Drawing Electron Dot and Lewis Diagrams

When drawing electron dot or Lewis Diagrams, always keep in mind that the nucleus is represented by the atomic symbol.

Dots represent electrons in electron dot diagrams.

Lines represent the bond of two electrons in Lewis Diagrams - also known as structural diagrams.

There are four orbitals - at the north, east, south, and west sides of the nucleus as shown.

Now in a covalent compound:

While hydrogen has one valence electron, chlorine has 7. When they are drawn like this, it means that hydrogen and chlorine are sharing the two electrons between them. These two electrons can also be represented as a horizontal line in a structural diagram.

Now in ionic compounds:

In the first example, Na has lost an electron and is now an ion with a +1 charge - it now has an empty valence shell and is stable. Cl has gained this electron and is now an ion with a -1 charge - it now has a full valence shell and is stable. These two ions are held together by their opposite attraction.


Diagrams for Polyatomic Ions

Because polyatomic ions have charges - electrons must be added or taken away. I find that it is easiest to do this after we have paired up the existing electrons. In the above example, we see CO3. After pairing up carbon's four valence electrons with those of the two oxygens on the left and on the right, we see that we need two more electrons in order to pair with the third oxygen on top. The 2- subscript means that the two electrons have been added.

Enjoy!

Periodic Trends

So what are periodic trends?
They are progress of increasing or decreasing with the charactierstics of certain elements.

They're are a lot of trends and we must know them all:
1)Metallic Properties
2)Atomic Radius
3)Ionization Energy
4)Electronegativity
5) Reactivity
6)Ion Cahrge
7) Melting and Boiling Point
8) Density

Here are some of the trends...let's start of with...
Metallic Properties

The elements change from left to right arosss the table and elelements that are more metallic moves down the table.

Atomic Radius:

The atom decreases aross a row from left to right and increases going down a family. When the atom is going from left to right the atomic number increases along with the nucleus.
When the atoms are going down a family the orbits are more compact and therefore the inner electons attract each other whereas the elctrons on th outer end repel each other.

Reactivity:

Metals and Non-metals have and show different trends. When ther metal moves down a family and right a row it is more reactive and when going up and left it is less reactive.
Non-metals are the opposite or metals, it increases or it's more reactive as an atom goes up a family.

Melting and Boiling Point:

The elements from the center has the highest melting point and the Noble Gases has the lowest melting point.

Ionization Energy:

Ionization Energy are enegies needed to remove one elctron from an atom. They increase as it does up and right in t he periodic table and all the noble gases have high ionization energy. The elctrons have these atoms remove eaily. When going left to right it has a higher attraction between the nucleus and the elctrons whereas the outer electrons decreases as it goes down a group.

Electronegativity:

Electornegativity is how much atom wants to gain electrons. When the atoms goes left to right and up they increase and decreases going down a group Atoms with high electornegativity are strognly attracts other elctrons from another atom or it may remove an electron from a neighbor.

Periodic Trends

In 1870, Dmitri Mendeleev was the first person to decide to put all the elements on a "table" so that it could be more organized and easier to study. Since then, it has the periodic table has been made better along with additional elements. 

To find out what the periodic trends are, we have to graph it and study them carefully.

Density: After graphing, you should be able to find that density usually increases along with the atomic number.

Melting and Boiling Point: After graphing, you should be able to find out that the melting and boiling points for metals usually increase from bottom to top of a group, and that for non-metals, they have an increase for melting and boiling points.

Ionization Energy: After graphing, you should be able to find out that it increases in a period (left to right), and it decreases as a group (top to bottom).
Electronegativity: After graphing, you should be able to find that it increases from left to right.

Atomic Radius: After graphing, you should be able to find that it decreases as a period and increases as a group.

How many valence electrons?

Valence electrons:
1) Exist on the outermost shell of an atom
2) Take part in chemical reactions
3) Are all electrons except those IN THE CORE, THOSE IN FILLED D- or F- SUBSHELLS

Here is some important terminology:
An open shell is a shell that has fewer than the maximum amount of electrons occupying it. In this picture, the 2nd or outermost shell of the oxygen atom is the open shell. While it can hold 8 electrons, only 6 occupy the shell.

A closed shell is a shell that contains the maximum number of electrons that it can hold. In this example, the closed shell is the first one. It contains two electrons - the most that it can hold.


Let's look at the core notation for SULPHUR
S = 1s^22s^22p^63s^23p^4 CORE: [Ne]3s^23p^4

From the core notation, we look at the portion that is not within brackets. We then add up the superscripts of the s and p orbitals and get 6. Thus, we can conclude that there are six valence electrons in sulphur. We don't add up the superscripts of the d- or f- orbitals unless they aren't full.

Some unrelated terminology on the periodic table.

Periodic Law - the properties of certain chemical elements that occur every so often when the elements are placed in their order.
Ianthanides - the first row under the table that starts with lanthanum.
Actinides - the row under Ianthanides.

Electronic Structure of the Atom!

What is the electronic structure of an atom? It's a notation that describes the orbitals in which electrons occupy and the total number of electrons in each orbit. 

Remember back to when we learned about Neils Bohr? The guy in that picture? He proposed that electrons exist in specific energy states and when electrons absorb/emit specific amount of energy, it instantaneously moves from one orbital to the next.


Now for some vocab:
Energy level- amount of energy which electrons in atoms can possess ( "n"= number of energy levels) 
Quantum of energy- energy difference between 2 particular energy levels
Ground state- when electrons of atoms are in their lowest possible energy level
Excited state- when 1 or more of an atom's electrons are in energy levels other than the lowest available level
Orbital- region of space occupied by an electron in particular energy level
Shell- set of all orbitals having the same "n" value
Subshell- set of orbitals of the same type

Orbitals are split into 4 different types: s, p, d, f
Each subshell consists of:
1: s-orbital
3: p-orbital
5: d-orbital
7: f-orbital


Pauli Exclusion Principle: maximum of 2 electrons can be placed in each orbit
This principle means that the maximum number of electrons in each subshell is:
2: s-subshell
6: p-subshell
10- d-subshell
14- f-subshell

This picture will help with writing the electronic configuration of atoms. For neutral atoms:
1) Always start with lowest energy level (Aufbau principle)
2) Figure out how many electron you have (neutral atom=atomic number)
3) Start at lowest energy level (1s) and keep adding on until none is left.

Each electron has an apposite spin designated by upward and downward arrows. An example is silicon. The last line of the picture shows the electronic configuration. Silicon has an atomic number of 14 thus it has 14 electrons at its neutral state. Notice that 2 electrons in 3p occupy separate orbitals and aren't paired? This is due to...
Hund's rule- when electrons occupy orbitals of equal energy, they don't pair up until they have to. 
So in written form: 1s^22s^22p^63s^23p^2
The exponents represent the number of electrons.




Now for writing electronic configurations for ions. 
For negative ions: 
Add electrons (equal to charge) to last unfilled subshell starting where neutral atom left off. An example is N -3
Nitrogen in its neutral state has 7 electrons but with the charge of -3 it is actually 10


For positive ions:
1) Start with neutral atom and remove electrons from outermost shell depending on the charge.
2) If there are electrons in both s and p orbital of the outermost shell, elextrons in the p-orbital are removed first.


Core Notation- way of showing electron configuration in terms of core and outer electrons
A set of electrons belonging to a given atom can be divided into 2 subsets: 
1) Core- set of electrons with configuration of the nearest noble gas (He, Ne, Ar, Kr, etc.) having an atomic number less than that of the atom being considered
2) Outer- consists of all electrons outside the core. Since core electrons, normally don't take part in chemical reactions.


So to find the core notation:
1) Locate the atom and note the noble gas at the end of the row and above the element
2) Write the electron configuration normally. BUT replace the part of the electron configuration corresponding to the configuration of the noble gas.
3) Write the noble gas in square brackets. Ex. [Ne]...
4) Follow the core symbol with the electron configuration of the remaining outer electrons.
Ex. Al: 1s^22s^22p^6 3s^23p^1      
      Ne: 1s^22s^22p^6 Bolded is the core. The rest is the outer. Therefore the core notation is: [Ne] 3s^23p^1


Exceptions!
Instead of: Cr --> [Ar] 4s^23d^4 (One electron short of half-filled subshell)
Cu--> [Ar] 4s^23d^9 (One electron short of a filled subshell)
In actuality:
Cr--> [Ar] 4s^13d^5 (Now both subshells are exactly half-filled)
Cu--> [Ar] 4s^13d^10 (4s^1 is exactly half-filled and 3d^10 is filled)
These 2 atoms indicate: A filled or exactly half-filled d-subshell is especially stable.


Link to awesome quiz to test your genius knowledge. :D

HAHAHA. SO FUNNY.

Modern Atoms

Modern atoms consist os electrons, protons and neutrons.

Remember the characteristics of each atom?
Proton ,  , mass = 1, charge = +1, located in the nucleus
Neutron,  ,mass is >1, charge = 0, located in the nucleus
Electron,  , does no = to 0, charge = -1, located around the nucleus

The # of protons = # of electrons

Atomic number (Z): number of protons found in a nucleus
  - the atomic number  = the number of protons = the number of electrons

Ions: atoms able to gain or lose electrons
  - # of protons - # of electrons = the charge

Electrically charged atoms:
 - negative charged ion are electrons added to a neutral atom
  - positive charged ion are electrons removed from a neutral atom

Mass number: the total number of protons and neutrons
  - atomic mass = # of protons + # of neutrons
  - number of neutrons = mass # - atomic #
  - mass # = # of protons + # of neutrons

Atomic mass: the average of an isotope

Isotopes: same number of protons and electrons but different number of neutrons

Natural Mixtures of Isotopes:
We have an element with 4 naturally occuring isotopes. Here are their atomic masses and percent abundance: a = 46 (25%), b = 47 (50%), c = 48 (15%), and d = 49 (10%). Find the atomic mass!

46 x 25% = 11.5

47 x 50% = 23.5

48 x 15% = 7.2

49 x 10% = 4.9

11.5 + 23.5 + 7.2 + 4.9 = 47.1 atomic mass units

You try some with the link below!
http://www.chem4kids.com/activities.html

Atomic Theory.

Aristotle: -believed that atoms were made from the 4 elements; earth, air, fire, and water Democritus: -in 300B.C, he believed that atoms were indivisible particles
Lavoisier (late 1700's) -stated earliest version of both the Law of Conservation and the Law of definite proportions
Proust (1799) - Proust believed that if a compound were to be broken down, it would still contain the same ratios as in the compound
Dalton (early 1800's) - believed that atoms are solid
-all atoms of given element are identical
-atoms of one element can combine with another to form chemical compounds
J.J. Thompson (1850) - proved existance of electrons
Rutherford (1905) - showed atoms had dense centres and electrons outside of them
-suggested atoms are mostly empty space
Niels Bohr (1885-1962) - studied gaseous samples of atoms
-proposed that electrons surrounds the nucleus in shells
The modern atom - an atom is the smallest particle of an element
-all atoms are made of three kinds of subatomic particles
         -electrons
         -protons
         -neutrons