Study Guide for Exam 2 for Chemistry I | LB 171, Study notes of Education Planning And Management

Download Study Guide for Exam 2 for Chemistry I | LB 171 and more Study notes Education Planning And Management in PDF only on Docsity! Increase Atomic Radii Increase Increase Ionization Energy Study Guide Exam #2 Cass Periodic Trends: All group one elements form a +1 cation; because they want to have a noble gas configuration-it is the most stable and least reactive. Group two elements form a +2 cation. Group 7 form a -1 ion. Group 6 form a -2 ion. Group 5 form a -3 ion. Atomic Radii: the size of the atom. The periodic trend is as you move down the periodic table size increases, because there are more energy level shells. Moving from the right to the left on the periodic table radius size increases as well. An increased number of electrons also increase the number of protons. More protons means the nucleus filled with protons pulls the electrons in tighter, therefore making the atom smaller. The more negative an atom the larger the radius, because the proton’s center does no exert a high enough pull. The more positive an atom the smaller the radius, because there is more of a pull toward its center. Zeff: The pull mentioned above. The higher the zeff value the higher the pull towards the center, meaning the smaller the atom. z is the nuclear charge or number of protons. s is the number of NON-valence electrons Anions: negatively charged elements. Electron Affinity: (EA) means the gaining of an electron; it is the desire of an atom to absorb an electron. Electron affinity is fueled by the desire to become a noble gas, and fill its octet. EA is measured in the energy, usually negative values, because energy is typically released when an electron is absorbed. A decrease in distance between two molecules or atoms increases EA. An increase in a positive charge also increases EA. To predict EA use the electrostatic equation: d is distance. Ionization Energy: The energy required to remove an electron. An increase in atomic size decreases ionization energy, because it makes harder for zeff to pull. Ionization energy increases as you move to the right and up on the periodic table. (the elements in group 6 have a slightly lower ionization energy than group 5, because losing an electron allows for those elements to have a half filled p orbital, and is therefore Formulas - Structural: shows connectivity between atoms. - Ball and stick model: Different color ball represents elements, sticks represent bonds - Space-fill model: atoms overlap each other. Naming - Why? To illustrate how atoms are connected. Need a universal way to represent and refer to compounds. Ionic naming 1. Cations named 1st, leave as is (no prefix or suffix) 2. Anions named 2nd, using suffix IDE (except for polyatomic ions, which are named as is) ie. NaI Sodium Iodide; Magnesium Chloride MgCl2 - No prefixes for ionic compounds because we know the ion it will make, so we also know how much of each ion is required to form a neutral charge (except hydrates). - Prefix and Hydrates: a hydrate is an ionic compound in which water co-crystalizes with it. Use the suffix only to represent the amount of water present. o ie. SnCl2* 6H2O Tin (II) Chloride Hexahydrate 1: mono 5: penta 2: di 6: hexa 3: tri 7: hepta 4: tetra 8: octa Roman Numerals: used to identify the charge of a cation that can form more than one type of cation. Group1 metals: +1 only Group2 metals: +2 only Beneath metalloid line: different charges ie. PbCl4 = Lead(IV) Chloride Cobalt(II) Chloride Hexahydrate = CoCl*6H2O Ionic “bonds”: Ionic compounds are a metal and a non-metal. Not really a bond because the electron is being transferred NOT shared. Just an electrostatic attraction. Valence electrons tell us what kind of ion can be formed, it can also tell how many electrons it can (or wants to) share. Octet Rule: each atom has co-ownership of 8 electrons total. C, N, O, F never violates the octet rule. Exceptions: Period 3 elements and bellow, because they have an empty d orbital. Study Guide Exam #2 Cass Polyatomic Ions: See the polyatomic ion sheet from angel! Lewis Dot Structures: All atoms want to fill their octet, and electro negativity tells whether an electron is transferred or shared. Gives a 2D picture of what bonds to what, bond polarities, and reactivity of a molecule. 1. Count number of valence electrons, they should be illustrated in the drawing 2. Place the LEAST electronegative element in the center (except Hydrogen) 3. Connect the outer atoms using single bonds (hydrogens never attach to center of polyatomic ion) 4. Fill the octet of outer atoms by placing electrons around the atom. (hydrogen follows the duet rule, only 2e-) 5. Any “left over” electrons get placed around the center atom. 6. Determine formal charge* of each atom 7. If the center atom has a positive formal charge and an outer has a negative formal charge the outer atom will donate an electron pair to form a multiple bond with the center atom. *Formal Charge: FC= (total valence e- in atom) – (# of bonds + # lone e-) O never goes in the middle unless it is paired with F. Resonance structures: illustrate movement of electrons around the molecule/ion. They do NOT rearrange which atoms are attached to which. Increased resonance increases the stability of an ion. They take turns being in a multiple bond or having a negative charge. (4) (5) Bond Order: BO= Increase bond order decreases bond length. Increased bond order increases bond energy (the amount of energy required to eliminate a bond). If a molecule has no resonance forms then the bond order is just the bond between the two atoms (so a double bonded molecule with no resonance would have a bond order of 2). The bond order for NO3- above would be 4/3, meaning the 1/3 of the time it a particular N-O bond is a double bond, and the other2/3 of the time it is a single bond. Electronegativity: causes unequal sharing of electrons. Linus Pauling’s electro negativity chart shows an atoms desire to attract electrons. As you move up and to the right on the periodic table electro negativity increases. Fluorine is the most electronegative. Bond Polarity: The difference in electro negativity determines the polarity of bonds. C-H bonds are always non-polar. 0.0-0.4 Non polar covalent 0.4-2 Polar covalent >2 Ionic VSEPR: Valence Shell Electron Pair Repulsion theory. Electrons clusters* around the center atom will arrange themselves as far from each other as possible. Negative charges do not want to be close to each other. *electron clusters are bonds (double and triple only count as one electron group), lone electron pairs, single electron. VSEPR helps predict shape (electron geometry) based on the number of electron groups around the center atom. Electron Geometry: helps to understand a lot about the reactivity based on a 3D structure. # of e- groups Example (1) Electron geometry Bond angle 2 BeH2 Linear 180 3 BF3 Trigonal planar 120 4 CH4 Tetrahedral 109.5 5 PF5 Trigonal bipyramidal 120, 90, 90