Chpt. 5: Chemical Bonding – Chemical

Formulas

This topic will be investigated under five main headings:

1. Chemical Compounds

2. Ionic Bonding

3. Covalent Bonding

4. Electronegativity

5. Shape of Molecules and Intermolecular forces

1. Chemical Compounds

Table Salt Water Dry Ice

In terms of elements can you think of a more scientific name for each of the above substances????

Chemical Compounds:A compound is a substance that is made up of two or more different elements chemically combined.

2H2(G) + O2 (G) → 2H20(L)

• The atoms of the elements in a compound are held together by attractive forces called chemical bonds

• Compounds can be broken down into their elements. If an electric current is passed through water (electrolysis) the compound splits into its elements of hydrogen and oxygen

*Note:

• Valence electrons – are the outer electrons

• Bonding involves only the valence electrons

The Octet Rule:The Octet Rule states that when bonding occurs, atoms tend to reach an electron arrangement with eight electrons in the outermost shell

• 1916 - Gilbert Lewis & Irving Langmuir proposed the Octet Rule

• Elements will try to loose, gain or share electrons to achieve eight electrons in their outer shell

• This outermost energy level is also known as the valence shell

Exceptions to the Octet Rule:

Transition Metals - can have more OR less than eight electrons in their outermost energy level

Elements near Helium e.g. Hydrogen, Lithium etc. tend to achieve electronic arrangement of Helium - 2 electrons in outer shell!!!!

Loss & Gain of electrons

Group number

No. of Electrons in outer shell

Electrons needed/to be lost for octet rule

Valency

1 1 e- 1 e- (lost) +1

2 2 e- 2 e- (lost) +2

3 3 e- 3 e- (lost) +3

4 4 e- 4 e- (shared) +/-4

5 5 e- 3 e- (gained) -3

6 6 e- 2 e- (gained) -2

7 7 e- 1 e- (gained) -1

8 8 e- 0 e- 0

Valency

IonsAn ion is a charged atom or group of atoms

Metals (left side of periodic table)

- lose valence electrons - achieve a stable valence shell (usually 8 e-)- gains a positive charge, i.e. a positive ion.

Loss of electrons (formation of cations):

Na 2,8,1 -> Na+ + e- 2,8 Li 2,1 -> Li+ + e- 2 K 2,8,8,1 -> K+ + e- 2,8,8 Mg 2,8,2 -> Mg2+ + 2e- 2,8Ca 2,8,8,2 -> Ca2+ + 2e- 2,8,8Al 2,8,3 -> Al3+ + 3e- 2,8.

Non-metals (right side of periodic table):

- gain valence electrons - achieve a stable valence shell (usually 8 e-)- gains a negative charge i.e. a negative ion

Gain outer electrons (formation of anions):F 2,7 + e- -> F- 2,8

Cl 2,8,7 + e- -> Cl- 2,8,8

Br 2,8,8,7 + e- -> Br- 2,8,8,8

I 2,8,18,18,7 + e- -> I- 2,8,18,18,8

O 2,6 + 2e- -> O2- 2,8

S 2,8,6 + 2e- -> S2- 2,8,8

N 2,5 + 3e- -> N3- 2,8

H 1 + e- -> H- 2.

1. Helium is a much safer alternative to hydrogen as it is stable and is therefore used in weather balloons and blimps.

2. Argon is used in electric light bulbs to prevent the tungsten filament from evaporating or reacting.

Uses of unreactive group – Noble Gases

Lewis Symbols – ‘Dot and Cross Diagrams’

Lewis symbols show the valence electrons as dots arranged around the atomic symbol.

hydrogen:

sodium: Na

chlorine: Cl

H ·

·

·· ·

·· · ·

Sodium & Chlorine

2 Ionic BondingIonic bonding is the force of attraction between oppositely charged ions in a compound.

• Remember:Ions are elements which have a positive or negative charge e.g.

Na has 11 e- (E.C.= 2,8,1) when Na gives away this one e- it now has more protons than electrons so it has an overall positive charge.

• Ionic bonds generally form between metals and non-metals.

Formation of Ionic Compounds Using Ionic Bonding

Formation of Table Salt (Sodium Chloride):

Bohr Type Circle Diagram

Dot and Cross Diagram

11p 11n 17p18n

Step 1: Draw Bohr diagrams of each atom

Na atom2,8,1 Cl atom

2,8,7

Example 1: Formation of Sodium Chloride – Bohr-type diagram

Step 2: Determine ions formed

11n11p 17p 18n

Na+ ion2,8

Cl- ion2,8,8

AND

Step 3: The formation of an Ionic Bond

The attraction between the positive sodium ion and the negative chlorine ion results in the formation of an ionic bond to give the ionic compound Sodium Chloride:

Example 1: Formation of Sodium Chloride – Dot-and-Cross Diagrams

Na · +

Step 1: Represent atoms using dot-and-cross diagram

Step 2: Determine ions formed:

Na

Clxx

xx

xxx

+and Cl

Na atom2,8,1 Cl atom

2,8,7

xx

xx

xxx·

__

Na+ ion2,8

Cl- ion2,8,8

Step 3: Formation of Ionic Bond:

Na · + Cl xx

xx

xxx ® Na+ [ ]Cl·

xx

xx

xxx

Sodium Chloride

Example 2:Show the formation of the ionic bond in magnesium fluoride, MgF2, by means of dot-and-cross diagrams

Try:Show the formation of the ionic bond in magnesium chloride by means of dot-and-cross diagrams

Writing Formulas of Ionic Compounds

A chemical formula is a way of representing a compound using symbols for the atoms present and numbers to show how many atoms of each element are present.

You must know how to write the formulas of ionic compounds of the first 36 elements.

When writing formula remember: an ionic compound has no net charge, overall it is neutral

*Note: transition metals will be discussed separately

*Note: • A compound which contains just two elements always ends in - ide

• A compound which contains oxygen as well as two other elements always ends in -ate

Working with Simple Ions:H+, H-

Na+ ,K+

Be2+, Mg2+, Ca2+

Al3+

O2- ,S2-

F-, Cl-, Br-, I-,

Write the formula of:

i) Potassium bromide

ii) Calcium Chloride

iii) Sodium Sulphide

iv) Aluminium Oxide

Potassium Bromide

+ Br

KBr

K1-

“To get correct formula just swop the valency numbers”

*Note: •The potassium ion is K+.

• The bromide ion is Br-

• Same number of positive and negative charges

Calcium Chloride2+ Cl

Ca

Cl2

Ca 1-

*Note: •The calcium ion is Ca2+.

• The chloride ion is Cl-

• In order to have the same number of negative charges as positive must use two Cl- ions

“To get correct formula just swop the valency numbers”

Sodium Sulphide

*Note: •The sodium ion is Na+.

• The sulphide ion is S2-

• In order to have the same number of positive charges as negative must use two Na+ ions

+ S

Na2

S

Na

2-

“To get correct formula just swop the valency numbers”

Aluminium Oxide

Al O3 + 2 -

Al2

O3

*Note:• Bring all charges up to their lowest common denominator – 6• To get six positive charges we need 2 Aluminium ions• To get six negative charges we need 3 Oxide ions

Working With Complex Ions:Must learn off by heart!!!!

Name Formula

Hydroxide IonNitrate IonHydrogen carbonate Ion

OH-

NO3-

HCO3-

Carbonate IonSulfate IonSulfite Ion

CO32-

SO42-

SO32-

Ammonium IonPhosphate Ion

NH4+

PO43-

Write the formula of the following:

i) Potassium Hydroxide

ii) Calcium Hydroxide

iii) Sodium Sulphate

iv) Ammonium Phosphate

Potassium Hydroxide

K OH+ 1-

K OH

Note: •The potassium ion is K+.

• The hydroxide ion is OH-

• Same number of positive and negative charges

Calcium Hydroxide

2+ OH

Ca (OH)2

Ca 1-

“When complex ions occur more than once brackets are needed”

*Note: •The calcium ion is Ca2+.

• The hydroxide ion is OH-

• In order to have the same number of negative charges as positive we must use two OH- ions

Sodium Sulphate

Na

Na SO42

+ SO42-

*Note: •The sodium ion is Na+.

• The sulphate ion is SO42-

• In order to have the same number of positive charges as negative charges we must use two Na+ ions

Ammonium Phosphate

NH4

PO4

(NH4)

PO43

1+ 3-

*Note: •The ammonium ion is NH4

+.

• The phosphate ion is PO43-

• In order to have the same number of positive charges as negative charges we must use three NH4

+ ions

Valency of Transition Elements• Transition elements have a variable valency (they loose electrons)

• The valency depends on the other elements the transition element is bonding to:

Transition Metal Ion ExampleIron (II) +2 FeCl2

Iron (III) +3 FeCl3

Copper (I) +1 Cu2O

Copper (II) +2 CuOChromium (III) +3 CrCl3

Chromium (VI) +6 Na2Cr2O7

Manganese (IV) +4 MnO2

Manganese (II) +2 MnSO4

Manganese (VII) +7 KMnO4

•The reason for this variable valency is because there is such a small energy difference between the 4s and 3d sublevels

*Note: You do not need to remember the actual valancies of any of the transition metals but you should be able to interpret the names as shown in the following examples

Working with Transition Metals1) Write the formula of:

i) Iron (II) carbonateii) Chromium (III) Chloride

2) Name the compound Cr2(SO4)3

Iron (II) Carbonate

Fe CO32+ 2-

FeCO3

*Note: •The Roman number (II) indicates Fe2+

• The carbonate ion is CO32-

• Valencies are balanced

Chromium (III) Chloride

Cr3+

Cr Cl3

Cl1-

*Note: •The Roman number (III) indicates Cr3+.

• The chloride ion is Cl-

• In order to have the same number of negative charges as positive charges we must use three Cl- ions

Name the compound Cr2(SO4)3

• One sulphate ion has a negative charge of 2- so three sulphate ions together must have a total negative charge of 6-

• Since overall charge must be zero the two chromium ions together must carry a charge of 6+

• Therefore each chromium ion must be Cr3+

Compound Name:

Chromium (III) Sulphate

Ionic Compounds and their Crystal Lattice Structure

Ionic bonds result in a three dimensional

crystal lattice structure.

Sodium Chloride has a cubic structure

Unit Cell

Uses of Ionic Substances in Everyday Life(Know at least two)

• Too little salt in diet = muscle cramps, so many athletes take salt tablets to replace salt lost in sweating.

• Food preservative (Brine)

• Salt used in manufacture of soap, leather, detergents • Salt is spread on roads during winter to help melt frost and snow

• Fluoridation of water supplies to prevent tooth decay

Properties of Ionic CompoundsContain positive and negative ions (Na+Cl-)

Solids such as table salt (NaCl(s))

High melting and boiling points

Strong force of attraction between particles

Separate into charged particles in water to give a solution that conducts electricity.

D- block elements and transition elements

All of the elements from scandium to zinc are the first row of the d-block

All of the elements from titanium to copper incl. are referred to as transition elements

Transition elements have the following characteristics:

• In d-block of table• Variable valency Note: Sc only forms Sc3+ ions and Zn only forms Zn2+ ions• Usually form coloured compounds Note: Sc and Zn only form white compounds• Widely used as catalysts Note: Sc and Zn show little catalytic activity• Have incomplete d sublevel Note: Sc3+ has empty d-sublevel and Zn2+ has full d-sublevel

A transition metal is one that forms at least one ion with a partly filled d-sublevel

Lets take a closer look:

Even though the 4s sublevel is filled before the 3d sublevel (as it is lower in energy) electrons are lost from the 4s first as location wise it is further from the nucleus and are therefore more easily removed:

4s electrons are lost before 3d electrons

Investigate e.c. of Sc, Sc3+, Fe, Fe2+ , Ti and Ti3+

3. Covalent BondingA covalent bond is the chemical bond formed by

sharing a pair of electrons

• Covalent bonds are typical of non-metal elements.

• A single bond has 1 shared pair of electrons.

•A double bond has 2 shared pairs of electrons.

•A triple bond has 3 shared pairs of electrons.E.G. H-H O=O N N

A molecule is a group of atoms joined together. It is the smallest particle of an element or compound that

can exist independently.

Examples of single covalent bonds:

Hydrogen Molecule (H2):

H H

Both hydrogen atoms want a full outer shell, but as they both only need one electron they each share an electron. Thus, each hydrogen atom achieves a stable electron structure like the noble gas Helium.

H H

X

X

Covalent bond

Examples of single covalent bonds:

1. The chlorine molecule

2. The water molecule

3. The ammonia molecule

4. The methane molecule

1. The chlorine molecule (Cl2):

2.The Water Molecule (H2O):

*Note: There are two pairs of electrons on the oxygen atom that are not involved in bonding. These are called lone pairs.

Lone Pairs – not bonded to another atomBond Pairs – pairs of electrons that are involved

in bonding

Thus, the oxygen atom in a water molecule has two lone pairs and two bond pairs

3. The Ammonia Molecule (NH3):

4. The Methane molecule (CH4):

* Note: a single bond is formed when one pair of electrons is shared between two atoms.

Valency

• Can predict number of covalent bonds formed around certain atoms.

• Hydrogen is said to be monovalent i.e. one hydrogen atom can only combine with one atom of any other element

The valency of an element is defined as the number of atoms of hydrogen or any other monovalent element with which each atom of the element combines.

e.g. Chlorine has a valency of 1 – HCl Oxygen has a valency of 2 – H2O Nitrogen has a valency of three – NH3

Double and Triple BondsIn some molecules the atoms share two or even three pairs

of electrons

• When two pairs of electrons are shared (between two atoms), we say that a double bond is formed

Examples of double covalent bonds:

1. The oxygen molecule2. The carbon dioxide molecule

1. The oxygen molecule (O2):

2. The carbon dioxide molecule (CO2):

• When three pairs of electrons are shared (between two atoms), we say that a triple bond is formed

Example of triple covalent bonds:

1. The nitrogen molecule

1. The nitrogen molecule (N2):

*Note: the number of electron pairs is the bond order

Combining single, double and triple covalent bonds

Draw electron dot structures showing the bonding in each of the following:

a) HCHO

b) CH3OH

c) COCl2

d) HOCl

Steps: 1. Determine number of bonds

required for each atom2. Draw skeleton structure with no

bonds (note: order of atoms)3. Complete the octet rule for each

atom by assigning appropriate bonds

Sigma ( σ ) and Pi ( π ) BondingIt is also possible to describe chemical bonding in terms of

atomic orbital’sSigma Bonding (σ):A sigma bond is formed when electrons are shared in line with the nuclei – a head on overlap of orbitals

example H2 molecule

1s orbital 1s orbital Molecular orbital H2

Example Cl2 molecule:E.C. : 1s2, 2s2, 2p6, 3s2, 3px

2, 3py2, 3pz

1

Two half filled pz atomic orbitals overlap head on to form a sigma bond

*ALL single bonds can also be called sigma bonds

Sigma Bonding:

Pi Bonding (π):A pi bond is formed when the shared orbitals are side on – not in line with the nuclei

example O2 molecule - E.C. 1s2,2s2 2px2, 2py

1 2pz1

Direct overlapBetween the two2py

1 of each atomSigma bond

Sideways overlap between The two 2pZ

1 of each atomPi bond

From our study of covalent bonding we know that the oxygen molecule is formed by a double covalent bond which we now know consists of 1 sigma and 1 pi bond

example N2 molecule: - E.C. 1s2, 2s2, 2px1, 2py

1, 2pz1

- 3 half filled p orbitals- 2px orbitals overlap head on – sigma bond- 2py orbitals overlap sideways – pi bond- 2pz orbitals overlap sideways – pi bond

• A covalent single bond is a sigma bond

• A covalent double bond consist of one sigma and one pi bond

• A covalent triple bond consists of one sigma and two pi bonds

*Note: Since there is less overlapping of orbitals in forming pi bonds these bonds are not as strong as sigma bonds

Characteristics of Ionic & Covalent Compounds

Remember ionic compounds usually consist of a network of ions whereas covalent compounds generally consist of individual molecules. This fact helps is understand many of the properties of ionic and covalent compounds

Characteristics of Covalent CompoundsProperty Explanation

Contain individual molecules

Solids, liquids, or gases at room temperature

(C6H12O6(s), H2O(l), CO2(g)

Soft Consist of molecules e.g. One crystal of iodine can be made up of millions of iodine molecules (I2) with only weak forces of attraction between one molecule and another

Low melting and boiling points During melting/boiling, molecules become separated.Forces of attraction between molecules are weak and little energy is required to separate them.

Property Explanation

Do not conduct electricity No mobile charged particlesMolecules not chargedElectrons tightly bound to atoms or shared by atoms in covalent bonds.

Shapes of Covalent Molecules

Ionic compounds consist of giant crystal lattices but covalent compounds consist

of separate molecules with each individual molecule having a particular

shape

Shapes of Covalent MoleculesThe Valence Shell Electron Pair Repulsion Theory

(VSEPRT)

The shape of a molecule depends on the number of pairs of electrons in the valence shell of the central atom (i.e. that lie around the central atom of the molecule)

• Since electrons are negatively charged, the electron pairs repel each other and arrange themselves in space so as to be as far apart as possible.

• The following shapes will arise:- Linear - Pyramidal- Triangular Planar - Planar/V-Shaped- Tetrahedral

Linear

Beryllium Chloride (BeCl2) has 2 bond pairs of electrons around the central atom.

These bond pairs (-ive) repel each other as far as possible resulting in a linear shape.

The bond angle is 180o.

Triangular (Trigonal) Planar

Boron trichloride (BCl3) has 3 bond pairs of electrons around the central atom.

The three pairs of electrons repel each other as far as possible resulting in a triangular planar shaped molecule.

The bond angle is 120o.

Tetrahedral (Tetrahedron)Methane (CH4) has 4 bond pairs of electrons around the central atom.The four pairs of electrons repel each other as far as possible resulting in a tetrahedral shaped molecule.The bond angle is 109.5o.

The wedge drawn indicates that this particular C-H bond is coming out of the plane of the paperThe dashed line indicates the

C-H bond is going behind the paper

C-H bond in plane of paper

In three examples given all pairs of electrons have been bond pairs i.e. no lone pairs present!!! However, some molecules have lone pairs present which must be taken into account.The VSEPR Theory has been adapted to account for lone pairs:

*Lone pair/lone pair > lone pair/bond pair > bond pair/bone pair*

i.e. Repulsion between two lone pairs in a molecule is greater than the repulsion

between a lone pair and a bond pair etc.

Using the VSPER Theory to predict shapes of molecules

1. Determine number of electrons around central atom.2. Work out the number of electrons contributed by the

other atoms to the bonds around the central atom3. Work out the total number of electron pairs around

the central atom.4. Taking lone pairs into account decide on the shape of

the molecule.

Pyramidal (Distorted tetrahedral)Using the VSEPR Theory deduce the shape of the ammonia molecule

Step 1: Nitrogen has 5 electrons in its outer shell.Step 2: Each hydrogen atom contributes one electron so a total of 3 electrons will be contributed to the bond pairsStep 3:

N = 5 electrons 3H = 3 electrons Total = 8 electrons

= 4 pairs of e-

suggests tetrahedral shape but ammonia has three bond pairs and 1 lone pair

Step 4: However, the lone pair of electrons in ammonia repel the bond pairs more than the bond pairs repel the lone pair, resulting in a pyramid shaped molecule.

The bond angle is 107o.Remember:

LP/LP>LP/BP>BP/BP

V Shaped (Distorted tetrahedral)Using the VSEPR Theory deduce the shape of the water molecule

Step 1: Oxygen has 6 electrons in its outer shell.Step 2: Each hydrogen atom contributes one electron so a total of 2 electrons will be contributed to the bond pairsStep 3:

O = 6 electrons 2H = 2 electrons Total = 8 electrons

= 4 pairs of e-

suggests tetrahedral shape but the water molecule has 2 bond pairs and 2 lone pairs

Since LP/LP>LP/BP>BP/BP the pairs of electron repel each other to result in a V-shaped/planar molecule.

The bond angle is 104.5o

Use the VSEPR Theory to deduce the shape of the following molecules:

i) SiCl4

ii) PBr3

Shapes of Molecules Summary

Number of Bond Pairs

Number of Lone

Pairs

Shape Bond Angle

Example

2 0 Linear 1800 BeCl2

3 0 Triangular Planar

1200 BCl3

4 0 Tetrahedral 109.50 CH4

3 1 Pyramidal 1070 NH3

2 2 V-Shaped 104.50 H2O

Electronegativity

Dividing chemical compounds into two groups, ionic and covalent, simplifies the subject of bonding to a considerable extent. Many compounds have bonding which is neither fully one nor the other, their properties suggest that they are partially ionic and partially covalent.

Electronegativity is a property which can be used to predict the type of bond formed when two elements combine

• In a covalent bond between identical atoms e.g. H2, Cl2 the pair of electrons is shared equally between the two atoms in the molecule.

• In bonds between different atoms, the pair of electrons is often more attracted to one of the atoms than the other e.g. HCl.

• In HCl the two electrons in the bond are more attracted to the chlorine than to the hydrogen. As a result this gives chlorine a slightly negative charge and leaves the hydrogen atom with a slightly positive charge.

H ---- Clδ+ δ-

Electronegativity is the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond i.e. chlorine is more electronegative than hydrogen

• Linus Pauling developed a scale of relative values of electronegativity:

From Pauling’s table of electronegativity values (pg 61/log tables) it can be seen that:

- the more electronegative elements are the very reactive non-metals – fluorine, chlorine etc.

- the least electronegative elements are the very reactive metals like potassium and sodium. These are said to be electropositive

Polarity:Polarity means that there is a positive and negative charge on the molecule and these two charges are separated by some distance.

• In a polar bond, electrons are shared unequally because of the difference in electron density:

e.g. HCl is a polar covalent molecule

HCl:

• In a non polar bond (purely covalent) the bonding electrons are shared equally:

*Note: in a polar covalent substance:

- the element with the lower electronegativity is δ+ and is written first

- the element with the higher electronegativity is δ- and is written second

e.g. HCl H --- Cl

H2,Cl2:

δ+ δ-

Uses of electronegativity values:

The difference in the electronegativity values of twojoined atoms can be used to tell

1. whether the bond between them is ionic or covalent

2. if covalent how polar that bond is

1. Is bond ionic or covalent???If predicting whether a bond is ionic or covalent thefollowing rules should be applied:

E.N. Difference = 0 → non polar bond (purely covalent)E.N. Difference ≤ 0.4 → bond is slightly polarE.N. Difference > 0.4 < 1.7 → bond is polar (polar covalent)E.N Difference > 1.7 → bond is ionic

*Note: • Polar covalent bonds are intermediate between

pure covalent bonds and ionic bonds and have therefore a certain amount of ionic character• When amount of ionic character is more than 50 %

bond is ionic

Example :Determine if potassium fluoride and methane are ionic or covalent compounds

KF

0.8 4.0E.N. Difference = 4.0 – 0.8 = 3.2 i.e. >1.7 → ionic bonding therefore this is an ionic compound

CH4

2.5 2.1E.N. Difference between carbon and hydrogen atom = 2.5 – 2.1 = 0.4 i.e. < 1.7 → covalent

bonding therefore this is a covalent compound

Exceptions to the rule:

LiHNaHKHCaH2

Using rule would indicate these molecules all contain covalent bonding when in fact they are all ionic compounds. Reason for this is the presence of the hydride ion H-

2. Predicting polarity of covalent bonds

Remember

The greater the electronegativity difference the more polar the bond

Investigate the polarity of the following covalent compounds:

H2O HClCH4

H2

(Note please leave space of half a page to complete this exercise)

The first 3 examples above have polar bonds and can therefore be considered polar molecules, however, some molecules have polar bonds but are considered non polar molecules

RememberPolarity means that there is a positive and negative charge on the molecule and these two charges are separated by some distance.

Let us look at the polarity of covalent compounds when taking their shape into account!!!!!

Water:H2O - E.N. Difference 1.4 = polar bonds

therefore a polar molecule- Shape V-shaped = centre of + and – charges are in different places. The centre of + is between 2 H’s while centre of – is in the O

H H

δ2-

δ+Centre of +ve

δ+

Centre of -ve

Carbon Dioxide:

CO2 - E.N. Difference: 1.0 = polar bonds therefore you would expect polar

molecule- Shape: Linear = centre of + and – charges are both in centre of C atom (i.e not separated by a distance ) - Since the centre of gravity of the partial + and – charges coincide carbon dioxide is considered non polar

O Oδ2+

δ-δ-

Centre of +ve

Centre of -ve

By considering the shape and the polarity of the bonds in the following covalent molecules decide whether they are polar molecules or not:

HFNH3

BCl3

CCl4

Mandatory Demonstration: Polarity Test for liquids

• Bring a charged rod towards a thin stream of liquid from a burette• If polar the stream will be attracted and bend towards the rod• If polarity of rod changed- stream will still bend towards it, but opposite side of molecule will be attracted• If non polar liquid used stream will not bend towards liquid

**Note: Diagram required

Importance of polarity:• One of the most important properties of water is that it is an excellent solvent – this property depends on the fact that it is a polar molecule

• Most ionic substances and most polar covalent substances dissolve in water e.g. NaCl (ionic substance)

The +ve sodium ion in NaCl is attracted to the –ve partial charge of oxygen in water molecule. Also, the –ve chlorine ion is attracted to the partial +ve charge of hydrogen in water molecule. As a result the crystal lattice structure of NaCl breaks up i.e. salt dissolves

• Rule for solubility: ‘like dissolves like’

Polar substances like HCl and NH3 readily dissolve in water (polar)

Most non polar substances do not dissolve in water

Non polar liquids such as methyl benzene and cyclopropane dissolve non polar substances such as wax and oil

Mandatory Demonstration: Testing solubility of ionic and covalent substances in different solvents

Intramoloecular and Intermolecular Bonding

1. Intramolecular Bonding - forces/bonds within molecules- ionic and covalent bonding are examples of intramolecular bonding

2. Intermolecular Bonding- forces/bonds that exist between one

molecule and another- 3 types of forces between molecules:

- van der Waals forces- Dipole-Dipole- Hydrogen Bonding

1. Van der Waal’s: These are the weakest forces caused by the

movement of e- within a molecule. The electrons move randomly within the bond so at 1 point in time they are nearer to 1 atom than the other. This induces a temporary dipole force.

Temporary dipoles will result in increased boiling points.

The greater number of e- in a molecule the greater the number of temporary dipoles.