3. Periodicity

• Decreases across a period • nuclear charge increases, while the electrons remain on the same valence shell • Increases down a group • Even though the nuclear charge is increasing, there are more electrons, which result in more electron shielding

• Cations get smaller • They lose their valence e-, thus the remaining electrons are pulled tighter to the nucleus • Anions get larger • Electron shielding weakens the force between electrons and the nucleus

• Increases across a period • nuclear charge is increasing despite the increase in electrons • Decreases down a group • electrons are further away from the nucleus in addition to electron shielding

A measure of the ability of an atom to attract bonded electron pairs to itself when in a covalent bond • Metals have small electronegativity values • Nonmetals have higher electronegativity values

The energy released when 1 mol of electrons is attached to 1 mol of neutral atoms or molecules in the gas phase • same increase/ decrease as Ionization Energy

• Low melting / boiling points • decreases down the group –> increase in shielding and less nuclear shielding • Very reactive • needs to lose only one electron to have a full shell • Large atoms • metals are soft and not dense

With water: • React vigorously with water to create a hydrogen and a base • Na(s) + H2O(l) –> NaOH(aq) +H2(g) With halogens: • React vigorously with halogens to form salts • 2Na(s) + Cl2(g) –> 2NaCl(s)

• Very reactive / electronegative • Needs to gain just one electron to have a full valence shell • High melting / boiling points • Increases down the group –> increased LDF between the simple covalent molecules and increased molecular weight • Down the group • becomes darker • decreases in reactivity • A halogen higher up the group will displace the one lower down

• Oxides of elements have increasing acidity across a period • Al is amphoteric • Metal oxides are basic • Nonmetal oxides are acidic

An element that possesses an incomplete d-sublevel in one or more oxidation states zinc is not a transition element because it has a full d-sublevel in its atoms/ ions Transition elements have an empty d orbital - the d orbitals splits into 2 energy sublevels and electrons moving between these gives them their properties • Cu / Cr - it is more energetically favorable to half-fill and completely fill the d sublevel, so they only contain one 4s e-

• produces colors and allows for complex ion formation • variable oxidation numbers - good catalysts • when transition metals lose electrons, they lose the 4s first • all transition metals have an oxidation state of +2 • occurs when they lose both s orbital e-s • most transition metals are diamagnetic (repelled) • only a few are paramagnetic due to unpaired electrons allowing spin in one direction to form poles

A neutral molecule / anion which contains a non-bonding pair of electrons, these form coordinate bonds with the metal ion to form complex ions - donates a pair of electrons to a central metal ion to form a covalent bond

Consists of a central atom, which is usually a metal atom/ ion and attached ligands Coordination number: the total number of points at which a central atom / ions attaches ligands • 2 - linear • 4 - square planar / tetrahedral • 8 - octahedral The charge on a complex ion is the sum of the charges of the metal ion and the ligands

Transition metal ions are colored due to the d-d electron transitions between d orbitals which are split in the electric field due to the presence of ligands • color obvserved is complementary to the color absorbed • ions with higher charge and ligands with greater charge density produce a greater split in the d orbitals • polydentate ligands form more than one coordinate bond with the metal ion