Chapter 4 Compounds and Their Bonds

4.4 Polyatomic Ions

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Polyatomic Ions

A polyatomic ion • is a group of atoms. • has an overall ionic charge.

Some examples of polyatomic ions are

NH4+ ammonium OH− hydroxide

NO3−

nitrate NO2−

nitrite

CO32− carbonate PO4

3− phosphateHCO3

− hydrogen carbonate(bicarbonate)

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Some Compounds with Polyatomic Ions

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Some Names of Polyatomic IonsThe 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 the prefix hydrogen (or bi).

HCO3− hydrogen carbonate (bicarbonate)

HSO3− hydrogen sulfite (bisulfite)

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Names and Formulas of Common Polyatomic Ions

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Naming Compounds with Polyatomic Ions

• 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

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Some Compounds with Polyatomic Ions

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Learning CheckMatch each formula with the correct name.

A. MgS 1) magnesium sulfite

MgSO3 2) magnesium sulfate

MgSO4 3) magnesium sulfide

B. Ca(ClO3)2 1) calcium chlorate

CaCl2 2) calcium chlorite

Ca(ClO2)2 3) calcium chloride

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Learning CheckName each of the following compounds:

A. Mg(NO3)2

B. Cu(ClO3)2

C. PbO2

D. Fe2(SO4)3

E. Ba3(PO3)2

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Writing Formulas with Polyatomic IonsThe 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 has the polyatomic ions in parentheses.

Mg2+ and 2NO3−

-> Mg(NO3)2

subscript 2 for charge balance

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Learning CheckSelect the correct formula for each.

A. aluminum nitrate

1) AlNO3 2) Al(NO)3 3) Al(NO3)3

B. copper(II) nitrate

1) CuNO3 2) Cu(NO3)2 3) Cu2(NO3)

C. iron(III) hydroxide

1) FeOH 2) Fe3OH 3) Fe(OH)3

D. tin(IV) hydroxide

1) Sn(OH)4 2) Sn(OH)2 3) Sn4(OH)

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Learning CheckWrite the correct formula for each.

A. potassium bromateB. calcium carbonateC. sodium phosphateD. iron(II) nitrite

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Flowchart for Naming Ionic Compounds

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Learning CheckName the following compounds:

A. Ca3(PO4)2

B. FeBr3

C. Al2S3

D. Zn(NO2)2

E. NaHCO3

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Learning CheckWrite the formulas for the following:

A. calcium nitrate

B. iron(II) hydroxide

C. aluminum carbonate

D. copper(II) bromide

E. lithium phosphate

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Chapter 4 Compounds and Their Bonds

4.5Covalent Compounds

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Covalent Bonds

Covalent bonds form

• when atoms share electrons to complete octets.

• between two nonmetal atoms.

• between nonmetal atoms from Groups 4A (14), 5A (15), 6A (16), and 7A (17).

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Hydrogen Molecule

A hydrogen molecule

• is stable with 2 electrons (helium).

• has a shared pair of electrons.

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Forming Octets in Molecules

In a fluorine, F2,, molecule, each F atom

• shares 1 electron.

• attains an octet.

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Carbon Forms 4 Covalent Bonds

In a CH4 (methane) molecule,

• 1 C atom shares electrons with 4 H atoms to attain an octet.

• each H atom shares 1 electron to become stable, like helium.

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Multiple Bonds

In a nitrogen molecule, N2,

• each N atom shares 3 electrons.• each N attains an octet.• the bond is a multiple bond called a triple

bond.• the name is the same as the element.

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Naming Covalent Compounds

In the names of covalent compounds, prefixes are used to indicate the number of atoms (subscript) of each element. (mono is usually omitted)

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Naming Covalent CompoundsWhat is the name of SO3?

1. The first nonmetal is S sulfur.

2. The second nonmetal is O, named oxide.

3. The subscript 3 of O is shown as the prefix tri.

SO3 -> sulfur trioxide

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

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Naming Covalent CompoundsName P4S3.

1. The first nonmetal, P, is phosphorus.

2. The second nonmetal, S, is sulfide.

3. The subscript 4 of P is shown as tetra.

The subscript 3 of O is shown as tri.

P4S3 -> tetraphosphorus trisulfide

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Formulas and Names of Some Covalent Compounds

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Learning CheckSelect the correct name for each compound.

A. SiCl4 1) silicon chloride2) tetrasilicon chloride3) silicon tetrachloride

B. P2O5 1) phosphorus oxide2) phosphorus pentoxide3) diphosphorus pentoxide

C. Cl2O7 1) dichlorine heptoxide2) dichlorine oxide3) chlorine heptoxide

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Learning CheckWrite the name of each covalent compound.

CO2 _____________________

PCl3 _____________________

CCl4 _____________________

N2O _____________________

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Example: Writing Formulas for Covalent Compounds

Write the formula for carbon disulfide.

STEP 1: Elements are C and S

STEP 2: No prefix for carbon means 1 C

Prefix di = 2

Formula: CS2

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Learning CheckWrite the correct formula for each of the following:

A. phosphorus pentachloride

B. dinitrogen trioxide

C. sulfur hexafluoride

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Learning CheckIdentify each compound as ionic or covalent, and

give itscorrect name.

A. SO3

B. BaCl2

C. (NH4)3PO3

D. Cu2CO3

E. N2O4

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Study Tip: Ionic or CovalentA compound is

• ionic if the first element in the formula

or the name is a metal or the polyatomic ion NH4+ .

K2O K is a metal; compound is ionic;

potassium oxide

• covalent if the first element in the formula or the name is a nonmetal.

N2O N is a nonmetal; compound is covalent;

dinitrogen oxide

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Learning CheckIdentify each compound as ionic or covalent and

give itscorrect name.

A. Ca3(PO4)2

B. FeBr3

C. SCl2

D. Cl2O

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Learning CheckDetermine if each is ionic (I) or covalent (C ), and

write the formula.

A. calcium nitrate

B. boron trifluoride

C. aluminum carbonate

D. dinitrogen tetroxide

E. copper(I) phosphate33

Chapter 4 Compounds and Their Bonds

4.6

Electronegativity and Bond Polarity

34Copyright © 2009 by Pearson Education, Inc.

Electronegativity

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.

• is high for the nonmetals, with fluorine as the highest.

• is low for the metals.

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Some Electronegativity Values for Group A Elements

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Low values

High values

` Electronegativity increases

` E

lectro

neg

ativity d

ecre

ase

s

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Nonpolar Covalent Bonds

A nonpolar covalent bond

• occurs between nonmetals.• is an equal or almost equal sharing of electrons.• has almost no electronegativity difference (0.0 to 0.4).

Examples: Electronegativity Atoms Difference Type of BondN-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

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Polar Covalent Bonds

A polar covalent bond

• occurs between nonmetal atoms.• is an unequal sharing of electrons.• has a moderate electronegativity difference (0.5 to

1.7).

Examples: Electronegativity

Atoms Difference Type of BondO-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

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Comparing Nonpolar and Polar Covalent Bonds

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Ionic BondsAn ionic bond • occurs between metal and nonmetal ions.• is a result of electron transfer.• has a large electronegativity difference (1.8 or

more).

Examples: Electronegativity

Atoms Difference Type of BondCl-K 3.0 – 0.8 = 2.2 IonicN-Na 3.0 – 0.9 = 2.1 IonicS-Cs 2.5 – 0.7 = 1.8 Ionic

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Electronegativity and Bond Types

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Predicting Bond Types

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Learning CheckUse the electronegativity difference to identify the type of bond [nonpolar covalent (NP), polar covalent (P), or ionic (I)] between the following:

A. K-NB. N-OC. Cl-ClD. H-Cl

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Chapter 4 Compounds and Their Bonds

4.7Shapes and Polarity of

Molecules

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VSEPRIn 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.

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Four Electron Groups

In a molecule of CH4,

• there are 4 electron groups around C.

• repulsion is minimized by placing 4 electron groups at angles of 109°, which is a tetrahedral arrangement.

• the shape with four bonded atoms is tetrahedral.

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Three Bonding Atoms and One Lone Pair

In a molecule of NH3,

• 3 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.

• with 3 bonded atoms, the shape is pyramidal.

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Two Bonding Atoms and Two Lone PairsIn a molecule of H2O,

• 2 electron groups are bonded to H atoms and 2 are lone pairs (4 electron groups).

• 4 electron groups minimize repulsion in a tetrahedral arrangement.

• the shape with 2 bonded atoms is bent.

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Shapes with 4 Electron Groups

Electron Pairs

Bonded Atoms

Lone Pairs

Molecular Shape

Example

4 4 0 Tetrahedral CH4

4 3 1 Pyramidal NH3

4 2 2 Bent H2O

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Learning CheckState the number of electron groups, lone pairs, and use VSEPR theory to determine the shape of the following molecules or ions.

1) tetrahedral 2) pyramidal 3) bent

A. PF3

B. H2S

C. CCl4

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Study TipTo determine shape, 1. draw the electron-dot structure.2. count the electron pairs around the central atom.

3. count the bonded atoms to determine shape.

4 electron pairs and 4 bonded atoms = tetrahedral 4 electrons pairs and 3 bonded atoms = pyramidal

4 electron pairs and 2 bonded atoms = bent

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Polar MoleculesA polar molecule • contains polar bonds.

• has a separation of positive and negative charge called a dipole, indicated with + and -.

• has dipoles that do not cancel. + -

• •

H–Cl H—N—H dipole

H dipoles do

not cancel

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Nonpolar MoleculesA nonpolar molecule • contains nonpolar bonds.

Cl–Cl H–H

• or has a symmetrical arrangement of polar bonds.

O=C=O Cl

Cl–C–Cl

Cl

dipoles cancel

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Determining Molecular Polarity STEP 1: Write the electron-dot formula. STEP 2: Determine the shape. STEP 3: Determine if dipoles cancel or not.

Example: H2O

. .

H─O: H2O is polar

│ H

dipoles do not cancel

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Learning CheckIdentify each of the following molecules as 1) polar or 2) nonpolar. Explain.

A. PBr3

B. HBr

C. Br2

D. SiBr4

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Chapter 4 Compounds and Their Bonds

4.8 Attractive Forces in Compounds

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Ionic Bonds

In ionic compounds, ionic bonds • are strong attractive forces.• hold positive and negative ions

together.

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Dipole-Dipole Attractions

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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 very electronegative atoms.

Dispersion Forces

Dispersion forces are• weak attractions between nonpolar

molecules.• caused by temporary dipoles that develop

when electrons are not distributed equally.

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Attractive Forces

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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 other dipole-dipole attractions.

• Dispersion forces are weak interactions and very little energy is needed to change state.

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Melting Points of Some Substances

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Learning CheckIdentify the main type of attractive forces for each:1) ionic 2) dipole-dipole 3) hydrogen bonds 4) dispersion

A. NCl3

B. H2O

C. Br-BrD. KClE. NH3

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