ChemMidterm#2studyguide.pdf CHEM 120
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Date Created: 11/02/15
Chem Midterm #2 Study Guide **Knowledge for both covalent and ionic bonding** Chemical bond: o Def: the attractive force that holds to atoms together in order to create a more complex unit. They are formed based on the interactions between the atoms electrons Lewis dot Structures o Def: the chemical symbol surrounded by dots that equal the number of valence electrons Valence electrons: the outermost full electron shell This number can be found form the periodic table set up Periodic table groups have the same number of valence electrons (because they have similar chemical propertied) Group number signifies the number of valence electrons it has (ex: group VIA has 6 valence electrons) 8 is the most electrons a shell can have o The Idea of Stability The idea that the electrons surrounding an atom will not undergo spontaneous change All noble gasses are considered stable because they have a full outer electron shell, meaning they do not need to transfer or share electrons to become more stable Octet rule So noble gases have 8 electrons, other elements want to have 8 electrons on their otter most shell that is why they bond and transfer electrons with other atoms (also known as trying to produce a noble gas configuration) This makes the compound more stable and less likely to change The exceptions…. o Sub octet Beryllium (Be) and Boron (B): Only needs three bonds o Expanded octet Needs to have at least a level 3 valence shell Phosphorus (P): has five bonds o Radicals Nitrogen (N): can have 7 electrons **Chapter 4: The Ionic Bond Model** Ionic bond Def: formed through the transfer of electrons, this will then form an ionic compound o Characteristics: High melting point Good conductors of electricity when in a molten state or a solution Bonds between a metal and a non metal Ion: an atom or group of atoms that is electrically charged as a result of loosing or gaining an electron o Gaining an electron = more negativity of the ion leading to a (-) overall charge o Loosing an electron = more positivity of the ion leading to a (+) overall charge Periodic table trends of loosing or gaining electrons in a transfer o Metals in group 1A, 2A, 3A tend to loose electrons Group 1A forms a net +1 charge Group 2A forms a net +2 charge Group 3A forms a net +3 charge o Nonmetals in groups 5A, 6A, 7A tend to gain electrons Group 5A form a net -3 charge Group 6A form a net -2 charge Group 7A form a net -1 charge o Fixed charges Ag (+) Zn, Cd, (2+) Al, Ga, (3+) Isoelectronic: an atom or group of atoms that have the same number and configuration of electrons o This is normally associated with compounds creating a noble gas configuration Chemical formulas for ionic compounds o The resulting compound always needs to be neutral o One needs to create the correct combination of parts Ex: K+ and S2- 2 (K+) = 2+, S2- = 2- Net charge of 0 o The symbol of the positivity charged ion always goes first o The charges are not written in the formula, they are inferred and should be known Chemical Naming o Binary compounds Contains one metal and one nonmetal Metal (+) and non metal (-) 1. Metal is named first in the full name 2. Nonmetallic stem with –ide ending EX: NaF Sodium fluoride If naming metals with variable charges you need to indicate the charge on the metal with roman numerals o Polyatomic compounds An ion is forced by a group of atoms They are normally very stable and they never occur alone and are always associated with ions of opposite charges When writing the chemical formula put parentheses around the polyatomic structure to indicate that they act as a unit Nitrogen Nitrate: NO3 (-) Ammonium: NH4 (+) Sulfur Sulfate: SO4 (2-) Hydrogen sulfate: HSO4 (-) Phosphorus Phosphate: PO4 (3-) Hydrogen phosphate: HPO4 (2-) Dihydrogen phosphate: H2PO4 (-) Carbon Carbonate: CO3 (2-) Hydrogen carbonate: HCO3 (-) Hydrogen Hydronium: H3O (+) Hydroxide OH (-) Cyanide: CN (-) **Chapter 5: The Covalent Bond Model** Covalent bond Def: formed through the sharing of electrons, this will then form a molecular compound. o Characteristics Low melting points Tend to be in a gas or liquid form A bond between similar atoms, usually non metals Lewis dot structures o The shared electrons are drawn as lines to indicate that they are bonded atoms One can have a single, double, or triple bond Single: one line that is equivalent to 2 electrons Double: 2 lines that is equivalent to 4 electrons Triple: 3 lines that is equivalent to 6 electrons o Steps for drawing Calculate the total number of Valence electrons present (use periodic table trends) Write out the chemical symbols next to each other with one bonded pair between each Add non bonded electrons around the atoms Check for the octet rule and shift some electrons into bonded pairs or lone pairs to achieve the correct number of valence electrons that you found earlier o Resonance The idea that the bonds between atoms can be drawn in different ways A double bond or triple bond can be shared between different atoms and still create an accurate LDS When this happens, the stability of the atom can be compromised, one drawing of the structure will be more stable than other resonance structures To find the stability you find the formal charge for each atom and the total formal charge closest to o is the most common form of the compound o Formal charge = (#valence electrons) – (number of lone electrons) – (1/2 the bonded electrons) Bond strength The average bond strength when there is a combination of single, double, or triple bonds a single bond = 1 a double bond = 2 a triple bond =3 add up all the bonds and divide them by the number of electron groups = the average bond strength o Lewis dot structures for polyatomic ions Has both ionic and covalent properties Covalent properties between the atoms in the ion Ionic in how it has a overall charge and will bond with ions as a whole unit Molecular geometry o Def: the 3D arrangement of atoms within a molecule o This helps define the physical and chemical properties of the substance o VSEPR (valence electron pair repulsion) Def: used the information present in a LDS structure to determine the molecular geometry Count electrons groups around the central atom A bonded pair of electrons whether it is a single, double, or triple bond count as 1 group A lone pair is also considered one 1 group 2 VSEPR electron groups Always called linear The angle between electron pair is 180 degrees 3 VSEPR electron groups No lone pair as around the central atom = trigonal planar 2 bonded groups and 1 lone pair = bent/angular The angle between electron pairs is 120 degrees 4 VSEPR electron groups 4 bonded electron groups = tetrahedral 3 bonded electron groups and 1 lone pair = trigonal pyramidal 2 bonded electron groups and 2 lone pairs = bent/angular The angle between electron pairs is 109 degrees o Molecules with more than one central atom You will have multiple VSEPR shapes Treat each central atom separately to find the molecular geometry around that atom This will then create bends in the molecule (this is getting into organic chem and bit and I don’t think multiple VSEPR questions will be asked extensively on the test) Coordinate covalent bonds o Def: when a covalent bond is formed but the bonded pair of electrons both come from one atom and non come from the other atom Once the bond is formed there is no way to distinguish that it is coordinate covalent by just looking at it. This concept is used to help explain bonding of polyatomic ions and some molecules because there electron arrangement would otherwise create problems with their electron arrangement Electronegativity o Def: the ability of different nuclei to have a relative attraction for shared electrons in a bond o More electronegative as you go across the periodic table from left to right More electronegative as you go up the table Bond polarity o Nonpolar = equal sharing between bonds o Polar = unequal sharing between bonds This create a net charge on the compound when originally it was a neutral compound o Fractional charge Used to determine if a bond is polar or not By using the trends of electron negativity from the periodic table, one can find the relative attraction between the atoms bonded The more electronegative atom will be denoted with a (-) delta The less electronegative atom will be denoted with a (+) delta The atom will more electronegativity will be pulling the bonded electrons more, giving the bond a charge in the direction of the more electron negative atom Naming molecular compounds o Binary compounds The naming of 2 non metals bonded together The full name of the less electronegative atom goes first Followed be the root of the other atom with –ide as the ending Numerical prefixes are used for both non metals 1 = mono 2 = di 3 = tri 4 = tetra 5= penta 6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca **Drawing the connection between ionic and covalent bonds** Bond classification o Most bonds have both ionic and covalent properties o Non polar covalent = electronegativity of .4 or less Very equal sharing of electron o Polar covalent = electronegativity between .4 and 1.5 Unequal sharing of electrons o Ionic or polar covalent= electronegativity between 1.5 and 2.0 Ionic if a metal is present Polar covalent when only non metals are present o Ionic= electronegativity of 2.0 or higher Electron transferring Both of these concepts of ionic and covalent bonds are “convenience concepts” meaning they were created by humans to make it easier for us to understand the interactions between atoms. A bond is not always 100% covalent or 100% ionic, that is why we create these models with the idea of purity in mind, but in real life nature is not 100% perfect all of the time. That is why there are so many exception and rules that are not always followed.
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