1

Name: ___________________________________________

Unit 6- Bonding and Nomenclature

Learn by 11/18!

Polyatomic Ion Name

NH4+

Ammonium

H3O+

Hydronium

NO3-

Nitrate

OH-

Hydroxide

CN-

Cyanide

C2H3O2-

Acetate

MnO4-

Permanganate

HCO3-

Bicarbonate (or Hydrogen Carbonate)

SO4-2

Sulfate

CO3-2

Carbonate

CrO4-2

Chromate

PO4-3

Phosphate

ClO3-

Chlorate

Day Page # Description IC/HW Due Date Completed

ALL 2 – 3 Warm-up IC

1 4 – 6 Bonding Notes IC

1 7 Hybridization IC

1 8 Lewis Structures Notes IC

1 9 VSPER Notes IC

1 10 VSPER Practice IC

2 11 – 13 Polarity IC

2 14 – 15 Carbon! IC

2 16 Bonding and Shapes Quiz Review HW

3 X Bonding and Shapes Quiz Thursday, November 17 and Friday,

November 18

3 17 – 18 Naming Notes IC

3 19 - 20 Polyatomic Ion Naming IC

4 21 Covalent Naming IC

4 22 Naming Acids IC

4 23 – 24 Organic Naming IC

5 25 - 27 Unit 6 Review HW

6 X Unit 6 Test MC on Moodle 12/5 – 12/9

In Class on 12/6 and 12/7

2

Warm-Up

Day 1

1. What is an ionic compound?

2. What is a covalent compound?

Day 2

1. Draw the Dot Diagram for PCl3

2. Draw the Dot Diagram for CaCl2

3. Complete the table below with the name of each shape:

Steric Number 0 Lone Pairs 1 Lone Pair 2 Lone Pairs

2

3

4

3

Day3

1. What two shapes are ALWAYS polar?

2. Give the formula for Boron and Chlorine.

Day 4

1. Complete the table

Polyatomic Ion Name

PO4-3

Chromate

OH-

ClO3-

Permanganate

Acetate

Ammonium

Day 5

1. Name the following compounds:

a. NaCl

b. BaBr2

c. CCl4

d. SO2

e. SnCl4

f. PF3

4

Bonding Notes

• There are three main types of bonding: ionic, metallic and covalent.

• Ionic Bonding occurs when the electronegativity difference between two atoms is greater than

__________.

• Metallic bonding occurs between two ____________________.

• Covalent bonding occurs when the electronegativity difference between two atoms is less than

___________.

Metallic Bonding

• Metallic bonds consist of positively charged metallic cations that donate electrons to the

__________________.

• The “sea” of electrons are shared by all atoms and can move throughout the structure.

• Properties:

o Thermal Conductivity

o Electrical Conductivity

o Malleability- the ability to be hammered down into thin sheets.

o Ductility- the ability to be drawn into a wire.

Ionic Bonds

• When a metal bonds with a nonmetal, an bond is formed.

• An ionic bond always involves the TRANSFER of electrons from the to the

.

• The cation and anion are held together by ___________________________________.

• An ionic compound does not consist of individual molecules. Instead, there is a huge network of

positive and negative ions that are packed together in a .

• Because their bonds are so strong, ionic compounds tend to have very melting

points.

• Ionic compounds are _____________________________, which means they can conduct electricity.

• A positive ion is called a . A negative ion is called an .

When forming ionic compounds the positive and negative charges must balance.

5

The Criss - Cross Rule

• When forming ionic compounds, to determine the ratio of cation to anion, take the absolute value

of the charge of the cation, and put it as the subscript for the anion. Take the absolute value of the

charge of the anion, and put is as the subscript for the cation. Simplify if necessary.

o EX: Aluminum and Oxygen

o EX: Barium and Oxygen

• Criss-Cross Practice:

o Lithium Iodide

o Strontium Chloride

o Aluminum Nitride

o Sodium Sulfide

o Potassium Oxide

• Here are some more compounds for practice. Fill in the table with the missing information:

Compound Name Cation Anion Binary Ionic Formula

lithium sulfide

calcium nitride

potassium oxide

strontium phosphide

iron(II) oxide

iron(III) oxide

cobalt(II) chloride

nickel(III) sulfide

6

Covalent Bonding

• Covalent bonds are between _____________________________________________.

• Covalent bonds are formed when electrons are _____________________ between two atoms. If

two atoms share 4 electrons, they form a ______________________________. If two atoms share

6 electrons, they form a ____________________________.

• There are two types of covalent bonds: polar and non-polar. Polar bonds have an electronegativity

difference between ___________________________. Non-polar bonds have an electronegativity

difference less than ___________________.

• In polar bonds, the electrons are shared ________________________________. In non-polar

bonds, the electrons are shared ________________________________.

• Covalent compounds can exist in any state (solid, liquid or gas). They have _____________ melting

and boiling points.

• Write the correct formulas for each covalent compound:

Compound Name Oxidation States Covalent Formula

water O (-2)

H (+1)

Carbon Dioxide C (+4)

O (-2)

Chlorine (Diatomic Element) Cl (-1)

Methane (5 total atoms) C (-4)

H (+1)

Ammonia (4 total atoms) N (-3)

H (+1)

Carbon tetrabromide (5 total

atoms)

C (+4)

Br (-1)

Phosphorous trichloride (4 total

atoms)

P (-3)

Cl (-1)

Diphosphorous trioxide (5 total

atoms)

P (-3)

O (-2)

7

Hybridization

• When atoms form bonds, they form two different types: sigma (σ) and pi (π). The sigma bond is the

single bond. A double bond would count as one sigma and one pi. A triple bond counts as one sigma

and two pi.

• We want all the sigma bonds to have the same energy (remember, s orbitals have less energy than p

orbitals). To achieve this, the orbitals will ________________________.

• When the s and three p orbitals hybridize, they form _____________sp3 orbitals. These orbitals

have 25% s character and 75% p character.

• We need one hybridized orbital for each ____________________ bond and each lone pair.

• CH4 has four sigma bonds, so it will need to form four sp3 orbitals.

• NH3 has three sigma bonds and a lone pair, so it will need to form four sp3

orbitals.

• CO2 has two sigma bonds and two pi bonds. It only needs two hybridized orbitals, so it will form two

sp orbitals. The pi bonds will just use regular p orbitals.

• CH2O has three sigma bonds and a pi bond. What do you think the hybridization for CH2O will be?

8

Lewis Structures Notes

• Draw the dot diagram for each atom. Make sure you place the electrons in the correct order.

• Draw the dot diagrams for Carbon, Nitrogen and Oxygen.

• Steps for drawing Lewis Structures:

o Order the atoms- the atom with the most unpaired electrons will be the central atom (Carbon

is always the central atom).

o Draw the dot diagram for each element in the compound.

o Pair up all unpaired electrons.

• Examples:

o Phosphorous and Chlorine (PCl3)

o Carbon and Bromine (CBr4)

• Ionic Dot Diagrams- the electrons are gained and lost, not shared. Put each atom in a bracket with

its balance electrons around it. The charge of the atom goes outside of the bracket.

o Ex: Sodium Chloride (NaCl)

9

VSPER- Valence Shell Electron Pair Repulsion

Bonds Lone Pairs Shape

Linear

Bent

Trigonal Planar

Trigonal Pyramidal

Tetrahedral

Steric number is the total number of ______________________________ AND ______________________.

10

VSEPR Practice

Complete the table with the requested information.

Molecule Structural Diagram Oxidation State of each

element

Geometry

CClF3

SF2

BF3

SiBr4

NH3

11

Polarity

Bond Polarity

Electronegativity

Ionic bonds have an electronegativity difference that is greater than 1.7. Covalent bonds have an

electronegativity difference less than (or equal to) 1.7. Electronegativity differences between 0 and 0.4

indicate non-polar covalent bonds. Electronegativity differences between 0.4 and 1.7 indicate polar

covalent bonds.

Polar Covalent Bond- a covalent bond in which the electrons are not shared equally because one atom

attracts them more strongly than the other.

Non-polar Covalent Bond- a covalent bond in which the electrons are shared equally.

Use the periodic table of electronegativities to answer the questions on electronegativity differences.

H

2.1

Li

1.0

Be

1.5 ELECTRONEGATIVITY

(electron attraction!)

B

2.0

C

2.5

N

3.0

O

3.5

F

4.0

Na

0.9

Mg

1.2

Al

1.5

Si

1.8

P

2.1

S

2.5

Cl

3.0

K

0.8

Ca

1.0

Sc

1.3

Ti

1.5

V

1.6

Cr

1.6

Mn

1.5

Fe

1.8

Co

1.9

Ni

1.9

Cu

1.9

Zn

1.6

Ga

1.6

Ge

1.8

As

2.0

Se

2.4

Br

2.8

Rb

0.8

Sr

1.0

Y

1.2

Zr

1.4

Nb

1.6

Mo

1.8

Tc

1.9

Ru

2.2

Rh

2.2

Pd

2.2

Ag

1.9

Cd

1.7

In

1.7

Sn

1.8

Sb

1.9

Te

2.1

I

2.5

Cs

0.7

Ba

0.9

La-Lu

1.0-1.2

Hf

1.3

Ta

1.5

W

1.7

Re

1.9

Os

2.2

Ir

2.2

Pt

2.2

Au

2.4

Hg

1.9

Tl

1.8

Pb

1.9

Bi

1.9

Po

2.0

At

2.2

Fr

0.7

Ra

0.9

Ac

1.1

Th

1.3

Pa

1.4

U

1.4

Np-No

1.4-1.3

Determine the type of bond that would form between the following two elements using differences in

electronegativity.

Example: Mg – O

O is 3.5 and Mg is 1.2, therefore, the difference is 3.5 – 1.2 = 2.3 IONIC

Example: Cl – Cl

Cl is 3.0. The difference is 3.0 – 3.0 = 0 NON-POLAR COVALENT

Bond Electronegativity Difference Bond Type

1. C – N

2. Li - F

12

Bond Electronegativity Difference Bond Type

3. N – Cl

4. Na - Cl

5. O – F

6. B – H

7. Ba – F

8. C – H

Molecular Polarity

Dipole moment- a property of a molecule whereby the charge distribution can be represented by a center

of positive charge and a center of negative charge.

Polar Molecule- a molecule that has a permanent dipole moment.

Determining if a molecule is polar.

1. If ALL of the bonds are non-polar, then the molecule is non-polar.

2. If some or all of the bonds are polar, you can consider the vectors. Vectors are arrows that point in

the direction of the negative charge (the direction the electrons are pulled). Examples:

O

POLAR

H H

a. If all the arrows point toward the central atom AND the central atom has lone pairs, then the

atom is polar.

H

NON-POLAR

H C H

H

b. If all the arrows point toward the central atom AND the central atom has no lone pairs, then

the atom is non-polar.

H F POLAR

c. If the arrow points toward one atom in a linear molecule, it is polar.

d. If the arrows all point away from the central atom AND there are no lone pairs on the central

atom, then the atom is non-polar.

13

e. If the arrows all point away from the central atom AND there are lone pairs on the central

atom, then the atom is polar.

3. Another way to determine if a molecular is polar or not, is to look at symmetry.

H

NON-POLAR

H C H O C O

H

a. If the molecule is symmetrical, then the molecule is non-polar.

H

H C H

F POLAR

H Cl

b. If the molecule is not symmetrical, then the molecule is polar.

c. If the molecule is bent or trigonal pyrimidal, then the molecule is polar (lone pairs on the

central atom mean that it is polar).

Draw the structural formula for each molecule, and determine if it is polar or non-polar.

Formula Structural Formula Polar/Non-polar

NH3

SCl2

CF4

PCl3

H2S

C2H2

14

Carbon!

Carbon is in the Carbon Group (14), so it has 4 valence electrons. Carbon tends to hybridize.

Hybridization of Carbon

Hybridization Single Bond (1 σ) Double Bond (1 σ, 1 π) Triple Bond (1 σ, 2π)

sp3

4 -- --

sp2

2 1 --

sp -- 2 --

sp 1 -- 1

Compounds containing carbon are considered organic. Organic compounds are necessary for life.

Examples you are familiar with are glucose (C6H12O6), aspirin (C9H8O4) and propane (C3H8). Carbon is part of

the backbone and the base pairs of DNA and RNA (remember biology).

Structure

Lewis Dot Structure for Carbon:

When carbon bonds with 4 hydrogens, it is called Methane, and has the following structure:

This structure is non-polar because it is symmetrical. The geometry is

tetrahedral because it has four sigma (σ) bonds and no lone pairs.

When C2H6 (ethane) forms, the carbons will form a single bond with hydrogens around the outside

(remember, always save the hydrogens for last).

This structure is non-polar because it is symmetrical. Each carbon has a

tetrahedral geometry because they each have four sigma (σ) bonds and no

lone pairs.

When C2H4 (ethene) forms, the carbons will form a double bond with hydrogens around the outside. You

would determine the Lewis dot structure in the following manner:

1. Put the carbons next

to each other and form

a bond.

2. Put equal number of

hydrogens on each carbon.

3. Each carbon has one electron left

over, so these electrons will also

bond. This gives you a double bond.

This will give you the following Lewis dot structure:

This structure is non-polar and the geometry around each

carbon is trigonal planar because each carbon has three sigma

(σ) bonds and one pi (π) bond. (Remember, do not count pi

bonds when determining geometry!)

15

When C2H2 (ethyne) forms, the carbons will form a triple bond with one hydrogen on each carbon. The

structure would look like this:

This structure is non-polar and the geometry around each

carbon is linear because each carbon has two sigma (σ) bonds

and two pi (π) bonds.

When CH3CH2OH (ethanol) forms, the carbons will form a single bond, and one carbon will bond with the

oxygen. The hydrogens go around the outside.

This structure is POLAR because it is asymmetrical. Each carbon has

a tetrahedral geometry because they each have four sigma (σ)

bonds. The geometry around the oxygen is bent because it has two

sigma (σ) bonds and two lone pairs.

When CH2O (formaldehyde) forms, the carbon will double bond to the oxygen, and form single bonds with

each of the hydrogens.

This structure is POLAR because it is asymmetrical. The geometry

around the carbon is trigonal planar because it has three sigma (σ)

bonds and one pi (π) bond. Formaldehyde has a permanent dipole.

When C6H6 (benzene) forms, the carbons will form a ring.

The structure is non-polar because it is symmetrical. The geometry

around each carbon is trigonal planar because it has three sigma

(σ) bonds and one pi (π) bond on each carbon. Benzene has a

resonance structure.

When CH3COOH (acetic acid) forms, the following structure occurs:

The structure is POLAR because it is asymmetrical. The geometry

around the first carbon (on the left) is tetrahedral because it has

four sigma (σ) bonds. The geometry around the second carbon (on

the right) is trigonal planar because it has three sigma (σ) bonds

and one pi (π) bond.

16

Bonding and Shape Quiz Review

Determine the formula for the compound formed by the two atoms and indicate if it is an ionic or

covalent compound

1. Calcium and Oxygen 2. Nitrogen and Fluorine

3. Sodium and Chlorine 4. Carbon and Oxygen

Draw the dot diagram for each of the IONIC compounds below

5. CaO 6. Na2S

7. SrF2 10. KI

Complete the table below.

Formula Electron Dot Diagram Bonding

Orbitals

Geometry Structural Formula Polar?

NCl3

CO2

H2O

C2H6

17

Naming Notes

Ionic Compounds

• To name ionic compounds:

o Name the _____________________ first.

o Name the anion second- change the ending to ________________.

o Al2O3

o BaCl2

o Ca3N2

o KF

• Ionic Compounds with transition metals:

o Metals can have multiple oxidation states, with the exception of:

� Group 1 (Alkali Metals) is always +1

� Group 2 (Alkali Earth Metals) is always +2

� Silver (Ag) is always +1

� Zinc (Zn) is always +2

� Cadmium (Cd) is always +2

� Aluminum (Al) is always +3

o Some elements, such as iron, form two or more cations with different charges. We use

__________________ to indicate the ion’s charge. For example, Fe+2

would be named ________ and Fe+3

would be named

__________ . If an element does not form more than one charge, then you do not use a

Roman numeral in its name.

� EX: Iron (III) Oxide

� PbO2

� Fe2S3

18

Practice

• Tin (II) Chloride

• Iron (III) Nitride

• Copper (I) Bromide

• PbCl2

• Co2O3

• SnS

• Circle the correct chemical formula for each compound below. Make sure the positive and negative

charges are balanced.

o calcium oxide CaO Ca2O CaO2

o magnesium fluoride MgF Mg2F MgF2

o sodium sulfide NaS Na2S NaS2

o barium nitride BaN Ba2N3 Ba3N2

o aluminum sulfide AlS Al2S3 Al3S2

Classical Names

• For transition metals, you can also use classical names (these come from the Latin name). The

ending –ous is less; -ic is more- so if you look at the possible charges on the periodic table, the

smaller charge will have the –ous ending and the larger charge will have the –ic ending

o Ferric (Fe+3

) and Ferrous (Fe+2

)

o Cupric (Cu+2

) and Cuprous (Cu+)

o Mercuric (Hg+2

) and Mercurous (Hg+)

o Stannic (Sn+4

) and Stannous (Sn+2

)

o Plumbic (Pb+4

) and Plumbous (Pb+2

)

19

Polyatomic Ion Naming

• Ions formed from a single atom are known as ions.

• You wrote formulas for ionic compounds using monoatomic ions. Many ionic compounds found in

chemistry contain polyatomic ions, which are ions made up of

.

• The following polyatomic ions should be studied and memorized:

Polyatomic Ion Name

NH4+

Ammonium

H3O+

Hydronium

NO3-

Nitrate

OH-

Hydroxide

CN-

Cyanide

C2H3O2-

Acetate

MnO4-

Permanganate

HCO3-

Bicarbonate (or Hydrogen Carbonate)

SO4-2

Sulfate

CO3-2

Carbonate

CrO4-2

Chromate

PO4-3

Phosphate

ClO3-

Chlorate

• How would you write the formula for calcium hydroxide? Is there a

difference between CaOH2 & Ca(OH)2 ? Circle the correct formula.

• When more than one polyatomic ion is present, the formula for the polyatomic ion is surrounded by

.

• Practice writing ionic formulas using polyatomic ions:

BaCO3

Zn(ClO)2

Cobalt (III) Nitrate

Silver Chlorate

Al2(SO4)3

Pb(C2H3O2)2

Ammonium Chloride

Barium Phosphate

You have memorized several polyatomic ions, but there are some you don’t know, but can figure out:

20

Use chlorate (ClO3-) as an example

• If the ion has 1 more oxygen atom than the base ion (ClO3-), it is named by a prefix per- and a suffix

–ate.

o ClO4- is perchlorate

• If the ion has 1 less oxygen atom then the base ion (ClO3-), then it is named by the suffix –ite.

o ClO2- is chlorite

• If the ion has 2 less oxygen atoms than the base ion (ClO3-), then it is named by the prefix hypo- and

a suffix –ite.

o ClO- is hypochlorite

Name the following:

1. SO32-

2. PO33-

3. SO22-

4. CO22-

5. PO53-

6. CrO32-

21

Covalent Naming

• Binary covalent compounds are characterized by having two nonmetals. Naming these compounds

involves the use of numerical prefixes:

Prefix Number Prefix Number

1 6

2 7

3 8

4 9

5 10

• If there is only ONE atom of the first element, you DON’T need a prefix. The FIRST element is named

as a normal element. The SECOND element has an –IDE ending.

o N2O4

o XeF4

o N2O5

o CO

o CBr4

o Diarsenic pentoxide

o Phosphorous pentabromide

o Carbon tetraiodide

o Trisilicon tetranitride

o Tetraphosphorous decoxide

22

Naming Acids

• If the compound begins with Hydrogen, it is an acid. If the acid does not contain a polyatomic ion,

write the prefix hydro-, then name the second element and change the ending to –ic.’

o HCl

o HBr

o H2S

Naming Acids with Polyatomic Ions

The polyatomic ions you have memorized have –ate as the ending, so you name the polyatomic ion and

change the ending to –ic.

Use sulfate (SO42-

) as the example

• H2SO4 is sulfuric acid

• If the ion has one more oxygen atom than the base (SO42-

), then the ion is named by adding the

prefix per- and the suffix –ic

o H2SO5 is persulfuric acid

• If the ion has one less oxygen atom than the base (SO42-

), then the ion is named with the suffix –ous.

o H2SO3 is sulfurous acid

• If the ion has two less oxygen atoms than the base (SO42-

), then the ion is named with the prefix

hypo- and the suffix –ous.

o H2SO2 is hyposulfurous acid

Name the following:

1. H2CO3

2. H3PO2

3. HClO4

4. H3PO3

23

Organic Naming

Remember, organic compounds are compounds that contain carbon. Organic compounds

that contain hydrogen and carbon are called hydrocarbons. We’re going to focus on three

different types of hydrocarbons: alkanes, alkenes and alkynes.

Hydrocarbons

• Alkanes contain single bonds between carbon atoms. Know the following names and

formulas:

# of C Name Formula (CnH2n+2)

1 Methane CH4

2 Ethane C2H6

3 Propane C3H8

4 Butane C4H10

5 Pentane C5H12

6 Hexane C6H14

7 Heptanes C7H16

8 Octane C8H18

9 Nonane C9H20

10 Decane C10H22

• Alkenes contain a double bond between carbon atoms. The name of alkenes is very

similar to that of alkanes, just change the ending from –ane to –ene. For example:

H2C=CH2 has 2 carbons AND a double bond, so ethane becomes ethene

o If there are more than 2 carbons (propane – decane), you must number the

atoms.

1 2 3 4 5 6 7 8 9 10

CH3CH=CHCH2CH2CH2 CH2CH2CH2CH3

o Number the atoms so the carbon containing the double bond has as small of a

number as possible (so you can go right to left). Because the double bond

starts on the second carbon, this is called 2-decene.

24

Name the following alkenes

CH3CH2CH2CH2CH=CHCH2CH3

CH3CH=CH2

CH3CH2CH=CHCH2CH3

• Alkynes contain a triple bond between carbon atoms. These are named the same

way you name alkenes, except the ending is changed to –yne.

7 6 5 4 3 2 1

CH3CH2CH2C CCH2CH3

This is called 3-heptyne

Name the following alkynes

CH3C CCH2CH3

CH3C CCH2CH2 CH2CH2CH2CH3

CH3C CCH3

25

Unit 6 Test Review

Complete the table below

Draw the dot diagram for the IONIC compound below

5.. AlN 6. HgCl2

7. AlF3 8. NaF

Write the formula for the following compounds

9. Aluminum Bromide 10. Dinitrogen Tetroxide

11. Ammonium Fluoride 12. Iron (II) Sulfate

13. Copper (I) Chloride 14. Carbon Dioxide

Name the following compounds

15. NaHCO3 16. MgO

17. H2S 18. N2O3

19. Fe2(SO4)2 20. CuCl2

Formula Electron Dot Diagram Bonding

Orbitals

Shape Structural Formula Polar?

CH2O

H2O

CH3F

C2H4

26

Naming Review

Name the following compounds:

1. HBr 2. BaSO3 3. Al2(SO4)3

4. H2CO3 5. HI 6. Ca(HSO3)2

7. Pb(NO4)2 8. Zn3(PO2)2 9. Ca3(PO4)2

10. H2SO3 11. HC2H3O2 12. K2CrO3

13. H2SO4 14. Ra(C2H3O2)2 15. NiClO

16. HNO3 17. H3P 18. H3PO4

19. HNO2 20. Fe(ClO4)3 21. KH

22. C4H10 23. SnO 24. H2SO3

25. ZnSO2 26. Al(ClO)3 27. HI

28. Hg2Cl2 29. Fe(OH)2 30. Fe2(CrO4)3

31. Ba(ClO2)2 32. Li3PO3 33. KMnO4

34. N2O4 35. Cl2S7 36. PbO2

37. Rb2CO3 38. C7H16 39. Fe(OH)3

40. NaHSO4 41. SCl2 42. Ag2S

27

Write the chemical formula for the following compounds:

1. Hydrochloric Acid 2. Copper (II) Perchlorate 3. Strontium Nitrate

4. Sodium Hypochlorite 5. Phosphoric Acid 6. Plumbic Oxide

7. Sulfurous Acid 8. Stannous Fluoride 9. Potassium Permanganate

10. Aluminum Hyponitrite 11. 1-pentene 12. Sodium Dihydrogen Phosphate

13. Mercurous Sulfide 14. Sodium Sulfite 15. Ammonium Hydrogen Phospate

16. Stannous Fluoride 17. Ferrous Nitrate 18.Calcium Perchromate

19. Hyposulfurous Acid 20. Mercuric Sulfide 21. Titanium (III) Chlorate

22. Calcium Hydroxide 23. Sodium Bicarbonate 24. Aluminum Hypophospite

25. Lithium Hyposulfite 26. Colbalt (III) Sulfite 27. Stannic Nitrite

28. Nonane 29. Lead (IV) Acetate 30. Ammonium Phosphite

31. Ferric Oxide 32. Hydrosulfuric Acid 33. Mercuric Chromate

34. Potassium Cyanide 35. Ammonium Sulfate 36. 2-hexyne

37. Barium Hydroxide 38. Mercury (II) Sulfide 39. Silver Chlorate

40. Lead (II) Sulfate 41. Potassium Permanganate 42. Silicon Dioxide