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Chem 105 Exam 3 Help

by: Allie Evey

Chem 105 Exam 3 Help Chem 105

Allie Evey

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This will help you prepare for the Exam on April 14th. Test review problems will be posted sometime on Saturday
Chem 105
Study Guide
Chem, 105, Finnegan, notes, Chemistry, exam
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This 9 page Study Guide was uploaded by Allie Evey on Thursday April 7, 2016. The Study Guide belongs to Chem 105 at Washington State University taught by Finnegan in Spring2015. Since its upload, it has received 70 views. For similar materials see Chem 105 in Chemistry at Washington State University.


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Date Created: 04/07/16
Things to Know for Exam 3 Ø Electron configurations of ions Ø Ground state electron configurations of monatomic ions • The large numbers represent the energy level. • The letters represent the sublevel. • The superscript numbers indicate the number of electrons in the sublevel. ü Electron Configuration for Sodium (Na): ???????? ???????????????? ???????? ???? ???? ???? The above is a diagram to easily help you remember which elements are in which “block”, for example sodium is in the third row S block so all the electron shells prior to the 3s block are completely filled The s shell can hold 2 electrons(2 columns in the s block) the o shell can hold 6 electrons(6 columns) the d shell can hold 10 electrons(10 columns) Remember: The general format for writing electron configurations is ns, np, n- 1d(starts at 3 d in the 4 row) The f shell isnt super important,but it holds 14 electrons and starts in row 5 • Main group ions - elements form monatomic ions that have a noble gas configuration (???????? ???????? ) ▯ • Transition metal ions -(the middle of the periodic table) don’t normally obtain a noble gas configuration Ø Orbital diagrams • Just a box representation of the electron shells Above is the orbital box diagram of sodium. Compare this to the above electron configuration Notice how for Oxygen the 2p orbital has an arrow in each box, and then it starts adding the down facing arrow. For any Orbital Diagram this is the case. You fill all the boxes in a specific set 1s, 2s, 2p etc. With an up facing arrow first, and then you fill it with a down facing arrow, or vice versa (the direction (up or down) isn’t important just that the first arrows face the same direction) Ø Paramagnetic vs. Diamagnetic • Paramagnetic-If there is an unpaired electron in the electron configuration the element/molecule is paramagnetic • Diamagnetic- When there are no un-paired electrons Ø Periodic trends (definitions, trend direction, explanation based on electron configurations, irregularities due to full and half-full subshells) • Atomic Radii - In general, the size of an atom will decrease as you move from left to the right of a certain period. • Ionic Radii :Neutral atoms tend to increase in size down a group and across a period. When a neutral atom gains or loses an electron, creating an anion or cation, the atom's radius increases or decreases. • Ionization energy - As you move Right and up n the periodic table the ionization energy (IE) gets larger • Electron affinity - Smaller Atoms release more energy when you add an electron because the electron gets sucked in close to the nucleus. 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 Ø Lewis electron dot symbols (for atoms and monatomic ions) • Bonds are shown as lines • Lone pair electrons are pairs of dots • All atoms (except H) want 8 electrons(octet rule) • An atom usually forms one bond for each electron • Central Atom Premise: the central atom is the least electronegative(except hydrogen, it only has one electron) Important Rules • Molecules with multiple Carbon and Nitrogen will usually have multiple centers • Carbon is the ONLY element the forms long chains (ie. Sugars) • Avoid O—O and F—F bonds, NEVER EVER make chains (2 or more) of Oxygen and Fluorine! • Fluorine does not form double bonds, Ever • DO NOT make triangles or squares with your bonds. • Pentagons and hexagons are okay. • Remember, rules can be broken, but when you do that bad things start happening. The more rules you break the worse you compound is. Ø Ionic bonds: attraction between 2 or more ions. Electron transfer is necessary to form ions Ø Lattice energy: the energy required to separate a mole of an ionic solid into gaseous ions Ø Born-Haber cycles: graphic representation of Hess’s Law (the total enthalpy change (Δ ▯ ) for a reaction is the sum of all Δ???? ) The above example is for Aluminum dioxide, notice how the different enthalpies are combined in each part of this equation Ø Covalent bonds : A bond that requires the sharing of electrons between atoms • shorter = stronger = stable = high bond energy • To Compare bond length, bond strength, bond stability and energy, think of the bonds as pieces of wood. A long piece of wood is much easier to break than a short piece of wood, you could break the short piece of wood but it would require a lot of energy Ø 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 Ø Electronegativity is shown using arrows above the bond. The arrow points towards the more electronegative atom Ø Resonance (IT’S IMPORTANT) • They show the possible electron arrangements of a given ion • They really show the extreme. An electron that makes a double bond is actually switching back and forth. That’s hard to show in a picture so we have resonance structures • Must show resonance structures that have the optimal formal charge • There are structures that have non-optimal formal charge (necessary for later Chem classes) • IF MORE THAN ONE RESSONANCE STRUCTURE EXIXTS, ALL EQUIVALENT (BY FORMAL CHARGE) RESONANCE STRUCTURES MUST BE SHOWN LINKED BY DOUBLED HEADED ARROWS Ø FORMAL CHARGE • Total the electrons around the atom(bonding electrons are split evenly) • Subtract this from an atoms normal charge(from the periodic table) • Formal Charge is used to decide between different possible structures • The best structures have a 0 formal charge Example: ???????? ▯ Current Charge-Charge found on the periodic table=Formal Charge 6-6=0 4-4=0 6-6=0 Ø Valence Shell Electron Pair Repulsion Theory (VSEPR) • VSEPR theory assumes that each atom in a molecule will achieve a geometry that minimizes the repulsion between electrons in the valence shell of that atom Ø AXE • ???????? ???? ???? • n=the number of bonding groups on the central atom • m=the number of lone pairs on the central atom • ANY type of bond (single, double, triple) is one bonding group • n+m=the number of electron groups on the central atom Ø Electron Geometry • The number of electron groups determines this • There are only 5 • Determines the ideal bond angles-elements try and arrange themselves as far away from each other as possible because the electrons repulse themselves Ø Molecular Geometry (shape)-electron geometry and the number of lone pairs determine the molecular geometry • The molecular shape, the bon polarities, and the formal charge distribution determine the molecular polarity Ø Bond Angles • Electron Geometry determines the ideal bond angles (they are approximate) Linear: 180 degree angles Triganol planar: 120 degree angles Tetrahedral: 109.5 degree angles Trigonal bipyramidal: 90,120 degree angles Octahedral: 90 degree angles • Lone pairs take up more space than bonding pairs • Double bonds take up more space than a single bonds • Important Note: When dealing with resonance structures, the amount of space the “rotating” double n=bond takes up is negligible because it is technically all double bonds ▯ ???? ???? • Electron Geometry: Triganol Bipyramidal • ???????????????????????????????????? ????ℎ????????????:Linear due to the 180 degree angle How do we do the bond angles for this molecule? By using the idea that double bonds take up more space than a single, and lone pairs take more space than bonds ???????????? ???? • Octahedral Electron Geometry: ???????? ???? ▯ ▯ • Molecular shape: square planar, due to the 90 degree angle • Bond angles for F-Xe-F is 90 degrees Ø Polar and Non-Polar Molecules • Something is polar if it has a negative side and a non polar side • In molecules this is determined by symmetry around the central atom • If a molecule is symmetric (by polar bonds)around the central atom, the charges cancel eachother out • If the molecule is not semetric(by polar bonds) than the molecule is polar • A numeric measure of polarity is the Dipole moment(μ) • The more polar a molecule the larger the dipole moment • Is ???? polar or non-polar? Non-polar ▯ • Is ???????????? ▯ polar or non polar? Non polar Ø Titrations • Acid base (Experiments #4 & 5) • Redox titrations (Experiment #10) • Equivalence point vs. end point # indicators To review the above just recall the labs you’ve done, and the lab backgrounds that went with them Will be added after it is covered in Lecture • ! Valence bond theory o " hybrid orbitals o " ó-bonds and ð-bonds - delocation of ð-bonds Titrations • " Acid base (Experiments #4 & 5) • " Redox titrations (Experiment #10) - equivalence point vs. end point # indicators


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