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UAB CH 115 - Unit 1: Atoms, Isotopes and Ions

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CH 115 1st Edition Lecture 7Unit 1: Atoms, Isotopes, and IonsOutline of Last Lecture I. Shapes of Quantum Numbersa) Principalb) Subsidiaryc) Magnetic II. Energy Levels of Electronsa) Rulesi. Aufbau Principleii. Pauli Exclusion Principleiii. Hund’s Ruleb) Orbital Notationi. Box vs. Energy DiagramsOutline of Current Lecture III. Writing Electron Configurationsa. Expanded, Condensed, and Noble Gas Configurationsb. Atoms vs. IonsCurrent LectureWriting Electron ConfigurationsThese notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.- Important Information: Core electrons – electrons in an atom’s inner energy levels Valence electrons – electrons in an atom’s outermost energy levels Paramagnetic – substances that have unpaired electrons (the prefix “para” meansside by side, and when electrons are unpaired they look as if they are PARAllel to each other). Diamagnetic – substances that have filled subshells and all of their electrons are paired (the prefix “di” means 2, so there are 2 paired electrons) The strength of magnetic attraction is proportional to the number of unpaired electrons.- Different Electron Configurations: Expanded– In expanded form, write each orbital with electrons like this  Condensed – In condensed form, write eachsubshell with electrons like this  Noble Gas – In noble gas form, write the next smallest noble gas in brackets. Write valenceelectrons in condensed notation like this - There’s added stability when a subshell is filledor half‐filled. Atoms are always trying to be morestable.- Example: Nitrogen’s electronconfiguration is more stable thancarbon’s or oxygen’s, as shown here: - When filling transitionmetalelectrons, fill the s‐subshell first and then fill the d‐subshell.- Special cases do exist in which you would try to either fill or half-fill the d-subshell first.The reason for this is because the atom will be more stable if one of the subshells is filledor half-filled instead of both of them not being filled or half-filled at all. These special cases include Cr: [Ar]4s13d5; Mo: [Kr]5s14d5; Cu: [Ar]4s13d10; Ag: [Kr]5s14d10- Ions have electron configurations as well, but they differ from an atoms’ electron configuration.o Atoms – have no net charge. Atoms are neutral.o Ions – have either a positive or negative net charge. Ions are either positively or negatively charged.o A monatomic ion – formed when an atom either gains or loses electronso Cation – A positively charged ion.o Anion – A negatively charged ion.- When writing the electron configuration for an ion, find the electron configuration of its neutral atom first and then add or subtract electrons from its outer most subshell. For example, the electron configuration of Nickel (Ni) is [Ar] 4s23d8. The electron configuration for Ni2+ (remember, this is a cation of Ni) is [Ar] 4s23d6. - If you are trying to find the electron configuration of a cation, you are going to subtractelectrons from the electron configuration of the neutral atom.- If you are trying to find the electron configuration of an anion, you are going to add electrons from the electron configuration of the neutral atom.However remember the special cases as well!o It may seem counter intuitive that you would add electrons to the electron configuration of the neutral atom for anions and subtract electrons from electron configuration of the neutral atom for cations, but it comes from the fact that it is called an electron configuration.o The electron configuration tells us how many electrons are in each orbital/subshell of the atom in question.o Cations have a positive charge, meaning that they lost electrons.o If the atom is originally neutral it has the same number of electrons and protons, but if it loses an electron there are now more protons than electrons, giving the newly formed monoatomic ion a positive charge (cation).o Thus, if the electron configuration is 1s2 for the neutral atom, it would be 1s1 for a cation and 1s22s1for an anion.- Main group elements will gain or lose electrons to form ions that have the same electron configuration as the noble gas that is closest in atomic number to


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UAB CH 115 - Unit 1: Atoms, Isotopes and Ions

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