John Newlands - Idea of Repeating Octaves of Properties

John Newlands - idea of repeating octaves of properties

Dimitry Mendeleyev - arranged known elements according to atomic weights and properties. Made predictions which were later proven accurate.

Henry Moseley - utilised X-Ray. Said that periodic table should be based on atomic numbers of elements.

Dalton = atom as the basic unit of an element that can enter into a chemical combination// very small and indivisible// everything is made up of discreet building blocks, that is atoms.

J.J Thompson - electrons - cathode ray - studying the deflection of a ray of high-speed electrons fired at a fluorescent screen and deflected with the aid of magnetic fields and electric fields.

The charge of an electron is -1.60 * 10^-19 C // The mass of an electron is 9.09 * 10^-28 g.

Fe2+ is ferrous ion // Fe3+ is ferric ion // Silicide = Si4- // Nitride N3- // Phosphide P3- // Sulfide S2- // Carbonate (CO3)2- // Chromate (CrO4)2- // Dichromate (Cr2O7)2- // Permanganate (MnO4)- // Peroxide (O2)2- // Sulfide S2- // Sulfite (SO3)2- //

An example of electron configuration : 1s2,2s2,2p6,3s2,3p6,3d10,4s2,4p6,4d10,4f14....

The type of orbital of the first element in a period corresponds to the period number. For example, Lithium is the first element in period 2. It's valence electron is in the 2s^1 orbital.

The first ionisation energies increase from left to right across a period.

Atomic radii decrease from left to right across a period due to the increase of nuclear charge when additional protons pull additional electrons closer to the centre. // Inert gas elements have a slight increase in atomic radii because of the electron repulsion of the filled valence shell.

Ionic radius differs from atomic radius. // For metals, the electron(s) lost causes the radius to decrease because of the lost of negative charge// Non-metals tend to gain electrons. Therefore, the increase in negative charge causes the ionic radius to be higher than the atomic radius.

The ELECTRO-NEGATIVITY of an element is a number that measures the relative strength with which the atoms of the element attract valence electrons in a chemical bond. // Electronegativity increases as we move across the periodic table because of the decrease in atomic radius.

The most electronegative element is fluorine because of its small atomic and ionic radius.

Ionisation Energy pertains to the removing of electrons from an atom. The first ionisation energy is the amount of energy required to remove one electron from the valence shell of an atom. The more valence electrons are remove, the higher the ionisation energy is because of the imbalance of nuclear charges. The lowest ionisation energy is found with the least electronegative atom. Therefore, the more electronegative an atom is, the higher its ionisation energy is. // Noble gasses have the highest ionisation energies but it decreases as we go down the group because of the increase in atomic radius. // The lowest ionisation energy occurs when a lone electron fills the outer shell or the valence shell of an atom.

The Ionization potential-atomic number graph is jerky because dips happen when there is a lone electron in the valence shell (so although the number of electrons has increased by one, the ionisation energy is lower) AND small peaks happen when an orbital is filled (like when a 2s1 orbital is filled, or when a the 3 squares in a 2p shell is filled each with one electron. So, although the configuration isn't exactly like an octet, it is similar and therefore the ionisation energy increases slightly). The peaks and dips are always related to the state of filling of the orbitals involved and the distance of these orbitals from the nucleus.

Electron Affinity pertains to the ability to accept one or more electrons. It is the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. The sign of the electron affinity is opposite that of the ionisation energy. Large positive electron affinity means that the negative ion is stable.

Diagonal Relationships - The first elements of groups in the periodic table are somewhat differing in their characteristics compared to the their group members. This is usually attributed to the unusually small size of 1st elements; they have comparably similar charge densities with their diagonal partner.

Peroxides - compounds with oxygen in their elements. The oxides have a negative charge of 2-.

Superoxides - compunds with oxygen in their elements. The oxides have a negative charge of 1-.

The reactivity with water of the elements in the alkaline earth metal group increases as we go down the group.

The VSEPR model for predicting molecular geometry is relatively simple but insufficient. It does not consider the differences in the type of covalent bond, for example(whether single, double or triple). Models who are more complex but more complete are the valence bond theory and the molecular bond theory. Molecular bond theory goes a little further than the valence bond theory because it explains how some molecules which, according to the valence bond theory, should be diamagnetic but are actually paramagnetic. An example is the O2 molecule which displays paramagnetic characteristics.

In molecular orbital theory, we have bonding and antibonding orbitals - the former having lower energy and greater stability and the latter having greater energy and lower stability.

Bonding

Valence electrons shift when elements combine to form substances.

When electronegativity values of 2 reacting atoms differ by 1.7 or more, an ionic bond is formed. Otherwise, the electrons are shared(covalent), either as non-polar or polar bonds.

REMEMBER! Some nonpolar molecules contain polar covalent bonds. Because they are arranged symmetrically, the characteristics of polar bonds don't show and therefore, the molecules are not polar molecules.

Coordinate covalent bonds occur when only one of the atoms supplies the electrons needed to form a covalent bond. E.g.: NH4+ and H2SO4.

In most metals, one or more than one of the valence electrons become detached from the metal atom and become free electrons. There exists a bond between the atom and the electrons because of the nuclear positive charge which is called metallic bond.

All of the intermolecular forces can be referred

to as Van