BELLRINGER

Explain in complete sentences the

difference between various

Structures of chemical bonds

Form an ionic bond between the following

• Na and Cl• Cl and F• Na and Mg• Mg and Br• NH4

+ OH-

• NH4+ SO4

2-

Covalent BondingCovalent Bonding

The Octet Rule • The octet rule says that atoms tend to

gain, lose or share electrons so they have eight electrons in their outer shell.

Covalent Bonding Bonding• Instead of gaining or losing electrons atoms

can get stable by sharing electrons

• This is always between two non-metals.

• Two fluorine atoms, for example, can form a stable F2 molecule in which each atom has 8 valence electrons by sharing a pair of electrons.

• In covalent bonds they can share

more than two electrons

Single, Double and Triple Bonds

• With Covalent bonds the elements can share two or more electrons

• A Single Bond is when 2 electrons are shared they are represented by a single line in bond diagrams

• A Double bond is when 4 electrons are shared they are represented by two lines in bond diagrams

• A Triple bond is when 6 electrons are shared they are represented by three lines in bond diagrams

• Oxygen has six valence electrons and is seeking an octet. Therefore it combines with two hydrogen atoms, each with one valence electron.

Representing covalent molecules• Chemists use two types of diagrams to

represent molecules:

1. Electron dot formula (Lewis structure) – each outer shell (valence) electron is represented by a dot.

2. Valence structure – lines are used to represent electrons: one line represents one pair of electrons.

Lewis Dot StructureLewis Dot Structure• In order to draw covalent bonds you must be

able to write Lewis Dot structure• This shows the number of valence electrons

around the symbol • Put one electron on each side before doubling

up. • We can see that helium does not follow this

rule exactly.

                       

• First, we must determine whether CO2 is a covalent or ionic bond.

• Because nonmetals form covalent bonds, CO2 will be covalent.

• Ionic bonds form when combining both metals and a nonmetals.

• Begin by drawing the Lewis dot structure of each individual atom.

• How many electrons does the carbon want?• 4• How many electrons does each oxygen want?• 2

• Determine which is the central atom

• All three of these atoms want an octet.

• Carbon needs to “share” four electrons.

• Carbon will share two of its electrons with the oxygen on the right and two with the oxygen on the right.

• The carbon dioxide Lewis Dot Structure:The carbon dioxide Lewis Dot Structure:

:C::C:

:C:• Take note of the boxed in electrons below.

• Now each atom has eight electrons.

What type of bonds does CO2 have?

Double bonds

Draw and complete the following tableName/

Formulae

No. Valence electrons

E dot diagram

Bond diagram

Type of bond

HCl H = 1

Cl = 7

Methane CH4

C = 4

H = 1

Oxygen (02)

O =

Include the following compounds CS2, NH3 , CF4.

Naming Covalent formulas1. First name generally stays the same.

2. Second non metal name changes to –ide.

3. If there is one or more of the second element it will have prefix in front

e.g. Mono (means 1) or di (2) or tri (3) tetra (4) penta (5) e.g. CO2 Carbon Dioxide

3. If there are two or more of the first element it must have a prefix

e.g. di (2) or tri (3) tetra (4) penta (5) Dihydrogen sulfide (H2S)

Naming Covalent compounds

• What is dihydrogen monoxide?

• H2O

• Name CO• Carbon Monoxide

• Name P2S3

• Diphosphourous trisulfide

Valence electrons• In covalent bonds when electrons are shared the

electrons spend most of their times between the two nuclei of the atoms. The electrons are called localised electrons.

• The shared electrons are called the bonding electrons

• The valence electrons not involved in bonding are called the non-bonding electrons.

• Pairs of non-bonding electrons

are called lone pairs.

Lone pair

Pair of bonding electrons

Your turn

• How many bonding electrons does Cl2 have?

• How many non-bonding electrons does Cl2 have

• How many lone pairs does Cl2 Have?

Covalency• Chlorine needs to share one electron to

give it a stable outer shell. Therefore its has a covalency of one.

• The covalency of an atom is generally the number of electrons it shares when bonding with another non metal atom.

Draw and Complete the following tableName/

Formulae

No. bonding electrons

No. of non bonding electrons

No. of lone pairs

Covalency

Fluorine (F2)

Methane CH4

C =

H =

Oxygen (02)

Include the following compounds HCl, CS2, NH3 , CF4.

Charge Cloud Model

• This Model of the atom pictures the electrons moving around the nucleus in a region called an electron cloud.

• The electron cloud is a cloud of varying density surrounding the nucleus. The varying density shows where an electron is more or less likely to be.

• an 'electron cloud' may be a single, double or triple bond, or a lone pair of electrons

What is the Electron Cloud Model?Diagram 1:

What is the Electron Cloud Model?Diagram 2:

Shapes of molecules• The shape of a molecule describes the way

in which the atoms are arranged

• Molecules can be linear, angular v-shaped or bent, triangular pyramid and tetrahedral.

• There are more shapes but you only need to the know the above

http://www.youtube.com/watch?v=FhVkCH9COZo

http://www.youtube.com/watch?v=q6g7BUruUK0&feature=related

http://www.youtube.com/watch?v=ra0A8Ni1DAA&feature=related

Shape of Molecule

Name of Shape

No. of Lone pairs

Bond angle

Examples

Linear 0 180 CO2, HCN

Bent/

Angular

2 <109.5 H2O, SCl2

Triangular pyramid

1 <109.5 NH3, PCl3

Tetrahedral 0 109.5 CH4, SiF4

Activity

Trigonal Planar 0 120° sulfur trioxide

The molecule is all in the same plane

and is two dimensional.

• Use Balloons and string to make the shape of a molecule you have been assigned

• Use the molecular models to make all the other shape molecules

• x

VValencealence

SShellhell

EElectronlectron

PPairair

RRepulsionepulsion

TheoryTheory

Trigonal Planar

Tetrahedral

Trigonal bipyramidal

Octahedral

The valence shell electron pair repulsion (VSPER) model

• VSPER is used for predicting the shapes of individual molecules, based upon the extent of electron-pair repulsion (to push away).

• a lone pair of electrons is a non-bonding pair of electrons

• 'electron clouds' are negatively charged since the electrons are negatively charged, so electron clouds repel (push away) from one another and try to get as far away from each other as possible.

The valence shell electron pair repulsion (VSPER) model

RULES 1. Lone electron pairs in the outer shell of an atom in a

molecule will repel each other the most due to their negative charge

2. lone pairs of electrons exert a greater repelling effect than bonding pairs do

3. lone pair-bonding pair repulsion is greater than bonding pair-bonding pair repulsion

lone pair-lone pair repulsion > lone pair-bonding pair repulsion > bonding pair-bonding pair repulsion

How to determine shapes of molecules1. Decide which atom is the central atom

2. Determine the number of valence electrons for each atom

3. Arrange the electrons to include bonding

4. Draw the bonding electrons and lone pairs so they are as far apart as possible

5.Count the total number of electron pairs around the central atom, including both the bonding pairs and lone pairs. Double and triple bonds are counted as on bonding pair.

6. Use the VSEPR geometry table to predict the geometry of the molecule (it can only be one of 5 shapes!!).

How to determine shapes of molecules

• Sometimes the molecules are represented by AXY, A is the central atom, X is the 2nd atom and Y is the # of peripheral atoms

• AX2 = linear or bent• AX3 = trigonal planar or trigonal pyrimadial• AX4 = tetrahedral (tetra = 4 faces)• AX5 = trigonal bipyramidal (2 pyramids)• AX6 = octahedral (octa = 8 faces)

Which shapes do you recognise?

• x

Your turn

• Use Lewis structures, VSPER and the molecular models to help you determine the shape of the following molecules.

• Carbon tetrachloride

• Silicon disulfide

• Dicarbon dihydride

• Boron trihydride

• Phosphorous trifluoride

What is Electronegativty? • Electronegativity: is a measure of an

elements ability to attract electrons

• The electronegativity of elements tends to increase from left to right across a period.

• Electronegativities decrease from top to bottom within a group.

• Fluorine is the most electronegative atom

Electronegativity Values• The electronegativity values can be found in

the periodic table• The higher the value the higher the

electrotnegativity• The Pauling scale is used to measure

electronegativity. It is a relative scale running from 0.7 to 4.0 (hydrogen = 2.2).

• The units for electronegativity are Pauling units.

• x

There are two types of covalent bonding

1. Non-polar bonding with an equal sharing of electrons.

2. Polar bonding with an unequal sharing of electrons. The number of shared electrons depends on the number of electrons needed to complete the octet.

1. Non Polar Bonds• What does polarised/polarity mean?

• Opposites

• Non polar bonding results when two identical non-metals equally share electrons between them.

• One well known exception to the identical atom rule is the combination of carbon and hydrogen in all organic compounds.

• Non polar bonding exists between carbon and hydrogen

• Iodine (I2) and oxygen (O2) forms a diatomic (di – two, atomic – atoms) non-polar covalent molecule.

• x

2. Polar Bonds• Polarised bonds occur in covalent molecules.

One atom has greater electronegativity and therefore has a partial (slight) negative charge because it attracts the bonding electrons more strongly (remember electrons have a –ive charge).

• The other atom has a slight

positive charge.

Partial positive is represented by this symbol

Partial negative is represented By this symbol

• x

Which is more electronegative H or Cl?

• Cl has greater electronegativity

• The bonding electrons therefore spend more time near the Cl atom

• So the Cl side of the molecule becomes slightly negative (partial negative

charge) and the H side becomes

slightly positive making a

polarised bond

• What about HF?

Polar molecules• Polar molecules are those containing polar

bonds that do not cancel each other out• Molecular polarity is dependent on the

difference in electronegativity between atoms in a compound and the asymmetry of the compound's structure.

• For example, a molecule of water is polar because of the unequal sharing of its electrons in a "bent" structure, whereas methane is considered non-polar because the carbon shares the hydrogen atoms uniformly.

Polar

• The covalent bond between HCl is said to be polarised and the molecule is called a polar molecule

• It is also said to be a dipole

• because is has two charged ends

Non polar molecules• Non polar molecules are those that

contain only non-polar bonds

OR

• Molecules in which polar bonds cancel each other out

Quiz

• What happens to the electrons in a polar bonds?

• How do you know if a bond is polar?

• What is a non polar molecule?

• What is a dipole?

• What is the symbol for slightly negative?

• Can non polar molecules have polar bonds?

How to determine polar moleculesThere are two important factors

1. The polarity of the individual bonds in the molecule;

2. The shape or geometry of the molecule.

Steps to take

a)Determine if a given individual bond is polar, Look at the difference between electronegativity of the atoms in the perioidc table. If the difference is:

0.3 < non polar

0.3 -1.7 = polar

b) Determine the shape of molecule. Then there are 3 option

i) if all bonds are non-polar, then the whole molecule is non-polar regardless of its shape.

ii) If there is symmetry in the molecule so that the polarity of the bonds cancels out, then the molecule is non-polar. (symmetry arround the central atom)

iii) If there are polar bonds but there is no symmetry the overall molecule is polar.

Which molecules are polar?

Determine if BF3 is polar or non polar

• B-F bonds are polar but they are symmetrically arranged around the central fluorine atom. No side of the molecule has more negative or positive charge than another side, and so the molecule is nonpolar.

Predict whether the following molecules are polar or nonpolar:

(a) BrCl; (b) SO2; • SOLUTION (a) Chlorine is more electronegative

than bromine. Consequently, BrCl will be polar with chlorine carrying the partial negative charge:

• (b) Because oxygen is more electronegative than sulfur, the molecule has polar bonds.

• For each of these, the VSEPR model predicts a bent geometry. Because the molecule is bent, the bond dipoles do not cancel and the molecule is polar

Which of the following molecules contains polar bonds but is nonpolar?

a) H2O

b) NH3

c) F2

d) CCl4e) CH2Cl2

• In CCl4, there are 4 bonds arranged symmetrically around the central carbon atom.

• Each bond is polar, you can tell due to the electronegativity differences between carbon and chlorine.

• Because the 4 atoms of chlorine are arranged symmetrically, the polarity of the bonds cancel each other out, making a non-polar molecule.

Determine if OF2 is polar or non polar

• The electronegativities of oxygen and fluorine, 3.44 and 3.98, respectively, produce a 0.54 difference that leads us to predict that the O-F bonds are polar. The molecular geometry of OF2 is bent. Such an asymmetrical distribution of polar bonds would produce a polar molecule.

• When there are no polar bonds in a molecule, there is no permanent charge difference between one part of the molecule and another, and the molecule is nonpolar.

What is the difference between a polar bond and a polar molecule?

• A polar bond is a charged bond. There is a positive and negative end that do not cancel each other out

• A polar molecule contains polar bonds that do not cancel each other out.

ACTIVITY

Discuss in a group and write a report about a molecule or compound given to you. Present your report to class.

Distribution of work: person 1 – geometric structure of molecule or compound, 2 – properties of molecule or

compound, 3 – properties if element 1, 4 - properties of elements 2 and 3.

Distribution for groups:Group I – H2O, NaCl; II – CO, O3; III – KBr, NH3;

IV – LiF, SO2; V – HCl, O2; VI – SO, F2; VII – SO2,Cl2VIII – CO2, Br2