Periodic Table

The distant past

Aristotle (384-322 BC) was an ancient Greek philosophers who thought that everything was made up of just four elements: earth, water, air and fire.

Antoine-Laurent de Lavoisier

In 1789, a French chemist Antoine-Laurent de Lavoiser preoduced the first modern chemical textbook which he compiled the first list of elements; oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc and sulfur.

Lavoisier, ‘the father of modern chemistry’

Jons Jakob Berzelius

In 1828, the Swedish chemist Jons Jakob Berzelius published a table of atomic weights, determined the composition by mass of many compounds, and introduced the letter-based symbols for elements.

Dobereiner’s triads

In 1892, Dobereiner discovered that strontium had similar chemical properties to calcium and barium, and that its atomic weight fell midway between the two.
Nature contains triads of elements where the middle element has properties that are an average of the other two members of the triad when ordered by the atomic weight.

Dobereiner’s triads

John Newlands

The English chemist John Newlands was the first person to devise a Periodic Table of the elements in order of their relative atomic weights. In 1865, he put forward his ‘law of actaves’-
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Any given element will exhibit analogous behaviour to the eight element in the table above
• Newlands’ arrangement showed all known elements in seven ‘octaves’
• The elements are ordered by the atomic weights that were known at the time.

Dmitri Mendeleev
Mendeleev’s table
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In Mendeleev’s table:
• Elements with similar properties were arranged in vertical columns
• Gaps were left for not found elements
• The order of elements was rearranged where their properties did not fit

Mendeleev’s periodic law:
• The elements, if arranged according to their atomic weights, exhibit an apparent periodicity of properties
• Elements which are similar, as regards to their chemical properties, have similar atomic weights or increase regularly
• The elements which are the most widely diffused have small atomic weights
• The magnitude of the atomic weight determines the character of the element
• Certain characteristic properties of elements can be foretold from their atomic weights

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C1 1.2 intermolecular forces and bondings

Strength of bonds and forces
Ionic and covalent bonds are strong, but the intermolecular force (forces between different molecules) are weak.

There are three common types of intermolecular forces:
-hydrogen bonds
-permanent dipole-dipole forces
-can set Waals’ forces

Permanent dipole-dipole interactions
Polar molecules (molecule that has an overall dipole across the bonds) have a permanent dipole, so the molecules are attracted to each other to form permanent dipole-dipole force.

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Van der Waals’ forces
Van der Waals’ forces are caused by the movement of electrons in the shells. So at any moment, there will be an instantaneous dipole across the molecule. The instantaneous dipole indices a dipole in neighbouring molecules, and the small induced dipoles cayses attraction between the molecules to form weak intermolecular forces (Van der Waals’ forces).

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Hydrogen bonding
A hydrogen binding is a strong dipole-dipole attraction between:
-an electron-deficient hydrogen atom on one molecule; and
-a lone pair of electron on a highly electronegative atom an a different molecule.

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A hydrogen bond is formed by attraction between δ+ and δ- charges on different water molecules. 

Linus Pauling

Linus Pauling

Linus Pauling (1901-1994) was an American chemist, biochemist, peace activist and a educator who won the Novel Price twice – Novel Price in Chemistry (1954) and Novel Peace Price (1962) – for the first and last time, yet.

Linus Pauling with his Novel Peace Prize

Linus Pauling became a peace activist strongly influenced by his wife’s pacifism. He joined Emergency Committee of Atomic Science which was to acknowledge people about the danger of nuclear weapons. His passport got denied from US Department for this but resorted soon before he won a Novel Prize in Stockholm in 1962.

Introduction of electronegativity

Electronegativity is a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond.

As shown in the table above, electronegativity increase as it goes up on the periodic table and along to the right on the periodic table.

Example:
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HCl is an example of a covalently bonded molecule and
• The Cl atom is more electronegative
• The Cl atom has a greater attraction for bonding pair of electrons.
• The bonding electrons are closer to the Cl atom.
There is now a small charge difference across the H Cl bond.
This small charge difference is called a permanent dipole and is shown by:
• a small positive charge on H atom, δ+
• a small negative charge on Cl atom, δ-

A polar molecule has an overall dipole, when you take into account any dipole across the bonds.

C1 1.2 Electrons and the Periodic Table

Sub shells and the periodic table

The periodic table is structured in the blocks of 2,6,10 and 14, linked to sub-shells.

There is s-orbital which contains 2 electrons;

And there is p-orbital which contains 6 electrons;

looks like a dumbell so the electrons are as far away as possible from each other.

There also is d-orbital;

And finally, there is f-orbital;

Sub-shells and energy levels

Within each shell orbitals of the same typr are grouped together as a sub-shell.
Each sub-shell is made up of one type of atomic orbital only. So there are s, p, d and f sub-shells.

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Electron energy levels
The sub-shells within a shell have different energy levels.

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C1 1.2 Chemical bondings

There are three types of chemical bondings –

IONIC BONDING:

  • The bonding between non-metal and metal
  • Happens between oppositely charged ions
  • Electrons are transfered
  • Strong intermolecular forces

COVALENT BONDING

  • The atoms share a pair of electrons
  • There is a electrostatic force which is very strong
  • The intermolecular force is very weak
  • Happens between non-metal and non-metal

METALLIC BONDING

  • Happens between metal and metal
  • There is a sea of electrons, electrons are free to move around

1.1 14 Oxidation number

An oxidation number is a measure of the number of electrons that an atom uses to bond with atoms of another element. Oxidation numbers are derived from a set of rules.

Oxidation numbers rules

Uncombined element – C; Na; O2; P4 – the oxidation number is 0

Combined oxygen – H2O; CaO – the oxidation number is -2
compound
                                   SO4
                    O              -2
                    O              -2
                    S           +4
                 overall      0

Combined hydrogen – NH3; H2S – the oxidation number is +1

Combined ion – Na+,+1; Mg2+, +2 – the oxidation number is charge on an ion

Combined flourine – NaF; CaF2 – the oxidation number is -1

Ionisation energy

Ionisation energy – means the energy required to remove 1 e from every atom in a mole of gaseous atoms to from a mole of mono-positive atoms.

or X(g) = X+(g) + e

Factors affecting ionisation enegy

Atomic radius : The greater atomic radius, the smaller the nulear attraction experienced by the outer electrons

Nuclear charge : -The greater the nuclear charge, the greater the attractive force on the outer electrons

Electron shielding or screening : -Inner shells of electrons repel the outer-shell electrons
-This repelling effect is called electron shielding or screening
-The more inner shells there are, the larger the shielding effect and the smaller the nuclear attraction experienced by the outer electrons.

Successive ionisation energies –

Successive ionisation energies are a measure of the energy required to remove each electron in turn. For example, the second ionisation energy is a measure of how easily a 1+ ion loses an electron to form a 2+ ion.

Example(First 3 ionisation energies for lithium)
: Li(g) – Li+(g) + e1stI.E. = +520kjmol-1
Li+ – Li2+(g) + e 2ndI.E. = +7298kjmol-1
Li2+ – Li3+(g) + e 3rdI.E. = +11815kjmol-1

Each successive ionisation energy is larger than the one before.

Lap report – Titration

Titration

Equipments/materials used:
– Burret
– 25ml pippett
– Pippett filler
– Cornical flask
– HCl
– 25ml unknown base solution
– Phenolphthalein
– Standing clamp

Procedure:
1.Set up the equipments as shown below;

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2.Put 100ml of HCl into the burret and measure out 25ml of unknown solution using pippett and the pippett filler and put into the cornical flask.

3.Put 3-4drops Phenolphthalein into the unknown solution.

4.Place the unknown solution under the burrett and put HCl until the colour of the unknown solution changes.

5.Repeat the titration until you either get 2sets of same results or 4 sets of results.

Result:
Rough-28ml, 1st-26.5, 2nd-26.6, 3rd-26.6

Lab report

Finding the relative atomic mass of an unknown metal by collecting gas

Equipments/materials used:
– 1.00moldm-3 of HCl
– Weighing scale
– 250cm3 measuring cylinder
– 25cm3 measuring cylinder
– Cornical flask
– Delivery tube
– Spatula
– Bowl
– Unknown metal, X

Procedure:
1.Set up the apparatus as shown below;

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2. Place between 0.15-0.20g of the unknown metal in the cornical flask.
*when using weighing scale, make sure the weighing scale is balanced!

3. Measure out 25.0cm3 HCl and add directly to the cornical flask as quickly as possible (because most of gas is given out in the beginning of the reaction) connecting the delivery tube to the flask immediately afterwards.

4.Ensure that all the metal has reacted.

5.Collect the gas and record the final volume if hygrogen.
*To increase the realiability, repeat the experiment several times.

The result:
Unknown salt(g)   Volume of gas(cm3)
    0.21                  146
    0.18                  123
    0.18                  130

No of moles; 0.006083, 0.005125, 0.005416

So, RMM; 34.56, 35.12, 33,23
The average of these is 34.30.

Which means, according to my result, the unknown metal is most likely to be Calcium.

C1.1-5 Types of formula

Empirical formula; is used as the simplest way of showing a formula.

Example:
0.6075g of Mg and 3.995g of Br
Ar: Mg, 24.3; Br, 79.9

Molar ratio- Mg; 0.6075/24.3 Br; 3.995/79.9
So, the ratio is 0.025:0.050 = 1:2
Finally, the empirical formula is MgBr2

Molecular formula tells the number of each type of atom that make up a molecule.

Example:
Empirical formula of CH2 and a relative molecular mass, Mr, of 56.0
Answer:
empirical formula mass of CH2 = 12.0 + (1.0*2) = 14.0
number of CH2 units in a molecule = 56/14 = 4
Molecular formula: (4*CH2) = C4H8