Topic 4 Bonding.

SL + HLIonic bond

Covalent bondIntermolecular forces

Hydrogen bondDipole-dipole attraction

van der Waals’ forcesMetallic bond

4.1 Ionic bond

Ions = charged particlesIonic bond= the electrostatic bond between positively and negatively charged ionsIonic compound= a compound built from ions, i.e. a salt

• Atoms want to get a “Noble gas electron configuration”. One way to get that structure is to throw away the valence electron or steal some electrons to get a full outer shell.

• Sodium gives one electron to chlorine and both will have noble gas configuration. Which noble gas?

• Sodium ion: positive electric charge. Cation.Chloride ion: negative electric charge. Anion.

Na + Cl Na+ + Cl-

2,8,1   2,8,7   2,8   2,8,8

Ionic crystals• + charge particles and – charge particle

attracts each other in three dimension and builds up a lattice/crystal. Strong electrostatic forces in three dimensions. Each cation is surrounded by anions and vice versa.

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-

Naming compounds

• Positive ions have the name of the atom: Sodium ion

• Negative ions have the name of the atom (or almost) + the ending –ide: Chloride ion

• Sodium + chlorine Sodium chloride• Lithium + oxygen Lithium oxide

Formula of ionic compounds

Na+ Mg2+ Al3+

Cl-

O2-

N3-

Write the chemical formula of the compounds formed between the positive and negative ions above.Write the name of the ionic compounds

Transition metals

• Transition metals can often form more than one ion, e.g.:– Fe2+ and Fe3+

– Cu+ and Cu2+

• HL: coloured because of d-orbitals

Some common polyatomic ions

• Nitrate NO3-

• Hydroxide OH- • Sulphate SO4

2-

• Carbonate CO32-

• Hydrogen carbonate HCO3-

(Bicarbonate)

• Phosphate PO43-

• Ammonium NH4+

Ca(OH)2

Formula of ionic compounds-2

Na+ Mg2+ NH4+

OH-

SO42-

PO43-

Write the chemical formula of the compounds formed between the positive and negative ions above. Write the name of the ionic compounds

When can we expect an Ionic bond?

• The quick rule: Metal + non-metal => Ionic compounds

(Salts)

CuSO4 (s)

Electronegativity

• Electronegativity is a measure of an atoms power to attract electrons.

• On the right side in the periodic table (group 7,6,5) the atoms have high values, they attract electrons readily. Best is Fluorine, e-neg 4.

• On the left side the values are low. Low ability to attract electrons.

Electronegativity values

FONClBrISCH

Ionic bond or not- calculate difference in electronegativity

• If you want to do a more “precise” estimation you can calculate it with the help of electronegativity values.

• If the difference in electronegativity > 1,7 then you can say it's an Ionic bond.

Ionic bond or not?Use electronegativity values in the Chemistry Data booklet

• NaCl• MgO• AlCl3

• SiO2

• Ca3N2

Typical properties of Salts

Hard, brittle, • Conduct electricity in solution or melted • High melting points => Strong bond• Hydration of Ion in Water solution

4.2 Covalent bond

Electron pair bond, molecular bond

• If the De-neg < 1,7 the bond is considered to be covalent

• Often between non-metals• Polar or non-polar

• In a covalent bond the atoms share electrons with

each other to get a “Noble gas electron configuration”

• The bond has a direction (one atom to an other atom)

• One bond consists of two electrons, an electron pair

OH H

• 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

• all valence electrons marked by dots or lines. Draw Lewis structures for:

• F2

• NH3

• CO2

• N2

• C2H4

Number of bonds bond lengths and bond strengths

• As the number of shared electrons increase (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

-COOH Single 143 358

-COOH Double 123 745

Non-polar covalent bond

• In, for example, 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

• Which molecules contains at least one polar covalent bond?

F2 HF NF3 SiF4

• Which bond in the pairs below have the highest polar character?

a) C-O, N-Ob) H-O, S-Oc) H-O, H-Sd) Se-S, Se-F

Ionic, polar or non-polar covalent bond?

• % ionic character of a bond: 0-90%(there are no 100% ionic compounds)

Na+ Cl-H-ClCl-Cl

Dative covalent bond

• In a “normal” covalent bond both atoms contribute with electrons to the bond.

• Sometimes only one atom contributes with both electrons (the electron pair) to the bond

• Then the covalent bond is called a dative covalent bond

Examples of dative bonds

• H2O + H+ H3O+

• NH3 + H+ NH4+

• C (4 ve-)+ O (6 ve-) CO

VSEPR

• Valence Shell Electron Pair Repulsion

• Shape and bond angles • Determine the molecules structure, the shape in 3

dimensions• Structure around a given atom is determined principally

by minimising electron-pair repulsion

• Bonding or non-bonding pairs will be as far apart as they can.

Linear

• Two negative centers• 180o • E.g. CO2

Trigonal planar (flat)

• Three negative centers• 120o • E.g. BF3

Tetrahedral arrangement

• Four negative centers• 109.5 o

Method

1. Draw Lewis structure2. Count electron pairs, minimise the repulsion3. Positions of the atoms4. Name of the structure

Ammonia, NH3 Tetrahedral, Trigonal planar 107o, one non-bonding electron pair

• Non-bonding pair (lone pair) takes more space => reduce bond angels

Water, H2O Tetrahedral, Non-linear (bent) 104o Two non-bonding electron pairs

Non-bonding pair (lone pair) takes more space => reduces bond angles

Non-polar and polarmolecules (dipoles)

• Based on bond polarity and molecular shape • May predict how a molecule will behave with

other compounds• Polar molecule = Dipole

A polar molecule (a dipole)• Must have 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 (NOT a dipole)

Dipole or not?

• H2O

• CO2

• NH3

• CH4

• CH3OH

Allotropes

• Some elements can be found in different forms

• E.g: – Carbon: Diamond, graphite, C60fullerene– Oxygen, Ozone

See PPT: Carbon allotropes

Silicon

• Metalloid, Semiconductors, non-metallic structure

• Similar structure as diamond

Silicon dioxide• SiO2 Silica, giant structure similar to diamond, quarts

• Silicates, NaSiO4, tetrahedrical, silicon-oxygen single bond

4.3 Intermolecular forces • Holds molecules together in liquids or solids (No

Intermolecular forces between gaseous particles)

• Weaker than covalent and ionic bonds

• Hydrogen bond (Quite strong)• Dipole-dipole (Middle weak)• van der Waal’s forces (~ London dispersion forces)

(Very weak)

• Accounts for differences in aggregatio state and solubility

van der Waal’s forces• “Vibrations” in the electron cloud => Temporary

dipoles.

• A temporary dipole in one molecule can induce a temporary dipole in another molecule

• Exist between all molecules

• The strength increases with molar mass of the molecule/atom

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

van der Waal’s forces, cont.

Dipole-dipole bond• Electrostatic attraction between molecules

with permanent dipoles• Stronger than van der Waals bond

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

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

Hydrogen bond• In molecules that contain Hydrogen bonded to

Oxygen, Nitrogen or Fluorine (high electronegativity and non-bonding electron pair)

• Stronger than dipole-dipole bonds

• Interaction of the non-bonding electron pair in one molecule and hydrogen (with high positive

charge) in another molecule.

Which intermolecular bond?• H2O b.p.= 100oC H2S b.p.= -61oC

• NH3 b.p.= -33 oC PH3 b.p.= -88oC

Which intermolecular bond?

• C2H6 b.p. CH3CHO C2H5OH b.p. -89 oC 20 oC 78 oC

Boiling points of hydrogen compounds

What kind of intermolecular bonds can be expected to dominate between the molecules/atoms?a) NH3(l)b) C5H12(l) c) Br2(l) d) HI(I) e) Ar(l) f) PCl3(l) g) CH3OH(l) h) CHCl3(l)

4.4 Metallic bonding

• 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 it’s location the electron will follow. The bond is between the ion and the electrons not between the ions.

Summary• The different kind of bonds is very important

for the behaviour of a compound, solution or a mixture

4.5 The strength/type of the bond affects:

• Melting points (impurities lower the melting point)

• Boiling points• Volatility (how easy a compound will convert

to gas)• Electrical conductivity• Solubility and miscibility

Fe (s) Fe (l) Fe (g)

NaCl (s) NaCl (l) NaCl (g)

Which bonds are broken?Which bond is strongest?

Changes in state (1)1538 ºC 2861 ºC

801 ºC 1413 ºC

H2O (s) H2O (l) H2O (g)

HCl (s) HCl (l) HCl (g)

CH4 (s) CH4 (l) CH4(g)

Which bonds are broken?Which bond is strongest?

Changes in state (2)0 ºC 100 ºC

-85 ºC -61 ºC

-162 ºC-184 ºC

• Molecules with van der Waals bonds

Solubility

• Unpolar compounds • Polar compounds• Ionic bonds• Molecules with

hydrogen bonds • Molecules with

dipol-dipol bonds

Intermolecular bonds in solutions• Ion-hydrogen bonds or ion-dipole bonds

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