Chem: Periodicity

• Atomic radius


  Not how big the atom is, but measured as half the distance between neighbouring nuclei
• As you go from left to right


  The effective nuclear charge increases because as you are adding protons, you are still in the same energy level so the inner electrons are not increasing
• Ionic radius


  The atomic radius of an ion in ionic crystal structures
• First ionization energy


  The energy required to remove one mole of electrons from one mole of gaseous atoms
• The small dips in the general trend across the period are caused by the sub-orbitals. The d-orbital fills one electron in each orbital first then they fill the top level after. Removing the bottom level requires more energy than the top
• Electron affinity


  The energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions
• Electronegativity


  The measure of the tendency of an atom in a molecule to attract a shared pair of electrons towards itself
• Metallic character


  Metals are characterised by having low ionisation energies, which means that they lose their valence electrons relatively easily to form positive ions
• Metal oxide


  A metal bonded with oxygen
• Metal oxides reacting with water


  An oxide is very strongly basic ion and due to small size and high charge it forms hydroxide ions when it reacts with water
• Acidity/basicity of an oxide

  More electropositive central atom = More basic, More electronegative central atom = More acidic
• Alkali metals


  Located in group 1 of the periodic table, a very reactive group of metals with properties quite different from those like iron and copper
• Halogens


  The elements in group 17, known as the halogens, are a very reactive group of non-metal elements
• The periodic table is organized by atomic number, which determines the number of protons (and electrons) in an atom.
• Isotopes are atoms with different numbers of neutrons but the same number of protons.
• Elements within a given period have similar chemical properties because they have the same number of electron shells.
• Atomic mass is determined by adding up the masses of all subatomic particles in an atom.
• Periods increase in energy level as we move down the periodic table.
• Elements can be classified into groups based on their chemical properties.
• Electron configuration refers to the arrangement of electrons around the nucleus of an atom.
• Atomic radius decreases across a period because there are more protons pulling the electrons closer together.
• Periodicity refers to the repetition of trends within periods or rows of the periodic table.
• Relative atomic mass is calculated based on the relative abundance of naturally occurring isotopes.
• Atoms can gain or lose electrons to form positive or negative ions.
• Group 1 elements are called alkali metals and have one valence electron that easily loses to form positive ions.
• Ionic compounds form when one or more metal atoms lose electrons to become positively charged cations, while nonmetal atoms gain electrons to become negatively charged anions.
• Group 1 elements are called alkali metals and share common physical and chemical characteristics such as low density, softness, reactivity, and high melting points.
• Ions can be represented using symbols such as Na+, Cl-, O2-, etc.
• Alkali metal compounds dissolve easily in water to form hydroxides that neutralize acids.
• Group 2 elements are called alkaline earth metals and have two valence electrons that also lose easily to form positive ions.
• Covalent bonds occur between two nonmetal atoms sharing pairs of electrons.
• Ionization Energy (IE) is the amount of energy required to remove one mole of gaseous atoms and convert them into one mole of positive ions.
• The atomic radius increases going down a group due to increased shielding from inner shells, allowing outer shells to expand further away from the nucleus.
• The number of valence electrons determines the group number of an element.
• Metallic bonding occurs when positively charged metal ions attract delocalized electrons, resulting in strong metallic bonds.
• The first ionization energy increases from left to right due to increasing nuclear charge, which attracts valence electrons more strongly.
• Metals have low ionization energies compared to non-metals.
• Metals have high melting points, low densities, and good conductors of heat and electricity.
• Group 2 elements are called alkaline earth metals, with two valence electrons that also readily lose to form positive ions.
• Ionization energy is the amount of energy required to remove one mole of gaseous atoms' outermost electron(s).