Polyatomic ion

A polyatomic ion, is a charged ion composed of two or more atoms covalently bonded or of a metal complex that can be considered as acting as a single unit in the context of acid and base chemistry or in the formation of salts. The prefix "poly-" means "many," in Greek, but even ions of two atoms are commonly referred to as polyatomic.

For example, a hydroxide ion is made of one oxygen atom and one hydrogen atom: its chemical formula is (OH)⁻. It has a charge of −1. An ammonium ion is made up of one nitrogen atom and four hydrogen atoms: its chemical formula is (NH₄)⁺. It has charge of +1.

Nomenclature

There are two "rules" that can be used for the learning the nomenclature of polyatomic ions. First, when the prefix bi- is added to a name, hydrogen is added to the ion's formula and its charge is increased by 1, the latter being a consequence of the hydrogen ion carrying a +1 charge. An alternate to the bi- prefix is to use the word hydrogen in its place: the anion derived from H⁺ + CO₃²⁻, HCO₃⁻ can be called either bicarbonate or hydrogen carbonate.

The second rule looks at the number of oxygens in an ion. Consider the chlorine oxoanion family:

oxidation state	−1	+1	+3	+5	+7
anion name	chloride	hypochlorite	chlorite	chlorate	perchlorate
formula	Cl−	ClO−	ClO2−	ClO3−	ClO4−

First, think of the -ate ion as being the "base" name, in which case the addition of a per- prefix adds an oxygen. Changing the -ate suffix to -ite will reduce the oxygens by one, and keeping the suffix -ite and adding the prefix hypo- reduces the number of oxygens by two. In all situations, the charge is not affected. The naming pattern follows within many different oxyanion series based on a standard root for that particular series. The -ite has one less oxygen than the -ate, but different -ate anions might have different numbers of oxygen atoms.

These rules will not work with all polyatomic ions, but they do work with the most common ones (sulfate, phosphate, nitrate, chlorate).

Examples of common polyatomic ions

The following tables give examples of commonly-encountered polyatomic ions. Only a few representatives are given, as the number of polyatomic ions encountered in practice is very large.

Anions
Acetate (ethanoate)	CH3COO− or C2H3O−2
Benzoate	C6H5COO− or C7H5O−2
Bicarbonate (hydrogen carbonate)	HCO−3
Carbonate	CO2−3
Cyanide	CN−
Hydroxide	OH−
Nitrate	NO−3
Phosphate	PO3−4
Sulfate	SO2−4

Cations
Ammonium	NH+4
Hydronium	H3O+
Mercury(I)	Hg2+2

Research Topics

Here is a list of compounds; choose one compound and research its chemical formula, molecular structure, bonding type, everyday uses and impact on our lives. Also, what makes it important, and is it safe or dangerous; harmful, such as a pollutant; or beneficial, such as a requirement for maintaining our health. Consider the following when choosing a compound (Remember choose your groups with three members in a group):

• Chemical formula
• Molecular Structure
• Type of bonding
• Physical properties of the compound
• Chemical properties of the compound
• Uses of the compound
• Harms of the compound

The list of the compounds:

1.     Potassium Chlorate

2.     Ascorbic Acid

3.     Borax (Sodium Tetraborate Decahydrate)

4.     Boric Acid

5.     Butane

6.     Calcium Carbonate

7.     Calcium Chloride

8.     Calcium Hydroxide

9.     Calcium Oxide

10.  Calcium Sulfate

11.  Carbon Dioxide

12.  Copper(II) Sulfate and Copper Sulfate Pentahydrate

13.  Magnesium Sulfate

14.  Naphthalene

15.  Propane

16.  Silicon Dioxide

17.  Potassium Chloride

18.  Sodium Acetate

19.  Sodium Bicarbonate

20.  Sodium Hydroxide

21.  Sucrose (Saccharose)

22.  Sulfuric Acid

                                       Chemical Bonding: Covalent Bonding

·      Covalent bonding occurs when two (or more) elements share electrons.

·      Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons).

·      This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons.

·      Covalent Compounds:

1. Gases, liquids, or solids (made of molecules)
2. Low melting and boiling points
3. Poor electrical conductors in all phases
4. Many soluble in nonpolar liquids but not in water.

·      A good example of a covalent bond is that which occurs between two hydrogen atoms. Atoms of hydrogen (H) have one valence electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to pick up a second electron. Hydrogen atoms will react with nearby hydrogen (H) atoms to form the compound H₂. Because the hydrogen compound is a combination of equally matched atoms, the atoms will share each other's single electron, forming one covalent bond.

·      Unlike ionic compounds, covalent molecules exist as true molecules. Because electrons are shared in covalent molecules, no full ionic charges are formed.  Thus covalent molecules are not strongly attracted to one another.  As a result, covalent molecules move about freely and tend to exist as liquids or gases at room temperature.  

·      For every pair of electrons shared between two atoms, a single covalent bond is formed. When two oxygen atoms form the compound O₂, they share two pairs of electrons, (4 electrons) forming two covalent bonds.   

·      Lewis structures can also be used to show bonding between atoms. Each dash represents one pair ( 2 electrons) of electrons, or one bond

·         Elements which are close together in electron negativity tend to form covalent bonds and can exist as stable free molecules. Carbon dioxide is a common example.

·          Polar and non-polar covalent bonding: two subtypes of covalent bonds.

·      Non-polar Covalent bonding:

1.      The H₂ molecule is a good example of the first type of covalent bond, the non-polar bond; because both atoms in the H₂ molecule have an equal attraction (or affinity) for electrons.

2.       The bonding electrons are equally shared by the two atoms, and a non-polar covalent bond is formed.

3.      Whenever two atoms of the same element bond together, a non-polar bond is formed.

·     Polar Covalent bonding:

1.      A polar bond is formed when electrons are unequally shared between two atoms.

2.      Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons.

3.       A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.

                                                    Chemical Bonding: Ionic Bonding

• Compounds are formed when two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.

·         Ionic bonding: In ionic bonding, (It is called an ionic bond because the atoms become ions, a charged atom that has either lost an electron or has extra electrons) electrons are completely transferred from one atom to another. In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions. The oppositely charged ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond.

For example, during the reaction of sodium with chlorine: Notice that when sodium loses its one valence electron it gets smaller in size, while chlorine grows larger when it gains an additional valence electron. This is typical of the relative sizes of ions to atoms. Positive ions tend to be smaller than their parent atoms while negative ions tend to be larger than their parent.

sodium (on the left) loses its one valence electron to chlorine (on the right),

resulting in

A positively charged sodium ion (left) and a negatively charged chlorine ion (right).

Ionic compounds share many features in common:

• Ionic bonds form between metals and nonmetals.
• In naming simple ionic compounds, the metal is always first, the nonmetal second (e.g., sodium chloride).
• Ionic compounds dissolve easily in water and other polar solvents.
• In solution, ionic compounds easily conduct electricity.
• Ionic compounds tend to form crystalline solids with high melting temperatures. The fact that ionic compounds are solids, results from the intermolecular forces (forces between molecules) in ionic solids.  

Each sodium ion is attracted equally to all of its neighboring chlorine ions, and likewise for the chlorine to sodium attraction. The concept of a single molecule does not apply to ionic crystals because the solid exists as one continuous system. Ionic solids form crystals with high melting points because of the strong forces between neighboring ions.

	Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1 Cl-1 Na+1
Sodium Chloride Crystal	NaCl Crystal Schematic

·         Elements from opposite ends of the periodic table will generally form ionic bonds. They will have large differences in electron negativity and will usually form positive and negative ions. The elements with the largest electron negativities are in the upper right of the periodic table, and the elements with the smallest electron negativities are on the bottom left. If these extremes are combined, such as in RbF, the dissociation energy is large. As can be seen from the illustration below, hydrogen is the exception to that rule, forming covalent bonds.

                                         Model Answers for Worksheets Grade 9

CHAPTER 2: SECTION 1, 2 AND 3

Concept Reviews

SECTION: CLASSIFYING MATTER

1.     a. E               d. E

        b. O               e. O

        c. E                f. O

2. An atom is the smallest unit of an element that maintains the properties of that element, while a molecule is the smallest particle of a substance that has the chemical properties of that substance. Molecules consist of one atom or two or more atoms bonded together.

3. A pure substance is made up of matter that has a fixed composition and definite properties. Although a homogeneous mixture is uniformly mixed, it is a combination of more than one pure substance and does not necessarily have a fixed composition.

4.  a. C               d. E

     b. E                e. C

     c. C                f. E

5. Elements are pure substances because each has a fixed composition of protons, neutrons, and electrons and particular characteristic properties. Compounds are pure substances because each has a fixed composition of atoms and definite properties.

6.   a. M              d. S

      b. S               e. M

      c. S                f. M

7. Copper and carbon are classified as elements because they are made of only one kind of atom and cannot be broken down by chemical means.

SECTION: PROPERTIES OF MATTER

1.     a. physical

        b. chemical

        c. physical

        d. physical

        e. chemical

2.     a. physical

         b. chemical

         c. physical

         d. physical

         e. physical

3. Aluminum is light and durable (physical properties) and nonreactive (chemical property).

4. 136.3 g silver

m = DV

= (10.49 g/cm3)(12.99 cm3) = 136.3 g

5. 2.3 g/cm3

D = m/V

= (820 g)/(350 cm3) = 2.3 g/cm3

Chapter Assessment

TEST A

1. b                  11. c

2. b                  12. b

3. d                  13. properties

4. c                   14. oxygen

5. c                   15. Cl2

6. c                   16. miscible

7. d                   17. reactive

8. b                   18. less dense

9. d                19. physical

10. b

20. Density affects how heavy something is, as well as whether something floats in a particular liquid. Boats float because the air inside the vessel is less dense than the water the vessel sits in. Cream can be separated from milk because it is less dense than milk and floats to the top.

CHAPTER6: SECTION 1 AND 2

SECTION 1 COMPOUNDS AND MOLECULES

Review

1. Different models show different characteristics of the compound’s structure. One model cannot show every characteristic.

2. The ball-and-stick model shows bond angles and bond lengths. The ball-and-stick model uses balls to represent atoms and sticks to represent bonds.

3.

4. The particles in liquids are more strongly attracted to one another than particles in gases. Thus, molecules in rubbing alcohol are more strongly attracted to one another than molecules in methane gas.

5. The attractions between particles are relatively weak.

Concept Reviews

SECTION: COMPOUNDS AND

MOLECULES

1. Compounds are held together by chemical bonds.

2. A ball-and-stick model gives you a better idea of bond lengths and bond angles. A space-filling model gives you a better idea of the space occupied by atoms.

3. The model would be more accurate because bonds can bend, stretch, and

rotate without breaking.

4. Substances with network structures have strong bonds holding the atoms or ions together. Much energy (a higher temperature) is needed to break these

bonds.

5. Table salt has a strong network structure consisting of very tightly bonded sodium cations and chloride anions. Table sugar is made of individual molecules. The bonds within each molecule are strong, but there are no bonds (just slight attractions) between molecules.

6. Because the boiling point of the compound is relatively low, the compound is likely to be in the form of individual molecules.

SECTION 2 IONIC AND COVALENT BONDING

Review

1. No, Na and K are both metals. They would not form ions with opposite charges. Thus, they would not form an ionic compound.

2. Atoms of the same element that are bonded together share electrons equally. Atoms of different elements that are bonded together do not share electrons equally.

3. Top row: 4; double; polar bottom row: 2 per bond; single; polar

4. The bond between O and H is covalent. The bonds between the two OH^- ions and Ca2⁺ are ionic.

5. Aluminum foil and dissolved KOH; aluminum is a metal, and metals can conduct electric current as solids. KOH is an ionic compound. Ionic compounds can conduct electric current when they are dissolved.

Concept Review

SECTION: IONIC AND COVALENT

BONDING

1. Atoms join to form bonds so that each atom has a more stable electron

configuration.

2. An ionic bond is formed by the attraction between oppositely charged ions. A covalent bond is formed when atoms share one or more pairs of electrons.

3. A triple bond is stronger than a double bond because more energy is required to break a triple bond with 3 pairs of shared electrons as opposed to 2 pairs of shared electrons in a double bond.

4. A compound that contains one or more polyatomic ions has both ionic and covalent bonds. The atoms making up the polyatomic ion are covalently bonded. The polyatomic ion forms an ionic bond with an oppositely charged ion.

5. Gold has metallic bonds in which the electron clouds of the gold atoms overlap. This overlap allows electrons to be transferred from atom to atom easily; therefore, gold is a good conductor of electricity.