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IB Chemistry Summary- By Paul Li & Silvia Riggioni

Table of Contents

Atomic Theory

The Electromagnetic Spectrum

  • Electromagnetic radiation is a form of energy.
  • The smaller the wavelength the higher he frequency => higher energy of the wave
  • Radio waves, microwaves, infrared, visible light, ultraviolet, x-ray, gamma rays

←                  Increases                ←

  • Velocity of waves = frequency x wavelength ()
  • Electromagnetic radiation comes in packages called quanta or photons.

Atomic Emission Spectra

  • White light is made up of all the colours of the spectra.
  • When it passes through a prism a continuous spectrum is obtained.
  • When energy is applied to specific (individual) elements they emit a spectrum which only contains emissions of particular s.
  • A line spectrum is not continuous. Each element has its own characteristic line spectrum.
  • Hydrogen spectrum- it consists of discrete lines that converge towards the high energy end of the spectrum. The lines converge as the shells are getting closer together. Energy levels increase because we get a higher frequency and a smaller wavelength. ()

Explanation for the Emission spectra

  • Electrons can only exist at specific energy levels
  • When energy is supplied to an atom e- are exited from the lowest (ground) state to an exited state.
  • When e- drop from a higher level to a lower level they emit energy (a photon of light). This energy corresponds to a particular and shows up as a line spectrum.
  • Jumps to the n=1 have the highest  and the smallest.

  • The emission spectra can also be used to find the ionization energy. This is done using the Rydberg equation.
  • First Ionization Energy→ the energy required to remove one electron from each atom of a mole of atoms in the gas state, to form one mole of cations in the gas phase, under s.t.p

Subatomic Particles

  • Almost all the mass of the atom is concentrated in the nucleus which has a very small radius.
  • Much of the atom is empty space
  • Electrons and protons are deflected by and electric field. Neutrons aren’t.

Atomic Numbers

  • It is the number of protons in the nucleus of an atom.
  • Defines which element the atom belongs to and consequently its position in the Periodic Table.

→ Z is the Atomic Number, X the symbol.

Mass Numbers

  • It is the sum of the number of protons plus the number of neutrons in an atom or ion.

→ A is the Atomic Mass, X the symbol.

  • The relative atomic mass is the measure of the average mass, taking into account the various types of isotopes. E.g. RAM of chlorine is 35.5, because Cl-35 is three times as abundant as Cl-37.


  • Two or more atoms of the same element which have the same number of protons (Z) but different number of neutrons.
  • All isotopes react in the same way. However, the different masses will affect the physical properties such as density and the rate of diffusion of both elements and compounds.

Calculating RAM by example – Lead (Pb)

                                        Abundance in %





Average of 100 atoms:

                =               306

        =          4,861.6

        =          4,678.2

        =        1,0878.4





                                                207.242  → RAM of Pb

If the abundance is not given as a percentage then divide by the total abundance.


  • Solute        →        Substance that is going to be dissolved.
  • Solvent        →        The liquid where we are going to dissolve the solute.
  • Solution        →        Solute + Solvent



  • It is a pure substance which can’t be made simpler by any chemical method and is made up of atoms, all of which have the same atomic number.
  • A pure substance is made up of only 1 type of atomic number.
  • Arranged in the periodic table by increasing atomic number.
  • The majority are metals.

Physical Properties

  • A covalent radius is half the4 minimum distance between the nuclei of 2 atoms of the same element covalently bonded in a diatomic molecule.
  • A Van der Waal’s radius is half the minimum distance between the nuclei of two atoms in the same element, which are NOT chemically bonded.

  • Atomic radius:         radius of an atom. Half the distance between the nuclei of adjacent atoms.
  • Ionic radius:         radius of an anion or cation.

Atomic and Ionic Radii in the Periodic Table

  • The Alkali Metals (Group I)
  • As you go down the group, both the atomic and the cationic radius increase.
  • Cations have a smaller atomic radius than the parent atom they come from because they have lost an outer shell electron.
  • The ions get smaller as they have the same electronic structure but a greater nuclear attration.
  • The Halogens (Group VII)
  • As you go down the group, the anionic ionic radius increases.
  • Anions have a larger atomic radius than the parent atom they come from, because they gain an electron in the outer shell. The electrostatic repulsion makes the outer shell expand.
  • There are more e- for the same nuclear charge, so each is held less strongly and thus, can be further away from the nucleus for the same energy.
  • Across a Period
  • The atomic radius decreases because the electrons are being pulled closer to the nucleus due to the increase in charge.  E.g. Na → Ar
  • The ions across the period have the same ionic structure, but an increase in the number of protons increases the electrostatic forces between the protons and electrons. Thus, the radius decreases.
  • Ions are isoelectronic, meaning that they have a similar electronic structure which resembles noble gas structures. E.g. Nitrogen’s is similar to Neon’s.
  • Down a Group
  • Down any group the ionic radius increases as there are more electrons in more shells, which are further from the nucleus and with more shielding. E.g. Li+, Na+, K+
  • As we go down the groups, the outermost electron is in a higher energy level, which is further from the nucleus, so the radius increases.

Electronegativity (Pauling’s)

  • It is the ability of an atom in a covalent bond to attract electrons to itself. It is a relative measure; hence it doesn’t have any units.
  • Electronegativity depends on 3 things:
  • Real nuclear charge, Z
  • Number of screening electrons (repulsion electrons)
  • Atomic Radius

  • If two elements have similar electronegativity, they bond covalently. If the difference is of 2ish or more, they will form ionic bonds.
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Electronegativity in the Periodic Table

  • The Alkali Metals (Group I)
  • These elements have low values for electronegativity
  • As you go down the group, electronegativity decreases, even thought the number of protons is increasing.
  • The Halogens (Group VII)
  • These elements have relatively high electronegativity values.
  • As you go down the group, electronegativity decreases, even thought the number of protons is increasing.
  • Across a Period
  • Electronegativity increases across the Periodic Table from left to right.
  • This is due to the fact that the nuclear charge increases and the atomic radius decreases slightly.
  • Down a Group

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