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: