CHEM1011 Lecture Notes - Lecture 6: Covalent Bond, Lead(Ii) Iodide, Lead(Ii) Nitrate

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TOPIC 6.

CHEMICAL REACTIONS AND IONIC EQUATIONS.

Reactions involving ionic compounds.

As discussed earlier, ionically bonded compounds consist of large aggregations of

cations and anions which pack together in crystal lattices in such a way that the

electrostatic attractions between oppositely charged ions are maximised and repulsions

between like charged ions minimised. When an ionic crystal is placed in water, in

many cases the solid dissolves, releasing the component ions into the SOLVENT to

form a SOLUTION. Such compounds are said to be SOLUBLE and the substance

that dissolves is called the SOLUTE. A well known example of a soluble ionic

compound is table salt or sodium chloride. The process of dissolving can be best

represented by an equation which is slightly different from the formula equations used

in Topic 5. Instead, an IONIC EQUATION is used to show the ions released into the

solution as follows.

NaCl(s)  Na+ + Cl–

There is an upper limit to the amount of solute that can be dissolved in a solvent at a

given temperature. When no more solute can be dissolved, the solution is said to be

SATURATED. The maximum solubility of compounds increases with temperature

because more energy is available at higher temperatures, allowing the ions to escape

from the attractive forces in the crystal lattice.

Ionic Equations.

An ionic equation is able to show the physical states of all the reactants and products

unambiguously by including any dissolved ionic species as ions.

The same rules apply as for formula equations in that all species shown on the left

must also be present on the right hand side of the equation. In addition, notice that the

electrical charge present on both sides of the equation must also balance.

Thus the equation for another ionic solid, barium chloride, dissolving in water would

be as follows

BaCl2(s)  Ba2+ + Cl– + Cl–

which is usually written as

BaCl2(s)  Ba2+ + 2Cl–

Because the Cl– ions are separate individual species, they are represented as 2Cl– and

not as Cl2

2–, which would mean two Cl atoms bonded together and bearing a 2

negative charge, (Cl−Cl)2–.

VI - 1

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VI - 2

Why do ionic solids dissolve in water?

When ions are released into water solution, they all experience attractions to water

molecules which form spheres around them. The reason for this attractive force

between water molecules and ions is the ability of the oxygen atoms in water

molecules to attract the electrons in their O–H bonds to a greater extent than do the

hydrogen atoms. The O atom is said to be more ELECTRONEGATIVE than the H

atom. This results in a slight negative charge on the oxygen atom and a slight positive

charge on each hydrogen atom. The O–H bond is an example of a POLAR BOND

and the water molecule, being angular in shape, has a non-symmetric distribution of

charge and is a POLAR MOLECULE.

In the dissolution of ionic solids such as sodium chloride, the oxygen atoms of water

molecules are attracted to the positive charge on cations (Na+ in this example) while

its hydrogen atoms are attracted to the anions (Cl– in this example). This results in

considerable energy being released as individual ions become surrounded by the

attracted water molecules as shown below.

Consequently, ions in water solution are said to be AQUATED and sometimes the

suffix (aq) is used to emphasise this point. The large amount of energy released by the

process of aquating ions may result in the crystal lattice breaking down.

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VI - 3

Thus the two equations above might also be written as

NaCl(s)  Na+(aq) + Cl–(aq)

BaCl2(s)  Ba2+(aq) + 2Cl–(aq)

Initially the (aq) suffix will be used here, but later it will be assumed that all ions in

water solution are aquated and the (aq) suffix will be omitted. Some other examples

of ionic equations for ionic solids dissolving follow.

K2CO3(s)  2K+(aq) + CO3

2–(aq)

(NH4)3PO4(s)  3NH4

+(aq) + PO4

3–(aq)

Many text books use the (aq) symbolism (incorrectly) to indicate a solution of a

substance in water by attaching (aq) to the formula of that solid. Clearly, there can

be no such species as an aquated ionic substance because if it has dissolved, it is

totally present as ions. Therefore equations showing species such as NaCl(aq) are

most misleading and should be ignored.

Not all ionic solids will dissolve in water to a significant extent. For example, the

amount of the ionic solid silver chloride, AgCl, which will dissolve in water is so

small that it is classed as insoluble. Insoluble ionic compounds of common metals

include three chlorides, about five sulfates, most carbonates, most phosphates and

most sulfides. All nitrates are soluble in water.

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Document Summary

When an ionic crystal is placed in water, in many cases the solid dissolves, releasing the component ions into the solvent to form a solution. Such compounds are said to be soluble and the substance that dissolves is called the solute. A well known example of a soluble ionic compound is table salt or sodium chloride. The process of dissolving can be best represented by an equation which is slightly different from the formula equations used in topic 5. Instead, an ionic equation is used to show the ions released into the solution as follows. There is an upper limit to the amount of solute that can be dissolved in a solvent at a given temperature. When no more solute can be dissolved, the solution is said to be. The maximum solubility of compounds increases with temperature because more energy is available at higher temperatures, allowing the ions to escape from the attractive forces in the crystal lattice.

Related Questions

What is the main purpose of this experiment?

What are some of the major assumptions of this experiment?

What are some major safety precautions for this experiment?

Qualitative Analysis of Group I Cations

Figure 1. Snack Mix

Are There any Cheese Flavored Snacks in the Mix?

You may have seen bowls of mixed snacks such as the one in Figure 1 at parties or in offices. Which one is your favorite? Some people prefer the salty pretzels, and some like the honey nuts, whereas others prefer the cheese flavored chips.

How do you tell if the bowl contains your favorite flavor snack? You will probably start by looking at the mix and see if you can spot it. If you do, then your search is done. You may however spot a similar snack but still be unsure about the flavor. You could then try to smell it to see if it smells familiar or go straight for the ultimate test: the taste test! Suppose you do not find the snack you are looking for, but you see another one that appeals to you. You will probably follow a similar pattern to identify the flavor of the snack. Similarly, qualitative analysis is used to check for the presence of ions in a mix or to identify unknown ions.

Qualitative Analysis of Ions

Qualitative analysis is used to separate and detect cations (positively charged ions) and anions (negatively charged ions) in a sample substance. The tests that are performed are not concerned with the amount of a substance in the sample as opposed to quantitative analysis, but only with a qualitative answer to the question of whether the substance is present in the sample or not.

The positively charged cations are divided into analytical groups. The groups are defined as sets of substances that can be precipitated with the addition of a single reagent. These groups are not to be confused with periodic system groups. Each cation group can be analyzed with a set of reactions to identify the presence or absence of each member of the group.

Group I Cation Separation and Identification

Qualitative analysis is made easier by the development of a standard scheme that can be followed step by step until all of the cations have either been detected or ruled out. For each group of cations, a flowchart of the scheme is prepared, and one can simply follow the steps and the decision branches in order. The flowchart for the group I cations is shown in Figure 2 below.

Figure 2. Flowchart for Separating and Identifying Group I Cations

Group I cations are silver (Ag+), lead (II) (Pb+2) and mercury (I) (Hg2+2). They can be precipitated from a solution as chlorides in a reaction with HCl. Among these three insoluble chlorides, lead chloride dissolves in hot water. Subsequently, to test whether the lead cation is present, potassium chromate (K2CrO4) is used. Lead cations react with chromate ions to form lead chromate (PbCrO4) as a yellow precipitate. Detection of the mercury (I) cation requires the addition of ammonia solution. This causes mercury (I) chloride to precipitate as the ammonium complex HgNH2Cl, and solubilizes silver chloride as an ammonium complex, [Ag(NH3)2]Cl. Last, the addition of nitric acid (HNO3) confirms the presence of silver ion as it precipitates as AgCl.

About This Lab

In this lab, you will test the identification reactions on individual samples of all three group I cations. You will then use your observations to detect the presence of the Group I cations in two unknown samples.

Chemistry 162 Lab Activity

Introduction Water has an amazing ability to adhere (stick) toitself and to other substances. Hydrogen bonds form when hydrogenatoms covalently bonded to nitrogen (N), oxygen (O), or fluorine(F) in the form of covalent compounds such as ammonia (NH3), water(H2O) and hydrogen fluoride gas (HF). In these molecules, thehydrogen atoms do not pull as strongly on the shared electrons asthe N, O, or F atoms. Therefore, the molecules are polar; thehydrogen atoms become positively charged and are able to formhydrogen bonds to negative ions or negatively charged parts ofother molecules (such as the N, O, and F atoms that becomenegatively charged in these compounds). Hydrogen bonds are not truebonds like covalent bonds or ionic bonds. Hydrogen bonds areattractions of electrostatic force caused by the difference incharge between slightly positive hydrogen ions and other, slightlynegative ions. These attractions are much weaker than true ionic orcovalent bonds, but they are strong enough to result in someinteresting properties.

In the case of water, hydrogen bonds form between neighboringhydrogen and oxygen atoms of adjacent water molecules. Theattraction between individual water molecules creates a bond knownas a hydrogen bond (Fig. 1).
Fig. 1: Hydrogen bonds shown as the dotted lines between watermolecules.
Cohesion Molecules of pure substances are attracted to themselves.This sticking together of like substances is called cohesion.Depending on how attracted molecules of the same substance are toone another, the substance will be more or less cohesive. Hydrogenbonds cause water to be exceptionally attracted to each other.Therefore, water is very cohesive.
Adhesion Adhesion is similar to cohesion, but it involves unlike(i.e. different) substances sticking together. Water is veryadhesive; it sticks well to a variety of different substances.Water sticks to other things for the same reason it sticks toitself â because it is polar so it is attracted to substances thathave charges. In this lab activity, you will be asked to use thematerials and procedures provided to explore the properties ofwater that involve both cohesion and adhesion. Be sure to writedown your observations and results as you progress. sults as youprogress.

Fig. 1

Hydrogen bonds shown as the dotted

lines between water molecules.

Part 1:

Exploring Water

Procedure

How many drops of water can you fit on the surface of a pennywithout the water spilling over?

Predict how many drops of water will fit on a penny. Record yourprediction.

Use the plastic pipette to drop water on the surface of a penny.Drop car

efully and slowly to fit as

many drops as possible on the penny. Count each drop until thewater spills over the side of the

penny and record your results.