1. IB DP1 Chemistry Bonding
What makes atoms join together to make compounds?

2. Topic 4: Bonding (12.5 hours)
4.1 Ionic bonding Describe the ionic bond as the electrostatic attraction between oppositely charged ions Describe how ions can be formed as a result of electron transfer Deduce which ions will be formed when elements in groups 1, 2 and 3 lose electrons Deduce which ions will be formed when elements in groups 5, 6 and 7 gain electrons State that transition elements can form more than one ion 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 State the formula of common polyatomic ions formed by non- metals in periods 2 and Describe the lattice structure of ionic compounds. 4.2 Covalent bonding Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei Describe how the covalent bond is formed as a result of electron sharing Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom State and explain the relationship between the number of bonds, bond length and bond strength 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 Predict the relative polarity of bonds from electronegativity values 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) Predict whether or not a molecule is polar from its molecular shape and bond polarities Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C60 fullerene) Describe the structure of and bonding in silicon and silicon dioxide. 4.3 Intermolecular forces Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules Describe and explain how intermolecular forces affect the boiling points of substances. 4.4 Metallic bonding Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons Explain the electrical conductivity and malleability of metals. 4.5 Physical properties Compare and explain the properties of substances resulting from different types of bonding.

3. Ionic Bonding

4. Crystals: 7 ‘perfect’ crystal shapes

6. Halite- rock salt- sodium chloride

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

8. Ions
charged particles with electrostatic attraction between them
Cl-
Na+

9. Sodium and chloride ions formed when electrons transferNa + Cl 
Na+
Cl-
2,8,1
2,8,7
2,8
2,8,8

10. 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-

11. Which is the smallest ion?
Na+ Al+3 Cl- P3-

12. Two or more electrons can be transferred
Different sized atoms give different mineral structures as they pack in a different way
Two or more electrons can be transferred
Hexagonal Beryl crystal; Image Wikipedia

13. What is the formula of iron (III) oxide?
Fe2O FeO Fe3O2 Fe2O3

14. Polyatomic ions: charge distributed over more than one atom
For example phosphate, PO4-3
can be found in products of reactions of phosphoric acid

15. Some common polyatomic ions
Nitrate NO3-
Hydroxide OH-
Sulphate SO42-
Carbonate CO32-
Hydrogen carbonate HCO3-
(Bicarbonate)
Phosphate PO43-
Ammonium NH4+

16. 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+
Common Anions
Common Name
Formula
Alternative name
Simple Anions
Chloride
Cl−
Fluoride F−
Bromide
Br−
Oxide
O2−
Polyatomic anions
Carbonate
CO32-
Hydrogen carbonate
HCO3−
bicarbonate
Hydroxide
OH−
Nitrate
NO32-
Phosphate
PO43-
Sulfate
SO42-
Anions from Organic Acids
Ethanoate
CH3COO−
acetate
Methanoate
HCOO−
formate
Ethandioate
C2O4−2
oxalate
Cyanide
CN-

17. 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

18. What is the formula of ammonium sulphate?
NH4SO4
(NH4)2SO4
NH4(SO4)2
SO4(NH4)2

19. 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

20. Covalent bonding

21. 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
Show wooden electronegativity chart

22. How ionic is an ionic compound?
bigger difference in electronegativity  more ionic
(‘ionic’ usually De-neg> 1.8 difference)
usually metal + non-metal

23. 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’?

24. ‘Sharing’ electrons De-neg < 1,7 covalent bonding forms molecules
Often between non-metals

25. Covalent bond formation- valence electrons

26. 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)

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

28. skeletal formula for complex organic molecules

29. Condensed formula
propanol CH3CH2CH2OH

30. Coordinate covalent bond (dative bond)
both electrons in the bond from the same atom
once formed, is the same as any other covalent bond

31. 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
Double
134
614
Triple
120
839
-COOH
Single
143
358
Double
123
745

32. Which bond has the highest bond polarity, δ
H-H Cl-Cl Al-F Al-Br

33. 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.

34. 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.

35. Bonds 100% Covalent bond  Polar covalent bond  Ionic bond
% ionic character of a bond: 0-90%
(there are no 100% ionic compounds)

36. Molecular shapes

37. 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

38. 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

39. Build molecules from plasticine and straws
bond: 3cm length of straw
atom: 1cm diameter plasticine ball
unbonded pair of electrons 1cm straw length

40. Shapes of simple molecules
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

41. Methane, Water and Ammonia
greater repulsion between non-bonding pairs
smaller bond angles than predicted

42. Intermolecular forces
Why do molecules stick together to form liquids and solids?

43. Intermolecular forces hold molecules together, affecting physical properties
Melting and boiling points
Strength
Flexibility
Viscosity
Deflection in electric field
Volatility (how easy a compound will convert to gas)
Electrical conductivity
Solubility
What properties are affected by Van der Waals´forces

44. Intermolecular forces
Hydrogen bond strong Dipole-dipole weaker van der Waal’s forces weakest

45. 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)
A dipole is a overall charge imbalance in a molecule.
PhET sim states of matter- where are there forces between molecules?
Not ions... But there is an electrostatic attraction between positive and negative

46. Induced dipoles in all molecules (van der Waal’s forces)
Movements in electron cloud  Temporary dipoles.
Temporary dipole in one molecule can induce a temporary dipole in another.
Image:

47. van der Waals forces
The strength increases with molar mass of the molecule.
He b.p 4K
Xe b.p. 165K.
Only effective over short range so the molecule “area” is also important.
Pentane, C5H12, b.p. 309K
Dimethylpropane, (CH3)4C b.p. 283K

48. Trends in physical properties
Plot one graph showing melting point and boiling point (in Kelvin) against molar mass for the halogens
Describe the pattern (2 sentences)
Explain the pattern (2 sentences)
melting point /C
boiling point /C
Flourine
-220
-188
Chlorine
-102
-34
Bromine -7 59
Iodine
114
184
Astatine
302
337
Show samples of I2, Br2 and Cl2 video of F2 at (Flourine periodic table of elements)
Data:

49. Is a molecule polar? A polar molecule has polar covalent bonds. AND
Is there a difference in electronegativity? (FONClBrISCH)
AND
has an asymmetric shape according to charge distribution.
Otherwise it is a non-polar molecule.

50. Molecular polarity HF NH3 H2O
Images:

51. Molecular polarity

52. 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 K

53. Induced dipole
Image:

54. Hydrogen bonding
H bonded to a highly electronegative element- F, O or N
proton  unbonded pair
important in water
Image:

55. Examples H2O b.p. 373K H2S b.p 212K NH3 b.p. 240K PH3 b.p 185K
C3H8 bp20 oC
CH3CHO bp42 oC
C2H5OH bp78 oC

56. Ice
Image:

57. Polar and non-polar liquids are immiscible
Image:

58. Allotropes

59. Allotropes: different structural forms of the same element
Oxygen
O2 diatomic oxygen
O3 ozone

60. Allotropes of Carbon

61. C allotropes: Diamond Hard, colourless, insulator
Tetrahedral, giant structure
Covalent bonds  sp3 orbitals.
Image:

62. C allotropes: Graphite
Slippery, black conductor
Layers of fused six-membered rings
Each carbon surrounded by 3 others in a trigonal planar arrangement  sp2 + p-orbital
p-orbital perpendicular to layers and gives close-packed p-orbitals
Delocalized electrons electrical conductor
Image:

63. C: Allotropes: Fullerene, C60
Spherical molecule
12 pentagons and 20 hexagons.
Image:

64. C allotropes: Carbon nanotubes
Image:

65. Silicon solid at room temperature
high melting and boiling points of 1414 and 3265 °C
conducts heat well
grey color and a metallic luster
strong, very brittle
crystallizes in a diamond cubic crystal structure
Images:

66. Silicon dioxide SO2 Silica giant structure similar to diamond
silicates, SiO4, tetrahedral, silicon- oxygen single bond
Image:

67. Metals

68. Metallic bond Metals have low electronegativity.
Atoms packed into a lattice.
Calence electrons delocalised
valence electrons have no “home”
atoms positive ions in a sea of electrons that keep them together.
Image:

69. Metallic properties
Electrical conductivity: electrons float around. Put one in, one is pushed out.
Malleability and Ductility: if the atom moves, the electron follows. Bond is between ion and electrons, not between ions.

70. Patterns in bonding and properties

71. 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:

72. Ionic salts Hard, brittle, Conduct electricity in solution or melted.
High melting points and boilign points
Ions hydrated in aqueous solution

73. Covalent compounds low mp and bp
poor conductors of electricty and heat

74. Summary of properties Structure type Property Giant Metallic Ionic
Covalent
Molecular
Hardness and malleability
Variable hard-ness, malleable rather than 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
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

75. Investigation

76. 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

77. Links
Ionic bonding /
Covalent bonding d/

78. Polarity links
Viscosity
States of matter of-matter

79. Polarity links
forces.shtml
States of matter
_13_5a.html
Notes:
Snowflakes:
Ice crystals

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

81. Teaching notes