QCE Chemistry U1, U2, U3, and U4

• The atomic number is the number of protons in the nucleus of an atom. The number of protons is equal to the number of electrons in a neutral atom
• Mass number is the sum of the number of protons and the number of neutrons in the nucleus of an atom.
• There is a 95% chance of finding electrons in the s, d, p, f electron orbitals.
• Each orbital can hold a maximum of two electrons
• The Aufbau principle - electrons fill the lowest energy sublevels first.
• Hund’s Rule – in each sublevel, orbitals half fill first, before containing an electron pair
• Pauli exclusion principle – paired electrons have opposite spin. This reduces mutual repulsion between the electrons.
• For the electron configuration of transition metal ions, take electrons first from 4s2.
• Chromium – 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1
• Copper – 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s1
• When electrons are excited to a higher energy level, and then return to a lower energy level, they release a photon of a specific energy, as shown by a specific frequency of light.
• Electrons moving back to the lowest energy states and over the longest distances release the highest E (short λ).
• Group – vertical columns of the periodic table which contain elements having similar chemical and physical properties
• Nuclear charge – due to number of positive protons in the nucleus.
• Valence electron shell – outer electron shell, hence these are the electrons that mainly determine physical and chemical reactions.
• Electron shielding – the interference of attraction of electrons to the nucleus by other electrons.
• Atomic Radii: decrease across a period,

  increase down a group.
• Atomic radii increase down the group as there is an added electron shell.
• Atomic Radii decrease across a period (except added subshell anomalies) due to increasing nuclear charge within the same valence electron shell, drawing the outer electrons closer to the nucleus.
• Ionic Radii: Cations get smaller because they lose electrons to have a full valence shell that is of lower energy
• Ionic Radii: Anions get larger because they gain electrons to fill their current valence electron shell
• First Ionization Energy – The energy required to remove one mole of electrons from a mole of atoms or ions in the gaseous phase
• X(g) ---> X+(g) + e-
• As the energy of the electron increases, the electron is farther away from the nucleus.  As a result, the force of attraction between the nucleus and the electron decreases.
• IE increases across a period because the nuclear charge increases, attracting the electrons and the number of electrons in the shell are increasing.
• IE decreases down a group because the electrons further away from the nucleus and electrons in lower shells are blocking the attraction causing electron shielding.
• As the size of the nuclear charge increases the force of the attraction between the negatively charged electrons and the positively charged nucleus increases. 
• Ionization Energy Graph: Large jumps are changes in energy levels, medium jumps are changes in sub-levels, and small jumps are due to mutual repulsion within a sublevel.
• As electrons are removed, the charge difference is greater. So the relative number of protons to electrons increases and so it takes more energy to remove the next electron.
• Electronegativity is a measure of the ability of an atom to attract bonded electron pairs to itself when in a covalent bond
• The Balmer series includes the lines due to transitions from an outer orbit n > 2 to the orbit n' = 2.
• Ionic Bond – electrostatic attraction between oppositely charged ions
• The degree of polarity can be determined by looking at the difference in electronegativity. 0.0-0.4 is nonpolar covalent bond, 0.4-1.8 is polar covalent bond. >1.8 is ionic bond.
• Polyatomic ions are covalent molecules with a charge that act as a unit in ionic compounds.
• Metallic bonds are the electrostatic attraction between a lattice of positive ions and delocalized electrons.
• A covalent bond is the electrostatic attraction between a shared pair of electrons and positively charged nuclei
• Covalent compounds are formed by sharing electrons. To name these compounds you must add prefixes to show how many atoms are in the formula. You know it is covalent because it will be a non-metal combined with a non-metal.
• For alkali metals, reactivity increases down the group as electron is easier to remove
• For halogens, reactivity increases up the group as electron is easier to gain
• Isotopes are atoms of the same element with the same atomic number, but different mass numbers, i.e. they have different numbers of neutrons.