Learning Outcomes :

04/09/2018

•Using your data book and using your knowledge of chemistry comment on the use of titanium metal in the aerospace industry.

CfE Higher Chemistry

Unit 1: Chemical Changes and Structure

Trends in the Periodic Table and Bonding

04/09/2018

Learning Outcomes :

Arrangement of Elements in the Periodic Table

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•I can predict the properties of an element using its position in the Periodic Table

•I can identify groups and periods in the Periodic Table.

•I can explain why certain elements have similar properties

•I can identify the alkali metals, halogens, noble gases and transition elements in the Periodic Table.

Lesson Starter: Data Book Task

. Use your iPad to join the Higher Chemistry class on iTunesU.

Enrolment key for the iTunesU course is:

FKH-NPF-BXL

. Download a copy of the Higher Chemistry Data book either from the SQA website or the iTunesU resource folder.

. Using your data book and using your knowledge of chemistry comment on the use of titanium metal in the aerospace industry.

04/09/2018

The Periodic Table • On March 6th 1869 Dmitri Mendeleev, a Russian chemist,

published his Periodic Table of the Elements. He arranged the known elements in order of increasing atomic masses. This was eventually changed to atomic number.

• Elements with similar chemical properties were arranged in groups. He left gaps for elements yet to be discovered

• In the years that followed his ideas were modified and new elements were discovered until we arrived at the modern Periodic Table.

04/09/2018

Trends in the Periodic Table The elements in the periodic table have different properties

The table is set up in such a way that these properties vary periodically

across a period,

or down a group

The properties are both physical and chemical

The chemical properties of an element stem from its physical properties:

Density, melting points and boiling points, atomic size, ionisation enthalpy, attraction for bonding electrons

Notes:

04/09/2018

After discussion with your group, make sure you can identify where all the following groupings are in the Periodic Table and what their properties are. Mark these on the periodic table handout and stick it into your notebook.

Metals Non-metals Alkali metals Transition metals Halogens Noble gases The diatomic elements The radioactive elements

Notes:

04/09/2018

Metals Non-metals

Alkali metals

Transition metals

Halogens

Noble gases The diatomic

elements

Notes: There are variations in the physical properties of the elements

across a period and down a group.

(i) Density

• Copy and complete the table below for the first twenty elements, using the information in the data booklet.

• What do all the elements in Group A have in common?

• What do all the elements in Group B have in common?

• How does the density of the elements change across a period?

• How does the density of the elements change down a group?

04/09/2018

Physical properties of the elements

Variation of density (g cm-3) with atomic number

period 2 (Li - Ne) maximum at boron (B) - group3

period 3 (Na - Ar) maximum at Aluminium (Al)- group 3

Variation of density (gcm-3) with atomic number

In general in any period of the table, density first increases from group 1 to a maximum in the centre of the period, and then decreases again towards group 0

2nd 3rd

4th

5th

Variation of density (g cm-3) with atomic number Adapted from New Higher Chemistry E Allan J Harris

down a group gives an overall increase in density

The melting and boiling points of elements give an indication of the forces that hold the atoms or molecules together

The higher the melting and boiling point the stronger the forces

The trend is similar for both melting and boiling so we’ll just look at melting

Melting and Boiling points

Notes: (ii) Melting and boiling points

• Using the information in the data booklet, how do the melting

and boiling points of the elements change across a period?

• What is the trend in melting and boiling points down Group 1?

• What is the trend in melting and boiling points down Group 7 and Group 0?

04/09/2018

• The melting point starts off low, gradually increases to a peak (at group 4) then gradually decreases to a very low value (at group 0 or 8)

• To explain this trend we must think about the strength of the forces between the molecules

• In group 1 the atoms are held together by metallic bonds • In group 4 the atoms are held together by many very strong covalent bonds

(covalent network) • In group 8 the atoms are held together by very weak bonds (monatomic

gases) • We will look at the different types of bonding later in the unit

Learning Outcomes :

Bonding and Structure of the first twenty elements

04/09/2018

•I can explain how a covalent bond is formed.

•I can describe the behaviour of outer electrons in metallic bonding.

•I can explain the difference between covalent network and covalent molecular.

•I can give examples of metallic, covalent molecular, covalent network and monatomic elements.

Lesson Starter: N5 2014 PP Question

04/09/2018

Bonding in the first 20 elements

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04/09/2018

Demo 1.6

1. Metallic Bonding e.g. Li, Na, K, Be, Mg, Ca, Al

• Strong electrostatic forces exist between the positive nuclei and the outer shell electrons.

• These electrostatic attractions are known as metallic bonds.

Positive nucleus (core)

Delocalised electron

+ + + +

+ + + +

• The outer shell in metals is not full and so metal electrons can move between these partially filled outer shells.

• This creates what is sometimes called a ‘sea’ or ‘cloud’ of delocalised electrons.

Physical properties of metals 1. Metals are malleable and ductile

2. Conduction of electricity

Metal atoms can ‘slip’ past each other because the metallic bond is not fixed and it acts in all directions.

The ‘sea’ of delocalised electrons can move and carry the charge

+

applied force

M.p.’s are relatively low compared to the B.P’s. This is because in a molten metal the metallic bonding is still present. B.p.’s are much higher as you need to break the metallic bonds throughout the metal lattice.

3. Change of state

Metal b.p.’s are dependant on (i) How many electrons are in the outer shell (ii) How many electron shells there are.

2. Covalent Networks

Giant lattice of covalently bonded atoms.

e.g. B, C, Si

Giant molecules held together by covalent bonds, resulting in high mpt and bpt.

Boron: M.pt.= 2573K= 2300oC

Silicon: M.pt.= 1683K= 1410oC

Model 1.8

Two Forms of Carbon

1. Graphite (covalent network)

Weak Van der Waals forces between layers

Strong covalent bonds between atoms

• 3 electrons from each C atom used in bonding.

• 1 electron from each C is delocalised so graphite conducts electricity.

• Stacked hexagonal layers of C atoms with only weak Van der Waals between layers so layers can slip and slide- graphite is soft and can be used as a lubricant.

2. Diamond (covalent network)

• Hardest natural substance very strong covalent bonds.

• Electrical insulator (no delocalised electrons)-all outer electrons are localised in covalent bonds.

• Tetrahedrally bonded carbon atoms

We will discuss the 3rd form of

carbon later

3. Discrete Covalent Molecules • These molecules have known numbers of

atoms (discrete molecules).

e.g. Oxygen M.pt.= 55K= -218oC

Sulphur M.pt.= 386K= 113oC

• Low Mpts indicate that weak Van der Waals forces are present.

a) Diatomic molecules

H – H

O = O

N ≡ N

F – F

Cl – Cl

Remember HON 7!

All gases due to weak Van der Waals forces.

Diatomic

b) P and S

Covalent solids held together by Van der Waals which are stronger due to higher molecular masses.

c) Carbon Buckminster Fullerene

• Very large, C60.

4. Monatomic

Group 8 elements e.g. He, Ne, Ar

Non-bonded atoms.

Only weak Van der Waals forces between atoms.

Noble gases have full outer shells, they do not need to combine with other atoms.

Noble gases

b.p / oC

B.p.’s increase as the size of the atom increases

This happens because the Van der Waals forces increase

-280

-260

-220

-200

-180

-160

-140

-120

-100

Helium

Neon

Argon

Krypton

Xeon

Practice Question PP Question 2004

04/09/2018

Bond Strengths

Bond Type Strength (kJ mol –1)

Metallic 80 to 600

Ionic 100 to 500

Covalent 100 to 500

Van der Waals forces

1 to 40

Bonding and melting point

substance ionic or covalent

molecular did it melt?

copper sulphate

salol

potassium nitrate

sodium chloride

paraffin wax

04/09/2018

Expt 1.9 The properties of a compound depend on the type of bonding present. In the following experiment you will investigate melting points of ionic and covalent molecular solids. Place the test tubes provided in a beaker of boiling water for a few minutes

Find out the actual melting point of these compounds.

Ionic compounds have ……………. melting points.

• Explain this in terms of arrangement and movement of particles as well as attraction between particles.

Covalent molecular solids have …………………… melting points.

• Explain this in terms of attraction between and movement of particles.

04/09/2018

Covalent network compounds Silicon dioxide has the formula SiO2. Silicon carbide has the

formula SiC. • What type of bonding would you expect to exist in these compounds?

• Would you expect these compounds to have high or low melting points?

• Find out the actual melting points.

• How do they fit with your prediction?

• Examine models of these compounds.

• Explain melting point and boiling point in terms of bonding and movement of particles

• Consider the valances of C and Si and use this information to work out the exact structure of silicon carbide.

• Silicon carbide (SiC) is widely used as an abrasive as it is an extremely hard material. Explain this is terms of its structure.

04/09/2018

Success Criteria:

Next Lesson:

I can explain how a covalent bond is formed.

I can describe the behaviour of outer electrons in metallic bonding.

I can explain the difference between covalent network and covalent molecular.

I can give examples of metallic, covalent molecular, covalent network and monatomic elements.

Patterns in the periodic table: Covalent Radius

Learning Outcomes :

Periodic Trends in Ionisation Energy and Covalent radius

04/09/2018

•I can use covalent radius to describe the changes in the size of atoms across a period and down a group.

•I can explain the change in covalent radius in terms of changes in the number of occupied shells or the nuclear charge.

•I can state what is meant by first, second and third ionisation energies.

•I can write state equations to represent first, second and third ionisation energies.

•I can use atomic size and screening effect to explain the change in ionisation energies down a group.

•I can use atomic size and nuclear charge to explain the change in ionisation energies across a period.

Covalent Radii of Elements

The size of an atom is measured by it’s covalent radius, the

distance between the nucleus and it’s outer electrons.

Values for covalent radii can be found in

the data book

nucleus

energy

levels

covalent

radius

Looking down a group

The single electron in the outermost energy level is much further from the

nucleus in caesium.

Cs Li

-

-

This causes the caesium atom to have a much larger covalent radius.

The caesium atom also has many more electrons between the single outer

electron and the nucleus.

This screening effect counteracts the attraction from the greater nuclear

charge.

- -

-

- -

- -

- -

-

- -

-

- -

-

- -

-

-

-

-

-

- - -

-

-

- -

-

-

-

- -

- - - -

-

-

- - -

-

- - - -

-

- -

- -

- -

Looking across a period

Across a period we can see the covalent radius decreasing.

So, from lithium to fluorine:

3+

-

9+

-

-

-

- -

-

-

Lithium Atom Fluorine Atom

As we move left to right we are adding a proton to the nucleus and an

electron to the outermost energy level.

Looking across a period

The lithium atom has a

smaller nuclear charge

than neon and so a larger

covalent radius

Fluorine’s greater nuclear

charge pulls the outer energy

level in closer.

3+

-

radius = 134pm radius = 71pm

9+

-

-

-

- -

-

-

9+

-

-

-

- -

-

-

Atomic Size Summary

Decreasing Atomic Size

Across a period from left to right atomic size decreases

This is because of the atom having more electrons & protons and therefore

a greater attraction which pulls the atom closer together hence the smaller

size.

Atomic Size Summary In

cre

asin

g A

tom

ic S

ize

Down a group atomic size increases

This is because of the extra outer energy levels and the screening effect of

the outer electrons.

Decreasing Atomic Size

Ionisation Energy

The ionisation energy is the energy required to remove

one mole of electrons from one mole of atoms in the

gaseous state.

The first ionisation energy of magnesium:

Mg (g) Mg+ (g) + e- 744 kJmol-1

Values for ionisation energies can be found in

the data book

Ionisation Energy

The third ionisation enthalpy shows a massive increase because it

requires an electron to be removed from magnesium’s second

energy level.

Mg2+ (g) Mg3+ (g) + e- 7750 kJmol-1

Mg+ (g) Mg2+ (g) + e- 1460 kJmol-1

The second ionisation energy of magnesium:

Looking across a period From lithium to neon the first ionisation energy increases. Why?

Li (g) Li+ (g) + e- 526 kJmol-1

Ne (g) Ne+ (g) + e- 2090 kJmol-1

Li Be B C N O F Ne

An atom of Lithium

The lithium atom has 3 protons inside the nucleus

Li (g) Li+ (g) + e- 526 kJmol-1

3+

-

The outer electron is attracted by a relatively

small nuclear charge

An atom of Neon

The neon atom has 10 protons inside the nucleus

10+

-

-

-

-

- -

-

-

Ne (g) Ne+ (g) + e- 2090 kJmol-1

Each of neon’s eight outer electrons is attracted by a

stronger nuclear charge

Looking down a group

The first ionisation energy decreases down a group in the periodic table.

Why?

Li (g) Li+ (g) + e- 526 kJmol-1

Cs (g) Cs+ (g) + e- 382 kJmol-1

1. More Energy Levels As we saw with atomic size, the single electron in the outermost energy level is

much further from the nucleus in caesium than in lithium.

Li

-

Caesium’s attraction for its outer electron is lowered by the screening

effect caused by all its other electrons.

Cs

-

-

-

- -

- -

- -

-

- -

-

- -

-

- -

-

-

-

-

-

- - -

-

-

- -

- -

-

- - - -

-

-

- - -

-

- - - -

-

- -

- -

- -

-

-

2. Screening Effect

Ionisation Energy Summary

Increasing Ionisation Energy

Across a period from left to right ionisation energy increases

This is due to the increase in atomic charge having a greater pull on the

electrons and therefore more energy is required to remove electrons.

Ionisation Energy Summary D

ecre

asin

g I

onis

ation E

nerg

y

Down a group ionisation energy decreases

This is due to the outer electrons being further away from the nucleus and

so the attraction is weaker and they are more easily removed.

Increasing Ionisation Energy

Screening Effect

• Down a group there is a shielding (or screening) effect from the extra energy levels, and increased distance from the nucleus makes it easier to remove an electron

04/09/2018

Atomic Size The atomic size is just like it sounds, the size of an atom

The covalent radius is defined as half the distance between the centres of covalently bonded atoms

The size of an atom will depend on two things:

The number of energy levels

The nuclear charge pulling the electrons in

Atomic Size Across a period from left to right the atomic size decreases. Down a group the atomic size increases from top to bottom.

Why?

Across a period an increasing nuclear charge (+) pulls the outer electrons (-) closer in towards the nucleus

Down a group there is an increase in the number of energy levels surrounding the nucleus

Ionisation Enthalpies Ionisation enthalpy is the quantity of energy required to remove 1 mole of electrons from 1 mole of atoms in the gaseous state.

ie. The energy required to remove an electron from an atom e.g.

Mg(g)

Mg+(g) + e-

What would affect the ionisation enthalpy?

How much attraction the outer electrons feel from the nucleus

Ionisation Enthalpies Across a period ionisation enthalpy increases from left to right

Down a group the first ionisation enthalpy decreases from top to bottom

What affects how much nuclear charge an electron feels?

Across a period increase in nuclear charge and decrease of atomic size makes it more difficult to remove an electron

Down a group there is a shielding (or screening) effect from the extra energy levels, and increased distance from the nucleus makes it easier to remove an electron

The first ionisation enthalpy is large for noble gases

There is a very large increase in ionisation enthalpy after the ion has achieved the noble gas configuration

Why?

It requires considerably more energy to remove an electron from a completely full shell (very stable) which is nearer the nucleus

Ionisation Enthalpies

Ionisation Enthalpies

The 1st ionisation energy is the energy needed to remove the first mole of electrons and the 2nd ionisation energy is the energy needed to remove the second mole of electrons, etc. e.g. the ionisation energies for magnesium are:

1st Mg(g) Mg+(g) + e- ∆H= 744 kJmol-1

2nd Mg+(g) Mg2+(g) + e- ∆H= 1460 kJmol-1

3rd Mg2+(g) Mg3+(g) + e- ∆H= 7750 kJmol-1

Why the big jump from 2nd to 3rd ionisation energy for Mg?

Ionisation Enthalpies There is a large jump in ionisation energy when the electron to be removed

comes from a new shell, closer to the nucleus.

Examples

Use your Data Book to calculate the energy required for the following changes:

a) Ca(g) Ca2+(g) + 2e-

b) Al(g) Al3+(g) + 3e-

• The total energy to remove more than 1 mole of electrons is equal to the sum of each mole added together (as above).

Learning Outcomes :

Periodic Trends in Electronegativity 04/09/2018

•I can explain the effect of electronegativity on electrons in a covalent bond.

•I can describe, using a data book, the change in electronegativity down a group.

• I can use atomic size and nuclear charge to explain the change in electronegativity across a period.

•I can use atomic size and screening effect to explain the change in electronegativity down a group.

Electronegativity

Electronegativity is a measure of an

atom’s attraction for the shared pair of

electrons in a bond

e

e

C H

Which atom would have a greater

attraction for the electrons in this bond

and why?

Linus Pauling Linus Pauling, an American chemist (and winner of two Nobel

prizes!) came up with the concept of electronegativity in 1932 to

help explain the nature of chemical bonds.

Today we still measure

electronegativities of elements using the

Pauling scale.

Since fluorine is the most

electronegative element (has the

greatest attraction for the bonding

electrons) he assigned it a value

and compared all other elements to

fluorine.

Values for electronegativity can be

found in the data book

Electronegativities

Looking across a row or down a group of the

periodic table we can see a trend in values.

We can explain these trends by applying the

same reasoning used for ionisation energies.

Looking across a period Increasing Electronegativity

Across a period electronegativity increases

The charge in the nucleus increases across a period.

Greater number of protons = Greater attraction for bonding

electrons

What are the

electronegativities of

these elements?

1.0

F C B N O Li Be

1.5 2.0 2.5 3.0 3.5 4.0

Looking down a group

4.0

3.0

2.8

2.6

F

Cl

Br

I

Decre

asin

g E

lectr

onegativity

Down a group electronegativity decreases

Atoms have a bigger radius (more electron shells)

The positive charge of the nucleus is further away from the bonding

electrons and is shielded by the extra electron shells.

What are the

electronegativities of

these halogens?

Attraction for bonding electrons Different elements have different attractions for bonded electrons

The relative ability of an element to attract electrons is called its Electronegativity

Those elements that require just one or two electrons to fill an energy level can attract electrons more easily

Electronegativity values increase across a period from LtoR

Electronegativity values decrease down a group

Therefore the most electronegative element is Fluorine

The smaller the atom the easier it is to capture an electron since they will feel a greater “pull” from the nucleus

Learning Outcomes :

Bonding Continuum 04/09/2018

•I can explain the relationship between differences in electronegativity and type of bonding.

•I can use data from properties such as conductivity, melting point and boiling point to deduce type of bonding and structure.

•I can list exceptions to the statements: “Compounds formed between non-metals only are covalent. Compounds formed between a metal and a non-metal are ionic”.

Electronegativity & Bonding

The difference in the ability of elements to attract electrons tells us about the type of bonding we can expect between them

No difference = pure covalent bonding

Small difference (<1.5) = polar covalent bonding

Large difference (>1.5) = ionic bonding

Revision - Bonding in elements

• There are two main bonding types found in elements: metallic and covalent.

Revision - Metallic Bonding

• Metallic bonding consists of the atoms losing their outer electrons to a common ‘pool’ of delocalised electrons.

• The atoms become positively-charged ions. • The charged metal ions are now attracted to the pool of electrons. • The attraction of opposite charges is called “electrostatic attraction” • The electrons are free to move so, metals conduct

electricity.

Revision - Covalent Bonding

• In Covalent Bonding Covalent bonding there is also electrostatic, but this time the atoms are held together by the attraction between their positive nuclei and negatively-charged shared pairs of electrons.

Ionic Bonds • Ionic bonds are formed between atoms with a large

difference in electronegativities.

• (A table of electronegativity values can be found in the data book.)

• They are often (though not always) between metals and non-metals.

Bonding in Compounds

• Compounds can show three different types of bonding between the atoms in the compound

• Ionic

• Pure Covalent

• Polar Covalent

• The type of bonding depends on the electronegativity values of the atoms in the compound

Ionic Bonds

• For example, in potassium bromide, the difference in electronegativities is so large that potassium will lose an electron and form a positive ion.

Reacting Elements:

Electron Arrangement:

During Reaction:

New Electron

Arrangement:

Ions Formed:

K Br

2,8,8,1 2,8,7

loses 1e- gains 1e-

2,8,8 2,8,8

transfer of

an electron

e-

K Br - +

• The electrostatic force of attraction between the oppositely charged ions is called the ionic bond

• Ionic compounds form a LATTICE STRUCTURE. • Millions of oppositely charged ions are held together in a

very stable arrangement.

Br - K +

Pure Covalent Bonding

e

e

H H

• This gives rise to a Pure Covalent Bond

• A pure covalent bond has no ionic character at all.

• If the electronegativities of both atoms are identical, the bonding electrons are evenly shared between both atoms.

2.2 2.2

Electronegativities

Polar Covalent Bonding • If there is a small difference between the

electronegativities of both atoms, the bonding electrons are pulled more closely to the more electronegative atom.

• The atom with the greater share of electrons will end up with a slight negative charge by comparison with the other atom.

• The symbols δ+ and δ– mean ‘slightly positive’ and ‘slightly negative’.

Polar Covalent Bonding

• This produces a Polar Covalent Bond.

• A polar covalent bond has some ionic character.

e

e

P Cl

δ- δ+

2.2 3.0 Electronegativities

Ionic Bond

• An ionic bond exists when the difference in electronegativities is so great that the movement of the bonding electrons between the two atoms is complete.

Li F

1.0 4.0

Electronegativities

+ - Li +

F -

• There is no sharing of the electrons and oppositely charged ions are formed.

e

e

Bonding Continuum

• We can place each of the types of intramolecular bond (bonds between different atoms) in one continuous series based on how much ionic character the bond displays.

Increasing ionic character

• This depends on the difference in electronegativities between the two atoms.

• This is called a Bonding Continuum

Bonding Continuum - Pure Covalent Bonding

e

H H e

Increasing ionic character

• The electronegativities of both atoms are identical. • A pure covalent bond has no ionic character at all

Bonding Continuum - Polar Covalent Bonding

e

H H e

Increasing ionic character

• There is a small difference between the electronegativities. • There is some ionic character

e

P Cl e

δ- δ+

Bonding Continuum – Ionic Bonding

e

H H e

Increasing ionic character

• There is a large difference between the electronegativities. • True ions are formed

e

P Cl e

δ- δ+ Li F

+ -

Bonding Continuum

e

H H e

e

P Cl e

δ- δ+

Increasing ionic character

Li F + -

Pure Covalent

Bond

Polar Covalent

Bond

Ionic Bond

To judge the type of bonding in any particular

compound it is more important to look at the

properties it exhibits rather than simply the names of

the elements involved.

Ionic or covalent bonding?

• In general, metal compounds tend to be ionic and non-metals bonded to non-metals tend to be covalent, but this is not always the case. • Tin(IV) iodide is an example of a metal compound

which has polar covalent, rather than ionic, bonds.

• In order to decide whether a compound is ionic, polar covalent or pure covalent, we must look at the properties of the compound.

Ionic bonding

• Ionic compounds do not conduct electricity when solid but do conduct when the ions are free to move e.g. when molten or in solution

• Ionic compounds are hard solids at room temperature due to the strong electrostatic bonding between the oppositely charged ions

• Ionic compounds are soluble in ionic solutions

Covalent bonding

• Covalent compounds do not conduct electricity as solids, melts or solutions

• The melting point of covalent compounds varies enormously as covalent compounds can exist as huge network structures with very high melting points( over 1,000oC) or as covalent molecules with melting points of less than 100oC • Covalent compounds are soluble in covalent liquids

(like dissolves like).

Trends in Periodic Table Summary

• An increase in the nuclear charge causes the outer electrons to be more strongly attracted to the nucleus.

• An increase in the number of electron shells causes the outer electrons to be screened from the nucleus.

04/09/2018

Practice Questions

1. Which of the following elements has the smallest electronegativity?

A Lithium

B Caesium

C Fluorine

D Iodine

Practice Questions

2. Which of the following atoms has the least attraction for bonding electrons?

A Carbon

B Nitrogen

C Phosphorus

D Silicon

Practice Questions

3. Which of the following chlorides is likely to have the least ionic character?

A LiCl

B CsCl

C BeCl2

D CaCl2

Practice Questions

4. Which of the following compounds has the greatest ionic character?

A Caesium fluoride

B Caesium iodide

C Sodium fluoride

D Sodium iodide

Practice Questions

5. Decide whether each of the following covalent molecules is polar or non-polar.

a) Oxygen

b) Hydrogen chloride

c) Carbon monoxide

d) Hydrogen

Practice Questions

6. Using hydrogen fluoride as an example, explain how a polar covalent bond arises.

Practice Questions

7. Which bond is more polar H-F or H-Cl ?

Explain your answer.

Practice Questions

8. Explain why ammonia, NH3, has polar covalent bonds yet Nitrogen, N2, does not.

Practice Questions

9. Which of the following chlorides is most likely to be soluble in tetrachloromethane (CCl4)?

A Barium chloride

B Caesium chloride

C Calcium chloride

D Phosphorus chloride

Practice Questions

10. Predict the type of bonding that would be present in the following substances:

a) melts at 6 oC and boils at 80 oC which does not conduct electricity

b) melts at 1044 oC and which conducts when molten but not as a solid

c) melts at 962 oC and conducts electricity when a solid

d) melts at 2,300 oC and does not conduct electricity