Chemistry 111 1st Edition Exam # 2 Study GuideChapter 3BThere are four main atomic orbitals which are s, p, d and f.These orbitals can hold a maximum number of electrons: s=2, p=6, d=10, f=14Aufbau Principle (“building up”)- Electronic configuration which consists of adding electrons one by one, always placing the electrons on the lowest available energy orbitals.Orbital Diagrams-A representation of the number of electrons in an atom by using boxes to symbolize different orbitals. These diagrams will use arrows to represent the electrons in each of the orbitals which must have opposite spins (Pauli Principle).Hund’s Rule: “fill up before pair up”. All orbitals must first be filled with unpaired electrons before pairing them up if possible. This may be done by breaking apart paired electrons or taking electrons from other incomplete orbitals in order to fully complete another. Chromium and Copper are both exceptions to this rule which is called anomalous configurations as they are more stable with two half-filled shells rather than a single completelyfilled one.Formation of Ions - The gain or loss of valence electrons to achieve a stable state or a complete shell. Isoelectronic Atoms-Different atoms with the exact same electron configuration like Na+, O2-, and Mg2+Atomic Radius: the distance between half nuclear centers in a molecule. The more protons an element has, the stronger the effective nuclear charge will be, thus the atom’s electrons will more closely attracted to the nucleus, making the atomic radii smaller.Ionization Energy (IE): refers to the amount of energy required to remove an electron from an atom. The more electrons an atom has, the easier it is to remove an electron and the less energy required. Helium (He) will have higher ionization energy than Magnesium (Mg).Note: When ionization energies are successive, the removal of each electron becomes harder than the previous and requires more energy. IE (2) > IE (1) When we are given successive ionization energies such as:IE (1) → 2150 kJ/molIE (2) → 2450IE (3) → 5300IE (4) →5800We are able to determine the element to which this data belongs to by looking at the biggest jump in energy. In this case, the energy value has the highest jump from IE (2) → IE (3) which means that this element will have 2 valence electrons. Electron Affinities (EA): the gaining of an electron. Does not include noble gases because they are already stable. EA values become more negative moving up and to the right in the periodic table. F + e- → F-Chapter 4ATypes of Chemical BondsIonic Bond: cation + anionCovalent bond: sharing of the outermost electron pairsMetallic bond: metal atoms surrounded by shared electronsBinary Ionic CompoundsName of the cation + name + suffix –ideTheir formulas must always be neutral*NaF – sodium fluorideK2O – potassium oxideFor transition metals which have different charges, we add a Roman numeral to indicate the charge of the cation.FeCl2: iron (II) chloride FeCl3: iron (III) chlorideRemember the Roman numeral refers to the charge, not the amount*Polyatomic AtomsTwo or more atoms joined by covalent bonds which possess a chargeKnow the following polyatomic atoms and their formulas: ammonium hydroniumAcetateCarbonateBicarbonatePerchlorateChloratechloritehypochloritechromatedichromatecyanidenitratenitritepermanganatephosphatephosphitesulfatesulfitehydroxideOxoanionsPolyatomic atoms containing oxygen with other elementsBinary Molecular CompoundsCompounds consisting of two nonmetalsName of nonmetal + prefix (number of atoms) + name of nonmetal + suffix –idePrefixes1-mono, 2-di, 3-tri, 4-tetra, 5-penta, 6-hexa, 7-hepta, 8-octo, 9-nona, 10-decaSO3 – sulfur trioxideNO2 – nitrogen dioxideBinary AcidsHydrogen + Monoatomic AnionHydro + name of anion+ suffix –ic + acidHBr – hydrobromic acidHCl – hydrochloric acidOxoacidsIf oxoanion name ends in –ate, the corresponding acid ends in –ic + acidIf oxoanion name ends in –ite, the corresponding acid ends in –ous + acidH2CrO4 – (chromate) → chromic acidHNO3 – (nitrate) → nitric acidHNO2 – (nitrite) → nitrous acidI ate a stic and ite was delicious*Know how to draw and understand Lewis Structures.Chapter 4BAllotropesDifferent molecular forms o f the same elementsO2 and O3 (ozone)Resonance StructuresTwo or more Lewis structures can be drawn for the same compound with different bonding arrangements like NO3 and O3.Bond OrderNumber of bonds / Number of outer elementsFormal ChargeValence electrons – (Unshared electrons + ½ (number of electrons in bonding