Electronegativity & Polarity!

Group

Period

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

1

H 2.1

He 0

2

Li 0.98

Be 1.57

B 2.04

C 2.55

N 3.04

O 3.44

F 3.98

Ne 0

3

Na 0.93

Mg 1.31

Al 1.61

Si 1.9

P 2.19

S 2.58

Cl 3.16

Ar 0

4

K 0.82

Ca 1

Sc 1.36

Ti 1.54

V 1.63

Cr 1.66

Mn 1.55

Fe 1.83

Co 1.88

Ni 1.91

Cu 1.9

Zn 1.65

Ga 1.81

Ge 2.01

As 2.18

Se 2.55

Br 2.96

Kr 0

5

Rb 0.82

Sr 0.95

Y 1.22

Zr 1.33

Nb 1.6

Mo 2.16

Tc 1.9

Ru 2.2

Rh 2.28

Pd 2.2

Ag 1.93

Cd 1.69

In 1.78

Sn 1.96

Sb 2.05

Te 2.1

I 2.66

Xe 2.6

6

Cs 0.79

Ba 0.89

La 1.1

Hf 1.3

Ta 1.5

W 2.36

Re 1.9

Os 2.2

Ir 2.2

Pt 2.28

Au 2.54

Hg 2

Tl 2.04

Pb 2.33

Bi 2.02

Po 2

At 2.2

Rn 0

7

Fr 0.7

Ra 0.89

Ac 1.1

Rf

Db

Sg

Bh

Hs

Mt

Uun

Uuu

Uub

Lanthanides

Ce 1.12

Pr 1.13

Nd 1.14

Pm 1.13

Sm 1.17

Eu 1.2

Gd 1.2

Tb 1.1

Dy 1.22

Ho 1.23

Er 1.24

Tm 1.25

Yb 1.1

Lu 1.27

Actinides

Th 1.3

Pa 1.5

U 1.38

Np 1.36

Pu 1.28

Am 1.3

Cm 1.3

Bk 1.3

Cf 1.3

Es 1.3

Fm 1.3

Md 1.3

No 1.3

Lr

Chemical bonds are formed when the electrons of one atom are attracted by the nucleus of another atom. They are also formed with various degrees of sharing of electrons between two atoms.

RECALL:

If the electron are shared equally between the two atoms, it is called a COVALENT BOND.

If the electron of one atom are given away completely to another atom, oppositely charged ion are then held together through an electrostatic attraction, and the force holding the two ions together is called an IONIC BOND.

There are various degrees of sharing which goes on between two atoms, depending on the amount of attraction each atom has for receiving another atom.

The attraction an atom has for the shared pair of electrons in a chemical bond is called an atom's ELECTRONEGATIVITY. It is the difference in electronegativity between two atoms that determines the degree of electron sharing which occurs between the two atoms.

Some fun facts:

• the atom with the highest electronegativity on the periodic table is Fluorine. It has a value of 4.0. This high value indicates that the atom readily pulls on electrons in a bond.
• the atom with the lowest electronegativity on the periodic table is Francium. It has a value of 0.7. This low value indicates that it doesn't readily pull on electrons within a bond.

Right about now, you're probably wondering why this is at all important.
And the answer is... well it probably isn't very important.  But we need to know it because it helps us determine the polarity of a compound.
Remember, there are 2 type of covalent bonds: non-polar and polar. 

You know you have a non-polar covalent bond if the difference between the electronegativities of the two elements comprising the bond have a difference of less than 0.5.

If the electronegativity difference is between 0.5 and 1.8, then it is a POLAR covalent bond. 

Once in a while, you get these weird situations where you have a metal and a non-metal bonding that has an electronegativity difference of less than 1.8, and you're wondering WHAT THE HECK MY TEACHERS HAVE BEEN LYING TO ME ABOUT IONIC BONDS.  I thought this too.  But actually, it's just the electronegativity difference that makes them seem not ionic.  They are still ionic compounds because they bond by transferring electrons, unlike the typical sharing of covalent bonds.

When you have a compound, you may be asked to draw out a diagram showing the movement of electrons (the electron shift). 

When you have a compound where there is a shift in electrons, that means one element gave away an electron and the other element gained one.  That leads to the conclusion that the element that gave away the electron becomes partially positive [it lost a -1 charge --> -(-1) = +1], and the element that gained an electron becomes partially negative.

For example... As you can see below, I kindly drew out for you what Lithium and Flourine would look like.  Since Fluorine has 7 valence electrons, it desperately wants 1 more to fill its shell.  On the other hand, Lithium only has 1 valence electron that it is dying to get rid of.  So Lithium kindly donates its electron to the cause, and allows Flourine to gain its full valence shell.  And everybody's happy! yay! So since Lithium gave away an electron, it becomes partially positive.  That squiggly S like thing means partially.  Cool beans.

The names Bond.

NOPE. WRONG BOND.

It's actually this guy...


Kidding again!

Another joke of the day:
Why did Carbon marry Hydrogen? 
I don't know, why? 
They bonded well from the minute they met. HAH


So today we will be talking about BONDING! Specifically, chemical bonding, which involves the transfer and sharing of valence electrons to gain stability (closed shell).


There's an existing force of attraction or repulsion  that exists between two charged particles and that is the electrostatic force. This force is acted upon in all directions.


Few things to note on:

• Opposite charges attract, and like charges repel
• The greater the distance, the smaller the attractive force
• The greater the charge, the greater the attractive force

Ionic Bonds

• the transfer of valence electrons from one atom to another
• formed by attraction of a metal to a non-metal
• creates a positive and a negative ion 

Atoms lose or gain enough electrons to attain a closed cell

Ionic bonds are arranged in a crystal lattice structure where there is an equal number of alternating positive and negative charges

Example of NaCl:

Ionic bonds are very strong, and only an enormous of energy can break the bonds. So if they need a LOT of energy, then the melting point must be high.

Electronegativity: measure of tendency of an atom to attract electrons from a neighbouring atom

(Non-Polar) Covalent Bonding

Non-polar covalent bonding is based on the octet rule.  That is, non-metals share electrons in order to obtain a closed shell.  Wow sharing?!  If my sister and I were a bond, we would be ionic.  As in I would steal her electrons.  We don't share, as a rule.  Sharing means I take stuff, say I will share, and then don't.  She usually forgets, so it's a very effective method of winning.  Elements have a lot to learn from me.


Here is an example of a covalent bond:

The electrons are shared equally among themselves, and each atom satisfies the octet rule. This nonpolar covalent bond is symmetrical, which is another factor of determining what type of covalent bond it is (we'll get to that later).

Anyways, I'm sorta like a non-metal, since non-metals have high levels of electronegativity, which means they attract electrons very strongly.  But, that being said, none of them are up to giving away any.  Greedy? I think so.  So these non-metals have come up with a bit of a compromise - they share electrons with other non-metals in order for both elements to be satisfied.  See, that's where humans (me...) and elements differ.  Elements are willing to share electrons to achieve a win-win, but for humans - a win isn't a WIN unless there's a loser.  Just saying.

IMPORTANTTERMSTHATYOUSHOULDNOTFORGET
omg i held the shift button for that entire thing instead of using caps lock... dummy.

INTRAmolecular forces (aka covalent bond) - found in the INTERIOR of the molecule - holds the molecule together
and the weaker INTERmolecular forces - found BETWEEN molecules, and holds 2 SEPARATE molecules together 


example:

Drawing Lewis Dot Structures

Here's a chart that shows the Lewis dot structures for main elements that we will use this year. The Lewis dot structures only show the number of valence electrons. Remember to follow the octet rule when bonding! (atoms tend to lose/gain and share electrons until they are stable)

COVALENT:
Covalent bonds SHARE electrons, so thus they are bonded like the diagram below.
Example: Hydrogen has one valence electron and Chlorine has 7. When they share their unpaired electrons, each of the atoms form a stable shell.

Note that the horizontal line represents the covalent bond (the sharing of electrons).

For covalent bonds, there can be more than one bond. Sometimes, when an atom has more bonding electrons - which are unpaired electrons - it can form double to triple bonds.

IONIC:

Ionic bonds transfer their electrons, thus a metal will give away their electrons and the non-metals will then gain them.

Example: Sodium has one valence electron to give away to Chlorine, who needs one more to become stable. 

The only difference between drawing covalent and ionic bonds are that ionic bonds are drawn with square brackets around each atom. We also indicate how many electrons are lost and gained by writing the charge outside of the brackets.

SO. We have a test on Tuesday the 24th.  Do not forget or you will fail.  Do not be THAT PERSON who walks in asking why everyone is studying and then has a minor heart attack and faints when they discover there is a test.

Yes, you. There's always gotta be someone...

History and Trends of the Periodic Table!

Hello fellow students, today I will be informing you about the history of the periodic table and also the trends on the periodic table. 

So here is a picture of a typical periodic table. But believe it or not, it did not look like this when it was first created. (I know right! Shocking!)

Here are some facts you might find interesting!

• By 1817, there were 52 elements discovered and 62 by 1863! 
• The first attempt to make a periodic table was in the 1820's. 



1863-1866 John Newlands began ordering the known elements by their masses (Law of Octaves - every 8th element shared a common set of properties)

In 1857, William Odling separated the known elements into 13 different groups based on their physical & chemical properties.

1869 Russian Dimitri Mendeleev published a method if organizing the elements according to both their masses & properties (can I get a wootwoot for mr. Mendeleev?)

1914 Henry Moseley discovered the relationship between element's atomic number and frequency of its x-rays

Now in a modern periodic table you have a series of rows and columns organized by their atomic number and not their atomic massses. 

In the pictures above, it shows that a period is the set of all elements in a given row going across the table. Whereas a group (or aka family) is the set of all elements in a given column going down the table.

There are different chemical families that are present in the periodic table. There is the:

Now let's talk a little bit about metals, non-metals, and semiconductors.

Metals reflect light when its polished (like silverware). It becomes shiny and has a metallic luster. Metals are usually opaque and good conductors of heat and electricity. They are generally flexible in sheet form and are usually solid at room temperature (expect for mercury). Metals lose electrons whereas non-metals gain.

Now non-metals are the white boxes from the image above. They are gases, liquids, or brittle solids at room temperature. They are poor conductors of heat and electricity compared to metals and they are opaque to translucent and dull - lustrous in appearance. Non-metals can be divided into 2 different groups: very low electrical conductivity & fair to moderate conductivity.

Semiconductors are metalloids/semimetals that have properties which resemble metals more than non-metals. Metal conductivity decreases with increasing temperature where as electrical conductivity increases with increasing temperature.

WOOHOO! Now that we're done with the history, lets move on to the trends on the periodic table.

There are several trends such as:

1. Metallic Properties
2. Atomic Radius
3. Ionization Energy
4. Electronegativity
5. Reactivity
6. Ion Charge
7. Melting/Boiling Point
8. Density

Here are some notes on these trends. I'm only going to do a few so that you don't fall asleep.

Metallic Properties

- properties of the element change from metallic to non-metallic going from left to right across the table.

- elements become more metallic going down a group in the periodic table.

Atomic Radius

- atomic radii of an atom decreases going across a row & increases doing down a group.

- if you go from left --> right on a given period, the atomic number (protons) increases & the positive charge on the nucleus increases also.

- increase in atomic number = increase in the number of electrons.

- all electrons in a given shell can be assumed to have the same average distance from the nucleus.

- as the number of protons increase, there is a greater force of attraction for electrons in the shell & the distance between the electrons and the nucleus decreases.

Ionization Energy

- energy needed to completely remove an electron from an atom.

- increases going up and to the right.

- all noble gases have high ionization energy (highest = Helium; lowest = Francium).

- opposite trend from atomic radius.

- measured in kJ/mol.

- increases going across a row.

- decreases going down a row.

Electronegativity

- how much atoms want to gain electrons.

- same trend as ionization energy.

- tendency of an atom to attract (completely remove) electrons from a neighboring atom.

- atoms with high electronegativity strongly attract their own valence electrons (difficult to remove when atoms have high electronegativity)

- atoms with low electronegativity have little attraction to the electrons & little tendancy to remove an electron from a neighbour.

- top right corner of the periodic table has the highest electronegativity (excluding noble gases). (Flourine is the MOST electronegative element).

- electronegativity increases as you go across a period.

- electronegativity decreases as you go down a group.

Ion Charge

- the ion charge changes as you go across the period on a periodic table.

- it goes from +1 (Alkali Metals), +2 (Alkaline Earth Metals), variable charges (Transition Metals), +3 (group that starts with Boron), +/-4 (group that starts with Carbon), -3 (group that starts with Nitrogen), -2 (group that starts with Oxygen), -1 (Halogens), 0 (Noble Gases).

Melting/Boiling Point

- elements in the centre of the table have the highest melting point.

- noble gases have the lowest melting point.

- from L --> R, melting point increases (until the middle of the table).

Some main things to remember is:

• When you are going down the periodic table, it means going to the outer shells so the radius increases.
• When you are going across the periodic table, it mean adding electrons to the same shell & increasing the positive charge of the nucleus, so the radius decreases as the shell is pulled in.

VALENCE ELECTRONS!

Anyway, today's blog is about predicting the number of valence electrons. (DIGRESSION: speaking of predictions, I predicted that Canucks would atleast get through round 1 and kill the Kings...but NOPE we lost.. so obviously predicting is just a bad thing to do... at least for me)

But since we have to PREDICT the number of valence electrons, this is how it's done.

First, let's start off with some important terms...

The core electrons are not valence electrons. They are the inner electrons that are in an atom and therefore don't do the action of bonding.

The outer core is where the valence electrons take place, excluding the filled d- and f- subshells.

Now when we talk about shells, we do not mean these shells:

An openshell is the shell that is not completely filled with the max. number of electrons and a closed shell is when the shell is completely filled with electrons to its maximum extent

•  Valence electrons are the outermost electrons in an open shell of an atom 
•  Valence electrons include all electrons EXCEPT those in the core and d & f subshells
•  AKA the reactable electrons, important part of chemical reactions and bonding                            
• ALL noble gases have ZERO valence electrons because they already have a stable octet meaning it has exactly 8 electrons on it.

Examples: 

Sulfur: [Ne] 3s2 3p4 ---> 6 valence electrons because you only count the ones in the outer core
Rubidium: [Kr] 5s1 ---> 1 valence elecrton
Neon: [Ne] ---> 0 valence electrons because it's a noble gas!! However, you can also say 8 electrons :-)
 
Tata

Ms Chen the pyromaniac

^That was pretty much Ms Chen last class...
hahahhah


Yes... that would be us.


Hahahaha it was a fun class. Ms Chen basically set a bunch of elements mixed with methanol into flames and each turned into a different color. It was pretty cool.

COPPER WAS MY FAVOURITE!

So on a slightly less exciting note, class continued with us viewing light throught spectroscopes.  We then drew out the line spectra shown in our magical little device.  I dunno about you guys, but I seemed to have some trouble viewing the actual spectral lines.  Maybe I needed to get a little closer, but I was rather frustrated that I could SORTA see the lines in the spectroscope, but the were in my peripheral vision so I couldn't QUITEEE get a proper glimpse.  Angry face.

...You know you were all thinking of Ms. Lehmann while filling out that worksheet... ;)

I was definately humming under my breath while working.

So the important thing from the worksheet - remember good ol ROY G BIV?

Well that little thingymabob will help you remember the order of wavelengths in the visible spectrum. 

Red waves have the shortest wavelengths and lowest frequency because they are very low energy.

Violet waves, on the other hand, have very short, frequent waves, which indicates high energy.

So that's about it... I'm pretty sure.

See ya later, alligator :)

Electronic Configuration 101

Hello fellow followers! Let's start the lesson with a fun chemistry joke, shall we?


Outside his buckyball home, one molecule overheard another molecule saying, "I'm positive that a free electron once stripped me of an electron after he lepton me. You gotta keep your ion them." 


Bahahaha, get it?!?! We all need some cheesy chemistry jokes once in a while :-)


Moving on now...

We had a very perplexing class today. Brace yourselves, for we will be covering a difficult topic: ELECTRONIC CONFIGURATION. So pay attention carefully!

What is it?

Electronic configuration of an atom is a notation that describe the orbitals where the electrons occupy and the total number of the electrons in each orbital.

Neils Bohr, the same guy that invented the Bohr diagram, proposed that electrons only exist in specific energy states. Electrons can move from one orbital to another when they absorb or emit a specific amount of energy.

Energy levels are the amount of energy that an electron can have. The lowest energy level is in the innermost orbit that forms a shell around the nucleus. Each orbit is assigned a quantum number (n) which indicates the number of the energy level. As n increases, the energy of the electron increases. The energy difference between two particular energy levels is called the quantum of energy.

Atoms with all electrons in the lowest energy levels are in a ground state. When electrons move up in the energy levels, they become less stable and more "excited", thus are in an excited state.

As mentioned before, electronic configuration revolves around orbitals. So what are they? An orbital is the actual region of space occupied by an electron in a particular energy level. The set of all orbitals having the same n-value or energy level is the shell. A subshell is a set of orbitals of the same type.

There are many different types of orbitals and each have their own fundamental shape. They are designated by letters s, p, d,and f. Each orbit holds a maximum of 2 electrons. (Pauli Exclusion Principle)

Specific types of orbitals are possible for a given value of "n".

             n=1 : s-type ONLY

             n=2 : s- and p-types

             n=3 : s- , p- , and d-types

             n=4  : s- , p- , d- and f-types

S-type subshell

• consists of 1 s-orbital
• has a spherical shape
• max # of electrons: 2

P-type subshell

• consists of 3 p-orbitals
• shaped like a dumbbell
• max # of electrons: 6

D-type subshell

• consists of 5 orbitals
• max # of electrons: 10

F-type subshell

• consists of 7 f-orbitals
• max # of electrons: 14

____________________________________________________________________________

Writing Electronic Configurations

(for neutral atoms)

STEPS:

1. Figure out how many electrons you have (if neutral atom, just look at the atomic number)
2. Follow the direction of the arrows and start with the lowest energy level first (Aufbau principle)
3. Starting from 1s, keep adding the notations until you get to the designated number (of electrons). Note that each orbital has a diff maximum of electrons! (eg. s-types can only go up to 2)

Each electron has an opposite spin designated by upward and downward arrows.

Example: Silicon, in its neutral atom, has 14 electrons. Thus, there 2 electrons in 1s, 2 in 2s, 6 in 2p, and 2 more electrons in 3s.

The letters "A" , "B", etc shows the steps taken to write the electron configuration for silicon. Note that in the 3p subshell, there are two electrons that occupy separate orbitals without being paired. This is because of Hund's Rule - when electrons occupy orbitals of equal energy, they don't pair up until they have to. So one would fill up each orbital with the "up" arrows first before adding "down" arrows.

You would write it as: 

­1s²2s²2p⁶ 3s²3p^2­­

FOR IONS:

Negatively charged ion:

• starting from the number of electrons in a neutral atom (equal to atomic #), add electrons according to the ion charge
• with the new number of electrons, write the electronic configuration
• remember that if there were half-filled orbitals, fill the remaining down arrows in!

Eg. Oxygen ion - O^-2 

Number of electrons = 10

↑↓   ↑↓   ↑↓ ↑↓ ↑↓ 

1s       2s         [  2p  ] 

Therefore: 1s²2s²2p⁶

Positively charged ion:

• starting from neutral number of electrons, remove electrons (from the outermost shell  first working backwards) according to charge
• work your way backwards with the orbitals, so if you have electrons in both s- and p-orbitals, remove the ones in the p-orbital first
• then write the notation the same way you would for the neutral atom but with a different number of electrons!

Core Notation

• Set of electrons in an atom can be divided into 2 subsets: core and outer electrons. Thus core notation shows the electronic configuration in terms of core and outer electrons
• CORE: set of electrons with the configuration of the first noble gas that comes BEFORE it, or the way I like to think of it, the noble gas that's a row above the element
• OUTER: the remaining electrons outside of the core

How do you write it? Well, you will need to refer to your periodic table!

1. Find the element. 
2. Then find the noble gas on the row above it.
3. Replace the first part of the configuration, that equals the same amount of electrons as the noble gas, with the symbol of the noble gas in square brackets.
4. Leave the remaining electronic configuration of the outer electrons outside the brackets.

eg.  Oxygen ion - 10 electrons

       1s²2s²2p⁶

^{On the periodic table, Helium is the first noble gas that comes before it. Helium has 2 electrons, so you would replace 1s^2.} 

^{Rewrite the configuration and you would get:}

[ He ]2s²2p⁶

^{The red indicates the core and the blue indicates the outer electrons!}

Also there are TWO exceptions you need to memorize!! 

1. Copper (Cu) gains stability by having a full d-subshell, so therefore its configuration would be [Ar]4s3d_­¹⁰
2. Chromium (Cr) gains stability by having a half-full d-subshell so it would be written as  [Ar]4s3d_­⁵

^{Here is an additional periodic table with some of the configurations already! Your welcome! :)}

1A	2A	3A	4A	5A	6A	7A	8A
1 H 1s1							2 He 1s2
3 Li 1s2 2s1	4 Be 1s2 2s2	5 B 1s2 2s22p1	6 C 1s2 2s22p2	7 N 1s2 2s22p3	8 O 1s2 2s22p4	9 F 1s2 2s22p5	10 Ne 1s2 2s22p6
11 Na [Ne] 3s1	12 Mg [Ne] 3s2	13 Al [Ne] 3s23p1	14 Si [Ne] 3s23p2	15 P [Ne] 3s23p3	16 S [Ne] 3s23p4	17 Cl [Ne] 3s23p5	18 Ar [Ne] 3s23p6

Oh my.... looking back this was a VERY LONG post. I'm sorry for those readers that had to endure all this way!

HAHAHAHA. In no way take this seriously! I apologize in advance!