IB DP1 ChemistryBonding

What makes atoms join together to make compounds?

Topic 4: Bonding (12.5 hours)

4.1 Ionic bonding4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions.4.1.2 Describe how ions can be formed as a result of electron transfer.4.1.3 Deduce which ions will be formed when elements in groups 1, 2 and 3 lose electrons.4.1.4 Deduce which ions will be formed when elements in groups 5, 6 and 7 gain electrons.4.1.5 State that transition elements can form more than one ion.4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values.4.1.7 State the formula of common polyatomic ions formed by non- metals in periods 2 and 3.4.1.8 Describe the lattice structure of ionic compounds.4.2 Covalent bonding4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei.4.2.2 Describe how the covalent bond is formed as a result of electron sharing.4.2.3 Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom.4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength.4.2.5 Predict whether a compound of two elements would be covalent from the position of the elements in the periodic table

or from their electronegativity values.4.2.6 Predict the relative polarity of bonds from electronegativity values4.2.7 Predict the shape and bond angles for species with four, three and two negative charge centres on the central atom using the valence shell electron pair repulsion theory (VSEPR).4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities.4.2.9 Describe and compare the structure and bonding in the

three allotropes of carbon (diamond, graphite and C60 fullerene).

4.2.10 Describe the structure of and bonding in silicon and silicon dioxide.4.3 Intermolecular forces4.3.1 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how theyarise from the structural features of molecules.4.3.2 Describe and explain how intermolecular forces affect the boiling points of substances.4.4 Metallic bonding4.4.1 Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons.4.4.2 Explain the electrical conductivity and malleability of metals.4.5 Physical properties4.5.1 Compare and explain the properties of substances resulting from different types of bonding.

Ionic Bonding

Crystals: 7 ‘perfect’ crystal shapes

Halite- rock salt- sodium chloride

Sodium chloride is an ionic compound with ions arranged in a lattice

Ions

charged particles with electrostatic attraction between them

Na+ Cl-

Sodium and chloride ions formed when electrons transfer

Na + Cl Na+ + Cl-

2,8,1   2,8,7   2,8   2,8,8

Ions

Group 1: H+, Li+, Na+, K+, Rb+, Cs+, Fr+

Group 2: Be2+, Mg2+, Ca2+, Sr2+, Ba2+

Group 3?/13: B3+, Al3+, Ga3+

Group 6?/16: O2-, S2-,

Group 7?/17: F-, Cl-, Br-, I-

Which is the smallest ion?

Na+

Al+3

Cl-

P3-

Two or more electrons can be transferred

Different sized atoms give different mineral structures as they pack in a different way

Hexagonal Beryl crystal; Image Wikipedia

What is the formula of iron (III) oxide?

Fe2O

FeO

Fe3O2

Fe2O3

Polyatomic ions: charge distributed over more than one atom

For example phosphate, PO4-

3

can be found in products of reactions of phosphoric acid

Some common polyatomic ions

Nitrate NO3-

Hydroxide OH-

Sulphate SO42-

Carbonate CO32-

Hydrogen carbonate HCO3-

(Bicarbonate)

Phosphate PO43-

Ammonium NH4+

Common Anions Common Name Formula Alternative

name Simple Anions Chloride Cl− Fluoride F− Bromide Br− Oxide O2− Polyatomic anions Carbonate CO3

2- Hydrogen carbonate

HCO3− bicarbonate

Hydroxide OH− Nitrate NO3

2- Phosphate PO4

3- Sulfate SO4

2- Anions from Organic Acids Ethanoate CH3COO− acetate Methanoate HCOO− formate Ethandioate C2O4

−2 oxalate Cyanide CN-

Common Cations Common Name Formula Alternative

name Simple Cations Aluminium Al3+ Calcium Ca2+ Copper(II) Cu2+ cupric Hydrogen H+ Iron(II) Fe2+ ferrous Iron(III) Fe3+ ferric Magnesium Mg2+ Mercury(II) Hg2+ mercuric Potassium K+ kalic Silver Ag+ Sodium Na+ natric Polyatomic Cations Ammonium NH4

+ Hydronium H3O+

Careful with...

name of atom can change when ion is formed chlorine atom (Cl) chloride ion (Cl-)

-ate is often a polyatomic ion with oxygen eg sulphate, phosphate, etc.

different ions often have similar names... nitrate NO3

-

nitrite NO2-

nitride N-3

What is the formula of ammonium sulphate?

NH4SO4

(NH4)2SO4

NH4(SO4)2

SO4(NH4)2

d-block (transition elements) can have variable valencies

Mn2+ manganese(II)

Mn3+ manganese(III)

Mn4+ manganese(IV)

Ni2+ nickel(II)/nickelous

Ni3+ nickel(III)/nickelic

Pb2+ lead(II)/plumbous

Pb4+ lead(IV)/plumbic

Cr2+ chromium(II)/chromous

Cr3+ chromium(III)/chromic

Cu1+ copper(I)/cuprous

Cu2+ copper(II)/cupric

Fe2+ iron(II)/ferrous

Fe3+ iron(III)/ferric

Hg2+ mercury(I)/mercurous

Covalent bonding

Define electronegativity

Electronegativity is the tendency of an atom to attract electrons towards itself. The atoms with higher values attract electrons more strongly.

Highest flourine (and rest of groups 7,6,5)

FONClBrISCH

Wikipedia table

How ionic is an ionic compound?

bigger difference in electronegativity more ionic

(‘ionic’ usually De-neg> 1.8 difference)

usually metal + non-metal

Which aluminium compounds will be ionic?

atom Al F O Cl Br

electronegativity

1.5 4.0 3.5 3.0 2.8

Formula of aluminium compound

De-neg ‘Ionic’ or ‘covalent’?

‘Sharing’ electrons De-neg < 1,7covalent bonding forms molecules

Often between non-metals

Covalent bond formation- valence electrons

2, 4 or 6 electrons?

Single bond: the two atoms share two electrons (1 pair)

Double bond: the two atoms share four electrons (2 pairs)

Triple bond: the two atoms share six electrons (3 pairs)

Lewis structures (dot structures) show valence electrons in pairs as dots, crosses or lines

skeletal formula for complex organic molecules

Condensed formula

propanol CH3CH2CH2OH

Coordinate covalent bond (dative bond)

both electrons in the bond from the same atomonce formed, is the same as any other covalent bond

Bond lengths and Bond strengths

As the number of shared electrons increases (single to triple) the bond lengths shortens and the bond energy increase

Bond Bond type Lengths (pm) Energy (kJ/mol)

CC Single 154 347

CC Double 134 614

CC Triple 120 839

Which bond has the highest bond polarity, δ

H-H

Cl-Cl

Al-F

Al-Br

Non-polar covalent bond

In, H2 the two electrons in the bond are shared equally between the two hydrogen atoms.

H-H De-neg =0.

The electron distribution is symmetrical.

Polar covalent bond

If two different atoms form a covalent bond there will be a difference in De-neg.

The atom with highest electronegativity will have the electrons closer; they don’t share equally.

Unsymmetrical electron distribution.

Bonds

100% Covalent bond Polar covalent bond Ionic bond

% ionic character of a bond: 0-90%

(there are no 100% ionic compounds)

Molecular shapes

What shape are molecules?

VSEPR theory (Valence shell electron pair repulsion)

pairs of electrons repel and sit as far away as possible from each other

double and triple bonds count as a pair

VSEPR: electron repulsion molecular shape

Structure of molecule given by pairs of electrons arranging around an atom to be as far apart as possible

non-bonded pairs repel more than bonded pairs

double and triple bonds count as one

Build molecules from plasticine and straws

bond: 3cm length of straw

atom: 1cm diameter plasticine ball

unbonded pair of electrons 1cm straw length

Number of charge centres

Name of shape Bond angles (s)

Example

2 linear 180 BeCl2

3 trigonal planar 120 BF3

4 tetrahedral 109.5 CH4

5 trigonal bipyramidal

90, 120, 180

6 octahedral 90, 180

Shapes of simple molecules

http://en.wikipedia.org/wiki/Phosphorus_pentafluoridehttp://en.wikipedia.org/wiki/Sulphur_hexafluoridehttp://en.wikipedia.org/wiki/Boron_triflouride

Methane, Water and Ammonia

greater repulsion between non-bonding pairssmaller bond angles than predicted

Intermolecular forcesWhy do molecules stick together to form liquids and solids?

Intermolecular forces hold molecules together, affecting physical properties

Melting and boiling points

Strength

Flexibility

Viscosity

Intermolecular forces

Hydrogen bond strong

Dipole-dipole weaker

van der Waal’s forces weakest

Why do molecules attract each other to make liquids and gases?

Intermolecular forces: electrostatic attraction between

permanent dipoles (polar molecules)

permanent dipole and a temporary dipole (induced polarity)

temporary diploes (induced polarity)

Why do molecules attract each other?

electrostatic attraction between…

permanent dipoles (in polar molecules)

temporary diploes

A dipole is a overall charge imbalance in a molecule.

Which of the following molecules are polar?

Induced dipoles in all molecules (van der Waal’s forces)

Image: http://www.uwec.edu/boulteje/Boulter103Notes/11December.htm

Movements in electron cloud Temporary dipoles.

Temporary dipole in one molecule can induce a temporary dipole in another.

van der Waals forces

The strength increases with molar mass of the molecule.

E.g. He b.p 4 K : Xe b.p. 165 K.

Only effective over short range so the molecule “area” is also important.

E.g: Pentane, C5H12, b.p. 309 K

Dimethylpropane, (CH3)4C b.p. 283 K

Is a molecule polar?

A polar molecule

Has polar covalent bonds.Look at the difference in electronegativity (FONClBrISCH)

AND

Unsymmetrical shape according to charge distribution.

Otherwise it will be a non-polar molecule.

Molecular polarity

Images: http://en.wikipedia.org/wiki/Molecular_polarity

HF

H2O

NH3

Molecular polarity

http://phet.colorado.edu/en/simulation/molecule-polarity

Dipole-dipole

Electrostatic attraction between molecules with permanent dipoles.

Stronger than vdW.

Hydrogen chloride M= 36,5 g/mol b.p. 188 K

Fluorine M= 38 g/mol b.p. 85K

Induced dipole

Image: http://www.uwec.edu/boulteje/Boulter103Notes/11December.htm

Polar and non-polar liquids are immiscible

Image: http://en.wikipedia.org/wiki/Petroleum

Hydrogen bonding

H bonded to a highly electronegative element eg F, O or N

proton strongly attracts electronegative element in another molecule

important in water

Image: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force

Hydrogen bond

In molecules that contain Hydrogen bonded to Oxygen, Nitrogen or Fluorine (high electronegativity and non-bonding electron pair).

Interaction of the non-bonding electron pair in one molecule and hydrogen (with high positive charge) in another molecule.

Examples

H2O b.p.=100oC H2S b.p.= -61oC

NH3 b.p.= -33 oC PH3 b.p.= -88oC

C3H8 b.p. CH3CHO C2H5OH

b.p. 20 oC 42 oC 78 oC

Examples

H2O b.p.=100oC H2S b.p.= -61oC

NH3 b.p.= -33 oC PH3 b.p.= -88oC

C3H8 b.p. CH3CHO C2H5OH

b.p. 20 oC 42 oC 78 oC

Ice

Image: http://en.wikipedia.org/wiki/Ice

Trends in physical properties

How strong are the forces between molecules?

Bond type Dissociation energy (kJ/mol)

Covalent 1600

Hydrogen bonds 50–70

Permanent dipoles 2–8

Induced dipoles <4

Data: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force

Trends in physical properties

melting point /C boiling point /C

Flourine -220 -188

Chlorine -102 -34

Bromine -7 59

Iodine 114 184

Astatine 302 337

Plot one graph showing melting point and boiling point (in Kelvin) against molar mass for the halogensDescribe the pattern (2 sentences)Explain the pattern (2 sentences)

Data: http://en.wikipedia.org/wiki/Halogen

How strong are the forces between molecules?

Bond type Dissociation energy (kJ/mol)

Covalent 1600

Hydrogen bonds 50–70

Permanent dipoles 2–8

Induced dipoles <4

Data: http://en.wikipedia.org/wiki/Induced_dipole#Debye_.28induced_dipole.29_force

Allotropes: different structural forms of the same element

http://catalog.flatworldknowledge.com/bookhub/4309?e=averill_1.0-ch18_s04

OxygenO2 diatomic oxygenO3 ozone

Allotropes of Carbon

Diamond

Hard, colourless, insulator

Tetrahedral, giant structure

Covalent bond => sp3 orbitals.

Graphite

Slippery, black, conductor

Layers of fused six-membered rings. Each carbon surrounded by three others in a planar trigonal arrangement => sp2 + p-orbital

The p-orbital is perpendicular to the layer and give close packed p-orbitals

stabilise the layers

Delocalisation of electrons => electrical conductivity

Fullerene, C60

Spherical molecule. Looks like a football. 12 pentagons and 20 hexagons.

Bonds: C60 –hydration C60H60

(C2H4 + H2 C2H6 ; 1 H2 / double bond)

Each carbon has a double bond

Silicon

Metalloid, Semiconductors, non-metallic structure

Similar structure as diamond.

Silicon dioxide

SO2 Silica, giant structure similar to diamond

 

Silicates, SiO4, tetrahedrical, silicon-oxygen single bond

Physical properties

Melting points (impurities lower the melting point)

Boiling points

Volatility (how easy a compound will convert to gas)

Electrical conductivity

Solubility

Properties

Structure typeProperty

GiantMetallic

GiantIonic

GiantCovalent

MolecularCovalent

Hardness and malleability

Variable hard-ness, malleable rather than brittle

Hard and brittle Hard and brittle Usually soft and malleable unless hydrogen bonded

Melting and boiling points

Variable dep. On No of valence e-

High Very High Low

Electrical and thermal conductivity

Good in all states

Not as solids, conduct in (aq) or (l)

No No

Solubility 

Insoluble, except as alloys

In Water mostly Insoluble Often more soluble in other than water except if H-bonded

Examples Iron, copper NaCl, Na2SO4 Diamond,SiO2 (Sand)

CO2, Cl2,

ethanol, sugar

Ionic salts

Typical properties

Hard, brittle,

Conduct electricity in solution or melted.

High melting points => Strong bonds

Hydration of Ion in Water solution

Metallic bond

Metals have low electronegativity.

The atoms are packed close together in a lattice.

The valence electrons are delocalised among all atoms. The valence electron have no “home” The atoms can be seen as positive ions in a see of

electrons that keep them together.

This can explain the metallic properties

Electrical conductivity: electrons float around. If you put in one, one will fall out.

Malleability (smidbarhet) and Ductility (sträckbarhet): if the atom is pushed from its location the electron will follow. The bond is between the ion and the electrons not between the ions.

Investigate a physical property of a mixture related to intermolecular forces

Quantitative independent variable (cause)

Quantitative dependent variable (effect) viscosity, deflection by charged object, or other physical property

Links

Ionic bonding http://www.teachersdomain.org/asset/lsps07_int_ionicbonding/

Covalent bonding http://www.teachersdomain.org/asset/lsps07_int_covalentbond/

Polarity links

http://phet.colorado.edu/en/simulation/molecule-polarity

Viscosity http://www.youtube.com/watch?v=3KU_skfdZVQ

States of matter http://phet.colorado.edu/en/simulation/states-of-matter

Polarity links

http://phet.colorado.edu/en/simulation/molecule-polarity

http://antoine.frostburg.edu/chem/senese/101/liquids/faq/h-bonding-vs-london-forces.shtml

States of matter http://phet.colorado.edu/en/simulation/states-of-matter

http://employees.oneonta.edu/viningwj/modules/CI_dipoleinduced_dipole_forces_13_5a.html

Notes: http://www.uwec.edu/boulteje/Boulter103Notes/11December.htm

Snowflakes: http://www.its.caltech.edu/~atomic/snowcrystals/class/class.htm

Ice crystals http://www.edinformatics.com/interactive_molecules/ice.htm

Links

http://phet.colorado.edu/en/simulation/molecule-shapes

http://en.wikipedia.org/wiki/Phosphorus_pentafluoride

http://en.wikipedia.org/wiki/Sulphur_hexafluoride

http://en.wikipedia.org/wiki/Boron_triflouride

Teaching notes