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W6: Chemical Principles notes (9/26, 9/28, 9/30)

by: Olivia Lange

W6: Chemical Principles notes (9/26, 9/28, 9/30) sch 100 01

Marketplace > Seton Hill University > Chemistry > sch 100 01 > W6 Chemical Principles notes 9 26 9 28 9 30
Olivia Lange
Seton Hill University

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week 6's notes
Chemical Principles
Professor Flowers
Class Notes
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This 7 page Class Notes was uploaded by Olivia Lange on Monday October 3, 2016. The Class Notes belongs to sch 100 01 at Seton Hill University taught by Professor Flowers in Fall 2016. Since its upload, it has received 4 views. For similar materials see Chemical Principles in Chemistry at Seton Hill University.

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Date Created: 10/03/16
9/26/16 Homework Online Conflicts: 3.47 1s^2 is the answer in format fo1s 2 2+¿ 3.42  X¿ See “Chemistry in Action” Table in textbook for reference 3.7 Ionic Bonds +¿ −¿ ● The arrangement of  Na ¿  and Cl¿  ions in a sodium chloride crystal ● The crystal is held together by ionic bonds. ● Ion­transfer reactions of metals and nonmetals form products unlike either  element. ● Because opposite electrical charges attract each other, the positive ion and  negative ion are said to be held together by an ionic bond. ● There are many ions attracted by ionic bonds to their nearest neighbors. These  crystals are ionic solids. ● Compounds of this type are referred to as ionic compounds. Remember: Covalent bonds share, ionics completely give up. 3.8 Formulas of Ionic Compounds ● Since all chemical compounds are neutral, it is easy to figure out the formulas of  ionic compounds. ● All chemical compounds are neutral. ● Once the ions are identified, decide how many ions of each type give a total  charge of zero. ● The chemical formula of an ionic compound tells the ratio of anions and cations. Remember: Cation is positive, Anion is negative. ● If the ions have the same charge, one of each is needed. +¿¿  +  −¿¿ = KF K F ● If the ions have different charges, unequal numbers of anions and cations must  combine to have a net charge of zero. +¿¿ −2 2 K   +  O  =  K 2 Ca+2 + 2  −¿¿  =   CaCl 2 Cl ● Table 3.3 on page 83 ● When the two ions have different charges, the number of one ion is equal to the  charge on the other ion. ● The formula of an ionic compound shows the lowest possible ratio of atoms and  is known as a simplest formula. ● There is no such thing as a single neutral particle of an ionic compound. ● Formula Unit: The formula that identifies the smallest neutral unit of an ionic  compound. +¿ −¿ ○ For NaCl, the formula unit is one Na ¿ ion and one  Cl ¿  ion ○ For CaF2, the formula unit is one  Ca+2  ion and two  −¿¿   F ions. ● Once the numbers and kinds of ions in a compound are known, the formula is  written using the following rules: ○ List the cation first and the anion second. ○ Do NOT write the charges of the ions. ○ Use parentheses around a polyatomic ion formula if it has a  subscript. Worked Example 3.6 ● Formula for compound formed by calcium ions and nitrate ions. ● 2+¿ and  NO −3 Ca¿ NO ❑ ● Ca(¿¿3) 2 ¿ 3.9 Naming Ionic Compounds ● Ionic compounds are named by citing first the cation and then the anion, with a  space between words. ● There are two kinds of ionic compounds: ○ Type I ionic compounds contain cations of main group elements. ○ Type II ionic compounds contain metals that can exhibit more than one charge. ● These require different naming conventions (­ous and ­ic). ● Type I ionic compounds contain cations of main group elements. ○ These charges on these cations do not vary. ○ Do not specify the charge on the cation. NaCl is sodium chloride. MgCO 3 is magnesium carbonate. ● Type II ionic compounds contain metals that can exhibit more than one charge. ○ Specify the charge on the cation in these compounds with either  the old (­ous, ­ic) or the new (roman numerals) system. FeCl 2  is iron (II) chloride or ferrous chloride. FeCl 3  is iron (III) chloride or ferric chloride. ● Do NOT name  FeCl 2 iron dichloride oreCl 3  iron trichloride. ● Once the charge on the metal is known, the number of anions needed to yield a  neutral compound is also known. ● Charges do not need to be included as part of the compound name. See chart on page 89. 9/28 Worked Example: Name Ionic Compounds KF = Potassium Fluoride Mg Cl2 = Magnesium Chloride Au Cl3 = Gold (III) Chloride Fe O = Iron (III) Oxide/ Ferric Oxide 2 3 Cr2 O 3 = Chromium (II) Oxide 3.11 H+ and OH­ Ions: An Intro. To Acids and Bases ● Two of the most important ions are ● Hydrogen cation (H+) ● Hydroxide anion (OH­) ● Hydrogen cation ● Simply a proton ● Whenever an acid dissolves in water, the proton attaches to a molecule of  water to form a hydronium ion +¿+¿ +¿ ¿ H 2 =  H3 ¿ H O ● Chemists use….  ● Hydroxide anion ● Polyatomic ion ● Oxygen atom is covalently bonded to a hydrogen atom ● The importance of the H+ cation and the OH­ anion is that they are  fundamental to the concepts of acids and bases ○ Acid: a substance the provides H+ ions in water ○ Base: a substance that provides OH­ anions in water ● See table 3.5  ● Different acids can provide different numbers of H+ ions per acid  molecule. +¿ ○ Hydrochloric acid, HCl, provides one H ¿  ion per acid  molecule. ○ Sulfuric acid,  H SO❑ , can provide two +¿¿  ions per  2❑❑ 4 H acid molecule. ○ Phosphoric acid,  H 3 PO 4 , can provide three +¿¿   ions H per molecule. ● Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are bases. ● When these compounds dissolve,  −¿ ¿ anions go into solution along  OH with the metal cation. ● Different bases can provide different numbers of  −¿¿  ions per formula  OH unit. −¿ ○ Sodium hydroxide provides one  OH ¿  ion per formula unit. ○ Barium hydroxide,  Ba(OH) 2  can provide two  −¿¿  ions  OH per formula unit. Chapter 4 4.1 Covalent Bonds (SHARE ELECTRONS) ● Covalent bond­ a bond formed by sharing electrons between atoms. ● Molecule­ a group of atoms held together by covalent bonds. ● Main group elements undergo reactions that leave them with 8 valence electrons  (or two for hydrogen), so that they have a noble gas electron configuration. ● Nonmetals can achieve an electron octet by sharing an appropriate number of  covalent bonds. ● Covalent bonding in hydrogen ( H 2 ): ○ Spherical 1s orbitals overlap to give an egg­shaped region. 2 ○ Two electrons between the nuclei, providing  1s  configuration  of helium. ● Covalent bonding in hydrogen ( H 2 ): ○ H­H or H:H or  H  all represent a hydrogen molecule. 2 ● Bond length is the optimum distance between nuclei in a covalent bond ○ Atoms too far apart ■ Attractive forces are small ■ No bond exists ○ Atoms too close ■ Repulsive interaction between nuclei is strong ■ Atoms are pushed apart ● Optimum point ○ Net attractive forces are maximized ○ Molecule is most stable ● H 2 molecule ○ Optimum distance between nuclei is 74 pm (picometers) ● Chlorine also exists as a diatomic molecule due to the overlap of 3p orbitals. ● See diagram on page 103 ● There are 7 diatomic elements: nitrogen, oxygen, hydrogen, fluorine, chlorine,  bromine, and iodine. 9/30/16 Remember Covalent shares and Ionic bonds give. 4.2 Covalent Bonds and the Periodic Table ● A molecular compound is a compound that consists of molecules rather than  ions. ● See visual on bottom of page 103 ● In these compounds, each atom shares enough electrons to achieve a noble gas configuration or filled octet. ● See visual chart “Figure 4.3” on page 104 ● Numbers of covalent bonds typically formed by main group elements to achieve  octet configurations. ● Exceptions to the Octet rule: ○ Boron has only 3 electrons to share, and forms compounds with  six electrons. ○ Elements in 3rd row and below in the periodic table have vacant  “d” orbitals that can be used for bonding. ● Remember: bonds occur in valance shell (outer) See Worked Example 4.1 & 4.3 ● Bonding in some molecules cannot be explained by sharing of only two electrons between atoms. ● These molecules’ atoms can only have outer­shell electron octets by sharing  more than two electrons. ● This results in the formation of multiple covalent bonds. ● See diagram at top of page 106. These are unstable because they still have gaps in their octets. Looking at the visual at the bottom of page 106, they will share more than  one electrons (a double covalent bond!). ● So octet = stable 4.3 Multiple Covalent Bonds ● Single bond ○ Sharing one electron pair ○ Represented by a single line: H­H ● Double bond ○ Sharing two electron pairs ○ Represented by a double line: O=O ● Triple bond ○ Sharing three electron pairs ○ Represented by a triple line N ≡ N ● Multiple covalent bonding ○ Particularly common in organic molecules ○ Consist predominantly of carbon ● Note that in compounds containing multiple bonds: ○ Carbon still forms four covalent bonds ○ Nitrogen still forms three covalent bonds ○ Oxygen still forms two covalent bonds See worked example 4.4 & 4.5 4.4 Coordinate Covalent Bonds ● A coordinate covalent bond is the covalent bond that forms when both  electrons are donated by the same atom. ● See colorful graphic on page 108 ● Once formed, no different from any other covalent bond ○ Result in unusual bonding patterns ○ Nitrogen with 4 covalent bonds ○ Oxygen with 3 bonds ( H +¿¿ ) 3 O ● See graphic at very bottom of page 108 Remember: metals form ionic compounds, non­metals form covalent


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