Chemistry

Key unit competence

Describe how properties of ionic compounds and metals are related to the nature of their bonding

Learning objectives

By the end of this unit, I will be able to:

•Explain why atoms bond together;

•Explain the mechanisms by which atoms of different elements attain stability;

•Explain the formation of ionic bonds using different examples;

•Represent ionic bonding by dot-and-cross diagrams;

•Describe the properties of ionic compounds based on observations;

•Perform experiments to show properties of ionic compounds;

•Assemble experimental set up appropriately and carefully;

•State the factors that influence the magnitude of lattice energy ;

•Relate the lattice structure of metals to their physical properties;

•Describe the formation of metallic bonds;

•State the physical properties of metals and forces of attraction that hold atoms of metal.

People like to bond with each other for many reason such as: to unite their forces and be stronger, to exchange idea and produce big things, to found a family, etc. we cannot live in isolation. People can have strong connection .Similarly some atoms can also have strong bonds between them. Some atoms have weak connections, just like two people can have connections.

Some atoms may not need to bond with others; they are self-sufficient as some people, a small number, may be self-sufficient.

Connections between atoms are called chemical bonds. Solids are one of the three fundamental states of matter. In molecules, atoms or ions are held together by forces called chemical bonds.There are 3 types of chemical bonds: Ionic, Covalent and Metallic bonds.

The type of a bond in molecules is determined by the nature and properties of the bonding atoms. However, in this unit we will only emphasize on ionic and metallic bonding.

3.1. Stability of atoms and why they bind together

Activity 3.1

1. In pairs discuss and write electronic configuration of sodium , neon, argon, magnesium, aluminium, oxygen and chlorine

2. What happens when oxygen and chlorine gain electrons?

3. What happens when sodium, magnesium and aluminium lose electrons?

4. Discuss on how atoms of elements can gain their stabilities by either loosing or gaining electron(s) on the valence shells and show with evidence that an atom is stable?

5. How does the formation of an ionic bond between sodium and chlorine reflect the octet rule?

Like people always relate and connect to others depending on their values, interests and goals so does unstable atoms. They are combine together to achieve stability. We know that noble gases are the most stable elements in the periodic table. The noble gases are extremely unreative: they do not tend to form compounds or combine to themselves

What do the noble gases have in common? They have a filled outer electron energy level. When an atom loses, gains, or shares electrons through bonding to achieve a filled outer electron energy level, the resulting compound is often more stable than individual separate atoms. Neutral sodium has one valence electron. When it loses this electron to chlorine, the resulting Na+ cation has an outermost electron energy level that contains eight electrons.

It is isoelectronic (same electronic configuration) with the noble gas neon. On the other hand, chlorine has an outer electron energy level that contains seven electrons. When chlorine gains sodium’s electron, it becomes an anion that is isoelectronic with the noble gas argon.

Below are examples of how magnesium bonds with oxygen and calcium with chlorine:

Example 1

Example 2

As you can see, the noble gases have the most stable electronic configuration, characterized by the presence of 8 electrons in the outermost shell: this structure is called the “Octet Structure” or “Octet Rule”; octet meaning eight in Latin.Any atom that does not have that structure, will combine with another atom (of the same element or different element) to achieve the octet structure.

Checking Up 3.1

Write the shorthand electronic configuration for copper.Predict the ions that will be formed from atoms in the table, name the isoelectronic noble gas and establish the electronic configuration (E.C) of the ion formed.

3.2. Ionic bonding

Atoms have many ways of combining together to achieve the octet structure, and one of them is the formation of an ionic bond.

In an ionic bond, electrons are transferred from one atom to another so that they form oppositely charged ions; in other words, one atom loses electron(s), the other gains electrons(s).The resulting strong force of attraction between the oppositely charged ions is what holds them together. Ionic bonding is the electrostatic attraction between positive and negative ions in an ionic crystal lattice.

3.2.1. Formation of ionic bond

Activity 3.2

Draw diagrams to illustrate the formation of ionic compounds in magnesium oxide, magnesium chloride, sodium peroxide, and sodium sulphide.

The transfer of electrons from one atom to another followed by attraction between positive and negative ions is called ionic bonding. This type of bonding occurs between metals and non-metals. The compounds formed are called ionic compounds. As stated previously, metals try to lose their outer electrons while non metals look to gain electrons to obtain a full outer shell. When metals lose their outer electrons they form positively charged ions called cations. When non-metals gain electrons they form negatively charged ions called anions. An example is shown below:

The curved arrow between sodium and fluorine represents the transfer of an electron from a sodium atom to a fluorine atom to form opposite ions. These 2 ions are strongly attracted to each other because of their opposite charges. A bond is now formed and the resulting compound is called Sodium Fluoride

Another example of ionic bonding is the bonding of Beryllium and Fluorine to form beryllium fluoride.

The transfer of electrons from beryllium results in the formation of an ionic bond. Beryllium now has a positive (+2) charge and Fluorine now has a negative charge. The resulting compound is called Beryllium Fluoride (BeF2).

Other examples showing the formation of ionic compounds using dot and cross diagrams.

a. Formation of calcium chloride

b. Formation of magnesium oxide

c. Formation of bonds in sodium fluoride

Checking Up 3.2

1. For each of the following ionic bonds: Sodium + Chlorine, Magnesium + Iodine, Sodium + Oxygen, Calcium + Chlorine and Aluminium + Chlorine

a. Write the symbols for each element.

b. Draw a Lewis dot structure for the valence shell of each element.

c. Draw an arrow (or more if needed) to show the transfer of electrons to the new element.

d. Write the resulting chemical formula.

e. Write the electron configurations for each ion that is formed. Ex. H1+ = 1s0

2. Solid sodium chloride and solid magnesium oxide are both held together by ionic (electrovalent) bonds. a. Using s,p and d notation write down the symbol for and the electronic configuration of (i) a sodium ion; (ii) a chloride ion; (iii) a magnesium ion; (iv) an oxide ion.

b. Explain what holds sodium and chloride ions together in the solid crystal

c. Sodium chloride melts at 1074 K; magnesium oxide melts at 3125 K. Both have identical structures. Why is there such a difference in their melting points?

3.2.2. Physical properties of ionic compounds

Activity 3.3(a)

Determination of Relative Melting Point of different substances

Procedure:

1. Cut a square of aluminum foil that is about5 by 5 cm

2. Set up a ring stand with an iron ring attached.

3. Place the aluminum square on the iron ring, as shown at right in Figure 3.6

4. Obtain a small pea-sized sample of NaCl. Place the sample on the aluminum foil, about 5cm from the center of the square.

5. Obtain a small pea-sized sample of table sugar. Place the sample on the aluminum foil, about 1 cm from the center of the square, but in the opposite direction from the salt.

6. Your square of aluminum foil should look like in Figure 3.7.

7. Light the Bunsen burner and adjust the flame height so that the tip of the flame is just a cm or so below the height of the aluminum foil.

8. Observe as the two compounds heat up.

9. Set up another sheet of aluminum foil and determine the relative melting points (low vs. high) of the four unknowns.

10. Record your results in the table 3.1 belowCaution: if the compounds burn with sparks do not panic.

Study question:

1. Which compound melts first?

2. Giving reasons compare the melting points of the two compounds.

Table: 3.1 relative melting points of different substances

Conclusions:

The melting points of ionic compounds are higher than those of covalent compounds; this is due to strong electrostatic forces between opposite charges in the ionic substances compared to the week forces of attraction between molecules in covalent substances . This also explains why all ionic compounds are solid at room temperature.

Activity 3.3(b)

Conductivity in Solution

Procedure:

1. Dissolve a spoonful of NaCl in water.

2. Connect the apparatus as shown in figure 3.8

3. Make an observation and record your results as in table 3.2 below

4. Repeat the procedure 1 to 3 above using sugar solution, ethanol and copper(II) sulfate solution

5. Record your results in the table below.

Table 3.2: relative conductivity of different substances

Study questions:

1. Give reasons for your observations above.

2. Solid sodium chloride does not conduct electricity whereas an aqueous solutionof sodium chloride does. Explain

Conclusion:

Based on our tests with salt and sugar, the ability to conduct electricity in solution of ionic compounds is much higher than in covalent compounds

Activity 3.3(c) Solubility test

Procedure:

1. Using forceps, place 5-8 crystals of each of sodium chloride, magnesium chloride, copper sulphate, calcium carbonate, copper carbonate, sodium sulphate (a small pinch) of the compound into one of the test tubes in test tube rack.

2. Half-fill the test tube with distilled water and stir with a clean stirring rod.

3. Observe if the crystals dissolve in water.

4. Record your findings in a suitable table.

Table 3.3: Solubility of different substances

Study questions:

1. Classify substances as soluble and slightly soluble in water.

2. Giving reasons, suggest an explanation for your observation.

Conclusion:

Water is a good solvent for many ionic compounds but not a solvent for covalent compounds, apart few exceptions (you will learn about later on).

Shattering: Why are Ionic compounds generally hard, but brittle?

It takes a large amount of mechanical force, such as striking a crystal with a hammer, to force one layer of ions to shift relative to its neighbour. However, when that happens, it brings ions of the same charge next to each other (Figure3.9). The repulsive forces between like-charged ions cause the crystal to shatter. When an ionic crystal breaks, it tends to do so along smooth planes because of the regular arrangement of the ions.

Checking Up 3.3

1. The diagrams below show the electric conductivity of distilled water, solid sodium chloride and a solution of sodium chloride respectively. Use the diagrams to explain the observations from the set up.

i) no light is given out by bulb in A

ii) no light is given out by bulb in B

iii) light is given out in C

2.Why are ionic compounds brittle?

3.Why do ionic compounds have high melting points?

4.What happens when an electric current is passed through a solution of an ionic compound?

3.2.3. Lattice energy

Activity 3.4

By using information in this student’s chemistry book and other books from the school library, attempt to answer the following questions.

1. Define lattice energy

2. Explain how the lattice energy is used to describe high melting points of ionic compounds.

3. What is the bonding force present in ionic compounds?

4. Why is the melting temperature of magnesium oxide higher than that of magnesium chloride, even though both are almost 100% ionic?

5. How is lattice energy of ionic compounds related to their high melting points?

It is a type of potential energy that may be defined in two ways. In one definition, the lattice energy is the energy required to break apart an ionic solid and convert its component ions into gaseous ions (Endothermic process). On the other hand lattice energy is the energy released when gaseous ions bind to form an ionic solid (Exothermic process). Its values are usually expressed with the units’ kJ/mol.

Lattice Energy is used to explain the stability of ionic solids. Some might expect such an ordered structure to be less stable because the entropy of the system would be low. However, the crystalline structure allows each ion to interact with multiple oppositely charge ions, which causes a highly favourable change in the enthalpy of the system. A lot of energy is released as the oppositely charged ions interact. It is this that causes ionic solids to have such high melting and boiling points. Some require such high temperatures that they decompose before they can reach a melting and/or boiling point.

Checking Up 3.4

Define lattice energy

Which one of the following has the greatest lattice energy?

MgO or NaCl

LiCl or MgCl2

Which one of the following has the greatest Lattice Energy?

NaCl or CaCl2

AlCl3 or KCl

How do the ionic radius and ion charge affect the lattice energy of an ionic substance?

Factors affecting lattice enthalpy

Activity 3.5

1. Basing on the information obtained from the definition of lattice energy, suggest the factors that affect the lattice energy.

2. Which has the larger lattice energy: NaCl or CsI?

Basing on the information obtained from the definition of lattice energy, suggest the factors that affect the lattice energy.

Which has the larger lattice energy: NaCl or CsI?

There are two main factors that affect lattice enthalpy.

a) The charges on the ions

Sodium chloride and magnesium oxide have exactly the same arrangements of ions in the crystal lattice, but the lattice enthalpies are very different.

From the above diagram the lattice enthalpy of magnesium oxide is much greater than that of sodium chloride. This is because in magnesium oxide, +2 ions are attracting -2 ions; in sodium chloride, the attraction is only between +1 and - 1ions.

b. The radius or the size (volume) of the ions

The lattice enthalpy of magnesium oxide is also increased relative to sodium chloride because magnesium ions are smaller than sodium ions, and oxide ions are smaller than chloride ions.

It means that the ions are closer together in the lattice, and that increases the strength of the attractions.

For example, as you go down Group 17 of the Periodic Table from fluorine to iodine, you would expect the lattice enthalpies of their sodium salts to fall as the negative ions get bigger - and that is the case:

As the negative ion gets bigger in size, the distance between the centres of oppositely charged ions increase and attractions so decrease. For instance as you go down group 1 halides, the enthalpy decreases.

3.3. Formation of metallic bonds and physical properties of metals

Activity 3.6

The figure above shows materials commonly used at home. If you reflect back around your house/home you will see hundreds of objects made from different kinds of materials.

1. Observe the objects (in picture) and classify them according to the materials they are made of.

2. Have you ever wondered why the manufacturers choose the material they did for each item?

3. Why are frying saucepans made of metals and dishes, cups and plates often made of glass and ceramic?

4. Could dishes be made of metal? And saucepans made of ceramic and glass

3.3.1. Formation of metallic bond

Another way of combining is the combination between metal atoms to form metallic bond.

When metal atoms combine together, there is no transfer of electrons since the combining atoms are of the same nature, i.e. all are metals and no one is ready to give up or to capture electrons.In metallic bonding, all metal atoms put together their valence electrons in a kind of pool of electrons where positive metallic cations seem to bathe. This model is called “Elecron Sea Model” (Fig. 3.15)

Metals have a sea of delocalized electrons within their structure. These electrons have become detached and the remaining atoms have a positive charge. This positive charged is attracted to the delocalized sea of electrons due to electrostatic forces of attraction (forces which result from unlike charges), and as a result has a strong interaction. It is this interaction which makes the metals so hard and rigid. Figures 3.13 and 3.14 are representation of metallic bond.

3.3.2. Physical properties of metals

Activity 3.7: Looking at metals

1. Collect a number of metal items from your home or school. Some exam-ples are listed below: hammer, electrical wiring, cooking pots, jewellery, burglar bars and coins, nails,

2. What is the function of each of these objects?

3. Discuss why you think metal was used to make each object. You should consider the properties of metals when you are answering this question.

a. Electrical conductivity

Activity 3.8

Procedure

1. Take a dry cell/battery, a torch bulb/ bulb, connecting wires, crocodile clips and connect them. As in the figure 3.17

2. Repeat the experiment above using different metals

3. Record your results in a suitable table.

Study questions:

1. Compare the relative conductivity of the metals used in the above experi-ment.

2. Suggest the purpose of the resistor in the experimental set up.

Due to the mobile valence electrons of metals, electricity can pass through the metals easily. So they are conductors of electricity. Silver and copper are the best conductors of electricity.

Note: mercury is a poor conductor of electricity.Thermal conductivity

Activity 3.9

Experiment to demonstrate the ability of different substances to conduct heat

Apparatus/apparatus: two cups (made from the same material e.g. plastic); a metal

Procedure:

1. Pour boiling water into the two cups so that they are about half full.

2. At the same time, place a metal spoon into one cup and a plastic spoon in the other.

3. Note which spoon heats up more quickly.

4. Record your observations.

Study questions:

1. Which one heats faster plastic spoon or metallic spoon and why?

2. Why do we use plastic cups?

3. Why are cooking pots made of metallic materials not plastics?

Results: The metal spoon heats up more quickly than the plastic spoon. In other words, the metal conducts heat well, but the plastic does not.

Conclusion: Metals are good thermal conductors, while plastic is a poor thermal conductor. The reason is due to the mobility of electrons with transfer of kinetic energy between electrons. This explains why cooking pots are metallic, but their handles are often plastic or wooden. The pot itself must be metal so that heat from the cooking surface can heat up the pot to cook the food inside it, but the handle is made from a poor thermal conductor so that the heat does not burn the hand of the person who is cooking.

c. Malleability and ductility

Activity 3.10

Experiments to demonstrate the malleability and ductility of metals

Materials: wires, nails, hammer, piece of cloth.

Procedure:

1. Wrap the material to be test in a heavy plastic or cloth to avoid pieces flying from the material.

2. Place the material on a flat hard surface

3. Use a harmer to pound the material flat

4. Record your observations as malleable or non-malleable.

Metals can have their shapes changed relatively easily in two different ways i.e.

Malleable: can be hammered into sheets or

Ductile: can be drawn into rods and wires

As the metal is beaten into another shape the delocalised electron cloud continues to bind the “ions” together.

The positive ions in the lattice are all identical. So the planes of ions can slide easily over one another attractive forces in the lattice are the same whichever ions are adjacent and remain the same throughout the lattice

Some metals, such as gold, can be hammered into sheets thin enough to be translucent.

Note: If a material is pounded into a flat shape it is called malleable but if it breaks or does not change it is called non-malleable.

d) Metal appears hiny/lustrous

Activity 3.11

Demonstration of shininess in metals

Procedure:

1. Hold a small piece of sodium metal using forceps.

2. Place it on a hard surface and cut it into two parts

3. Observe the cut surface. What do you observe?

4. Look at the surface of aluminium sheets, how does it appear?

Study question:

Explain what makes metal surfaces appear shiny/luster.

Light is composed of very small packages of electromagnetic energy called photons. We are able to see objects because light photons from the sun (or other light source) reflect off of the atoms within the object and some of these reflected photons reach the light sensors in our eyes and we can see the objects.

When photons of light hit the atoms within an object three things can happen:

The photons can bounce back from the atoms in the object, can pass through an object such as glass or can be stopped by the atoms within the object.Objects that reflect many photons into our eyes make the objects appear shiny. Objects that absorb photons and reflect less photons appear dull or even dark black to our eyes.

Did you know? Of all of the metals, aluminium and silver are the shiniest to our eyes. Gold is also one of the more shiny metals. However, gold is not as shiny as silver and aluminium. Mercury, a liquid metal, is also shiny and special telescope mirrors have been made of mercury.

(e)Melting and boiling points

Activity 3.12:

1. Why do metals have variable melting points?

2. Why do metals have high melting points compared to non-metals?

Melting point is a measure of how easy it is to separate individual particles. In metals it is a measure of how strong the electron cloud holds the positive ions. The ease of separation of ions depends on the electron density and Ionic / Atomic size.

Melting point increases across the period (from Na---Al), the electron cloud density increases due to the greater number of valence electrons contributed per atom. As a result the ions are held more strongly. Hence more energy will be needed to separate the ions as shown in the table 3.4.

Table 3.4: Valence electron density as a factor that influence melting and boil-ing points of metal

Conversely, as the size of the atom/ion increases down the group the melting and boiling points decrease as shown in the table 3.5

Table 3.5. Atomic/ionic size as a factor that influences melting and boiling points of metals

(*) Sevenair & Burkett(1997), Introductory Chemistry, Investigating the Molecular Nature of Matter; WCB.

N.B: You may find, in another book, different values of atomic/ionic sizes, due to different methods used for the determination the sizes.

3.3.3. Factors affecting the strength of metallic bonds.

Activity 3.13

1. Explain different strengths of metallic bonds in different metals?

2. Compare the metallic strength of the following metals:

(i) Sodium and magnesium

(ii) Sodium and potassium

The three main factors that affect the metallic bond are:

•Number of protons/ Strength of nuclear attraction: The more protons the stronger the force of attraction between the positive ions and the delocalized electrons

•Number of delocalized electrons per atom: The more delocalized electrons the stronger the force of attraction between the positive ions and the delocalized electrons

Size of atom: The smaller the atom, the stronger the force of attraction between the positive ions and the delocalized electrons and vice-versa, the larger the atom, the weaker the force of attraction between the positive ions and the delocalised electrons.

The strength increases across a period from left to right because:

The atoms have more protons. There are more delocalized electrons per atom.Electrons are added to the same energy level. Group1 elements have 1 electron in their outer shells and so contribute 1 electron to the sea of electrons, Group 2 elements contribute 2 electrons per atom, and Group 3 elements contribute 3 electrons per atom

If the atoms/ions are smaller; there is therefore a greater force of attraction between the positive ions and the delocalized electrons.

In group 1 elements, the melting and boiling points decrease as the size increases hence attraction between the delocalized electrons and metal cations decreases down the group as shown table 3.6

Table 3.6: Variation of melting and boiling points of group 1 elements

Checking Up 3.6

1. Look at the table below, which shows the thermal conductivity of a number of different materials, and then answer the questions that follow:

The higher the number in the second column, the better the material is at conducting heat (i.e. it is a good thermal conductor). Remember that a material that conducts heat efficiently will also lose heat more quickly than an insulating material. Use this information to answer the following questions:

1. Name two materials that are good thermal conductors.

2. Name two materials that are good insulators.

3. Explain why:

a. cooler boxes are often made of polystyrene

b. Homes that are made from wood need less internal heating during the cold months.

c. Igloos (homes made from snow) are so good at maintaining warm temperatures, even in freezing conditions.

d. Houses covered by iron sheets and houses covered by tiles can be compared in their capacity of keeping the interior of the house hot of fresh during a sunny and hot day.

4. Magnesium has a higher melting and boiling point than sodium. This can be explained in terms of the electronic structures, the packing, and the atomic radii of the two elements.

a. Explain why each of these three things causes the magnesium melting and boiling points to be higher. b. Explain why metals are good conductors of electricity.

c. Explain why metals are also good conductors of heat.

5. Pure metals are usually malleable and ductile.

a. Explain what those two words mean. If a metal is subjected to a small stress, it will return to its original shape when the stress is removed. However, when it is subjected to a larger stress, it may change shape permanently. Explain, with the help of simple diagrams why there is a different result depending on the size of the stress.When a piece of metal is worked by a blacksmith, it is heated to a high temperature in a furnace to make it easier to shape. After working it with a hammer, it needs to be re-heated because it becomes too difficult to work. Explain what is going on in terms of the structure of the metal. Why is brass harder than either of its component metals, copper and zinc?

3.4. End Unit Assessment

1. Choose from a list of words and fill in the missing words in the text below:

List of words:

Conduct electricity, electrodes, electrolysis, electrostatic attraction, free electrons, good conductivity, great malleability, high density, high melting points, ionic bond,metal, negative ion, non-metal, positive ion,regular crystal shape and attractive forces.

Te x t :

Metals have a layered structure of ................... in fixed positions but between them are oppositely charged ............... that can move around at random between the metal atoms. There is a strong ...................................... between these oppositely charged particles which gives them ..........................The strong forces also give a ....................... making the average ........... heavier than an average ................ The presence of ........................... in the structure keeps the bonding intact when metals are bent or hammered giving them.......................................... Also, these ....................give metals ..............as regards heat and electricity.When electrons are transferred from (usually) ............ atom (e.g. sodium) to................. atom (e.g. chlorine) an ionic bond is formed. Sodium loses an electron to form a singly charged ........................... and chlorine gains an electron to form a singly charged negative ion. In an ionic compound, the ionic bond is the electrostatic attraction between the neighbouring positive ions and negative ions.

The strong forces holding this giant ionic lattice together give these ionic compounds............................ and...................................................................When ionic compounds are melted they are found to ................ in a process called ...........using electrical contacts called............ In this process, move to the negative electrode (cathode) and metalsare released. At the same time, ....................move to the positive electrode (anode) and ................ are formed. Research from the internet or text books to find out other physical properties of metals and ionic compounds that are not mentioned above.

Answer these questions by choosing the best alternative repre-sented by letters from A, B, C and D.

1. Metals lose electrons from their lattice to become

a. positive ions

b. negative ions

c. alkalis

d. non- metals

2. Neither ions nor electrons are free to move in

a. liquids

b. metals

c. ionic solids

d. All of the above

3. Attractive forces between metal ions and delocalized electrons can be weakened or overcome by

a. hammer

b. high temperature

c. water

d. All of the above

4. Metals are good conductors due to

a. ionic lattices

b. crystalline lumps

c. mostly solids

d. delocalized electrons

5. Most atoms adopt one of three simple strategies to achieve a filled shell. Which of the following is NOT one of these strategies?

a. They accept electrons

b. They share electrons

c. They give away electrons

d. They keep their own electrons

6. Which of the following is NOT a type of chemical bond?

a. Covalent

b. Metallic

c. Valence

d. Ionic

7. In metallic bonding...

a . One atom takes the outer shell electrons from another atom.

b. A couple of atoms share their electrons with each other.

c. Some electrons are shared by all the atoms in the material.

d. Bonding takes place between positively charged areas of one atom with a negatively charged area of another atom.

8. Which of the following is NOT a characteristic of metals?

A. Shiny /lustre

B. Brittle/Shatters easily

C. Conducts electricity

D. Malleable

9. When two or more metal elements are combined they form an...

a. bronze

b. alloy

c. Covalent bond

d. Brass

10. Sulphur is a solid non-metallic element at room tempera-ture, so it is?

a. A good conductor of heat

b. A substance with a low melting point

c. Easily bent into shape

d. A good conductor of electricity

11. Copper is a metallic element so it is likely to be a?

a. substance with a low boiling point

b. poor conductor of electricity

c. good conductor of heat

d. substance with a low melting point

12. Sodium chloride is a typical ionic compound formed by combining a metal with a non-metal. Sodium chloride will?

a. have a low melting point

b. consist of small NaCl molecules

c. conduct electricity when dissolved in water

d. not conduct electricity when molten

13. Copper is a metallic element so it is likely to be a?

a. Substance with a shiny surface

b. Poor conductor of electricity

c. Poor conductor of heat

d. Substance with a low melting point

14. When an ionic bond is formed between atoms of different elements?

a. Protons are transferred

b. Electrons are transferred

c. Protons are shared

d. Electrons are shared15. Sodium chloride has a high melting point because it has:

a. Many ions strongly attracted together

b. Strong covalent double bonds

c. A giant covalent 3-dimentional structure

d. Molecules packed tightly together

16. Which substance is likely to have a giant ionic structure:

a. Melts at 1400oC, insoluble in water, good conductor of electricity either when solid or molten

b. Melts at 2800oC, insoluble in water, non-conductor of electricity when molten or solid

c. Melts at 17oC, insoluble in water, non-conductor of electricity either when solid or molten

d. Melts at 2600oC, dissolves in water, non-conductor of electricity when solid,undergoes electrolysis in aqueous solution

17. Sodium chloride conducts electricity when:

a. Solid or molten

b. Solid or in solution

c. Molten or in solution

d. Non of the above

18. The structure of magnesium oxide is a

a. Giant covalent lattice

b. Giant ionic lattice

c. Simple ionic lattice

d. All the above19. What is the formula for magnesium chloride (contains Mg2+ and Cl? ions)?a. MgClb. Mg22Clc. MgCl2d. MgCl320. Why does sodium chloride have a lower melting point than magnesium chloride?

a. Its positive ions are smaller and have a smaller charge

b. Its positive ions are larger but have a smaller charge

c. Its positive ions are smaller but have a larger charge

d. All the above

21. Explain the conductivity of sodium chloride

a. It conducts electricity when molten because it contains free electrons

b. It conducts electricity when molten because sodium has metallic bonding

c. It conducts electricity when molten because its ions are free to move.

d. None of the above

Short and long answer questions

22.(a) Explain why the lattice dissociation enthalpy of NaBr is a bit less than that of NaCl.

(b) Explain why the lattice dissociation enthalpy of MgO is about 5 times greater than that of

NaCl

23.a) The table (using figures for lattice energies from gives experimental and theoretical values for the silver halides.(The values are listed as lattice dissociation energies.) compare the values and give a detailed explanation.

(b) For AgF, the experimental and theoretical values are very close. What does that show?