Functional Groups

Functional groups include compounds with elements other than C and H. This is the most reactive part of molecule. These compounds can be a single atom or a group of atoms, such as F or NO2.

For example, these groups include alcohols, nitro, aldehydes, ketones

Halides and Nitro:
These are simple hydrocarbons that are attached to alkane, -ene, or -yne

The main chain name will have a prefix, depending on the atom.
If :
F = fluoro
Cl = chloro
Br = bromo
I = iodo
NO2 = nitro


In here, we would say it is methane, but there are 3 chlorines.
Then, we would say trifluoromethane.

F is most unreactive, and I is most reactive.

Nitro is unreactive to chemical attack, but is explosive, such as in TNT, and nitroglycerine.

Alcohols:
In alcohols, an OH functional group is present. (Hydroxyl)
These are named by using the longest C chain containing the OH group.
All you have to do is replace the "e" ending with "ol"

For example:
      H   H
       |     |
H - C - C - OH      
       |     |
      H   H
The longest C chain is 2, and we have an OH functional group.
We would end up with ethane, except we replace the e with ol, ending up with ethanol.
Naming the alcohol functional group in different positions is also the same. Just include the lowest position number possible. If there is more than one OH, add prefixes such as -diol, -triol

All alcohols are poisonous to some degree.

Aldehydes and Ketones:
This carbonyl functional group features a double-bonded oxygen.
Aldehydes have the double bonded oxygen at the end of the chain, and the name ends in -al. But in ketones, the double bond is not on either end, and it ends in -one.



Aldehydes are very reactive, and ketones are unreactive.



Here's a video on halides:



And a video on general functional groups, such as alcohols.

Here's some worksheets for practice:
http://misterguch.brinkster.net/PRA051.pdf
http://www.arps.org/users/hs/thompsom/chemcom/unit_3/Functional_Groups_Worksheet.pdf

Alkenes and Alkynes

Carbons can form double or triple bonds with carbon atoms. This, in turn, would mean fewer hydrogens become attached. The naming rules are similar to alkanes.

Try to get the lowest number for indicating the double or triple bond position. An alkene has a double bond.



For example,
             
CH2 = CH - CH2 - CH3

would be butene. Notice that the name will end in -ene.

CH2 = CH2  would be ethene

The general formula is C(N)H(2N)

Some general rules for naming:
1) Find the longest chain. This will be the part that makes up the end of the name.
2) Number the carbon atoms to get the lowest number for the start of the double bond and place this number before the parent name.
3) Write the names and numbers for the side groups, assembling them alphabetically.

Some alkenes have the same structure, but a different geometry (or a different shape). These are called geometric isomers.

We use trans- or cis- to distinguish between these isomers.
If two adjacent carbon atoms are bonded by a double bond and have side chains on them, two chains are possible.

If the larger groups are both on the top or on the bottom, then it is cis.
If the larger groups are diagonal to each other, or across from each other, then it is trans.
If the larger groups are one on top of the other, then there are no isomers, and therefore no cis or trans.



Now, let's do alkynes. These are the same as alkanes or alkenes in terms of naming, but these just end in -yne.

Alkynes have one or more triple bond between carbons which lead to unsaturated hydrocarbons.
For example:
CH ≡ CH would be ethyne.

The general formula is C(N)H(2N-2)
There are no cis or trans forms since they are mostly the same.

Now, let's try a few examples.



The longest chain is 6, and the position of the triple bond is 2, counting from right to left. This is because we get the lowest number possible.

First, write down 2-hexyne. There is a methyl group on the 5th position (since we are already counting right to left, we must stay consistent)

Finally, we end up with 5-methyl-2-hexyne.

Here's a video on alkene.



Alkynes:



http://herh.ccrsb.ca/staff/FarrellL/chem11downloads/org/org2.pdf

Organic Chemistry

Now, we will start a new section on organic chemistry. Organic chemistry is the chemistry of carbon compounds. This is responsible for many of the everyday products that are used around the world, such as polyester, alcohol, cocaine.



Properties of organic compounds are low melting points, with low electricity. The carbon atoms can be linked in straight lines (which are really zigzag), circular structures, or in branches. As well, they can be linked in single, double or triple bonds.

First, we will start with alkanes. These are unbranched, straight chain compounds that only contain H and C. These are all hydrocarbons that are bonded together in single bonds. The names will end in "-ane" because they belong to the "alkane" group.



For the molecular formula C2H6, the full structure is:
      H   H
       |     |
H - C - C - H
       |     |
      H   H

As you see, the two carbons are bonded together, and the remaining bonds are with hydrogens.

The formulas for regular alkanes are:

Methane CH4              Hexane C6H14
Ethane C2H6               Heptane C7H16
Propane C3H8             Octane C8H18
Butane C4H10             Nonane C9H20
Pentane C5H12           Decane C10H22

For alkanes, the general formula is C(N)H(2N+2).

When hydrocarbons have side branches, they are called substituted hydrocarbons or branched hydrocarbons. (If you refer to the diagram above, this means that one of the Hydrogens surrounding the C is replaced with, say, a methane)

Carbon has 4 bonds. When a carbon is attached to another carbon, there are 3 hydrogens surrounding the   original carbon. When carbon is attached to two other carbons, there are 2 hydrogens. When attached to 3 carbons, there is 1 hydrogen left. When carbon is attached to 4 carbons, there are no hydrogens left.

For example, if a carbon is attached to another carbon, we would get:
                                              H
                                               |
  CH3                              H - C - H
       |                                       |
H - C - H               or      H - C - H      (the CH3 is the same as an added carbon, with 3 hydrogens
       |                                       |             branching out)
      H                                     H

The CH3 is supposed to be methane, but one H got bonded with the original H that was there, and so CH3 is left.

To name this, we would say this is ethane. We will go through this more later.

Suppose there is a name: 2-methylpentane



The methyl part is an alkyl group, which is an alkane which has lost one hydrogen atom.
The parent hydrocarbon is pentane (which is the longest straight chain, as there are 5 carbons).

The names of the alkyl groups end in "yl" because they are alkyl. If there are 2 or more of the same kind of alkyl group, use prefixes like di, tri, tetra, penta.



If more than one alkyl group (as in different kinds of alkyl groups) are present, list them alphabetically, and put the position number in front, with a dash between each group and number.

Here's a video:



http://www.arps.org/users/hs/thompsom/chemcom/unit_3/Naming_Alkanes_Worksheet_1.pdf
http://www.grossmont.edu/martinlarter/chemistry141/reference/AlkaneWorksheet.pdf

Chemical Bonding

As you know, chemical bonding only involves valence electrons, and compounds continue to gain/lose or share electrons until they have a full closed shell.

If electrons are shared equally, we say the covalent bond is non-polar. If they are shared unequally, a polar covalent bond is formed.

But when electrons are transferred, it is totally different, and it is an ionic bond.



Recall from the periodic table trends the term electronegativity. It is the tendency to attract electrons, with non-metals having a high number of electronegativity.

In non-polar bonding, equal sharing is observed to get a full shell. Electrons are attracted to the nuclei, and there are some in between the two atoms. The bonds are very strong, and require a high amount of energy to break them.

Intramolecular forces hold atoms of the molecule together. Intermolecular forces bond the molecules together. When melting occurs, the bond between the atoms is not broken, and only the intermolecular, weak bonds are affected.

In polar bonding, atoms with greater electronegativity pull electrons toward itself, so the shared electrons will actually be closer to this atom.



Just remember, atoms with the higher electronegativity form a PARTIAL NEGATIVE CHARGE, while the lower electronegativity atom will form a PARTIAL POSITIVE CHARGE.
Then, add an arrow to indicate which way the electrons will tend to move.  (in other words, draw the arrow pointing to the partial negative atom)

E.g. What is the bond between C and O?

To do this, we will use a Table of Electronegativities. If the difference is < 0.5, it is covalent. If the difference in EN is >0.5 and <1.8, it is polar covalent. If the difference is >1.8, it is ionic.



So, the EN of C is 2.55, and the EN of O is 3.44, giving a difference of 0.89, which is polar covalent.

This means our diagram would look like:

C δ+ ---------> O δ-


Here's a video:



On related videos, there are continuations of this video.

Here is some practice:
http://chemistry.sswiki.com/file/view/8.4+Review.pdf
http://www.dorjegurung.com/chemistry/IB_year1/worksheets/wkst_hybridization_shape_polarity.pdf

Test next day!

Today, we just had a review day for our test on Monday on Atomic Theory and Periodic Table and trends.

Good luck to all!

Electron Dot Diagrams

How to draw them

Drawing electon dot diagrams (or Lewis diagrams) are quite easy if you follow these steps/rules.

1. The nucleus is represented by the atomic symbol
2. Each electron is represented by a dot which are around the atomic symbol
3. There are four orbitals on each side of the nucleus. Each orbit can hold a maximum of 2 electrons
4. 8 electrons represent a closed shell, except for H, which needs 2 electrons to become stable
5. Determine the # of valence electrons

For example sodium is in group 1 therefore it has 1 valence electron

Here are some examples that include double/triple bonding

There are many types of electron diagrams, such as Bohr, Lewis or structural.

In a structural diagram, each bonded pair is represented as a line.
Here's an example:

Now, if we want to represent ionic bonds, we would look at the charge of the ion. For example, for NaCl:

You would put the negative ion in square brackets, and write the positive ion right beside it. So, you would write Na+ [Cl]- (but you would include dots around the four orbits to symbolize that the compound has satisfied the octet rule).

In Lewis structures, the number of bonds is shown, along with the lone paired electrons, which are electrons that do not participate in sharing.

But how do we determine which element is the central atom?
Here are a few rules to follow:
1. H and F are never in the middle
2. If there is a metal, the metal is always the central one
3. If a molecule contains only one atom of one element, and many of another, then the single atom is in the centre.
4. Atoms that need the most electrons to complete their valence shell is in the centre.

Let's try an example:
C2H2:

First, how many valence electrons do we have?

4x2 + 1x2 = 10

Now, we have two C's in the middle, with a bond in between them.

C - C

Then, we have two H's branching out from the C's.

H - C - C - H

That makes 6 electrons, so we have to distribute 4 more to the C's (since H's are already full)

If we distributed them evenly, we would have two lone electrons on each of the C's. (preferably one on each of the orbits)

Then, if we took one of the electrons and bonded it with an electron from the other C, we would have a double bond between the C's and one lone electron remaining on each C.

Now, if we take those ones and bond them together, we will successfully get our full shell, with a triple bond between the Carbons.




Here's some practice:
http://misterguch.brinkster.net/PRA017.pdf
http://www.arps.org/users/hs/thompsom/honors/chap_07/Lewis_Structure_Worksheet_1.pdf

Bohr Model Diagram

The Bohr model was approximately correct to what we know today, it was close in the theory of quantum mechanics. Bohr proposed that electrons occupied shells, otherwise known as orbitals. He believed that electrons had certain "energy levels", the lowest energy state being the |ground state. Bohr thought that the electron can "jump" to higher levels when they are "excited"and vice versa to when the electrons fall to the lower lessons. He thought that each "jump" or "fall" would give off a quantum of light energy, emitting a spectrum of light.

 

In the first shell, only 2 electrons are able to occupy it. The second shell there can be 8 electrons, and the third shell, there can also be 8 electrons, and this is called the octet. Bohr also suggested that electrons were not able to move freely in the atom and that when electrons are heated, they will give off a certain wavelength that is unique for each element.

Bohr also wrote something in the middle of the diagram. He wrote the number of protons and neutrons in the atom. For example, the diagram above is a representation of a Boron atom. The Atomic number is 5, meaning the proton is also 5... So in the middle it says P:5. The atomic number of this element is 10.8, and to get the neutrons you do 10.8-5  which equals 5.8, which is 6 when rounded. There fore you write N:6 in the middle. After drawing the thing in the middle, don't forget to draw the electrons that are in the orbitals. In this case, there are 5 electrons, so you put 2 electrons in the first shell, and 3 in the next.

and... BAYUM u just got yourself a Bohr Model Diagram! how easy was that! YAY! success ^^ WHOOP WHOOP!! :)

For more practice or more knowledge on Bohr Diagrams.. refer to these sites/videos:

http://needham.wikispaces.com/file/view/03+BohrModelPractice.pdf
http://pw.vsb.bc.ca/wyper/sci9/2-3_bohr_diagram_worksheet.pdf

Periodic Table Trends

In this blog, we will cover some patterns and trends in the periodic table.

Some of the patterns that we will discuss include
1. Metallic properties
2. Atomic Radius
3. Ionization energy
4. Electronegativity
5. Reactivity
6. Ion charge
7. Melting/boiling point
8. Density

First, let's start with metallic properties.
Although this one seems obvious, it is still good to point it out. When moving from left to right in the periodic table, the properties of the elements change from metallic to non-metallic.



Also, when going down a family (column), elements become more metallic, or better metals.

Secondly, let's study about the atomic radii.
When moving from left to right, the atomic radii decreases. This is because the atomic number and the positive charge increase. The increase in atomic number means an increase in both electrons and protons, making the force of attraction much stronger, and decreasing the distance between each other.



When going from up to down in a column, the radii will increase. This is because there will be more orbits taken up by the electrons. The inner electrons also repel the outer electrons, increasing the distance between the outermost electrons and the nucleus.

To recap atomic radii: left to right, decrease; up to down, increase.

Now, ionization energy. First, let's define the term. Ionization energy is the energy required to remove an electron from the neutral atom. Measured in kJ/mol, IE is basically the opposite of the atomic radius. Helium has the highest IE, and Francium has the lowest IE.



Usually, the outermost electron is removed, and so it should always be a valence electron being removed, unless it is a closed shell.

When moving left to right across a period, ionization energy increases. This is because the radii has been decreased, meaning a very strong attraction between the electrons and the nucleons (protons and neutrons). This means that it will be harder to remove an electron.

When moving up to down, the ionization energy decreases due to a larger radius. Attraction between the nucleus and the outer electrons is also decreased because there are more orbits in between blocking the way.

We can refer to removing the first electron as the first ionization energy, and removing the second as second ionization energy, etc.

Number 4, electronegativity. Definition: how much atoms want to gain electrons, or tendency to attract electrons from a neighbouring atom.



This means that if an atom has high electronegativity, it strongly attracts electrons from a neighbouring atom, and could almost "steal" an electron from its neighbour. As well, this means that the atom has a strong attraction with its valence electrons, so the electrons are harder to remove and thus has a higher ionization energy.

In the same way, lower electronegativity = lower ionization energy, because the atom does not have a strong attraction with its own valence electrons, then electrons are easier to remove.

The top right (Fluorine), except for noble gases, has the highest electronegativity.

To recap: left to right, increase in electronegativity; up to down, decrease in electronegativity.

5. Reactivity
In metals:
left to right, decreases
up to down, increases, because it is easier for electrons to be given away, meaning higher reactivity.

In non-metals, left to right, increase
up to down, decreases, because non-metals have higher electronegativity.

6. Ion Charge
The charges depend on its group.
For example,
group 1, +
group 2, 2+
group 13, 3+
group 14, 4+
group 15, 3-
group 16, 2-
group 17, 1-
group 18, 0
In the transition metals, the charges are variable.

7. Melting/boiling point
The noble gases have the lowest melting points, and the elements in the center have the highest. Melting point increases from left to right, except for in the middle.
In metals, going down the group decreases the melting and boiling points.
In non-metals, going down the group increases the melting and boiling points.



8. Density
As you go down a group in the periodic table, density increases. This is because, as you go down, atomic radius increases, and volume increases, so it will be more dense.

In summary,


http://mysite.oswego308.org/schools/uploads/files/3231/ws_periodic_table_and_trends.pdf
http://butane.chem.uiuc.edu/cyerkes/Chem102AEFa07/worksheets/Worksheet%2012.pdf

Quiz day and Periodic Table studying

Today, we had a short quiz on Atomic Theory and Electron Configuration. Then, we headed to the Curriculum Lab to study some trends in the Periodic Table.

We were given information about each element, such as density, melting/boiling points, ionization energy, electronegativity, and atomic radius, and we were to graph the properties and determine general trends found in the periodic table.

We will discuss these trends in our next blog! :)

Here is a good website to check out if you want to go ahead and look at some periodic trends!
http://www.chemguide.co.uk/atoms/propsmenu.html#top

Predicting Valence Electrons

Valence electrons are located in the outermost open shell of an electron. These take part in chemical reactions.

What we mean by open shell is that the shell contains less than the maximum number of electrons.
In contrast, a closed shell would be completely full.



The valence electrons are all the electrons in the atom except in core and filled d/f subshells. In other words, the valence electrons are the s, p types of the outer set, and ONLY unfilled d and f subshells.

For example, if we were given the configuration of Calcium as:

1s2 2s2 2p6 3s2 3p6 4s2

To see the valence electrons easier, put that configuration in Core Notation:

[Ar] 4s2

We see there are only 2 electrons in the outer set, and they meet the rules, since it is an s-subshell.

For a harder one, let's do Pb.

[Xe] 6s2 4f14 5d10 6p2

How many valence electrons are there?

If you got 4, you are correct. This is because there are 2 from the s-subshell, and 2 from the p-subshell. We do not count the f and d, because they are FILLED.

And that is all for this short section! :)

History of the Periodic Table

In 1863-1866, John Newlands found that every eighth element, starting with hydrogen first, shared a common set of properties.



In 1869, Dmitri Mendeleev listed the elements according to the masses and showed that certain properties recur, breaking the list into rows (periods) and columns (groups). He left gaps in his periodic table, and suggested those elements had yet to be discovered. He could also predict the properties and characteristics of these undiscovered elements.





Now, the modern periodic table is organized in atomic number. The Periodic Law suggests that properties of elements recur periodically when elements are arranged in atomic number.

The families (or groups or columns) include:
Alkali metal, alkaline earth metals, halogens, and noble gases.

The rows underneath the table are called Lanthanides (starting with Lanthanum) and Actinides (starting with Actinium). These rows are considered as inner transition metals. The main transition metals are in between the metals and non-metals, and the staircase starting in group 3 signifies metalloids, which have both metal and non-metal properties.

Metals are opaque and shiny. They are good conductors of heat and electricity, and can be hammered into sheets (malleable) or drawn into wires (ductile). They are usually solids at room temperature (with Mercury being the exception), and they tend to lose electrons.



Non-metals are gases and liquids at room temperature. They are poor conductors of heat and electricity. Some non-metals are brittle solids, and are dull in appearance and opaque.



Semiconductors are non-metals with electric conductivity. They are also called metalloids, and their properties resemble metals more than non-metals. The difference is that metal conductivity decreases with increasing temperature, but semiconductor conductivity increases with increasing temperature.



Here's another summary:
http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/metal.html

Writing Electronic Configurations of Neutral Atoms

We split this section into 2 parts, since it was getting very long.


Now, we will learn how to write configurations of neutral atoms.


1. Figure out how may electrons there are. (look at atomic number) Start at 1s and keep adding until there are none left.
2. When drawing arrows to show electron spins, always draw the upwards first, then downwards. You can have more upwards than downwards. (if this is confusing, read on and we will explain later)


For example, let's do Na.


Na has the 11th atomic number, and in a neutral atom, it has 11 electrons.


We always start with 1s. If we count to 11, we realize that there will be 2 in 1s, 2 in 2s, 6 in 2p, 1 in 3s, because 2+2+6+1 = 11.


If we want to visualize it:
1s   2s          2p             3s
↑↓   ↑↓       ↑↓ ↑↓ ↑↓         ↑


Each arrow represents one electron, and counting the arrows, there are 11.

Notice that the electron in 3s is not paired because there are not enough arrows. 
In written form, we would write the configuration as 
1s2 2s2 2p6 3s1 
The superscript just tells us how many electrons are in each subshell. Notice that the first 3 subshells are full, and the last one (3s) just has 1 electron. If the element was Magnesium, then it would be full, since magnesium has one more electron than sodium. 


Now, let's write some configurations for ions. This is pretty similar to above, but it just involves one extra step. For example, if we were given Cl-


The - just means that we need to gain an extra electron to make it full. First, just do the electronic configuration normally, as if the - charge did not exist. (in other words, do it for 17 electrons)




We would get 1s2 2s2 2p6 3s2 3p5. In arrows, it would be:
1s  2s         2p       3s        3p
↑↓  ↑↓   ↑↓ ↑↓ ↑↓   ↑↓    ↑↓ ↑↓ ↑↓
The 3p is not full, as there is one more spot available. Now, we look back at the charge. - means to add an electron, so we would add one to the last unfilled subshell. (in red)


For a + charged ion, we would still write it out as if it were neutral, but then removing electrons from the outermost. 


Now what if we were asked to do the configuration of an element with an extremely high atomic number? It would take forever! There's a shortcut called Core Notation.


The set of electrons can be divided into two subsets: core and outer. The core of the atom is the set of electrons with the configuration of the nearest noble gas with an atomic number lower than the element in question. 
The outer is the electrons outside the core. The core electrons normally take part in chemical reactions.



Step 1: locate the noble gas above the element. If we had Nickel, we would not use Krypton, even though it is closer. We would use Argon instead.
Step 2: Square bracket the noble gas configuration.
Step 3: Start from the noble gas and add electrons until you reach the element. 


For example, Chlorine. 
The noble gas above chlorine is Neon, which has 10 electrons. 
In written form, the configuration for chlorine would be: 


1s2 2s2 2p6 3s2 3p5. Now, we can replace the first 10 electrons with Ne. This means that 1s2 2s2 2p6 will be replaced. 


Core Notation for Chlorine: [Ne] 3s2 3p5
That seems so much simpler and cleaner than the entire written configuration, right?
However, you don't want to write it out fully, and then replace it. Instead, realize that Ne goes up to 2p6, and just write [Ne] first, then add the remaining electrons. 



There are two exceptions for electronic configurations. 
In copper, you would expect 1s2 2s2 2p6 3s2 3p6 4s2 3d9 or [Ar] 4s2 3d9. However, it is actually 1s2 2s2 2p6 3s2 3p6 4s1 3d10 or [Ar] 4s1 3d10. This is because copper likes to gain stability will a full d-subshell. 


In chromium, we expect [Ar] 4s2 3d4, but instead, we actually get [Ar] 4s1 3d5, because chromium likes to gain stability with a half-full d-subshell. 


Here's a worksheet:
http://www.chemteam.info/Electrons/WS-Configs&light.pdf
http://www.everettcc.edu/uploadedFiles/Student_Resources_and_Services/TRIO/Electron_configurations_wksht.pdf

And a video!

Electron Configuration, with Core Notation

In this lesson, we will introduce electronic configuration.

Electronic configuration is a notation that describes the orbitals in which electrons occupy, and it also gives the total number of electrons in each orbital.

In other words, it is just another way of showing which level electrons are in, just like Bohr's model.



For example, the electronic configuration for an element like carbon would be:

1s2 2s2 2p2

Now don't worry if that confuses you, because by the end of this blog, you'll know what they mean!

According to scientist Niels Bohr, electrons existed in specific energy states, and when an electron absorbs or emits energy, it can move from one level to another.

The scientific definition of an energy level is the maximum amount of energy an electron can hold.

When electrons are in "ground state", it means that all the electrons are in their lowest possible level. But when they are in an "excited state", it means that one or more of the electrons are in energy levels other than the lowest available level.

To write electron configurations, you should know these things:
1. Orbital - a region of space occupied by the electron
2. Shell - the set of all orbitals in the same energy level
3. Subshell - the set of orbitals of the same type.
4. There are four types of subshells (s, p, d, f)



s is like the starting point. In the very first energy level, only s-types exist. Since there are in energy level one, we can say they are 1s.
In the second energy level, s and p types exist. They are in energy level 2, so they are 2s and 2p.
In the third, s, p and d types exist (3s, 3p, 3d), and in the fourth, s, p, d, and f exist (4s, 4p, 4d, 4f).

In every orbital, there are 2 electrons.
An s-type consists of one orbital (2 electrons)
p consists of 3 orbitals (6 electrons)
d consists of 5 orbitals (10 electrons)
f consists of 7 orbitals (14 electrons)



The order of writing the levels are like the picture above. Arrange the orbitals like a descending staircase, with 1s at the top.

However, in a ground state, we must remember that we can only move on to the next level when the orbitals are full. Therefore, 1s2 should be at the top. The next step in the staircase would be 2s2 2p6.

Here is a good video of electron configuration. Unfortunately, the maker of the video disabled embedding, so here's a link.
http://youtu.be/xH1k1dtgiVY

Here's a good practice worksheet; however, you will be able to do the ones that ask for abbreviated configurations after you read the next blog!
http://misterguch.brinkster.net/PRA014.pdf

Atomic Structure

In the atom, there are 3 subatomic particles: neutrons, protons, and electrons.



Neutrons have a neutral charge, protons have a positive charge, and electrons have a negative charge. The neutrons and protons are in the nucleus, but the electrons are actually surrounding the nucleus in energy shells.

When an atom is neutral, it means that there is no net charge. This also means that the number of protons will equal to the number of electrons.

To find the number of protons, simply look at the atomic number of the element.



That means, if you add a proton to an element, it will become a totally different element.
Say, we have 4 protons in the nucleus of some element. This would be the element Beryllium. If we added a proton, then our element would change to Boron.

What if an atom does have some kind of charge? Then, it is called an ion. Atoms can gain or lose electrons by accepting or giving electrons to other atoms. If we are given the number of protons and the charge, we can find the number of electrons by

proton - charge

A negatively charged ion is called an anion. An anion has more electrons than protons, because non-metals tend to gain electrons.

In contrast, a positively charged ion is called a cation. This is when there are fewer electrons than protons. Metals are positively charged and tend to lose electrons.



The mass number is the total number of protons and neutrons, or it can also be found by rounding the atomic mass number to the nearest whole number. To calculate the number of neutrons in an atom, we can use the mass number and subtract from it the number of protons.

So: Mass number - atomic number (number of protons) = number of neutrons

Atomic mass is different from the mass number. The atomic mass is the average mass of an element's isotopes. They have decimal values because they are averages; however, the mass number is just a rounded number of the atomic mass.

If you add a neutron into the atom, you will get a heavier version of the element.

Now, let's try an example.
How many protons, electrons, and neutrons are in Co?
proton: 27
electron: 27
neutron: 59-27=32

An isotope of an element has the same atomic number, but a different number of neutrons or atomic mass.



Remember back in around December, when we used the atomic mass as the molar mass?
Well, now we can understand that this mass is really an average of value of a combination of isotopes.

Let's try to find the average atomic mass, given the data of naturally occurring isotopes.

Ne-20 (90.02%) Ne-21 (0.26%) Ne-22 (9.72%)
All we have to do, is just multiply the percentages to the masses of each of the isotopes and add them up.

So: 0.9002 X 20 + 0.0026 X 21 + 0.0972 X 22 = 20.197

We have 2 sig figs, so we will get the final answer of 20.

Now here are a few links for practice:
http://cmsweb1.loudoun.k12.va.us/52820831134912597/lib/52820831134912597/Atoms%20and%20Atomic%20Theory/Homework/ws.atomic.20structure.20set.pdf
http://www.scribd.com/doc/3370461/atomic-structure-worksheet

Here's a video:

Atomic Theory

In this next chapter, we will cover Atomic Theory and the Periodic Table and trends.

First, let's look at some history behind the Atomic Theory.

In the ancient times, Greek philosophers suggested that matter was made of atomos, which means the smallest pieces of matter.

At around 400 BC, Democritus was the first to propose this, and that atoms were indivisible particles.



Aristotle, a famous philosopher at the time, disagreed with this theory and thought that matter was made of earth, air, fire and water. This theory became true for 2000 years, but since it was only conceptual (much like Democritus's), scientists were able to disprove the theory later on.



In the 1700s, a French chemist named Lavoisier stated the Law of Conservation of Mass and the Law of Definite Proportions. These laws suggested that in a compound of say, H2O, there will always be 11% Hydrogen and 89% Oxygen.



In 1799, Joseph Proust experimentally proved Lavoisier's laws, and added that when a compound is broken down, products will exist in the same ratio as in the compound.



Then, in the early 1800s, John Dalton developed the basis of the modern Atomic Theory. He suggested that:

1. Elements were made of tiny indestructible spheres called atoms.
2. All atoms of an element were the same.
3. Atoms of a given element can be differentiated from another element by its relative atomic weights.
4. Atoms of one element will combine with atoms of other elements to create compounds.



In the 1850s, J. J. Thomson created an experiment known as the Raisin bun model (or the Plum Pudding). This model had solid, positively charged spheres, as well as negatively charged spheres. Using this model, he proposed that a unit over 1000 times smaller than the atom, known as the electron, existed.



A student of J. J. Thomson, Ernest Rutherford discovered that atoms have a positively charged, dense center with electrons surrounding it on the outside. He explained why electrons spun around the nucleus, but he could not explain why the electron did not fall into the nucleus and destroy the atom.





Luckily, a partner of Rutherford's, Niels Bohr found the solution. Bohr was studying gaseous samples of atoms at the time, and came to the conclusion that electrons surrounding the nucleus were in specific energy levels. When the electron was excited, it would jump to a higher level. When an electron came back down, it would release energy in the form of light. Each of these jumps gives off light in different wavelengths; therefore creating different colours, as the colours ROYGBIV all have different wavelengths.





In conclusion:
An electron surrounded the nucleus, and had a negative charge. If its mass was 1, then the mass of protons would be 1836, and the mass of neutrons would be 1837. Protons, positively charged particles, along with neutrons, neutral particles, were in the nucleus of an atom.

This is a really good website to check out:
http://chemistry.learnhub.com/lesson/3663-history-of-the-atomic-theory-i-ancient-times
Now a video to summarize!!