE ANSWER: n = 4, L = 2, m_L = -2, m_s = -1/2

 

Examples: What are the quantum numbers for the last electron of each element below?  Remember to have your regular periodic table handy.  Also if you need it have the quantum periodic table and the energy level periodic table ready.

Ca	n = 4, L = 0, mL = 0, ms = +1/2
Mo	n = 4, L = 2, mL = 1, ms = -1/2
In	n = 5, L = 1, mL = -1, ms = -1/2

 

PRACTICE PROBLEMS: What are the quantum numbers for the last electron of each element below?  Remember to have your regular periodic table handy.  Also if you need it have the quantum periodic table and the energy level periodic table ready.

Si	n = 3, L = 1, mL = -1, ms = -1/2
O	n = 2, L = 1, mL = 0, ms = +1/2
Mn	n = 3, L = 2, mL = 2, ms = -1/2
Rb	n = 5, L = 0, mL = 0, ms = -1/2
Th	n = 5, L = 3, mL = -2, ms = -1/2
Kr	n = 4, L = 1, mL = 1, ms=+1/2

 

 

What sections should I know before attempting to learn this section?

—> Electron Configuration Diagram Part 5

 

What does the m_s stand for in the quantum numbers?

The last quantum category is m_s (It is pronounced “M sub S”) represents what they call the spin. It is the easiest to figure out and the easiest to understand. Just like in the electron diagrams we used an up and a down arrow to represent the electrons traveling different directions within the same orbital. Here instead we use +1/2 and -1/2 to represent the up and down respectively. You can honestly choose these things at random so long as no two electrons in a row (within the same orbital) have the same m_s. You can put them in a specific order on the periodic table but it does not matter that much. I usually put +1/2 first. This is one example of a periodic table where the electron spin is illustrated.

 

Examples: Give the m_s for the last electron in the following elements. (These do not have definite answers but I give the answer I do because of the way I have explained them to you.)

Be	ms = -1/2
Pd	ms = -1/2
Al	ms = +1/2

 

PRACTICE PROBLEMS: Give the m_s for the last electron in the following elements. (These do not have definite answers but I give the answer I do because of the way I have explained them to you.)

Bi	ms = +1/2
F	ms = -1/2
Ni	ms = -1/2
Li	ms = +1/2

 

 

 

What sections should I know before attempting to learn this section?

—> Orbitals Part 1

—> Energy Level Part 2

—> Complete Electron Configuration Part 3

—> Quantum Numbers Part 1 (n)

—> Quantum Numbers Part 2 (L)

 

What does the m_L stand for in the quantum numbers?

The next quantum number category m_L (it is pronounced “M sub L”) stands for the specific orbital within the orbital types. This relates back to how many orbitals each orbital type has. The s orbitals have one orbital each. The p orbitals have 3 orbitals each. The d orbitals have 5 orbitals each. The f orbitals have 7 orbitals each. The m_L depends on the (L). The table below will show you best how m_l is dependent on (L).

If (L) is	Then mL can be…	Orbital types	# of electrons	# of orbitals
L = 0	0	s	2	1
L = 1	-1, 0, 1	p	6	3
L = 2	-2, -1, 0, 1, 2	d	10	5
L = 3	-3, -2, -1, 0, 1, 2, 3	f	14	7

Notice in the table above the two columns I highlight in red are related. On the bottom row there are 7 integers running from negative 3 to positive 3. Each one of those integers represents an individual orbital among that orbital type. Again you can reference the quantum periodic table to show you better where they are and how they are arranged. Keep in mind that different teachers like to display them in slightly different ways. The way I demonstrate on the periodic table seems to be the easiest to understand but it is not the only way to understand this concept.

 

Examples: Give the m_L for the last electron in the following elements.

Be	mL = 0
Pd	mL = 1
Al	mL =-1

 

PRACTICE PROBLEMS: Give the m_L for the last electron in the following elements.

K	mL = 0
S	mL = 0
Ge	mL = -1
Ti	mL = -2
Au	mL = 2
Lu	mL = 3

 

What sections should I know before attempting to learn this section?

—> Orbitals Part 1

—> Energy Level Part 2

—> Complete Electron Configuration Part 3

—> Quantum Numbers Part 1 (n)

 

What does the L stand for in quantum numbers?

The next quantum number category L stands for the type of orbitals. In electron configurations the types of orbitals were the s, p, d, f orbitals that were discussed in the orbitals section of this lesson. In the quantum numbers if L = 0 represents the s orbitals. If L = 1 represents the p orbitals. If L = 2 represents the d orbitals. If L = 3 represents the f orbitals. Another way to represent that is the table below.

 

L =	Orbital types
0 —>	s
1 —>	p
2 —>	d
3 —>	f

 

You can also see how this is represented on the quantum periodic table with this link.

 

Examples: Give the L (orbital type) of the last electron in each element.

Be	L = 0
Pd	L = 2
Al	L = 1

 

PRACTICE PROBLEMS: Give the L (orbital type) of the last electron in each element.

P	L = 1
Co	L = 2
Sn	L = 1
Ba	L = 0
Th	L = 3
W	L = 2

 

 

 

 

What sections should I know before attempting to learn this section?

—> Orbitals Part 1

—> Energy Level Part 2

—> Complete Electron Configuration Part 3

 

What are quantum numbers?

VIDEO Explanation of quantum numbers.

Quantum numbers and electron configurations are talking about the same thing. Each of them represent where the electrons are around the atom. However, the quantum numbers do it in a different way than the electron configurations do. The first difference is the quantum numbers only talk about one electron at a time.  The quantum numbers also divide into 4 different categories represented by letters. Those categories are n, L, m_L, m_s.

 

The quantum category n stands for the energy level. This is the same energy level that was discussed in the earlier energy level section of this lesson. This periodic table link can help you recognize the energy levels on the periodic table.

 

Examples: Give the n (energy level) of the last electron in each element.

Be	n = 2
Pd	n = 4
Al	n = 3

 

PRACTICE PROBLEMS: Give the n (energy level) of the last electron in each element.

He	n = 1
Fe	n = 3
Br	n = 4
Na	n = 3
U	n = 5
Hg	n = 5

 

 

 

What sections should I know before attempting to learn this section?

—> Orbitals Part 1

—> Energy Level Part 2

—> Complete Electron Configuration Part 3

—> Abbreviated Electron Configuration Part 4A

 

What are electron configuration diagrams?

VIDEO Explanation of electron configuration diagrams.

Electron configuration diagrams are how to represent the electron configurations in a picture form. Mainly what you need to know is how the picture below this paragraph is created. I will explain it in words as you are simultaneously looking at the picture below.  You might also want to bring up the energy level periodic table. The electron configuration picture is displayed by putting the orbitals closest to the nucleus (of the lowest energy) further toward the bottom of the picture. As the orbitals get further from the nucleus (higher in energy) they are displayed higher and higher in the picture. For example, the 3p orbitals are higher on an electron configuration diagram then the 2s orbital. The orbital types (s, p, d, f…) are also organized in a specific way. All the s orbitals are lined up furthest toward the left side of the picture. The s orbitals are also stacked one on top of the next. The p orbitals are slightly to the right of the s orbitals and slightly above the s orbitals in the same energy levels. For example, the 2p orbital is to the right of and slightly higher than 2s orbitals. However, the p orbitals like the s orbitals, are stacked one on top of another.

 

 

What is each box representing in the picture above? Each box represents one orbital. Each orbital can only hold 2 electrons maximum. DO NOT CONFUSE THIS WITH THE ORBITAL TYPES. If you take the maximum number of electrons that an orbital type can hold and divide it by 2, you get the amount of orbitals in that orbital type. For example, the d orbitals can hold 10 electrons. So, 10 divided by 2 equals 5. That is why the picture above has 5 boxes representing the 3d orbitals. If we are only able to put 2 electrons in each of the d orbitals, then we can put a total of 10 electrons in all the 3d orbitals. In the electron configuration diagram above, the electrons are represented by the up and down arrows. It does not matter if it is an up or down arrow. Both of them represent an electron. The only difference between the up and down arrow is the direction the electrons are traveling. The up versus down arrows are saying that the electrons are traveling in opposite directions relative to one another. I describe this movement of electrons by comparing it to cars on a racetrack. If I were representing cars on a racetrack by arrows, I would make them go in all the same direction (all up), because all cars go in the same direction during the race. However, what if I took half the cars in my race and made them go in the other direction? Besides a lot of potential car accidents happening, you could then show the cars running the races as half of them being up arrows and half of them being down arrows. This is just like electrons which race toward each other within the same track (orbital). Although the electrons do not collide with each other because they are both negatively charged, the cars around a racetrack versus electrons in an orbital is a very close analogy to what is actually happening in an atom or orbital.

 

To determine which orbitals are closest to the nucleus (lowest in energy), you have to break down your determination into two categories. The first is what energy level is that orbital on. The second is what type of orbital is it. Orbitals of a lower energy level will be closer to the nucleus. For example, the 2p orbitals will be closer to the nucleus than the 4s orbitals because the 2p orbitals is on the second energy level and the 4s are on the fourth energy level. Also, there is an order of orbitals within the same energy level. The s orbitals will be closer to the nucleus than the p orbital within the same energy level. The p orbitals will be closer to the nucleus than the d orbital within the same energy level. The d orbitals will be closer to the nucleus than the f orbital within the same energy level. For example, the 3p orbitals are closer to the nucleus than the 3d orbitals.

 

Examples: Which orbitals are CLOSEST (HAVE THE LOWEST ENERGY) to the nucleus?

3s or 1s	1s
4f or 2p	2p
5d or 5f	5d

 

PRACTICE PROBLEMS: Which orbitals are CLOSEST (HAVE THE LOWEST ENERGY) to the nucleus?

5f or 4f	4f
7p or 6s	6s
2s or 2p	2s
4d or 4p	4p
3d or 4p	3d
6f or 6s	6s

 

Now that we have a general understanding of the number of orbitals, orbital types, and the relationship of orbitals to the nucleus we should be able to start constructing electron configuration diagrams. How do we do that? We should first open up a periodic table with the energy levels and orbital types clearly displayed. Here is the same periodic table viewed in a different way. Just like the electron configures before, we want to start any electron configuration by starting at the top of the periodic table in the 1s orbitals or the spot where hydrogen is. Since it is very hard to explain electron diagrams in text, there are video examples below.

 

VIDEO Example 1 Nitrogen: How to draw an electron diagram.

VIDEO Example 2 Sulfur: How to draw an electron diagram.

VIDEO Example 3 Iron: How to draw an electron diagram.

 

PRACTICE PROBLEMS: Draw the electron diagram of the following elements or charged atoms. Use this periodic table if possible. If that does not work for you try the orbital periodic table. It is best to draw them from memory but you can refer to a template electron diagram if you wish.

B	Answer 1 link
F	Answer 2 link
Ca	Answer 3 link
Rh2+	Answer 4 link
Se2-	Answer 5 link
Ru	Answer 6 link