1.4 Bonding in Compounds Learning Intentions

Higher Chemistry Unit 1 – Chemical Changes and Structure Section 3 – Bonding in Compounds

1.4 Bonding in Compounds Learning Intentions
Learn how to explain differences in physical properties such as viscosity, melting point and boiling point in terms of differences in strength of intermolecular forces. I can describe a covalent bond in terms of electrons and positive nuclei. I can use electronegativity values to distinguish between pure covalent and polar covalent bonds. I can use the data book to assign + and - partial charges on atoms. I can distinguish between bonds which show a dipole and those that are non-polar. I can explain the relationship between differences in electronegativity and type of bonding. I can use data from properties such as conductivity, melting point and boiling point to deduce type of bonding and structure. I can list exceptions to the statements: “Compounds formed between non-metals only are covalent. Compounds formed between a metal and a non-metal are ionic”. I can explain the difference between intermolecular and intramolecular forces. I can name temporary and permanent Van der Waal’s forces. I can give examples of compounds with temporary and permanent Van der Waal’s forces. I can explain how London dispersion forces arise and I can state where London dispersion forces arise. I can relate the strength of London dispersion forces to the number of electrons within the atom or molecule. I can relate London dispersion forces to other types of bonds in terms of strength. I can explain how permanent dipole-permanent dipole interactions arise. I can use spatial arrangement of polar covalent bonds to predict whether a molecule is polar or non-polar. I can relate permanent dipole-permanent dipole interactions to other types of bonds in terms of strength. I can correctly describe hydrogen bonding as an example of intermolecular or intramolecular bonding. I can identify from the structural formula or the molecular formula compounds which have hydrogen bonds. I can relate hydrogen bonding to other types of bonds in terms of strength. I can explain the relationship between the type and strength of intermolecular bonds and melting point, boiling point and viscosity. I can predict the strength of intermolecular forces by considering the polarity of molecules and the number of electrons. I can explain patterns in melting and boiling points in terms of the strength of intermolecular forces. I can explain why the boiling points of ammonia, water and hydrogen fluoride are higher than expected given the number of electrons present in the molecules. I can explain how hydrogen bonding affects boiling points, melting points, viscosity and solubility (miscibility). I can explain why ice is less dense than water. I can state the type of solvent which tends to dissolve ionic compounds and polar molecular compounds. I can select groups within molecules which imply hydrogen bonding. I can identify molecules with a permanent dipole by considering the spatial arrangement of polar covalent bonds.

The Chemical Bond Chemical Bond Intermolecular Intramolecular
(Within) Intermolecular (between) Van der Waals The Chemical Bond Metallic Covalent Polar covalent Ionic Hydrogen bonding Permanent Dipole- Permanent Dipole interactions London’s forces

Ionic Compounds Ions - metals lose electrons and form positive ions
- non-metals gain electrons to form negative ions - electrons are transferred from metals to non-metals

Na atom + Cl atom Na+ ion + Cl- ion (2.8.1) (2.8.7) (2.8) (2.8.8)
transfer - + Cl Cl + Na Na Na atom + Cl atom Na+ ion + Cl- ion (2.8.1) (2.8.7) (2.8) (2.8.8)

Ionic Compounds The positive and negative ions are attracted (electrostatic bond )to each other. Ionic bond (electrostatic attraction) Na+ Cl- A giant lattice structure is formed. Each Na+ ion is surrounded by 6 Cl- ions. While each Cl- is surrounded by 6 Na+ ions. Ionic bonding is the electrostatic force of attraction between positively and negatively charged ions. This ionic network compound has many ionic bonds so ionic compounds have high m.p.s

Ionic Compounds - NaCl A giant lattice structure is formed when each Na+ ion is surrounded by 6 Cl- ions and each Cl- ion is surrounded by 6 Na+ ions. Sodium Chloride The formula of sodium chloride is NaCl, showing that the ratio of Na+ to Cl- ions is 1 to 1. The m.p. of NaCl is 801 0C The size of the ions will effect the strength of the ionic bond and how the ions pack together. e.g. NaF m.p. 1000oC, NaI 660oC

Molecular Ions, e.g. SO4 A single covalent bond.
Oxygen Sulphur S O A single covalent bond. 2 additional electrons e.g. Copper can donate the extra 2 electrons needed. Cu Cu e Copper sulphate contains the Cu2+ and the SO42- ions. There is, therefore, covalent bonding and ionic bonding in copper sulphate A solution of copper sulphate can conduct electricity. Molten ionic compounds can also conduct electricity.

Bond Strengths Bond Type Strength (kJ mol –1) Metallic 80 to 600 Ionic
Covalent Hydrogen 40 Dipole-Dipole 30 London’s forces 1 to 20

Covalent Molecular Compounds

Covalent Bonding Sharing electrons
takes place between non-metal and non-metal shared electrons count as part of the outer shell of both Atoms shared electrons attract the nuclei of both atoms this attraction is called the covalent bond

Hydrogen chloride H Cl (linear) HCl Cl H

Ammonia N (pyrimidal) NH3 H N H H H

Water O (angular) H2O H O H H

Draw electron dot cross diagrams for the following molecules and structural formula
SCl2 CO2 CH4 H H C H S Cl-S-Cl X O=C=O

Covalent Hydrogen 40 Dipole-Dipole 30 London’s Forces 1 to 20

Discrete molecules are formed when two or more atoms share electrons. The atoms are non-metal elements. An example is methane. Methane: CH4 H C C H Methane has strong intra-molecular and weak inter-molecular. It’s b.p. is -183oC

Non- metals elements can form double and triple covalent bonds. C H ethane C2H6 ethene C2H4 C H Double covalent bond Covalent molecular compounds have low m.p.’s because the weak forces holding the molecules together require only small amounts of thermal energy to break them.

Properties Low m.p.’s and b.p.’s., this increases with size of the molecule and the increasing number of atoms in the molecule. m.p.’s of the carbon halides 171 CF4 Temp / oC 90 CCl4 -23 CBr4 CI4 -183 m.p.’s increase because the strength of the London’s forces forces increase with the increasing size of the molecule. So more energy is needed to separate molecules.

Covalent Network Compounds

Silicon Carbide SiC Silicon, like carbon, can form giant covalent networks. Silicon carbide exist in a similar structure to diamond. Si C Covalent Bond Tetrahedral shape The 4 carbon atoms are available to bond with another 4 silicon atoms. This results in a COVALENT NETWORK COMPOUND

Silicon Carbide SiC Silicon carbide (carborundum) has a chemical formula is SiC. As this compound is linked by strong covalent bonding, it has a high m.p. (2730oC). It is a hard substance as it is very difficult to break the covalent lattice. SiC is used as an abrasive for smoothing very hard materials. Each Si is bonded to 4 C’s and each C is bonded to 4 Si’s. Hence the chemical formula, SiC

Silicon Dioxide SiO2 Silicon and oxygen make up nearly 75% of the Earth’s crust. They are therefore the most common elements in the Earth’s crust. They combine together to make a covalent network compound called silicon dioxide. This is usually found in the form of sand or quartz. Each Si atom is bonded to 4 O atoms, and each O atom is bonded to 2 Si atoms. Hence the chemical formula, SiO2 . Silicon dioxide (silica) also has a high m.p. (1600 oC) and like SiC, it is very hard and used as an abrasive. It is relatively un-reactive.

Periodic trends in electronegativity
Learning intention Learn the definition of electronegativity, and how to explain periodic trends in terms of nuclear charge, covalent radius and the screening effect of inner electrons.

Electronegativity Electronegativity is a numerical measure of the relative ability of an atom in a molecule to attract the bonding electrons towards itself.

Electronegativity Electronegativity is a measure of an atom’s attraction for the shared pair of electrons in a bond C H e Which atom would have a greater attraction for the electrons in this bond and why?

Values for electronegativity can be found on page 10 of the data book
Linus Pauling Linus Pauling, an American chemist (and winner of two Nobel prizes!) came up with the concept of electronegativity in 1932 to help explain the nature of chemical bonds. Today we still measure electronegativities of elements using the Pauling scale. Since fluorine is the most electronegative element (has the greatest attraction for the bonding electrons) he assigned it a value and compared all other elements to fluorine. Values for electronegativity can be found on page 10 of the data book

Electronegativities Looking across a row or down a group of the periodic table we can see a trend in values. We can explain these trends by applying the same reasoning used for ionisation energies.

Looking across a period
Increasing Electronegativity Li Be B C N O F 1.0 1.5 2.0 2.5 3.0 3.5 4.0 What are the electronegativities of these elements? Across a period electronegativity increases The charge in the nucleus increases across a period. Greater number of protons = Greater attraction for bonding electrons

What are the electronegativities of these halogens?
Looking down a group Decreasing Electronegativity F Cl Br I 4.0 3.0 What are the electronegativities of these halogens? 2.8 2.6 Down a group electronegativity decreases Atoms have a bigger radius (more electron shells) The positive charge of the nucleus is further away from the bonding electrons and is shielded by the extra electron shells.

Trends in electronegativity
Electronegativity increases across a period. Electronegativity decreases down a group Going across the period, the nuclear charge increases. This pulls the electron shells closer to the nucleus. As a results, the electronegativity increases. Going down the group, the nuclear charge increases but the number of electron shells also increases. As a result of ‘shielding’ and an increase distance the outer shell is from the nucleus, electronegativity decreases.

Chemical bonds: types of bonds Explores how different types of bonds are formed due to variations in the electronegativity of the bonded atoms. The distortion of the orbitals and the polarity of the bond is also displayed. Linus Pauling ( ) An account of the life and work of the Nobel Prize-winning chemist, Linus Pauling. Periodic Table of Data Visual database of the physical and thermochemical properties of the chemical elements which allows the user to plot graphs and tables, play games and view diagrams.

Polar covalent bonds Learning intention
Learn how differences in electronegativity between bonding atoms lead to the formation of polar covalent bonds.

Polar Covalent Bonds Non-polar covalent bond – electrons shared equally between atoms (same electronegativity) Polar covalent bond – electrons shared unequally between atoms (atom B is more electronegative)

Covalent Bonding A covalent bond is a shared pair of electrons electrostatically attracted to the positive nuclei of two atoms. Both nuclei try to pull the electrons towards themselves - + - + This is like a tug-of-war where both sides are pulling on the same object. It creates a strong bond between the two atoms. The atoms achieve a stable outer electron arrangement (a noble gas arrangement) by sharing electrons.

Covalent Bonding Picture a tug-of-war:
If both teams pull with the same force the mid-point of the rope will not move.

Pure Covalent Bond H e H This even sharing of the rope can be compared to a pure covalent bond, where the bonding pair of electrons are held at the mid-point between the nuclei of the bonding atoms.

Covalent Bonding What if it was an uneven tug-of-war?
The team on the right are far stronger, so will pull the rope harder and the mid-point of the rope will move to the right.

Polar Covalent Bond A polar covalent bond is a bond formed when the shared pair of electrons in a covalent bond are not shared equally. This is due to different elements having different electronegativities.

Polar Covalent Bond I δ- δ+ H e.g. Hydrogen Iodide e
If hydrogen iodide contained a pure covalent bond, the electrons would be shared equally as shown above. This makes iodine slightly negative and hydrogen slightly positive. This is known as a dipole. However, iodine has a higher electronegativity and pulls the bonding electrons towards itself (winning the tug-of-war)

Polar Covalent Bond C Cl δ+ δ-
In general, the electrons in a covalent bond are not equally shared. δ- δ+ e.g. C Cl 2.5 3.0 Electronegativities δ- indicates where the bonding electrons are most likely to be found.

Polar Covalent Bond Consider the polarities of the following bonds:
Electronegativities Difference C Cl 0.5 P H O H 1.3 C Cl δ- δ+ O H δ- δ+ P H Increasing Polarity Complete a similar table for C-N, C-O and P-F bonds.

Polar Covalent Bonds In the covalent bond between fluorine and hydrogen. The bonding electrons are not shared equally between the two atoms. Hydrogen Fluorine The fluorine nucleus has more protons and has a stronger pull on the electrons than the hydrogen nucleus..

- Thus the fluorine atom has a greater share of the bonding electrons and acquires a slight negative charge. + F H The hydrogen atom is then made slightly positive. The bond is a polar covalent bond and we use the symbols + and - to show this. The dipole produced is permanent. Fluorine is the most electronegative element. It is small atom compared to others and its nucleus is massive for its atomic size. Some other polar covalent bonds are O-H and N-H

Permanent dipole-permanent dipole interactions
Learning intention Learn about this additional intermolecular force of attraction which exists between polar molecules.

Dipole-Dipole Attractions
The differing electronegativities of different atoms in a molecule and the spatial arrangement of polar covalent bonds can cause a molecule to form a permanent dipole. H Permanent dipole Asymmetrical molecule + Cl Cl - Cl - - 3 polar covalent C–Cl bonds and 1 polar covalent C-H bond in CHCl3 POLAR molecule

Permanent Dipole-Dipole Interactions
The attraction is stronger than London's forces Hydrogen bonding is a particular example of dipole-dipole attractions. Molecules with permanent dipoles attract each other.

Polar Molecules and Permanent Dipoles
Both propanone and butane have the same formula mass of 58 however, butane boils at – 1 oC while propanone boils at 56oC Propanone is a polar molecule as it has a permanent dipole, so has polar-polar attraction as well as London’s forces between molecules. - + b.p. 56 o C Butane has no permanent dipoles, so only London’s forces between molecules. So has a lower boiling point. b.p. -1 o C

4 polar covalent C-Cl bonds in CCl4 tetrahedral shape
Symmetry CCl4 No permanent dipole Symmetrical molecule Cl - + Cl Cl - - Cl Tetrachloromethane has a symmetrical arrangement of polar bonds and the polarity cancels out over the molecule. - 4 polar covalent C-Cl bonds in CCl4 tetrahedral shape NON-POLAR molecule

2 polar covalent C=O bonds in CO2 linear shape
Symmetry CO2 No permanent dipole Symmetrical molecule + - - O O Carbon dioxide has a symmetrical arrangment of polar bonds and the polarity cancels out over the molecule. 2 polar covalent C=O bonds in CO2 linear shape NON-POLAR molecule

Bonding continuum Learning intention
Learn about the bonding continuum which stretches between pure covalent and ionic bonds, in terms of differences of electronegativity between bonding atoms.

Electronegativity Difference and Bond Type:
Difference Bond Example Covalent (nonpolar) H-H 0.0 Covalent (polar) H-Cl H20 0.7 Covalent (very polar) H-F 1.9  2.0 Ionic NaCl 2.1

The greater the difference in electronegativity the greater the polarity between two bonding atoms and the more ionic in character.

Bonding Continuum “Covalent compounds are formed by non-metals only”
IS NOT AN ABSOLUTE LAW! Some compounds break this rule….

Making Tin(IV)iodide First gently heat the tin and iodine in a small conical flask containing 10cm3 of tolulene on a hot plate. Then collect the yellow precipitate by filtration using Büchner filtration

Making Tin(IV)iodide Determine the melting point of the solid collected.

Making Tin(IV)iodide Melting point of tin(IV)iodide is 143oC.
Tin electronegativity of 1.8 Iodine has electronegativity of 2.6 Molecule contains polar covalent bonds, but the symmetry cancels out the dipoles, therfore only weak London’s forces so low melting an boiling point.

Titanium (IV) chloride
TiCl4 is a dense, colourless liquid. It is one of the rare transition metal halides that is a liquid at room temperature, This property reflects the fact that TiCl4 is ………….; that is, each TiCl4 ……………… is relatively …………… associated with its neighbours. Most metal chlorides are ionic. The attraction between the individual TiCl4 molecules is weak, these weak Van der Waals (intermolecular) interactions result in low melting and boiling points. TiCl4 is soluble in tolulene and dichloromethane, as are other non-polar species. covalent molecule weakly

TiCl4 Used in smoke grenades and for smoke screens

Hydrogen Bonding Learning intention
Learn about this strong type of intermolecular forces which exists between molecules containing N-H, O-H or F-H bonds.

Relating physical properties to intermolecular forces
Learning intention Learn how to explain differences in physical properties such as viscosity, melting point and boiling point in terms of differences in strength of intermolecular forces.

Intermolecular - Hydrogen Bonding
Consider the compounds formed between elements in group 4 of the Periodic table and hydrogen The group 4 hydrides are CH4, SiH4, GeH4, SnH4 They are all covalent molecular so have low melting points and boiling points.

The boiling point increases as you go down the group.

As you go down the group the central atom gets bigger.
There are more electrons so a greater chance of an uneven distribution of electrons within the atom. The London’s forces between the molecules gets stronger as you go down the group. More energy is needed to separate the molecules from each other.

Intermolecular – Hydrogen Bonding
A similar pattern would be expected in the other covalent molecular hydrides The group 5 hydrides NH3, PH3, AsH3 and SbH3 The group 6 hydrides H2O, H2S, H2Se and H2Te The group 7 hydrides HF, HCl, HBr and HI

NH3, has a higher boiling point than expected.

H2O has a higher boiling point than expected.

HF has a higher boiling point than expected.

H2O HF NH3 It is more difficult to separate NH3, H2O and HF molecules from each other than expected.

These compounds all have H atoms directly bonded to very electronegative atoms. In HF the H-F bond is polar covalent. The F has a much higher electronegativity than H. The pair of shared electrons in the covalent bond spend more time closer to the fluorine than the hydrogen. The H-F bond is polar. Hδ+ - Fδ-

The HF molecules can attract each other Hδ+ - Fδ- Hδ+ - Fδ- Hδ+ - Fδ- This is called hydrogen bonding. Hydrogen bonding is weak but is stronger than very weak London’s forces.

NH3 has H atoms directly bonded to very electronegative N atoms. There are Hydrogen bonds as well as London’s forces between the ammonia molecules. N- H+ H+ H+ H+ N- H+ N H+ H+ H+

H2O has H atoms directly bonded to very electronegative O atoms. There are Hydrogen bonds as well as London’s forces between the water molecules. O- H+ H+ H+O- H+ O H+ H+

Proteins consist of long chain atoms containing polar C=O and H-N bonds. Hydrogen bonds help give enzymes their shape.

Water O H Oxygen has 2 lone pairs of electrons which can form
- O H + + Oxygen has 2 lone pairs of electrons which can form a hydrogen bond with two hydrogen atoms. Each water molecule, in theory, could be surrounded by 4 hydrogen bonds

Hydrogen Bonding in Water

Water Density of water Water has its greatest density at a temperature of 4oC. When, as water cools further, the molecules start to move further apart, due to the hydrogen bonding, until a more open structure is formed at its freezing point. So ice floats!! New Higher Chemistry E Allan J Harris

Hydrogen bonding in ice
Hydrogen bonding in solid water gives rise to an open structure. This is why ice is less dense than liquid water.

Hydrogen bonding is also responsible for holding the two strands of nucleic acids together in DNA

Viscosity Viscosity is related to the molecular mass and the number of –OH present. Hydrogen bonding between the molecules will increase its viscosity. Density of water

Water Surface tension Water has a high surface tension. The molecules on the surface have in effect, hydrogen bonds. This has the effect of pulling the surface molecules closer together.

Behaviour in Electrical Fields

Nappies Cloth nappies cost between £100-£400 as opposed to disposable at £800-£1,200 for the 2.5 years of normal nappy use. 3 billion nappies are thrown away in the UK each year with 90% going to landfill. They can take up to 500 years to decompose. Disposables make up 4% of total household waste and up to 50% of that of families with one baby Disposable nappies use up to 5 times more energy to produce than cotton ones – that's including the washing process . Seven million trees are felled every year in Canada and Scandinavia to supply the pulp for disposables sold in the UK. 87

Sodium polyacrylate is a polymer with a molecular weight of over one million!
sodium carboxylate Chemical Background Groups called sodium carboxylate are attached along the backbone. 88

- Sodium poly(acrylate) absorbs 500 times its own mass of water. water
+ + Na+ -

- Sodium poly(acrylate) absorbs 500 times its own mass of water. + + +

Sodium poly(acrylate) absorbs 500 times its own mass of water.
- - - - - -

Predicting solubility from solute and solvent polarities
Learning intention Learn how the polarity of both the solute and solvent molecules influences solubility.

Solvent Action A liquid that a substance dissolves in is called a SOLVENT. Solvents can be either polar or non-polar molecules. Immiscible liquids do not mix, e.g. oil and water, however, non-polar liquids are miscible with each other. Polar solvents will usually dissolve polar molecules. Non-polar solvents will usually dissolve non-polar molecules. Water is a polar molecule so it is a polar solvent. Water has a polar covalent bonding between O and H. + + - H O - H +

Dissolving in Water Ionic Compound dissolving in water - + - +
- + - + - - + + - + Hydrated ions

Dissolving in Water

Dissolving in Water Pure Hydrogen chloride is polar covalent. When water is added it breaks to produce ions - + + - H Cl - + Cl- - + + - H+ Hydrated ions

Dissolving in Water Generally, covalent molecules are insoluble in water. However, small molecules like ethanol (C2H5OH), with a polar O-H functional group, will dissolve, Ethanol - + H2O + O H C -

Density change across a period
0.5 1 1.5 2 2.5 3 Sodium Magnesium Aluminium Density g/cm3 Silicon Phosphorus Sulphur Chlorine Argon Na Mg Al Si P S Cl Ar Na to Al the atom size decreases leading to greater packing in metal lattice. Si is a covalent network, tightly packed atoms in covalent lattice. P and S are covalent molecular solids with quite densely packed molecules. Cl and is a covalent molecular gas at room temperature. Ar and is a monomolecular gas at room temperature.

Covalent Hydrogen 40 Dipole-Dipole 30 Londons forces 1 to 20

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