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Study guide exam 2

by: Bridget House

Study guide exam 2 CHEM 120 001

Bridget House
GPA 3.48
General Chemistry
Deborah Weigand

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General Chemistry
Deborah Weigand
Study Guide
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This 6 page Study Guide was uploaded by Bridget House on Thursday October 29, 2015. The Study Guide belongs to CHEM 120 001 at University of Washington taught by Deborah Weigand in Fall 2015. Since its upload, it has received 44 views. For similar materials see General Chemistry in Chemistry at University of Washington.


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Date Created: 10/29/15
Study Guide Exam 2 Chapter 4 and 5 compounds are formed when atoms bind together Typesofbonds forms with electron transfer forms between two atoms of dissimilar elements usually a nonmetal and a metal consist of alternating positive and negative ions solid at room temperature has high melting point good conductors of electricity when the ionic solid if soluble in water forms an aqueous solution This is the formation of two ions A positive ion is called a and has fewer electrons An is a negative ion and has more electrons than protons Properties of an ionic bond a network of ions the attraction of positive charge to negative charge is very strong ionic bonds are stronger than covalent bonds ions come together to form a crystalline structure in solids 2 forms with electron sharing forms between like elements usually a nonmetal and a nonmetal consists of molecular compounds can be solid liquid or gas at room temperature low melting point does not conduct electricity bond results from two nuclei attracting the same shared electrons Two or more covalently bonded atoms form a molecule Can come in the form of bond in which two atoms share one pair of electrons bond in which two atoms share two pair of electrons These are twice as strong and twice as hard to break as a single covalent bond There must be at least two vacancies in an atoms valence electron shell prior to bonding in order for a double bond to be possible bond in which two atoms share three pair of electrons three times as strong and hard to break There must be at least three vacancies in an atoms valence electron shell for triple bonding to be ossible Thus rou 7A and Hydrogen cannot double or triple bond H a covalent bond in which both electrons in a given shared pair come from one of the two atoms involved in the bond Typically forms with metal ions pair of valence electrons that are shared between the atoms in a covalent bond pair of valence electrons in an atom that are not involved in sharing in a covalent bond c unit of a compound of covalently bonded atoms electrons in the outermost subshells of an atom These are only applied to representative and noble gas elements The valence electrons are only found in the s or p subshells Two important concepts 1 Not all electrons in an atom participate in the bonding with other atoms Valence electrons are the electrons in the atom that do participate in bonding 1 Certain arrangements of electrons in an atom make it more or less stable this concept is explained by the octet rule the lowest energy Does not readily or easily undergo change The noble gases are the most stable elements because they have no tendency to form bonds with other elements The electron configurations are full for noble gases besides Helium thus their s and p subshells are completely filled When compounds are formed the atoms of the element either lose gain or share electrons The goal of this process is to produce a noble as electron confi uration for each atom involved in the bond Some elements have less than an octet H has 2e B has 3e Be has 4e Some elements have more than an octet called an expanded octet P can have 10e S 12e Cl Br and I can have 14e Some elements can have an odd number of electrons some N compounds have 7 electrons Some compounds whose central atoms do not follow the octet rule include SF6 BeFZ C102 XeF4 a representation of the bonding patterns in a molecule 5 Count up all valence electrons for all atoms involved in the bond Write the chemical symbols for the atoms in the molecule in the correct order of bonding place a single bond two dots or a line in between each pair of atoms The central atom will usually be the atom that appears only once in the chemical formula Carbon will be central for most Hydrogen and Fluorine will never be the central atom Add the nonbonding electron pairs to the structure so that each atom that is bonded to the central atom has an octet Hydrogen is only 2 Place the remaining electrons on the central atom If there are not enough electrons to give the central atom on octet use on or more pairs of nonbonding electrons on the atoms bonded to the central atom thus forming double or triple bonds Count the total number of electrons and check number with step 1 a description of the 3D arrangement of atoms within a molecule a set of procedures used to predict the molecular geometry of a molecule using the Lewis structure and other information Central concept is that electron pairs in valence shells adopt an arrangement that minimizes the repulsion between the electron pairs This depends on the number of electron pairs Electron pair arrangements about the central atom 1 two electron pairs that are as far apart from one another as possible and are found at opposite sides of the nucleus at a 180 degree angle 2 three electron pairs found at opposite corners of an equilateral triangle at a 120 degree angle 0 If all three electron groups are bonding then the bond is still called Example HZCO o If one of the three bonds is nonbonding the bond is called Example 802 four sets of electron pairs makes a four sides solid with all sides of equal and identical equilateral triangles electron pairs separated by 109 degrees 0 If all electron groups are bonding it is a Exam le CH4 0 If one electron group is nonbonding it is a Example NH3 0 If two electron groups are nonbonding it is Example H20 a collection of valence electrons that are present in a localized region that is around the central atom in a molecule One group can consist of a single covalent bond 2 electrons a double covalent bond 4 electrons or a triple covalent bond 6 electrons Operational rules of VSEPR electron groups 1 Draw Lewis structure or molecule Identify the central atom which is the atom in which the geometrical information is desired 2 Determine the number of VSEPR electron groups present about the central atom this includes both bonding and nonbonding electron groups there is no distinction Also single double and triple bonds all count as one electron group 3 Predict the VSEPR electron group arrangements about the central atom by assuming the groups would orient themselves to minimize repulsion a measure of the relative attraction that an atom has for the shared electrons in a bond Electronegativity increases going up a group and across a period on the periodic table The higher the electronegativity the greater the attraction of that element for the shared electrons in a bond Non metals have higher electronegativity s tend to gain electrons in ionic bonds Metals have lower electronegativity s and tend to lose electrons in ionic bonds a measure of the degree of inequality in the sharing of electrons between two atoms in a bond a covalent bond in which there is equal sharing of electrons between two atoms two atoms of equal or very similar electronegativity s electronegativity difference between the two atoms of 04 or less covalent bond in which there is unequal sharing of electrons between two atoms this creates a fractional positive and negative on the atoms The electrons spend more time next to the more electronegative atom and less time near the less electronegative atom Electronegativity difference between the two atoms is greater than 04 but less than 15 Differences in electronegativity s of greater than 15 but less than 20 are either if there is a metal bonded to a nonmetal or if the bond is between two nonmetals measure of the degree of inequality in the attraction of bonding electrons to various locations in a molecule Depends on bond polarities and molecular geometry electronic charge molecule in which there is a nonsymmetrical distribution of molecule in which there is a symmetrical distribution of electronic charge when there are multiple correct ways to make a Lewis structure Example 03 the shorter the bond the stronger the bond for each atom take the number of valence electrons and subtract the number of electrons assigned in the molecule Guidelines for determining the formal charge 1 2 3 4 5 6 Both electrons in a lone pair belong to the atom in question Bonding electrons split evenly between bonded atoms The sum of the formal charges for all the atoms in the bond must equal the overall charge Carbon usually doesn t appear with a lone pair Oxygen usually doesn t appear with a positive formal charge When choosing which Lewis structure to chose if none of them have a formal charge of O chose the one with a formal charge closest to O the determination of the direction of bond dipole direction the arrow goes towards the atom with higher electronegativity delta goes to the atom with lower electronegativity delta goes to the atom with higher electronegativity The larger the electronegativity difference between two atoms the more ionic the bond a separation of charge representation Molecules with polar covalent bonds may also have an overall permanent dipole moment depending on the molecular shape A molecular dipole moment occurs when the electrons in a covalent bond are not shared equally by the atoms which is due to an electronegativity difference of the two atoms Example H20 If there is no net dipole moment dipole moment0 then the molecule is nonpolar compound in which two elements are present 1 2 3 4 Find the name and symbol of atoms Find the charge for each element using the periodic table Use the charge on each ion to determine what ratio is necessary to cancel the charges of each ion and make the overall charge zero making it neutral After this ste namin differs for ionic and covalent i an ionic compound in which one element is a metal and the other element is a nonmetal The metal is always the positive ion and the nonmetal is always the negative ion The presence of a nonmetal and metal in a compound is the key way to recognize that the compound is ionic Naming binary ionic compounds and creating formula no numerical prefixes are added to the name of binary ionic compounds The full name of the metal element is written first second is a separate word containing the stem of the nonmetal element name with the ending ide added on to the end Example sodium uoride two types of binary compounds 1 Main group elements the metal in the compound has a fixed charge This includes Group 1 and 2 and the following Agquot Znquot2 Cdquot2 Alquot3 Gaquot3 To name write the name of these compounds use the normal rule stated above 2 Transition group elements these metals are able to form more than one type of ion so they have more than one particle charge possible This includes all elements in the transition group When writing the names for compounds including a transition metal you must incorporate the charge on the metal ion into the name using roman numerals Then the nonmetal is written with the same rules above by adding an ide at the end a molecular compound in which two nonmetals are present use the same general rules above then add the prefix to both elements names to indicate how many atoms are present Mono one di two tri three tetrafour penta 5 hexa 6 Never use mono for the first element if it has only one Ions 1 2 an ion formed from a single atom through loss or gain of electrons Examples Cl Na Ca2 an ion formed from a group of atoms that are held together with covalent bonds through the loss or gain of electrons These are not molecules but instead charged pieces of compounds They require the presence of positive and negative ions and are neutral overall Example sulfate ion SO4quot2 Generalizations regarding the names and chemical formulas of polyatomic ions 1 Most not all ions carry negative charge that varies from 1 to 3 Exception to this includes H30 and NH4 Most not all include oxygen atoms Two ions do not CN and NH4 Most not all of the negative ions have names that end in ate Two ions that don t are OH named hydroxide and CN cyanide Two positive ions have names that end in ium Those include the two exceptions from 1 H30 hydronium and NH4 ammonium The number of pairs of ions exist where one member of the pair differs by having a Hydrogen atom present Example CO3quot2 carbonate differs from HCO3quot hydrogen carbonate In these pairs the charge on the ion containing Hydrogen is 1 less than the other ion 6 When there is more than one polyatomic ion present in a formula the polyatomic ion is put inside parentheses and the number of polyatomic ions is placed outside the parentheses Example FeOH3 To maintain the identity of the polyatomic ion the same elemental symbol may be used more than once in the chemical formula Example NH4N03 The rules for naming polyatomic ions is similar to the rules for naming binary ionic compounds differences include 1 If a polyatomic ion is present then the name of it is substituted for the metal name If a negative polyatomic ion is present then the name of it is substituted for that of the nonmetal stem and the ending is ide If both positive and negative polyatomic ions are present then both the name substitutions occurs and the name is just the names of both polyatomic ions When making Lewis structure for an ionic compound that contains polyatomic ions the positive and negative ions are treated separately to show they are individual and not linked in covalent bonds


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