Presentation on theme: "Chapter 5 Compounds and Their Bonds"— Presentation transcript:
1 Chapter 5 Compounds and Their Bonds 5.1 Octet Rule and Ions
2 Octet Rule An octet is 8 valence electrons is associated with the stability of the noble gasesdoes not occur with He; He is stable with 2 valence electrons (duet)Valence ElectronsHe 1sNe 1s22s22pAr 1s22s22p63s23pKr 1s22s22p63s23p64s23d104p
3 Ionic and Covalent Bonds Atoms form octetsto become more stableby losing, gaining, or sharing valence electronsby forming ionic or covalent bonds
4 Metals Form Positive Ions by a loss of their valence electronswith the electron configuration of the nearest noble gasthat have fewer electrons than protonsGroup 1A(1) metals ion 1+Group 2A(2) metals ion 2+Group 3A(3) metals ion 3+
5 Formation of a Sodium Ion, Na+ Sodium achieves an octet by losing its one valence electron.With the loss of its valenceelectron, a sodium ion has a 1+charge. Sodium atom Sodium ion11p p+11e– e–
6 Formation of Magnesium Ion, Mg2+ Magnesium achieves an octet by losing its two valence electrons.With the loss of two valenceelectrons magnesium forms apositive ion with a 2+ charge Mg atom Mg2+ ion12p+ 12p+12e– e–
7 Formation of Negative Ions In ionic compounds, nonmetalsachieve an octet arrangementgain electronsform negatively charged ions with 3–, 2–, or 1– charges
8 Formation of a Chloride Ion, Cl– Chlorine achieves an octet by adding an electron toits valence electrons.
9 Ionic Charge from Group Numbers The charge of a positive ion is equal to its Group number.Group 1A(1) = 1+Group 2A(2) = 2+Group 3A(3) = 3+The charge of a negative ion is obtained by subtracting 8 or 18 from its Group number .Group 6A(16) = – 8 = 2–or 16 – 18 = 2–
10 Group Number and Ionic Charge Ionsachieve the electron configuration of theirnearest noble gasof metals in Groups 1A(1), 2A(2), or 3A(13) have positive 1+, 2+, or 3+ charge.Of nonmetals in Groups 5A(15), 6A(16), or 7A(17) have negative 3–, 2–, or 1– charge.
11 Groups Numbers for Some Positive and Negative Ions
12 Chapter 5 Compounds and Their Bonds 5.2 Ionic Compounds
13 Ionic Compounds Ionic compounds consist of positive and negative ions have attractions called ionic bonds between positively and negatively charged ionshave high melting and boiling pointsare solid at room temperature
14 Salt Is an Ionic Compound Sodium chloride, or “table salt,” is an example of an ionic compound.
15 Ionic Formulas An ionic formula consists of positively and negatively charged ionsis neutralhas charge balancetotal positive charge = total negative chargeThe symbol of the metal is written first, followed by the symbol of the nonmetal.
16 Charge Balance for NaCl, “Salt” In NaCl,a Na atom loses its valence electrona Cl atom gains an electronthe symbol of the metal is written first, followed by the symbol of the nonmetal.
17 Charge Balance In MgCl2 In MgCl2, a Mg atom loses two valence electronstwo Cl atoms each gain one electronsubscripts indicate the number of ions needed to give charge balance
18 Writing Ionic Formulas from Charges Charge balance is used to write the formula forsodium nitride, a compound containing Na+ and N3−.Na+3 Na N3− = Na3N3(1+) (3–) =
19 Chapter 5 Compounds and Their Bonds 5.3Naming and Writing Ionic Formulas
20 Naming Ionic Compounds with Two Elements To name a compoundwith two elements,identify the cation and anionname the cation first, followed by the name of the anion
21 Examples of Ionic Compounds with Two Elements Formula Ions NameCation AnionNaCl Na+ Cl– sodium chlorideK2S K S2– potassium sulfideMgO Mg2+ O2– magnesium oxideCaI Ca2+ I– calcium iodideAl2O3 Al S2– aluminum sulfide
22 Transition Metals Form Positive Ions Most transition metals and Group 4(14) metals,Form 2 or more positive ionsZn2+, Ag+, and Cd2+ form only one ion.
23 Metals with Variable Charge The names of transition metals with two or more positive ions (cations) use a Roman numeral after the name of the metal to identify ionic charge.
24 Naming Ionic Compounds with Variable Charge Metals
25 Naming FeCl2 STEP 1 Determine the charge of the cation from the anion. Fe ion Cl– = Fe ion – = 0Fe ion = 2+ = Fe2+STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge.Fe2+ = iron(II)STEP 3 Write the anion with an ide ending. chlorideSTEP 4 Name the cation first, then the anion.iron(II) chloride
26 Naming Cr2O3STEP 1 Determine the charge of cation from the anion Cr ions + 3O2– = 2Cr ions + 3(2–)= 2Cr ions + 6– = 02Cr ions = 6+ Cr ion = 3+ = Cr3+STEP 2 Name the cation by the element name, and use a Roman numeral to show its charge.Cr3+ = chromium(III)STEP 3 Write the anion with an ide ending. oxideSTEP 4 Name the cation first, then the anion.chromium (III) oxide
28 Writing Formulas Write a formula for potassium sulfide. STEP 1 Identify the cation and anion.potassium = K+sulfide = S2−STEP 2 Balance the charges.K S2−K+2(1+) + 1(2–) = 0STEP 3 Write the cation first.2K+ and 1S2− = K2S1 = K2S
29 Writing Formulas STEP 1 Identify the cation and anion. Write a formula for iron(III) chloride.STEP 1 Identify the cation and anion.iron (III) = Fe3+ (III = charge of 3+)chloride = Cl−STEP 2 Balance the charges.Fe Cl−Cl−1(3+) + 3(1–) = 0STEP 3 Write the cation first.1Fe3+ and 3Cl− = FeCl3
30 Chapter 5 Compounds and Their Bonds 5.4 Polyatomic Ions
31 Polyatomic Ions A polyatomic ion is a group of atoms has an overall ionic chargeExamples:NH4+ ammonium OH− hydroxideNO3− nitrate NO2− nitriteCO32− carbonate PO43− phosphateHCO3− hydrogen carbonate(bicarbonate)
32 Some Names of Polyatomic Ions The names of common polyatomic anionsend in ateNO3− nitrate PO43− phosphatewith one oxygen less end in iteNO2− nitrite PO33− phosphitewith hydrogen attached use prefix hydrogen (or bi)HCO3− hydrogen carbonate (bicarbonate)HSO3− hydrogen sulfite (bisulfite)
34 Examples of Names of Compounds with Polyatomic Ions The positive ion is named first followed by the name of the polyatomic ion.NaNO3 sodium nitrateK2SO4 potassium sulfateFe(HCO3)3 iron(III) bicarbonateor iron(III) hydrogen carbonate(NH4)3PO3 ammonium phosphite
35 Writing Formulas with Polyatomic Ions The formula of an ionic compoundcontaining a polyatomic ion must have a charge balance that equals zero (0)Na+ and NO3− NaNO3with two or more polyatomic ions put the polyatomic ions in parentheses.Mg2+ and 2NO3− Mg(NO3)2subscript 2 for charge balance
36 Chapter 5 Compounds and Their Bonds 5.5 Covalent Compounds
37 Forming Octets in Molecules In a fluorine (F2) molecule, each F atomshares one electronacquires an octet
38 Diatomic ElementsThese elements share electrons to form diatomic, covalent molecules.
40 Single and Multiple Bonds In a single bond, one pair of electrons is shared.In a double bond, two pairs of electrons are shared.In a triple bond, three pairs of electrons are shared.
41 Electron-Dot Formula of CS2 Write the electron-dot formula for CS2.STEP 1 Determine the atom arrangement. The C atom is the central atom.S C SSTEP 2 Determine the total number of valence electrons for 1C and 2S.1 C(4e–) S(6e–) = 16e–STEP 3 Attach each S atom to the central C atom using one electron pair.S : C : S16e– – 4e– = 12e– remainingSTEP 4 Attach 12 electrons as 6 lone pairs.: S : C : S :
42 Electron-Dot Formula of CS2 (continued) To complete octets, form one or more multiple bonds.Convert two lone pairs to bonding pairs between C andS atoms to make two double bonds.
43 A Nitrogen Molecule has A Triple Bond In a nitrogen molecule, N2,each N atom shares 3 electronseach N atom attains an octetthe sharing of 3 sets of electrons is a multiple bond called a triple bond
44 Resonance Structures Resonance structures are two or more electron-dot formulas for the same arrangement of atomsrelated by a double-headed arrow ( )written by changing the location of a double bond between the central atom and a different attached atom
45 Writing Resonance Structures Sulfur dioxide has two resonance structures.STEP 1 Write the arrangement of atoms.O S OSTEP 2 Determine the total number of valence electrons.1 S(6e−) + 2 O(6e−) = 18e−STEP 3 Connect bonded atoms by single electron pairs.O : S : O e− used18e− – 4e− = 14e− remaining
46 Writing Resonance Structures (continued) STEP 4 Add 14 remaining electrons as 7 lone pairs.STEP 5 Form a double bond to complete octets. Two resonance structures are possible.
47 Formula from Ionic Charges Write the ionic formula of the compound containing Ba2+ and Cl.Write the symbols of the ions.Ba2+ ClBalance the charges.Ba2+ Cl two Cl needed ClWrite the ionic formula using a subscript 2 for two chloride ions that give charge balance.BaCl2
48 Chapter 5 Compounds and Their Bonds 5.6Naming and Writing Covalent FormulasNO nitrogen oxideNO2 nitrogen dioxideN2O4 dinitrogen tetroxide
49 Names of Covalent Compounds Prefixes are usedin the names of covalent compoundsbecause two nonmetals can form two or more different compoundsExamples of compounds of N and O:NO nitrogen oxideNO2 nitrogen dioxideN2O dinitrogen oxideN2O4 dinitrogen tetroxideN2O5 dinitrogen pentoxide
50 Naming Covalent Compounds STEP 1 Name the first nonmetal by its element name.STEP 2 Name the second nonmetal with an ide ending.STEP 3 prefixes to indicate the number (from subscripts) of atoms of each nonmetal. Mono is usually omitted.
51 Naming Covalent Compounds (Continued) What is the name of SO3?STEP 1 The first nonmetal is S sulfur.STEP 2 The second nonmetal is O, named oxide.STEP 3 The subscript 3 of O is shown as the prefix tri.SO3 → sulfur trioxideThe subscript 1(for S) or mono is understood.
52 Naming Covalent Compounds (Continued) Name P4S3STEP 1 The first nonmetal P is phosphorus.STEP 2 The second nonmetal S is sulfide.STEP 3 The subscript 4 of P is shown as tetra.The subscript 3 of O is shown as tri.P4S3 → tetraphosphorus trisulfide
53 Guide to Writing Formulas for Covalent Compounds
54 Writing Formulas of Covalent Compounds Write the formula for carbon disulfide.STEP 1 Elements are C and SSTEP 2 No prefix for carbon means 1 CPrefix di = 2Formula: CS2
55 Chapter 5 Compounds and Their Bonds 5.7 Electronegativity and Bond Polarity
56 Electronegativity The electronegativity value indicates the attraction of an atom for shared electronsincreases from left to right going across a period on the periodic tabledecreases going down a group on the periodic tableis high for the nonmetals, with fluorine as the highestis low for the metals
57 Some Electronegativity Values High valuesLow values
58 Nonpolar Covalent Bonds A nonpolar covalent bondoccurs between nonmetalshas an equal or almost equal sharing of electronshas almost no electronegativity difference (0.0 to 0.4)Examples:Atoms Electronegativity Type of Bond Difference _______________N–N – 3.0 = Nonpolar covalentCl–Br 3.0 – 2.8 = Nonpolar covalentH–Si – 1.8 = Nonpolar covalent
59 Polar Covalent Bonds A polar covalent bond occurs between nonmetal atomshas an unequal sharing of electronshas a moderate electronegativity difference(0.5 to 1.7)Examples:Atoms Electronegativity Type of BondDifference_________________________________________O–Cl – 3.0 = Polar covalentCl–C 3.0 – 2.5 = Polar covalentO–S – 2.5 = Polar covalent
61 Chapter 5 Compounds and Their Bonds 5.8 Shapes and Polarity of Molecules
62 VSEPR In the valence-shell electron-pair repulsion theory (VSEPR), the electron groups around a central atomare arranged as far apart from each other as possiblehave the least amount of repulsion of the negatively charged electronshave a geometry around the central atom that determines molecular shape
63 Shapes of Molecules The three-dimensional shape of a molecule is the result of bonded groups and lone pairs of electrons around the central atomis predicted using the VSEPR theory (valence-shell-electron-pair repulsion)
64 Guide to Predicting Molecular Shape (VSEPR Theory)
65 Two Electron Groups In a molecule of BeCl2, there are two electron groups bonded to the central atom, Be (Be is an exception to the octet rule): Cl : Be : Cl :to minimize repulsion, the arrangement of two electron groups is 180°, or opposite each otherthe shape of the molecule is linear
66 Two Electron Groups with Double Bonds In a molecule of CO2,there are two electron groups bonded to C (electrons in each double bond are counted as one group)repulsion is minimized with the double bonds opposite each other at 180°the shape of the molecule is linear
67 Three Electron Groups In a molecule of BF3, three electron groups are bonded to the central atom B (B is an exception to the octet rule)..: F:: F : B : F :repulsion is minimized with 3 electron groups at angles of 120°the shape is trigonal planar
68 Two Electron Groups and One Lone Pair In a molecule of SO2,S has 3 electron groups; 2 electron groups bonded to O atoms and one lone pair:O :: S : O :..repulsion is minimized with the electron groups at angles of 120°, a trigonal planar arrangementthe shape is bent (120°), with two O atoms bonded to S● ●
69 Four Electron Groups In a molecule of CH4, there are four electron groups around Crepulsion is minimized by placing four electron groups at angles of 109°, which is a tetrahedral arrangementthe four bonded atoms form a tetrahedral shape
70 Three Bonding Atoms and One Lone Pair In a molecule of NH3,three electron groups bond to H atoms, and the fourth one is a lone (nonbonding) pairrepulsion is minimized with 4 electron groups in a tetrahedral arrangementthe three bonded atoms form a pyramidal (~109°) shape
71 Two Bonding Atoms and Two Lone Pairs In a molecule of H2O,two electron groups are bonded to H atoms and two are lone pairs (4 electron groups)four electron groups minimize repulsion in a tetrahedral arrangementthe shape with two bonded atoms is bent (~109°)
72 Determining Molecular Polarity Determine the polarity of the H2O molecule. Solution: The four electron groups of oxygen are bonded to two H atoms. Thus the H2O molecule has a net dipole, which makes it a polar molecule.
73 Chapter 5 Compounds and Their Bonds 5.9 Attractive Forces inCompounds
74 Dipole–Dipole Attractions In covalent compounds, polar moleculesexert attractive forces called dipole-dipole attractionsform strong dipole attractions called hydrogen bonds between hydrogen atoms bonded to F, O, or N, and other atoms that are very electronegative
78 Melting Points and Attractive Forces Ionic compounds require large amounts of energy to break apart ionic bonds. Thus, they have high melting points.Hydrogen bonds are the strongest type of dipole–dipole attractions. They require more energy to break than do other dipole attractions.Dispersion forces are weak interactions, and very little energy is needed to change state.