History of Periodic Table

In the Beginning

• The first scientific discovery of an element occurred in 1649 when Hennig Brand discovered phosphorous.
• By 1869, a total of 63 elements had been discovered.
• As more and more elements were being discovered, scientists began to recognize patterns in properties.

• In 1863, John Newlands arranged chemical elements in order of their relative atomic masses and he arranged his elements in columns.



John Newlands

• He proposed the Law of Octaves in 1864
• This law stated that any element will show similar behaviours at the eighth element following it in the table.

Dmitri Mendeleev

• In 1864, Dmitri Mendeleev improved Newlands' idea and proved its effectiveness.
• Mendeleev is known as the 'father' of the periodic table.
• He organized the periodic table like no other chemist.
• He managed to organize them into groups possessing similar properties, not by atomic mass.
• He left gaps in the table, predicting that a new element would be discovered in place.

Mendeleev's Periodic Table

• In 1895, Lord Rayleigh discovered the gaseous element Argon.
• Did not fit with any of the periodic groups.
• In 1898, William Ramsey organize argon into the periodic table between chlorine and potassium.
• He made this decision even though argon's atomic weight was greater than potassium's.
• This group was called "zero" group due to the zero number of valence electrons in each element.
• He also predicted correctly properties such as neon.
• This groups is currently called the Noble Gases.

Modern Periodic Table

• Glenn Seaborg made the last major changes to the periodic table.
• Discovered plutonium in 1940
• Also discovered all elements from 94 to 102.
• Reconfigured table by moving the actinide series below the lanthanide series.
• Element 106 was name seaborgium (Sg) in his honor.
• The periodic table is completely organized by atomic number not by mass.
• Periodic Law: Elements recur periodically when arranged from lowest to highest atomic number.

Divisions in Periodic Table

• Period: horizontal rows
• Groups or Families: vertical columns
• Blocks: according to subshells (s-, p-, d-, and f-block)

Periodic Table Blocks

Chemical Families

• Alkali Metals
• Group 1
• Located in the s-block
• Highly reactive metals
• Alkaline Earth Metals
• Group 2
• Located in the s-block
• Share similar properties
• Relatively reactive metals
• Halogens
• Group 17
• Non-metal elements
• Contains all three familiar states of matter at a regular temperature and pressure.
• Highly reactive because the atoms are highly electronegative due to heir high nuclear charge.
• Reacts with water.
• Noble Gases
• Share similar properties
• All odorless, colorless, monatomic gases, with low chemical reactivity.
• Groups 18
• Outer shell full
• Lanthanide Series
• Includes 15 elements with atomic number 57 through 71.
• Are f-block elements
• Lanthanum and lutetium are labeled group 3 because they both have a single valence electron, therefor in the d-shell.
• Actinide Series
• Includes 15 elements from 89 to 103.
• Have similar properties.
• f-block elements
• Have a much more variable valence than the lanthanies.
• All radioactive and release energy during radioactive decay
• Can be used in nuclear reactors and nuclear weapons.

Modern Periodic Table

Electronic Structure of the Atom

electron configuration is the arrangement of electrons of an atom, a molecule


Shells
    -The electron shells are labelled K, L, M, N, O, P, and Q; going from innermost shell   outwards. 
    -Electrons in outer shells have higher average energy and travel farther from the nucleus than those in inner shells.

Subshells
    -Each shell is composed of one or more subshells, which are themselves composed of atomic orbitals. 

    - Example: The first (K) shell has one subshell, called "1s"; the second (L) shell has two subshells, called "2s" and "2p"; the third shell has "3s", "3p", and "3d"; and so on.

• Each s subshell holds at most 2 electrons
• Each p subshell holds at most 6 electrons
• Each d subshell holds at most 10 electrons
• Each f subshell holds at most 14 electrons
• Each g subshell holds at most 18 electrons

Notation
   -Notation consists of a sequence of atomic orbital labels (e.g. for phosphorus the sequence 1s, 2s, 2p, 3s, 3p) with the number of electrons assigned to each orbital (or set of orbitals sharing the same label) placed as a superscript.

   -Examples:

       Hydrogen :  1s¹       (has one electron in the 1s-subshell)                                                  

       Lithium    : 1s² 2s¹   (has two electrons) 

       Phosphorus: 1s² 2s² 2p⁶ 3s² 3p³   (has fifteen electrons)    

                                            

Valence Electron
    - the electrons in the last shell or energy level of an atom
    - the electrons of an atom that can participate in the formation of chemical bonds with other atoms
    -The valence electrons increase in number as you go across a period
    

Periodic table group	Valence electrons
Group 1 (I)	1
Group 2 (II)	2
Groups 3-12	See note *
Group 13 (III)	3
Group 14 (IV)	4
Group 15 (V)	5
Group 16 (VI)	6
Group 17 (VII)	7
Group 18	8**

* The general method for counting valence electrons is generally not useful for transition metals. Instead the modified d electron count method is used.

** Except for helium, which has only two valence electrons.

    -Example: Valence electron of Phosphorus(1s² 2s² 2p⁶ 3s² 3p³)   : 2+3 = 5                                        

    
     
Core Notation
    - Instead of writing out all of the electrons in the configuration, you can write out just the ones since the last noble gas
    -Example: electron configuration of Magnesium is 1s 2 2s 2 2p 6 3s 2
                                               ---->［Ne］3s 2


Ground State
    -The condition of an atom, ion, or molecule, when all of its electrons are in their lowest possible energy levels, not excited. 
    -When an atom is in its ground state, its electrons fill the lowest energy orbitals completely before they begin to occupy higher energy orbitals, and they fill subshells in accordance with Hund's rule 


Excited State
  -An excited state is any state with energy greater than the ground state

  -Excited states tend to have short lifetimes; they lose energy either through collisions or by emitting photons to "relax" back down to their ground states. 

Hund's Rule
    -Developed by the German scientist,
    -A rule of thumb stating that subshells fill so that the number of unpaired spins is maximized. Also known as therule of maximum multiplicity.
    - Electrons are added to the lowest available energylevel (shell) of an atom.


Pauli Exclusion Principle
    -it was formulated by the Austrian-born physicist Wolfgang Pauli in 1925. 
    -The Pauli exclusion principle states that each electron must have its own unique set of quantum numbers that specify its energy.
    -no two electrons can occupy the same quantum state at the same place and time 

Atomic Structure

Atomic Structure

Atoms are made up of 3 types of particles electrons , protons  and neutrons .  These particles have different properties. 


      Electrons: tiny, very light particles that have a negative electrical charge (-)


      Protons:  much larger and heavier than electrons and have the opposite charge,    protons have a positive charge


      Neutrons: large and heavy like protons, however neutrons have no electrical charge

Centrifugal Force
   -force that makes objects move outwards when they are spinning around something or travelling in a curve.
 
   -The protons and electrons stay together because just like two magnets, the opposite    electrical charges attract each other.

What keeps the protons and electrons from crashing into each other?

   -The electron spins around nucleus(the center of the atom).  The centrifugal force of the spinning electron keeps the two particles from coming into contact with each other.

Electron Cloud
   -Atoms are extremely small. One hydrogen atom, for example, is approximately 5 x 10^-8   mm in diameter.

 - Protons and neutrons behave like small particles, sort of like tiny billiard balls.

 - The electron has some of the properties of a wave.  In other words, the electron is     more similar to a beam of light than it is to a billiard ball.

  - Thus to represent it as a small particle spinning around a nucleus is slightly misleading.  In actuality, the electron is a wave that surrounds the nucleus of an atom like a cloud.

Hydrogen: a proton surrounded by an electron cloud 

 Neutron

  -Helium has the 2 protons in the nucleus have the same charge on them.

  -They would tend to repel each other, and the nucleus would fall apart.  

  -To keep the nucleus from pushing apart, helium has two neutrons in its nucleus.  

  -Neutrons have no electrical charge on them and act as a sort of nuclear glue, holding   the protons, and thus the nucleus, together.

Ion

 - An atom that carries an electrical charge

 - total number of electrons is not equal to the total number of protons

          ● cation : -positively charged ion

                          -has fewer electrons than protons

          ● anion : -negatively charged ion

                         - has more electrons than protons

                           

Listed below are three forms of hydrogen; 2 ions and the electrically neutral form.

H+ : a positively charged hydrogen ion	H : the hydrogen atom	H- : a negatively charged hydrogen ion

Isotope

  - two atoms with same number of protons but different number of neutrons

  - their chemical properties are almost identical since the chemical behavior of an atom is largely determined by its electronic structure.

  - some of them are radioactive

 two isotopes of hydrogen

Hydrogen Atomic Mass = 1 Atomic Number = 1	Deuterium  Atomic Mass = 2  Atomic Number = 1

Atomic Theory

Democritus: -in 300B.C, was the first to theorize the existence of atoms


 Aristotle  however disagreed with Democritus and believed that atoms were made from the 4 elements; earth, air, fire, and water
Lavoisier (late 1700's) -stated earliest version of both the Law of Conservation of mass and discovered combustion

Proust (1799) - Proust believed that if a compound were to be broken down, it would still contain the same ratios as in the compound; continued on with Lavoisier and proved the law of definite porportions

Dalton (early 1800's) - all matter is made of atoms
-all atoms of given element are the same
-atoms of one element can combine with another to form chemical compounds
-a chemical reaction is a rearrangement of atoms

J.J. Thompson (1850) - proved existence of electrons; uses the plum pudding model

Rutherford (1905) - showed atoms had dense centres and electrons outside of them
-suggested atoms are mostly empty space; proved Thompson's model to be incorrect

Niels Bohr (1885-1962) - studied gaseous samples of atoms
-proposed that electrons surrounds the nucleus in shells
-when exciting electrons of a certain element, it gives off light

All these people contributed to the model of the atom. However, the atom today consists of 3 subatomic particles: protons, electrons, neutrons. Protons and neutrons lies within a dense nucleus whilst the electrons surround the nucleus in shells.

Percent Yield & Percent Purity

PERCENT YIELD

Sometimes in a chemical process, not all the products are used up. Percent yield is there to help us calculate the % of the product recovered.

% Yield =   .      grams of actual product recovered             X 100%

                    grams of product expected from stoichiometry

The percent yield should never be higher than 100%. Where did all that extra stuff come from anyway?

Let’s do an example:

Consider this equation: 1 CrCl2 + 2 Ca à 2 CaCl + 1 Cr

If 15.0g of CrCl2 is reacted with excess Ca and 4.10g of Cr was produced, what is the % yield?

STEP 1: Check to see if the equation is balanced. If not, balance it.

It is balanced in this case.

STEP 2: Find the amount of Cr produced when 15.0g of CrCl2 reacts

15.0g CrCl2 X  1mol CrCl2  = 0.12195mol CrCl2    ßconvert to moles

                          123g CrCl2

0.12195mol CrCl2 X  1mol Cr .  = 0.12195mol Cr   ßuse mole ratio

                                   1mol CrCl2

0.12195mol Cr  X 52.0g Cr  = 6.34g Cr   ßconvert to grams

                                1mol Cr

STEP 3: Find % yield

% yield =  4.10g Cr obtained .  X 100%  = 64.7% yield

                   6.34g Cr expected

For every 100.0g of product predicted, only 64.7g of product is actually formed.

 

PERCENT PURITY

Sometimes the reactants that were used are not pure, meaning there are bits of other elements in the main element. % purity calculates how much is actually there.

% Purity =  mass of pure substance  X 100%

                      mass of impure sample

Again, the % purity should never go above 100%.

Let’s start with an easy example:

If a 10.0g sample of zinc ore contains 7.3g of zinc metal, what is the percent purity?

Find the percent purity

% Purity =  7.3g Zn metal (pure substance)  X 100%  = 73.0% purity

                      10.0g Zn ore (impure sample)

Here’s a harder example:

Consider this equation: Ca + H2O à CaOH + H2

If 6.2g of impure calcium, that is 82% pure, is reacted, how many moles of hydrogen gas is produced?

STEP 1: balance the equation

2 Ca + 2 H2O à 2 CaOH + 1 H2

STEP 2: calculate the mass of Ca

% Purity X mass of impure Ca = mass of pure Ca

82.0%  X 6.2g impure Ca  = 5.084g pure Ca

­100%

STEP 3: calculate moles of H2 produced by 5.1g Ca

5.1g Ca X  1mol Ca  =  0.1268mol Ca   ßconvert to moles

                   40.1g Ca

0.1268mol Ca X 1mol H2  = 0.063mol H2   ßuse mole ratio

                            2mol Ca

6.3 x 10^-2 mol of H2 is produced.