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UMass Amherst CHEM 110 - Chemistry Chapter 7

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Chemistry Chapter 7Key IdeaThe number and types of bonds an atom will form, can be derived from the number andarrangement of the atom’s electrons● Electron configurations are a list of all electrons in an element and the orbitals theyoccupy● Electron configurations allow us to predict physical and chemical properties of theelements● Orbitals: Single vs Multi-electron○ Similarities■ Similar in shape■ Identical in nodal structure.○ Differences■ Each multi-electron orbital is smaller and lower in energy than thecorresponding hydrogen atom orbital. Since they are lower in energy, theorbital gets closer to the nucleus.■ In multi-electron atoms, the orbital energy depends on the shell (n) andsubshell (l)● Aufbau Principle○ Lower energy levels should be filled before higher energy levels■ Example: A 4s orbital is higher in energy than a 3s orbital○ As ℓ increases, orbital energy increases■ Example: In n = 3 shell, 3s<3p<3d○ Electrons fill orbitals in order of increasing (n+l) values■ Example: The 4f ( 4+3=7 ) orbital is higher in energy than the 5s orbital (5+0=5 ) , despite its lower n value○ If two subshells have the same n+l, the subshell of lower n is filled first■ Example: 3p (3+1=4) orbital is filled first then the 4s (4+0=4) orbital● Pauli’s Exclusion Principle○ States that no 2 electrons within an atom can have the same set of four quantumnumbers (n, ℓ, mℓ , and ms)■ This means that an orbital can accommodate no more than 2 electrons○ Any two electrons occupying the same orbital must have opposite spins■ Allowed values of ms are + ½ or - ½● Hund’s Rule of Maximum Multiplicity○ Every orbital in a subshell is singly occupied with one electron before any oneorbital in that subshell is doubly occupied○ All electrons in a singly occupied subshell have the same spin (parallel spins)● Elements within a group have the following in common:○ They have the same number of valence electrons and similar electronconfigurations○ For the main group elements, the number of valence electrons is equal to thegroup number (with the exception of He)● Important note about d-block elements○ For transition metals s-orbitals fill before previous shells d-orbitals. Whentransition metals lose electrons, the s shell is lost first not the d shell.■ Example: 3d vs. 4s (n+l) à 3+2 = 5 vs. 4+1 = 4○ But…transitional metals generally lose the s-orbital electrons first!■ Example : Iron (Fe)○ Elements in 6B and 1B columns are exceptions to Aufbau’s principle. D orbital isstabilized rather than s orbital.■ Only one electron in s orbital and rest go to D orbital● Ex: Cr = [Ar] 4s1, 3d5● Types of Magnetic Materials○ Paramagnetic materials contain atoms, molecules, or ions with unpairedelectrons. It is attracted to a magnetic field.○ Diamagnetic materials have all electrons spin paired and the material does nothave a net magnetic field. Every electron is paired. It is slightly repelled to amagnetic field.○ Ferromagnetic: Has unpaired electrons that are aligned in a particular direction.It is strongly attracted to a magnetic field.● Isoelectronic Ions: Ions with the same number of electrons but different number ofprotons○ Example: Na+ and O2- ions● Periodic Trends: Atomic Radii○ Size of an atom is controlled by the size of its orbitals, but orbitals don’t have adefined boundary○ Chemist most often use covalent radius(bond distance) to determine the atomicsize○ Moving down a group the atomic radius increases with an increase in principlequantum number○ Moving across a period the atomic radius decreases since the Z* increases.○ Related to the metallic character of an element. The greater the atomic radius,the more metallic.● Effective Nuclear charge (Z*)○ Valence electrons experiences net charge of +1○ Z* for the highest-energy electrons (valence electrons) in an atom can becalculated by:■ Z* = Z (atomic number) – [number of core electrons]■ Z* increases moving left to right across the periodic table● Ionic Size○ Cations are smaller than the corresponding neutral atom○ Anions are larger than the corresponding neutral atom○ Anions are generally larger than cations○ Cation with the highest charge has the smallest radius○ Anion with the highest charge has the largest radius○ In an isoelectronic series, the ion with the most protons is smallest because thenucleus exerts a stronger force of attraction and the electrons are pulled closer tothe nucleus.○ The ion with the fewest protons is largest because the nucleus exerts a weakerforce of attraction and the electrons are held less tightly.● Ionization Energy (IE)○ It is the energy required to remove an electron from an atom in the gas phase○ It is always a positive value● Electron Affinity (EA)○ The energy change when a gaseous atom accepts an electron○ It can be a positive or a negative value○ Negative value indicates energy is


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