Chapter 7 chemical formulas and bonding

Chapter 7chemical formulas and bonding CHAPTER 7

CHAPTER 7:CHEMICAL FORMULAS AND BONDINGIONIC BONDING • DESCRIBE THE DISTINGUISHING CHARACTERISTICS OF AN IONIC BOND • DESCRIBE SOME PROPERTIES OF IONIC COMPOUNDS • EXPLAIN THE OCTET RULE • DRAW LEWIS DOT DIAGRAMS TO SHOW THE VALENCE ELECTRONS OF AN ATOM • DISTINGUISH AMONG ANIONS, CATIONS, AND POLYATOMIC IONS CHAPTER 7

IONIC BONDING • IONIC BONDS AND IONIC COMPOUNDS • STATIC ELECTRICAL ATTRACTION IS THE BASIS FOR IONIC BONDS • IN AN IONIC BOND, A POSITIVELY CHARGED ION IS ATTRACTED TO A NEGATIVELY CHARGED ION. • A COMPOUND THAT IS COMPOSED ENTIRELY OF IONS IS CALLED AN IONIC COMPOUND. • ALL IONIC COMPOUNDS CONSIST OF POSITIVELY CHARGED IONS, CALLEDCATIONS, AND NEGATIVELY CHARGED IONS, CALLEDANIONS. • IONIC COMPOUNDS ARE ELECTRICALLY NEUTRAL, SO THAT THE ELECTRICAL CHARGES OF THE CATIONS AND THE ANIONS MUST BALANCE CHAPTER 7

IONIC BONDING • PROPERTIES OF IONIC COMPOUNDS: • HIGH MELTING POINTS • BRITTLE • DISSOLVE IN WATER (SOLUBLE) • SEPARATED IONS MOVE ABOUT FREELY IN WATER, WHICH MAKES THEM GOOD CONDUCTORS OF ELECTRICITY • MOLTEN IONIC COMPOUNDS ARE GOOD CONDUCTORS OF ELECTRICITY • DO NOT CONDUCT ELECTRICITY IN THE SOLID STATE • IONS ARE HELD IN POSITION AND CANNOT MOVE FREELY CHAPTER 7

IONIC BONDING • THE OCTET RULE • ATOMS TEND TO GAIN, LOSE, OR SHARE ELECTRONS IN ORDER TO ACQUIRE A FULL SET OF VALENCE ELECTRONS. • MOST ATOMS HAVE 8 VALENCE ELECTRONS, OR AN OCTET, IN A FULL SET. • THE EXCEPTIONS INCLUDE HYDROGEN AND HELIUM, THE FIRST PRINCIPAL ENERGY LEVEL IS FULL WITH ONLY 2 ELECTRONS CHAPTER 7

IONIC BONDING A SODIUM ATOM READILY LOSES ONE OF ITS ELECTRONS - - Na Na+ Na [Ne] 3s1 Na+ [Ne] CHAPTER 7

IONIC BONDING THE CHLORINE ATOM READILY GAINS AN ELECTRON Cl Cl- - - - Cl [Ne] 3s2, 3p5 Cl- [Ne] 3s2, 3p6 CHAPTER 7

IONIC BONDING • THE OCTET RULE • ATOMS TEND TO GAIN, LOSE, OR SHARE ELECTRONS IN ORDER TO ACQUIRE A FULL SET OF VALENCE ELECTRONS. • MOST ATOMS HAVE 8 VALENCE ELECTRONS, OR AN OCTET, IN A FULL SET. • THE EXCEPTIONS INCLUDE HYDROGEN AND HELIUM, THE FIRST PRINCIPAL ENERGY LEVEL IS FULL WITH ONLY 2 ELECTRONS CHAPTER 7

IONIC BONDING Ne Na+ Cl- Ar - IN BOTH CASES THE ION THAT FORMS HAS THE ELECTRON CONFIGURATION OF A NOBLE GAS CHAPTER 7

IONIC BONDING • LEWIS DOT DIAGRAMS • GILBERT LEWIS (1875-1946) • VALENCE ELECTRONS ARE REPRESENTED AS DOTS PLACED AROUND THE ELEMENT SYMBOL (SMALL x’s AND o’s CAN ALSO BE USED INSTEAD OF DOTS) • THE DOTS ARE PLACED ALONE OR IN PAIRS AROUND THE ELEMENT SYMBOL. • TRADITIONALLY, THE DOTS ARE WRITTEN ALONG THE SIDES OF AN IMAGINARY BOX, AND NO MORE THAN TWO DOTS ARE PLACED ON ANY SIDE OF OF THE BOX • LEWIS DOT MAY ALSO BE USED TO ILLUSTRATE HOW ELECTRONS ARE REARRANGED DURING CHEMICAL REACTIONS CHAPTER 7

IONIC BONDING ns1 and ns2 np1 np3 Ar np4 np6 np2 np5 CHAPTER 7

LEWIS DOT DIAGRAMS CHAPTER 7

IONIC BONDING Na Na Cl Na Cl Cl CHAPTER 7

IONIC BONDING • TYPES OF IONS • MONOATOMIC IONS (“ONE-ATOM” IONS) • FORMED FROM ONE (1) ATOM • MONOATOMIC CATIONS • METALS THAT HAVE ONLY ONE KIND OF CATION • THESE METALS HAVE ONLY ONE POSSIBLE OXIDATION NUMBER AND THEREFORE CAN ONLY BOND IN ONE WAY • BECAUSE OF THIS THE NAMING PROCEDURE FOR COMPOUNDS WITH MONOATOMIC CATIONS IS “STRAIGHT FORWARD”: • THE METAL NAME IS ALWAYS THE FIRST WORD IN THE COMPOUND • METALS THAT HAVE MORE THAN ONE KIND OF CATION • THESE METALS HAVE MORE THAN ONE OXIDATION NUMBER AND THEREFORE CAN BOND IN MORE THAN ONE WAY • BECAUSE OF THIS THE NAMING PROCEDURE MUST REFLECT THE SPECIFIC CATION BEING USED. • POLYATOMIC IONS (“MANY- ATOMS”) • FORMED FROM MORE THAN ONE ATOM • THESE ARE GENERALLY ANIONS ! (THE EXCEPTION IS AMMONIUM NH4 ) CHAPTER 7

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IONIC BONDING • MONOATOMIC ANIONS • TO NAME THE MONOATOMIC….REPLACE THE SUFFIX OF THE ELEMENT NAME WITH THE SUFFIX “-IDE” • CHLORINE ATOM CHLORIDE ION • OXYGEN ATOM OXIDE ION • SULFUR ATOM SULFIDE ION CHAPTER 7

CHAPTER 7

Some elements commonly form more than one kind of cation. These elements include the transition metals, which do not Follow the octet rule in forming cations. COMMON METALS WITH TWO PRIMARY OXIDATION NUMBERS (“-ous” or “-ic” ending on the metal or Roman Numerals after the metal + nonmetal with “-ide” ending) CHAPTER 7

COMMON ELEMENTS WITH SEVERAL OXIDATION NUMBERS (“MONO-, DI-, TRI-, TETRA-, PENTA-, HEXA-, SEPTA-, OCTO-, NONA-, DECA-”) (Prefix is used on the metal and nonmetal (to designate the subscript) or use Roman Numerals after the metal) CHAPTER 7

IONIC BONDING • POLYATOMIC IONS • IONS THAT CONSIST OF MORE THAN 1 ATOM • BONDED TOGETHER BY COVALENT BONDS…BUT ACT AS A UNIT. • POLYATOMIC IONS FORM AN IONIC BOND WITH AN ION OF THE OPPOSITE CHARGE • ALMOST ALL POLYATOMICS ARE NEGATIVE IONS (EXCEPT AMMONIUM[NH4]+ ) CHAPTER 7

CHAPTER 7

IONIC BONDING • BINARY IONIC COMPOUNDS • CONTAIN THE IONS OF ONLY TWO ELEMENTS. • TYPES OF CHEMICAL FORMULAS • GENERAL FORMULAS • EMPIRICAL FORMULAS • MOLECULAR FORMULAS • STRUCTURAL FORMULAS LISTED IN ORDER OF INCREASING SPECIFICITY CHAPTER 7

FORMULAS • GLUCOSE • GENERAL FORMULA CnH2nOn • BASIC MATHEMATICAL STRUCTURE • EMPIRICAL FORMULA C H2 O • RATIO (CARBOHYDRATE) • MOLECULAR FORMULA C6H12O6 • SIX-CARBON MONOSACCHARIDE • STRUCTURAL FORMULA • -D-GLUCOSE CH2OH H H C O C C H H H C C OH OH OH CHAPTER 7 OH

IONIC BONDING • What is an ionic bond? • List the properties of ionic compounds. • State the octet rule. Use the octet rule to describe the reaction between chlorine and sodium. • Describe the Lewis dot diagram to show valence electrons of an atom. • Distinguish among anions, cations, and polyatomic ions. CHAPTER 7

COVALENT BONDING • A COVALENT BOND IS FORMED BY A SHARED PAIR OF ELECTRONS BETWEEN TWO ATOMS. • REMEMBER!!! THE OCTET RULE • MOLECULES AND THEIR FAMILIES • A GROUP OF ATOMS THAT ARE UNITED BY COVALENT BONDS IS CALLED A MOLECULE. • A SUBSTANCE THAT IS MADE OF MOLECULES IS CALLED A MOLECULAR SUBSTANCE CHAPTER 7

COVALENT BONDING • DESCRIBING COVALENT BONDS • LEWIS STRUCTURE: NH3 SINGLE COVALENT BOND COVALENT BONDS UNSHARED PAIR OF ELECTRONS H N H N H H H H SINGLE COVALENT BONDS COVALENT BONDS COVALENT BONDS OCTET RULE: Nitrogen = 8 e- Hydrogen = 2e- CHAPTER 7

COVALENT BONDING • MULTIPLE BONDS • SINGLE COVALENT BOND (SINGLE BONDS) • IN A SINGLE BOND, 2 ATOMS SHARE EXACTLY ONE PAIR OF ELECTRONS • DOUBLE COVALENT BOND (DOUBLE BOND) • CONSISTS OF TWO PAIRS OF SHARED ELECTRONS • TRIPLE COVALENT BOND (TRIPLE BOND) • CONSISTS OF THREE PAIRS OF ELECTRONS CHAPTER 7

COVALENT BONDING O H C H • DOUBLE BONDS • FORMALDEHYDE (H2CO) • YOU CAN REPLACE THE PAIRED ELECTRON DOTS WITH A REPRESENTATIVE DASH ( ). • ( ) = single bond • ( ) = double bond • ( ) = triple bond SINGLE BOND DOUBLE BOND H H C C O O H H CHAPTER 7

COVALENT BONDING TRIPLE BOND SINGLE BOND • TRIPLE BONDS • ETHYNE (C2H2) C C H H H C H C H C H C SINGLE BOND TRIPLE BOND CHAPTER 7

DRAWING LEWIS STRUCTURES • DRAWING LEWIS STRUCTURES 1. SUM THE VALENCE ELECTRONS FROM ALL ATOMS A. Don’t worry about keeping track of which electrons come from which atoms. Only the total number is important. 2. WRITE THE SYMBOLS FOR THE ATOMS TO SHOW WHICH ATOMS ARE ATTACHED TO WHICH, AND CONNECT THEM WITH A SINGLE BOND ( ). A. Atoms are often written in the order in which they are connected in the molecule or ion, as in HCN. When the central atom has a group of other atoms bonded to it, we usually write the central atom first, as in CO32- and SF4. CHAPTER 7

DRAWING LEWIS STRUCTURES 3. COMPLETE THE OCTETS OF THE ATOMS BONDED TO THE CENTRAL ATOM. A. Remember, however, that hydrogen can only have only two electrons. 4. PLACE ANY LEFTOVER ELECTRONS ON THE CENTRAL ATOM, EVEN IF DOING SO RESULTS IN MORE THAN AN OCTET. 5. IF THERE ARE NOT ENOUGH ELECTRONS TO GIVE THE CENTRAL ATOM AN OCTET, TRY MULTIPLE BONDS. A. Use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds. CHAPTER 7

DRAWING LEWIS STRUCTURES DRAWING THE LEWIS STRUCTURE FOR PHOSPHORUS TRICHLORIDE PCl3 P3, 4+, 5+ [Ne] 3s2, 3p3 = 5 Cl1- FIRST, SUM THE VALENCE ELECTRONS [Ne] 3s2, 3p5 = 7 x(3) 26 CHAPTER 7

DRAWING LEWIS STRUCTURES Cl P Cl SECOND, ARRANGE THE ATOMS TO SHOW WHICH ATOM IS CONNECTED WHICH, AND DRAW A SINGLE BOND BETWEEN THEM. Cl THERE ARE VARIOUS WAYS THE ATOM MIGHT BE ARRANGED. IN BINARY COMPOUNDS, THE FIRST ELEMENT LISTED IS GENERALLY SURROUNDED BY THE REMAINING ATOMS CHAPTER 7

DRAWING LEWIS STRUCTURES Cl P Cl Cl THIRD, COMPLETE THE OCTETS ON THE ATOMS BONDED TO THE CENTRAL ATOM. THIS ACCOUNTS FOR 24 ELECTRONS REMEMBER THAT THE BONDED ELECTRONS (DASH) ACCOUNT FOR TWO ELECTRONS! CHAPTER 7

DRAWING LEWIS STRUCTURES Cl P Cl Cl FOURTH, PLACE THE REMAINING TWO ELECTRONS ON THE CENTRAL ATOM, COMPLETING THE OCTET AROUND THAT ATOM AS WELL. CHAPTER 7

DRAWING LEWIS STRUCTURES CO2 4 C = CARBON HAS 8 O2 = 12 XX O C O 16 XX XX XX THAT’S SWEET DUDE !! THAT’S 20 ELECTRONS !! NOW YOU HAVE 16 ELECTRONS AND I STILL FOLLOW THE OCTET RULE !! OXYGEN HAS 8 CARBON HAS 8 WAIT…I’LL FIX IT !!! CHAPTER 7

DRAWING LEWIS STRUCTURES DRAWING LEWIS STRUCTURES FOR OXYIONS (POLYATOMICS) BrO31- 1- Br = 7 O O Br 18 O3 = 25 O NOTE: THE TOTAL NUMBER OF VALENCE ELECTRONS (Br=7)+ (O3=18)=25. TO SATISFY THE OCTET RULE THERE ARE 26 ELECTRONS (ONE MORE THAN REQUIRED). THE COUMPOUND IS PLACED IN BRACKETS AND THE EXTRA NEGATIVE CHARGE(S) IS/ ARE RECORDED AS A SUPERSCRIPT OUTSIDE THE BRACKETS. CHAPTER 7

DRAWING LEWIS STRUCTURES • DRAW THE LEWIS STRUCTURE FOR THE FOLLOWING COMPOUNDS • CH4 • HCN • SO42- • CF2Cl2 CHAPTER 7

RESONANCE STRUCTURES • WE SOMETIMES ENCOUNTER SUBSTANCES IN WHICH THE KNOWN ARRANGEMENT OF ATOMS IS NOT ADEQUATELY DESCRIBED BY A SINGLE LEWIS STRUCTURE. • CONSIDER OZONE (O3), WHICH MUST HAVE ONE DOUBLE BOND TO ATTAIN AN OCTET OF ELECTRONS AROUND EACH ATOM: CHAPTER 7

RESONANCE STRUCTURES INDICATESRESONANCE FORMS O O O O O O O O O THE TWO ALTERNATIVE STRUCTURES FOR OZONE ARE EQUIVALENT EXCEPT FOR THE PLACEMENT OF THE ELECTRONS. THE REAL MOLECULE IS DESCRIBED BY AN AVERAGE (BLEND) OF THE STRUCTURES. THE MOLECULE DOES NOT OSCILLATE RAPIDLY BETWEEN THE FORMS (THERE IS ONLY ONE FORM OF THE MOLECULE). CHAPTER 7

EXCEPTIONS TO THE OCTET RULE • ATOMS WITH LESS THAN AN OCTET • MANY COMPOUNDS OF BORON DO NOT FOLLOW THE OCTET RULE (BF3) • WE COULD COMPLETE THE OCTET BY FORMING A DOUBLE BOND CREATING POSSIBLE RESONANCE STRUCTURES. • THESE STRUCTURES FORCE FLUORINE TO SHARE ADDITIONAL ELCTRONS WITH THE BORON ATOM…HOWEVER….THIS IS INCONSTENT WITH THE HIGH ELECTRONEGATIVITY OF FLUORINE. CHAPTER 7

EXCEPTIONS TO THE OCTET RULE • ATOMS WITH MORE THAN AN OCTET • SOME ATOMS FOUND BEYOND THE 2ND PERIOD OF THE PERIODIC TABLE (MOST NOTABLY SULFUR AND PHOSPHORUS) SOMETIMES FORM BONDS THAT GIVE THEM MORE THAN AN OCTET OF ELECTRONS. • THE ADDITIONAL ELECTRONS FILL THE 3d ORBITALS OF THESE ATOMS • SF4 , AsF61-, PCl5, AND ICl41- CHAPTER 7

EXCEPTIONS TO THE OCTET RULE • MOLECULES WITH AN ODD NUMBER OF ELECTRONS • IN THE VAST MAJORITY OF MOLECULES THE NUMBER OF ELECTRONS IS EVEN, AND COMPLETE PAIRING OF ELECTRON SPINS OCCURS. • A MOLECULE WITH AN ODD NUMBER OF ELECTRONS CANNOT FOLLOW THE OCTET RULE. • ClO2, , NO , AND NO2 CHAPTER 7

BOND POLARITY AND ELECTRONEGATIVITY • THE CONCEPT OF BOND POLARITY IS USEFUL IN DESCRIBING THE SHARING OF ELECTRONS BETWEEN ATOMS • A NONPOLAR BOND IS ONE IN WHICH THE ELECTRONS ARE SHARED EQUALLY BETWEEN TWO ATOMS. • A POLAR COVALENT BOND, ONE OF THE ATOMS EXERTS A GREATER ATTRACTION FOR THE ELECTRONS THAN THE OTHER. • IF THE DIFFERENCE IN RELATIVE ABILITY TO ATTRACT ELECTRONS IS LARGE ENOUGH, AN IONIC BOND IS FORMED. CHAPTER 7

BOND POLARITY AND ELECTRONEGATIVITY • ELECTRONEGATIVITY: • THE ABILITY OF AN ATOM, IN A MOLECULE, TO ATTRACT ELECTRONS TO ITSELF. • THE GREATER THE ELCETRONEGATIVITY, THE GREATER ITS ABILITY TO ATTRACT ELECTRONS • RELATED TO ITS IONIZATION ENERGY AND ELECTRON AFFINITY • FLUORINE HAS THE HIGHEST ELECTRONEGATIVITY (4.0) • CESIUM HAS THE LOWEST ELECTRONEGATIVITY (0.7) CHAPTER 7

BOND POLARITY AND ELECTRONEGATIVITY GREY AREA BETWEEN 1.67 AND 2.0 0.0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 BETWEEN 0.5 AND 1.9 POLAR COVALENT > 2.0 IONIC < 0.4 NONPOLAR COVALENT < 0.4 NONPOLAR CHAPTER 7

BOND POLARITY ANDELECTRONEGATIVITY HF F2 LiF COMPOUND ELECTRONEGATIVITY DIFFERENCE 4.0 - 1.0 = 3.0 4.0 - 4.0 = 0 4.0 - 2.1 = 1.9 IONIC POLAR COVALENT TYPE OF BOND NONPOLAR COVALENT CHAPTER 7

BOND POLARITY ANDELECTRONEGATIVITY • IN F2 THE ELECTRONS ARE SHARED EQUALLY BETWEEN THE FLUORINE ATOMS, AND THE BOND IS NONPOLAR. F F CHAPTER 7

BOND POLARITY ANDELECTRONEGATIVITY • IN HF, THE FLUORINE ATOM HAS A GREATER ELECTRONEGATIVITY THAN THE HYDROGEN ATOM. • THE SHARING OF ELECTRONS IS UNEQUAL (THE BOND IS POLAR) • THIS IS REPRESENTED IN TWO WAYS +  - F F H H CHAPTER 7

BOND POLARITY ANDELECTRONEGATIVITY • THE + AND -ARE MEANT TO REPRESENT PARTIAL POSITIVE AND NEGATIVE CHARGES RESPECTIVELY. • THE ARROW REPRESENTS THE PULL OF ELECTRON DENSITY OFF THE HYDROGEN BY THE FLUORINE, LEAVING THE HYDROGEN WITH A PARTIAL POSITIVE CHARGE • THE HEAD OF THE ARROW POINTS IN THEDIRECTION IN WHICH THE ELECTRONS ARE ATTRACTED CHAPTER 7

Chapter 7 chemical formulas and bonding