1 Chapter 5 Compounds and Their Bonds 5.1 Octet Rule and Ions

1

Chapter 5Compounds and Their Bonds

5.1Octet Rule and Ions

2

An octet is 8 valence electrons is associated with the stability of the noble gases does not occur with He; He is stable with 2 valence

electrons (duet)Valence Electrons

He 1s2 2

Ne 1s22s22p6 8

Ar 1s22s22p63s23p6 8

Kr 1s22s22p63s23p64s23d104p6 8

Octet Rule

3

Ionic and Covalent Bonds

Atoms form octets to become more stable by losing, gaining, or sharing

valence electrons by forming ionic or covalent

bonds

4

Metals Form Positive Ions

Metals form positive ions by a loss of their valence electrons with the electron configuration of the

nearest noble gas that have fewer electrons than protons

Group 1A(1) metals ion 1+



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 ion 11p+ 11p+

11e– 10e–

0 1 +

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+ ion 12p+ 12p+

12e– 10e–

0 2+

7

Formation of Negative Ions

In ionic compounds, nonmetals achieve an octet arrangement gain electrons form 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) = 6 – 8 = 2–

or 16 – 18 = 2–

Group Number and Ionic Charge

Ions achieve the electron configuration of their nearest noble gas of 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.

10

Groups Numbers for Some Positive and Negative Ions

11

12

Chapter 5 Compounds and Their Bonds

5.2Ionic Compounds

13

Ionic compounds consist of positive and negative ions have attractions called ionic bonds between positively

and negatively charged ions have high melting and boiling points are solid at room temperature

Ionic Compounds

14

Salt Is an Ionic Compound

Sodium chloride, or “table salt,” is an example of an ionic compound.

15

An ionic formula consists of positively and negatively charged ions is neutral has charge balance

total positive charge = total negative charge

The symbol of the metal is written first, followed by the symbol of the nonmetal.

Ionic Formulas

16

Charge Balance for NaCl, “Salt”In NaCl, a Na atom loses its valence electron a Cl atom gains an electron the symbol of the metal is written first, followed by the

symbol of the nonmetal.

17

Charge Balance In MgCl2In MgCl2, a Mg atom loses two valence electrons two Cl atoms each gain one electron subscripts 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− = Na3N

Na+

3(1+) + 1(3–) = 0

19


5.3Naming and Writing Ionic Formulas

20

Naming Ionic Compounds with Two Elements

To name a compound with two elements, identify the cation and

anion name the cation first,

followed by the name of the anion

21

Formula Ions Name Cation Anion

NaCl Na+ Cl– sodium chloride

K2S K+ S2– potassium sulfideMgO Mg2+ O2– magnesium oxide

CaI2 Ca2+ I– calcium iodide

Al2O3 Al3+ S2– aluminum sulfide

Examples of Ionic Compounds with Two Elements

22

Transition Metals Form Positive Ions

Most transition metals and Group 4(14) metals, Form 2 or more positive ions Zn2+, Ag+, and Cd2+ form only one ion.

Metals with Variable Charge

The names of transitionmetals with two or morepositive ions (cations) use aRoman numeral after thename of the metal to identifyionic charge.

23

24

Naming Ionic Compounds with Variable Charge Metals

25

Naming FeCl2

STEP 1 Determine the charge of the cation from the anion.

Fe ion + 2 Cl– = Fe ion + 2– = 0 Fe 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 Cr2O3

STEP 1 Determine the charge of cation from the anion. 2Cr ions + 3O2– = 2Cr ions + 3(2–)

= 2Cr ions + 6– = 0 2Cr 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

27

Guide to Writing Formulas from the Name

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

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. Fe3+ Cl−

Cl− Cl−

1(3+) + 3(1–) = 0

STEP 3 Write the cation first.1Fe3+ and 3Cl− = FeCl3

30


5.4Polyatomic Ions

31

A polyatomic ion is a group of atoms has an overall ionic chargeExamples:

NH4+ ammonium OH− hydroxide

NO3−

nitrate NO2−

nitrite

CO32− carbonate PO4

3− phosphate

HCO3− hydrogen carbonate

(bicarbonate)

Polyatomic Ions

32

The names of common polyatomic anions end in ate

NO3− nitrate PO4

3− phosphate with one oxygen less end in ite

NO2− nitrite PO3

3− phosphite with hydrogen attached use prefix hydrogen (or bi)

HCO3− hydrogen carbonate (bicarbonate)

HSO3− hydrogen sulfite (bisulfite)

Some Names of Polyatomic Ions

Guide to Naming Compounds with Polyatomic Ions

33

34

The positive ion is named first followed by the name of the polyatomic ion.

NaNO3 sodium nitrate

K2SO4 potassium sulfate

Fe(HCO3)3 iron(III) bicarbonate

or iron(III) hydrogen carbonate (NH4)3PO3 ammonium phosphite

Examples of Names of Compounds with Polyatomic Ions

35

Writing Formulas with Polyatomic Ions

The formula of an ionic compound containing a polyatomic ion must have a charge balance

that equals zero (0)Na+ and NO3

− NaNO3

with two or more polyatomic ions put the polyatomic ions in parentheses.

Mg2+ and 2NO3−

Mg(NO3)2

subscript 2 for charge balance

36


5.5Covalent Compounds

37

Forming Octets in Molecules

In a fluorine (F2) molecule, each F atom shares one electron acquires an octet

38

Diatomic Elements

These elements share electrons to form diatomic, covalent molecules.

Guide to Writing Electron-Dot Formulas

39

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–) + 2 S(6e–) = 16e–

STEP 3 Attach each S atom to the central C atom using one electron pair.

S : C : S 16e– – 4e– = 12e– remaining

STEP 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 and S atoms to make two double bonds.

43

A Nitrogen Molecule has A Triple Bond

In a nitrogen molecule, N2, each N atom shares 3 electrons each N atom attains an octet the sharing of 3 sets of electrons is a multiple bond

called a triple bond

44

Resonance structures are two or more electron-dot formulas for the same

arrangement of atoms related by a double-headed arrow ( ) written by changing the location of a double bond

between the central atom and a different attached atom

Resonance Structures

45

Sulfur dioxide has two resonance structures. STEP 1 Write the arrangement of atoms.

O S O STEP 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 4e− used 18e− – 4e− = 14e− remaining

Writing Resonance Structures

46

STEP 4 Add 14 remaining electrons as 7 lone pairs.

STEP 5 Form a double bond to complete octets. Two resonance structures are possible.

Writing Resonance Structures (continued)

47

Write the ionic formula of the compound containing Ba2+ and Cl.

Write the symbols of the ions.Ba2+ Cl

Balance the charges. Ba2+ Cl two Cl needed

Cl

Write the ionic formula using a subscript 2 for two chloride ions that give charge balance.

BaCl2

Formula from Ionic Charges

48


5.6Naming and Writing Covalent Formulas

NO nitrogen oxideNO2 nitrogen dioxide

N2O4 dinitrogen tetroxide

49

Names of Covalent Compounds

Prefixes are used in the names of covalent compounds because two nonmetals can form two or more different

compoundsExamples of compounds of N and O:

NO nitrogen oxideNO2 nitrogen dioxide

N2O dinitrogen oxide

N2O4 dinitrogen tetroxide

N2O5 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

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 trioxide

The subscript 1(for S) or mono is understood.

Naming Covalent Compounds (Continued)

52

Name P4S3

STEP 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

Naming Covalent Compounds (Continued)

Guide to Writing Formulas for Covalent Compounds

53

54

Write the formula for carbon disulfide.

STEP 1 Elements are C and SSTEP 2 No prefix for carbon means 1 C

Prefix di = 2 Formula: CS2

Writing Formulas of Covalent Compounds

55


5.7Electronegativity and Bond Polarity

56

The electronegativity value indicates the attraction of an atom for shared electrons increases from left to right going across a period on the

periodic table decreases going down a group on the periodic table is high for the nonmetals, with fluorine as the highest is low for the metals

Electronegativity

57

Some Electronegativity Values

Low values

High values

58

A nonpolar covalent bond occurs between nonmetals has an equal or almost equal sharing of electrons has almost no electronegativity difference (0.0 to 0.4)

Examples: Atoms Electronegativity Type of Bond

Difference _______________N–N 3.0 – 3.0 = 0.0 Nonpolar covalentCl–Br 3.0 – 2.8 = 0.2 Nonpolar covalentH–Si 2.1 – 1.8 = 0.3 Nonpolar covalent

Nonpolar Covalent Bonds

59

A polar covalent bond occurs between nonmetal atoms has an unequal sharing of electrons has a moderate electronegativity difference

(0.5 to 1.7)

Examples: Atoms Electronegativity Type of Bond Difference_________________________________________

O–Cl 3.5 – 3.0 = 0.5 Polar covalentCl–C 3.0 – 2.5 = 0.5 Polar covalentO–S 3.5 – 2.5 = 1.0 Polar covalent

Polar Covalent Bonds

60

Comparing Nonpolar and Polar Covalent Bonds

61


5.8Shapes and Polarity of Molecules

62

VSEPR

In the valence-shell electron-pair repulsion theory(VSEPR), the electron groups around a central atom are arranged as far apart from each other as possible have the least amount of repulsion of the negatively

charged electrons have 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 atom is 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 other

the 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 arrangement

the 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 C repulsion is minimized by placing four electron groups at

angles of 109°, which is a tetrahedral arrangement the 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) pair repulsion is minimized with 4 electron groups in a

tetrahedral arrangement the 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

arrangement the shape with two bonded atoms is bent (~109°)

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.

72

73


5.9Attractive Forces in

Compounds

74

Dipole–Dipole Attractions

In covalent compounds, polar molecules exert attractive forces called dipole-dipole

attractions form strong dipole attractions called hydrogen

bonds between hydrogen atoms bonded to F, O, or N, and other atoms that are very electronegative

75

Dipole–Dipole Attractions (continued)

76

Dispersion Forces

Dispersion forces are weak attractions between nonpolar molecules caused by temporary dipoles that develop when

electrons are not distributed equally

77

Comparison of Bonding and Attractive Forces

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.

Documents

1 Chapter 5 Compounds and Their Bonds 5.1 Octet Rule and Ions