1. Periodic trends

2. Things to understand about the periodic table
Understand how the following properties relate to the element’s position on the periodic table
Atomic Radii
Distance between the nucleus and outer surface of the largest orbital
Ionization Energy
Energy required to remove an electron from an atom in the gas phase
Electron Affinity
The energy change when an electron is added to an atom.
Electronegativity
The relative strength of the nucleus to attract electrons.

3. Atomic Radius
The distance between the nucleus and the edge of the outermost orbital Measured by finding half the distance between two atom nuclei The “Atomic Radius” is quite different from the “Ionic Radius”

4. Atomic Radius
As n increases, atomic orbitals become larger and less stable. As Z increases, any given atomic orbital becomes smaller and more stable.

5. Radii trends
General decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus. Atomic radius generally increases as you move down a group The outermost orbital size increases down a group, making the atom larger.

7. Ionic Radii
Losing electrons reduces the atom size
Loss can leave orbital empty
Repulsion between electrons in valence shell is less, reducing size
Adding electrons increases atom size
adds new shells or orbitals
Increases e- - e- repulsion increasing size

8. Ionic radii
Positive and negative ions still increase down the group (electrons in heavier ions still occupy larger orbitals)

9. Isoelectronic atoms
Atoms and ions with the same number of electrons Example: O2–, F–, Ne, Na+, Mg2+ all have 10 electrons Trend in radius: O2– > F– > Ne > Na+ > Mg2+ Even if they have the same number of electrons a greater negative charge means more repulsion = larger ion Likewise a large positive charge means greater attraction between the electrons and the nucleus = smaller ion

10. Ionization energy
Energy required to remove an electron from a gas atom First electron removed requires the “first ionization energy” Electrons are removed in order from the highest n and l (either the s or p orbitals) e.g. if atom is [Ne]3s23p3, i.e. P; when given sufficient ionization energy it becomes: [Ne]3s23p2 or P+

12. Other ionization energies
Each additional ionization energy requires much more energy Each successive ionization in the same orbital type (spdf) will require a steady increase Successive ionizations of different orbital types require larger jumps in energy (because the orbitals are even closer to the core, e.g. 3d vs 4s) These patterns can be used to identify an element.

14. Ionization energy trend
As size increases electrons are less tightly held Generally decreases ionization energy More valence electrons increases ionization energy First ionization increases across a period The electrons of atoms in the same period have the essentially the same “n”, but more protons. This implies a greater attraction to nucleus, i.e. it is harder to remove electrons When we change rows, (increase n) we decrease the difficult to remove electrons, since they are now farther from the core

15. Electron Affinity
The energy released when an atom attracts an electron Tends to become more negative (more energy released) as we move from left to right along a period. E.g. Chlorine becoming Cl– releases a lot of energy (very favourable)

16. Electronegativity
The balance between “ionization” (loosing electrons) and “affinity” (gaining electrons) is known as “electronegativity”
Electronegativity indicates an atoms ability to attract electrons in a chemical bond.
The greater the difference in electronegativity the more ionic a bond.
Decreases down the group and increases to left to right
Larger atoms have lower electronegativity

17. Ionic vs Covalent bonds
A chemical bond forms when two atoms approach one another As the positive nuclei approach each other, the negative electrons may become attracted to the other atom The “atomic orbital” becomes spread out between the two atoms, creating a “molecular orbital” If the electronegativity difference between the two atoms is >2.0; then the electrons are transferred to the other atom, creating an “ionic bond” If the difference in electronegativity is < 2.0 the bond is “polar covalent”, the electrons are mostly shared. If the difference is < 0.4 it is a pure covalent bond