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Chem 105 Notes March 9th-March 23

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by: Allie Evey

Chem 105 Notes March 9th-March 23 Chem 105

Marketplace > Washington State University > Chemistry > Chem 105 > Chem 105 Notes March 9th March 23
Allie Evey

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This notes cover material that will be on exam 3 which is April 14th
Chem 105
Class Notes
Chem, 105, Finnegan, notes, Chemistry
25 ?




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"Yes please! Looking forward to the next set!"
Haylie Rowe

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This 4 page Class Notes was uploaded by Allie Evey on Tuesday April 5, 2016. The Class Notes belongs to Chem 105 at Washington State University taught by Finnegan in Spring2015. Since its upload, it has received 42 views. For similar materials see Chem 105 in Chemistry at Washington State University.


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Yes please! Looking forward to the next set!

-Haylie Rowe


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Date Created: 04/05/16
3/9/16-3/23/16 Electron Configuration • Main Group elements form monatomic ions that have a noble gas configuration(???????? ???????? ) ▯ • So Oxygen which is normally 1???? 2???? 2???? becomes ???? and has the same electron configuration as Neon • Calcium will try to attain the noble gas configuration by going from ???????? → ???????? Calcium would then have the same electron configuration as Argon • Important Note: Calcium can go back to Argon because it is much easier for it to lose electrons, than it is for calcium to gain them. Bromine, however, wouldn’t lose electrons and go back to Argon, instead it would gain an electron and have the same electron configuration as Krypton • Transition Metals (the middle of the periodic table) don’t normally obtain a noble gas configuration • Transition metals will ‘steal’ electrons from s-shell to try and get a stable d- shell. • d-orbitals-are stable when they are half full, and when they are full, so when they have 5 or 10 electrons. They will steal electrons from the s-orbital to obtain 5 or 10 electrons • Cr normally has a half filled 3d orbital (because it stole one electron from the 4s orbital), in ????????▯ the d-orbital only has 3 electrons. 1 electron was removed from the 4s orbital (the only electron the 4s orbital had) and the other two were removed from the 3d orbital. • Electrons in the transition metals are taken from the s-orbital first, followed by the d-orbital. Like in the above example • A ???? − 1 ???? electron configuration is very stable, so ions rarely destroy this configuration • Paramagnetic-If there is an unpaired electron in the electron configuration the element/molecule is paramagnetic • Diamagnetic- When there are no un-paired electrons Ionization Energy • The energy required to remove an electron (1mol) from an atom in the gas state • Yes, it does cost energy when you remove an electron from an atom, atoms try and hold on to all their electrons • Periodic Trend!!: As you move Right and up n the periodic table the ionization energy(IE) gets larger • The smaller the atom(recall atomic radii) the higher the IE! This is because the smaller the atom the more pull the nucleus has on the electrons, so the more the atom holds on to the electron. • Important Note: There are exceptions to this, if an atoms subshell is half filled, or full, It is not going to want to give up the electron that is making it stable, so the IE is going to be higher than the atom next to it that doesn’t have half or full subshells • Oxygen has a low first IE because if you take one electron from Oxygen its p subshell becomes half full. Boron is another example of this • As you remove more and more electrons the amount of energy required to take that electron is going to be higher. ???????? ???? ???????? ???? • When there is a HUGE jump in IE energy it is because it takes A LOT more energy to go from removing valence electrons to removing core electrons!! • What the above means is that you can tell how many valence electrons an atom has just by looking at an atoms ionization energies • For example, if at ???????? there is a huge jump in the amount of energy it takes to ▯ remove an electron, that atom has 5 valence electron (like in Nitrogen) Electron Affinity • The change in energy when an electron is added to an atom • The first electron affinity (EA) tends to decrease, because Δ???? becomes more negative • Smaller Atoms release more energy when you add an electron because the electron gets sucked in close to the nucleus • Important Note: Noble gasses and the 2A periods are the exception to this rule. A few others exist due to the stability of half and full subshells • The second(and subsequent) EA cost energy because the electrons are being repulsed by the negative charge on the atom CHAPTER 9 • The Octet rule is VERY important: each atom needs 8 electrons • In Potassium Iodide the Iodine takes Potassium’s extra electron to form an octet(stable), Potassium then has no valence electrons and is stable • Ionic Bond: the electrostatic attraction between 2 or more ions • Electron transfer is necessary to form ions ???????????????????????????????? ????ℎ????????????????∙???????????????????????????????? ????ℎ???????????????? • Coulomb’s law: ???? = ???????????????????????????????????? ∙ ???????????????????????????????? • Born-Haber Cycle:Is a graphic representation of Hess’s Law(the total enthalpy change(Δ ???? ) f▯r a reaction is the sum of all Δ???? ) For Aluminum Dioxide Covalent bonds • A bond that requires the sharing of electrons between atoms • A covalent bond requires a balance of forces: attractive forces (electron(-)- nucleus) and repulsive forces (nucleus-nucleus, electron(-)-electron(-)). • The distance between two atoms is the distance where the opposing forces are equal (one isn’t stronger than the other) • This is a position of minimum energy (energy is released when a bond is formed). • Covalent bonds are localized. They exist between a specific pair of atoms. Know the relationship between bond lengths, bond strength (bond energy) and bond order. • Polar covalent bonds: bonds where the electrons aren’t equally shared. One Atom is more electronegative (wants electrons more ex. Fluorine) so it hogs the electrons • For the following bonds which atom is more electronegative: C-F N-H H-Si O-Cl Fluorine is more electronegative than Carbon, Nitrogen is more electronegative than Hydrogen, silicon is less electronegative than hydrogen, and oxygen is slightly more electronegative than Chlorine • Electronegativity is shown using arrows above the bond. The arrow points towards the more electronegative atom • Ionic character: Polar bonds can be thought of as a combination of a covalent and an ionic bond. The more polar the bond, the higher the percentage of ionic bonding to covalent bonding.


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