Chem 105 Ch. 4 and 5
Chem 105 Ch. 4 and 5 Chem 105
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This 3 page Class Notes was uploaded by Anagrace Salem on Thursday October 6, 2016. The Class Notes belongs to Chem 105 at University of Mississippi taught by Dr. Gerald Rowland in Fall 2016. Since its upload, it has received 3 views. For similar materials see General Chemistry 105 in Chemistry at University of Mississippi.
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Date Created: 10/06/16
Degenerate: orbitals with the same energy Half d orbitals are signiﬁcantly more stable Valence electrons are the electron in the outermost shell (highest energy level) Group 6 and 11: the ones where they subtract an electron from the ground state orbitals and add one to a valence electron to have more stability Metalloids: semi conductivity, highly temp dependent, used in circuits, cellphones, etc. (Ex: Silicon) Metals: good conductors for heat and electricity, malleable, ductile, form cations Non-Metals, poor conductors for heat and electricity, brittle, form anions Atomic Radius: -increase when you go down a group, decrease when you go across a period from left to right -when the number of electrons increase and is added to the same shell, valence electrons are held tighter -valence shell is farther from the nucleus (higher energy level) -valence electrons occupy larger orbitals, which results in a larger size Increase in effective nuclear charge, because of adding additional protons Isoelectronic: the same electron conﬁguration as other elements. the more negative, the more larger Ionization Energy: energy required to remove an electron from an atom/ion from the gas state -Decreases when going down a group -Increase when going across a period -endothermic -valence electrons are easiest to remove because lowest ionization energy because they’re farther from the nucleus -metals are easiest group to remove because they require less ionization energy Electron Afﬁnity: the charge in energy of a neutral atom when an electron is added to the atom to form a negative ion. -Mostly exothermic -Halogens have the highest -The more negative, the more energy is being released -Increase from L to R and becomes more negative Chemical Bonds: interactions that keeps compounds together Ionic Bond: between a non-metal (anion) and metal (cation)/strongest chemical bonds Covalent Bond: between non-metals, forms molecular compounds/molecules, sharing electrons Empirical Formula: small number ratio Molecular Formula: exact number of each type of atom in a compound Structural Formula: shows how the atoms are connected to one another, lines represent covalent bonds and a pair of electrons Ionic Formula: charges must cancel out, neutral charge formula, wants the smallest whole number ratio Ionic Compounds: when ions are held together by electrostatic forces (transition metals have only one charge in these compounds) Type 1: Metals can have one charge -when naming them add the cation + anion with the ending -ide -Ag, Al, Ba, Be, Ca, Cs, K, Li, Mg, Na, Rb, Sc, Sr, Zn Type 2: more than one charge (primarily transition metals) -when naming them, add cation + (roman numeral) + anion with the ending -ide -Br, Cl, F, I, N, O, P, S -Ag+, Co2+, Co3+, Cr2+, Cr3+, Cu+, Cu2+, Fe3+, Hg2+, Pb2+, Pb4+, Sn2+, Sn4+, Zn2+ Polyatomic Ions: groups of atoms with charges / an ion composed of 2 or more atoms with an overall charge Only Polyatomic Cation is Ammonium -1 Charge Polyatomic Ions: NO3 (Nitrate) NO2 (Nitrite) ClO3 (Chlorate) ClO2 (Chlorite) ClO (Hypochlorite) ClO4 (Perchlorate) C2H3O2 (Acetate) HCO3 (Hydrogen Carbonate) OH (Hydroxide) MnO4 (Permanganate) HSO3 (Hydrogen Sulfate) CN (Cyanide) -2 Charge Polyatomic Ions: SO4 (Sulfate) SO3 (Sulﬁte) CO3 (Carbonate) CrO4 (Chromate) Cr2O7 (Dichromate) HPO4 (Hydrogen Phosphate) H2PO4 (Dihydrogen Phosphate) O2 (Peroxide) -3 Charge Polyatomic Ions: PO4 (Phosphate) Some Metals that form cations with different charges: Cr, Fe, Co, Sn, Hg, Pb
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