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Survey to Chemistry 1; Chapter 4: Chemical bonds

by: Jessica Taflinger

Survey to Chemistry 1; Chapter 4: Chemical bonds Ch 1043

Marketplace > Mississippi State University > Chemistry > Ch 1043 > Survey to Chemistry 1 Chapter 4 Chemical bonds
Jessica Taflinger

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These notes cover all of chapter 4 of Survey to Chemistry 1
Survey to Chemistry 1
Laura Smith
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This 5 page Class Notes was uploaded by Jessica Taflinger on Tuesday September 13, 2016. The Class Notes belongs to Ch 1043 at Mississippi State University taught by Laura Smith in Fall 2016. Since its upload, it has received 34 views. For similar materials see Survey to Chemistry 1 in Chemistry at Mississippi State University.

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Date Created: 09/13/16
Chapter 4: Chemical Bonds The Art of Deduction: Stable Electron Configurations ● Fact: Noble gas, such as helium, neon, and argon, are inert; they undergo  few, if any, chemical reactions. ● Theory: The inertness of noble gases results from their electron  configurations. Each (except helium) has an octet of electrons in its outermost  shell. ● Deduction: other elements that can alter their electron configurations to  become like those of noble gases would become less reactive by doing so. ● Valence electrons  ○ Recall that they are the electrons in the outermost shell ● Core electrons ○ Those in all the other shells are lumped together.  ● Isoelectronic ○ Having the same numbers of electrons or the same  electronic structure. Lewis (Electron­Dot) Symbols ● Representing valence electrons as dots when writing equations ○ Consider that a chemical symbol has four sides ○ Starting on any side, place single dots until each side has  one dot before adding another ○ There should be no more than two dots on any given side of  the chemical symbol. ● Crystal of sodium chloride ○ Na+ and Cl­ repeatedly arranging themselves in an orderly  fashion ● Ionic Bonds ○ The forces holding the crystal together ­ the attractive forces  between positive and negative ions ● Atoms and Ions ○ The names and symbols of an atom and its ion may look alot alike, especially the positively charged ions, called cations, but the actual  entities are very different ○ There is less room for confusion when it comes to naming  the negativvely charged ions, call anions ○ The names of anions are made to sound different through  the use of suffixes such as ­ ide in chloride. ● Octet rule ○ In case of helium, a maximum of two electrons can occupy  its single electron shell. ● Binary ionic compounds ○ Simple ions of opposite charge can be combined ○ To get the correct formula 1. Write each ion with its charge 2. Swap the charge number  3. Confirm that the transposed original atomic  ration 4. Write them as subscripts ● Covalent bond  ○ The bond formed when atoms share electron ● Single bond  ○ When one pair of electrons in the chloride molecule ● Lone pairs ○ The electrons that stay on one atom and are not shared ● Double bond ○ The covalent bond that forms from two pairs of electrons  being shared between two atoms ● Triple bond ○ A covalent linkage in which two atoms share three pairs of  electrons being shared between two atoms ● The number of bonds between two atoms is known as the bond  multiplicity. The more bonds between the two atoms, the higher the bond  multiplicity. ● Hydrogen and chlorine react to form a colorless gas called hydrogen  chloride ● Electronegativity ○ Can be viewed as its ability to attract electrons in a molecule ● Polar covalent bond ○ An unequal sharing of an electron pair leads to bond polarity ● Nonpolar covalent bonds  ○ The bonds where the electron pairs are shared equally ● HONC will enable you to write formulas for many molecules ○ Hydrogen forms 1 bond ○ Oxygen forms 2 bonds ○ Nitrogen forms 3 bonds ○ Carbon forms 4 bonds ● Polyatomic ions  ○ Charged particles containing two or more covalently bonded  atoms Rules for writing Lewis Formula ­ First choosing your skeletal structure (tells us the order in which the atoms attached to one another) ­ After choosing a skeletal structure you 1. Determine the total number of valence electrons. This total is the sum of the valence electrons for all the atoms in the molecule or ion.   You must also account for the charges on a polyatomic ion. a. For a polyatomic anion, add to its total number  of valence electrons the number of negative charges. b. For a polyatomic cation, subtract the number of positive charges from the total number of valence electrons. Examples:  N 2 4 Has [2 x 5] + [4 x 6] = 34 valence electrons −¿ ¿ Has [1 x 5] + [3 x 6] +1 = 24 valence electrons. N O 3 +¿ N H 4¿ Has [1 x 5] + [4 x 1] ­ 1 = 8 valence electrons. 2. Write a reasonable skeletal structure and connect bonded  pairs of atoms by a dash (one dash per each shared electron pair). 3. Starting with the most electronegative atom in the polyatomic molecule or ion, place electrons in pairs around outer atoms so that each  has octet, except for hydrogen, which will have a duet. 4. Subtract the number of electrons assigned so far ( both in  bonds and as lone pairs) form the total calculated in step 1.  Any electrons that remain are assigned in pairs to the central atoms. 5. If a central atom has fewer than eight electrons after step 4,  one or more multiple bonds are likely.  Move one or more lone pairs from  an outer atom to the space between the atoms to form a double or triple  bond.  A deficiency of two electrons suggests a double bond, and a  shortage of four electrons indicates either two double bonds or a triple  bond to the central atom. ● Valence­shell electron pair repulsion (VSEPR) theory ○ To predict the arrangement of atoms about a central atom in  a molecule ● Determining the shapes of many molecules 1. Draw a Lewis formula in which a shared electron pair  (bonding pair) is indicated by a line.  Use dots to show any unshared pairs (lone pairs) of electrons. 2. To determine shape, count the number of electron sets  around the central atom. Recall that a multiple bond counts only as one  electron set.  3. Using the number of electron sets determined in step 2,  draw a shape as if all the sets were bonding pairs and place these  electron sets as far apart as possible 4. If there are no lone pairs, the shape from step 3 is the shape of the molecule. If there are lone pairs, remove them, leaving the bonding  pairs exactly as they were (This may seem strange, but it stems from the  fact taht all the sets determine the geometry, even though only the  arrangement of bonded atoms is considered in the shape of the molecule.) 5. The presence of lone pairs of electrons on the central atom  decreases the angles between the bonds, as lone pairs require more  space than bonding pairs of electron. Likewise, double and triple bonds  command more space and push remaining bonds away. While somewhat  distorted, most molecules maintain their predicted arrangement of electron sets. The actual angles must be measured experimentally and you will not be asked to memorize or figure the out! ● Dipole  ○ A molecule with a positive end and a negative end  ● Polar molecule ○ Separate centers of positive and negative charge H ­ H and Cl ­ Cl H ­ Cl or H ­ Cl  Nonpolar                       Polar


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