Electronegativity and Polarity

Hello all, and welcome to one of the less interesting blogs that we have to offer you.  Unfortunately it is important information so we'll try to make it as short and sweet as possible.  This is an electronegativity chart showing the charges for how easy/hard it is for each of these elements to loose an electron(s).  For instance, elements like Cesium is much more likely to give away its electron to another element (such as fluoride).  Make sense so far?  I hope so.














These charges can be used to determine what kind of bond you are looking at.  So, for example, you were looking at what kind of bond sodium and chlorine would make.  You would find their electronegativity difference.  


Na= 0.9                                              3.0-0.9= 2.1
Cl= 3.0
(if you get a negative answer, then think of what you have as having absolute value bars around it, just make it positive)
Based on this chart...and on prior knowledge, we know that this is an ionic bond.  However, when you do the same process with say, Nickel and oxygen, you will see that it instead comes up as a covalent bond.  That is something that will be explained more next year.

Here is the promised video to help keep the difference between the three rememberable.  If you can't deal with bad corny science songs, then don't click play.  Otherwise, I hope it helps!
http://www.youtube.com/watch?v=oNBzyM6TcK8

Part 2: knowing how to draw these.

Luckily, for the most parts, it is drawn as a regular lewis dot structure; which you already know how to do right? Right.  The only difference, is you need to indicate which atoms would become positive (loose electrons)  and which would be come negative (gain electrons)  Then draw an arrow pointing toward the negative one.  Have no idea what is being said, well, here is a nice long video that takes you back to periodic trends and explains it a bit more.

Best of luck :)

Chemical Bonding

Here starts the tedious introduction blog of the last unit we will do this year.  Many things that are mentioned here will come up later so if you don't understand this one....take the time to figure it out.

Electrostatic Force

Is a force existing as a result of the attraction or repulsion (two positive or two negative particles repel each other but one of each attract each other) between two charged particles.  It is a important to remember all bonds are based on the electrostatic relationships formed between particles.  When these particles are farther from each other, the attraction is greater.  If there is more charge, then there is more force.  This lovely force that occurs acts in all directions which is why that in both theory and practice, bonds can be made on any electron in the valance shell.  

Ionic Bonds

Ionic bonds are formed between positive and negative ions (most often medal and non-medal).  In an ionic bond, the electrons are transferred from what atom to another.  If you need a good way to remember that, just look at the chemistry cat on the side of our blog.  These bonds are always made to have a closed shell.  Ionic bonds are:

• very strong
• have high boiling points
• high electronegativity
• ionization energy

After the bond, all atoms become ions.  The positive ions are called cations and the negative ions are called anions.  

Then there is the crystal lattice...

In a crystal lattice, the ions in the bonds line up so the positive side of one molecule line up with the negative side of the next.  This is what allows them to connect and make things such as table salt. 

[Non-Polar] Covalent Bonding

These bonds, unlike ionic bonds, share their electrons amongst that atoms in the molecule.  As a result of this, the electrons are stuck in the middle between the two atoms.  These bonds tend to be very strong with lower melting points.  

Individual Molecules

Individual molecules are actually quite strong.  During melting, it is not the molecules themselves breaking apart, but the different molecules breaking apart from each other.  They are held together only by the intermolecular force.  This is especially true for covalent bonds.

The next blog will have a video that will help to remember the differences between the three.  (Yes, there is one more; covalent bonds aren't always non-polar...)

A Brief History of the Atom

Democritus: Everything is made up of infinitely small particles. Titled it atomos, meaning indivisible.

Dalton: Created modern atomic theory. All atoms are hooked together, and form specific ratios. Did not realize that atoms are made up of different particles

Thomson: Discovered electron with cathode ray tube. Applied a magnetic force, and noticed that the ray deviated, due to opposing magnetism. Also created a new idea of appearance, titled plum pudding formation.

Rutherford: Fired alpha particles through gold, some rays deviated from center, while others bounced right back. This was because of the positively charged nucleus. He also realized that a lot of the volume of the atom is empty space.

Bohr: Worked in Rutherford's lab, and realized if the charge on the inside was positive, that the negative and positive charges would attract. Created the Bohr model on the assumption that electrons followed specific paths.

Schroedinger: claimed that electrons weren't particles but were in fact waves.

Chadwick: Discovered the neutron by disintegrating atoms through bombardment of  alpha particles. First to create artificial nuclear reaction
Heisenberg: Confirmed both of Bohr and Schroedinger's hypothesis. The electron is indeed a particle, that travels in a wave like motion, which is described by quantum theory.

History of the Periodic Table

Since the 5th century, a few elements had been discovered, including gold, tin, silver, lead and mercury, as these were easy to find. The first discovery was by Henning Brand, a German man, on a wild chase for the philosophers stone. In 1669, he discovered phosphorus, but kept it a secret, and the rock was later publicized in 1680 by Robert Boyle.
A hundred years later, a chemist named Antoine Lavoisier lived. He created a text book of elements that could not be broken down, which also included light, and caloric, which were then believed to be material substances.
Lavoisier's elements

Gases New names (French) Old names (English translation) Lumière Light Calorique Heat Principle of heat Igneous fluid Fire Matter of fire and of heat Oxygène Dephlogisticated air Empyreal air Vital air Base of vital air Azote Phlogisticated gas Mephitis Base of mephitis Hydrogène Inflammable air or gas Base of inflammable air

Metals New names (French) Old names (English translation) Antimoine Antimony Argent Silver Arsenic Arsenic Bismuth Bismuth Cobolt Cobalt Cuivre Copper Étain Tin Fer Iron Manganèse Manganese Mercure Mercury Molybdène Molybdena Nickel Nickel Or Gold Platine Platina Plomb Lead Tungstène Tungsten Zinc Zinc

Nonmetals New names (French) Old names (English translation) Soufre Sulphur Phosphore Phosphorus Carbone Pure charcoal Radical muriatique Unknown Radical fluorique Unknown Radical boracique Unknown

Earths New names (French) Old names (English translation) Chaux Chalk, calcareous earth Magnésie Magnesia, base of Epsom salt Baryte Barote, or heavy earth Alumine Clay, earth of alum, base of alum Silice Siliceous earth, vitrifiable earth

 By 1809, 47 elements had been discovered. Chemists noted pattern in reactions, and attempted to classify them. 

In 1862, the first form of element organization is published by a French chemist, Alexandre-Émile Béguyer de Chancourtois, that appeared to be a spiral, with increasing mass downward.

1864 John Newland proposes the Law of Octaves, which he found by assigning a mass of 1 to helium, and then ordering the rest by mass. This lead to the discovery that every 8th element had similar properties, 

in 1869, Mendeleev published the first official periodic table. He ignored certain previous ideas, and instead grouped certain elements together according to properties, vs the previous method of weight organization. Mendeleev predicted atomic weights of yet to be discovered elements, and left empty sections in his table to accommodate them

In 1911, Ernest Rutherford publishes a paper that explained nuclear charges. Later that year, Antonius Van Den Broek publishes a paper relating the atomic weight of an element to the charge of the atom. This became the basis of organization of the table, the atomic number.

Two years later, Henry Moseley proposes that the wavelengths of x-ray emissions of elements is directly proportional to that of the atomic number.

The latest changes made were by Glen Seaborg, who discovered plutonium, and elements 94-102. It was his doing that resulted in the lanthanide and actanide series' being placed below the table.

Periodic Table Trends

The periodic table is not organized randomly, it follows certain trends.

There are some trends we need to know.

Here is a youtube video that pretty much covers EVERY trends we need to know.

Below are the trends summarized in point form :D

1) Metallic properties

• From left to right across the table, the elements change from metal to non-metal
• From top to bottom down the table, elements become more metallic, or better metals

2) Atomic radius

• From left to right across a row, the radius of atom decreases because increase in number of protons brings the electrons closer to the nucleus
• From top to bottom down a group, the radius increases

3) Atomic size

• From left to right across a period, the size decreases
• From top to bottom down a group, the size increases

4) Reactivity

• For metal:
• From top to bottom down a group, the elements become more reactive
• in transitional metals, the middle part of the table is the least reactive and to the left or to the right become more reactive
• the most reactive metal is Francium
• For non-metal (excluding noble gas):
• From top to bottom down a group, the elements become less reactive

5) Ion charge

• Metals tend to have positive charges
• Non-metals tend to have negative charges
• the transitional metals have variable charge

6) Melting & boiling point

• elements in the centre of the table have the highest melting and boiling point
• Noble gases have the lowest M & B point
• from left to the right across the table, the M & B point increases until the middle of the table

7) Ionization energy

• this is the energy required to completely remove an electron from an atom
• From bottom to up, left to right cross the table, the energy increases
• All noble gases have high ionization energy
• Helium has the highest ionization energy while Francium has the lowest
• Note: this trend is opposite from atomic radius trend

8) Electronegativity

• Electronegativity is basically how much atoms want to gain electrons
• Note: this trend is the same as the trend for ionization energy
• Noble gases are excluded from the trend as they already have full shells (so they don't want to gain or loss electrons)

Writing Electronic Configurations

Neutral Ions

-always start with the lowest level first (Aufbau Principle)

-Figure out how many electrons you have (Neutral atom = Atomic #)

-start with the lowest energy level (1s) and keep adding until you have no electrons left

-Each electron has an opposite spin designated 

Ex. Sillicon: 14 electrons

2 electrons in 1s, 2 electrons in 2s, 6 electrons in 2p, 2 electrons in 3s, 2 electrons in 3p

A-E shows the steps taken to write the electron configuration

The 2 electrons in the 3p subshell aren't paired because of Hund's rule: when electrons occupy orbitals of the same energy, they can't be paired up until they have to.

written as: 1s²2s²2p⁶3s²3p²

Ions

negative charged ions: add electrons to the original number of electrons according to its charge

ex. P3- : 15+3

15 being its original # of electrons and 3 added because of its charge

total electrons = 18

written as: 1s²2s²2p⁶3s²3p³

positive charged ions: remove electrons from its original number of electrons according to its charge

ex. Ba2+: 56-2

total electrons: 54

written as: 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶5s^24d105p6

Core Notation

-set of electrons that can be divided into two subsets: core electrons and outer electrons

core: set of electrons with the configuration of the nearest noble gas that comes before it

outer: consists of all electrons outside the core (normally takes part in chemical reactions)

ex. Ca

1)the noble gas before calcium is Argon, so you put argon in square brackets

[Ar]

2)then add the remaining electrons

[Ar]4s²

2 EXCEPTIONS: Copper (Cu) and Chromium (Cr)

Cu:  [Ar]4s3d_­¹⁰

^{Cr: [Ar]4s3d5}

Electronic Structure of the Atom

-the electronic figuration of an atom is a notation that describes the orbits in which the electrons occupy and the total number of electrons in each orbital
ex. Manganese's electron configuration: 1s²2s²2p⁶3s²3p⁶4s²3d⁵


energy level: the amount of energy which an electron in an atom can posses (n)
quantum of energy: energy difference between 2 particular energy levels
ground state: when all the electrons of an atom are in their lowest possible energy levels
excited state: when one or more of an atom's electrons are in energy levels other than the lowest available level
orbital: the actual region of space occupied by an electron in a particular energy level
shell: the set of all orbitals having the same n-value
subshell: set of orbitals of the same type

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
Pauli Exclusion Principle: A maximum of 2 electrons can be placed in each orbit

S-type subshell: ONE s-orbital and 2 max # of electrons

P-type subshell: THREE p-orbitals and 6 max # of electrons
D-type subshell: FIVE d-orbitals and 10 max # of electrons
F-type subshell: SEVEN f-orbitals and 14 max # of electrons

Valence Electrons

- valence electrons are all electrons EXCEPT those in the core or in the filled d- or f- subshells

Valence Electrons: in the outermost (energy level) open electron shell of an atom
Open Shell: contains less than its maximum number of electrons
Closed Shell: exactly its maximum number of electrons

ex. Br:  [Ar]4s²3d¹⁰4p⁵  =  7 valence electrons

Ca: [Ar]4s2  =  2 valence electrons

Xe: [Xe]  =  8 valence electrons

Atoms, Ions and Isotopes

Review

Properties of sub-particles

• neutral atom has no overall net charge
• number of protons = number of electrons

The atomic Number (number of protons)

• protons found in nucleus
• IF no electric charge
• atomic number=number or protons=number of electrons
• if a proton is added to an elements nucleus, a new element will be produced

Ions

• most atoms can gain or loose electrons
• few elements (ex. hydrogen) can gain or loose both
• atoms that gain or loose electrons care called Ions
• ion is an electricity charge atom
• negatively charged normally non-metal
• positively charge tend to be metals

Mass Number

Mass number:  total number of protons and neutrons or atomic mass number

• number of neutrons= mass number-atomic number  OR atomic mass - atomic number
• atomic number(pro.)= mass number-electrons
• atomic mass=average of all isotopes
• mass number= round number of atomic mass
• MASS NUMBER DOES NOT EQUAL ATOMIC MASS

Isotopes

• neutron is added to an elements nucleus which makes a heavier version of the same element
• same element with bigger atomic mass
• same number of protons and electrons
• different number of neutrons so different atomic mass

For example, if you had Calcium-41, how many neutrons would there be?

Calcium-41 = Calcium with an atomic mass of 41

Calcium       = atomic number is 20 (20 protons)

                       atomic mass= protons + neutrons

                    = n + 20 = 41

                       so there are 21 neutrons

                             there is no charge listed so assume 

                                 number of electrons= number of protons

                                 Your done :)

Review Isotopes!

NOW HERE IS A VIDEO THAT CONTAINS EVERYTHING YOU NEED TO KNOW!