3.1.3 Bonding

Define what an ionic bond is and what is it caused by

a bonding between a metal and a non-metal due to the electrostatic force of attraction between oppositely charged ions formed by electron transfer

Give examples of ionic compounds

• sodium chloride

• calcium sulfide

• lead nitride

• aluminium iodide

Explain why an ionic compound is a solid at room temperature/ has high m.p and b.p.

• it is a giant ionic lattice structure

• which has strong electrostatic forces of attraction between oppositely charged ions

• in all directions

• that require lots of energy to overcome

Explain why ionic compounds can only conduct electricity when molten and not when solid

• when solid the ions are held in a fixed position and so are not free, thus cannot move around and carry electric charge

• when motlen the ions are free to move and carry electrical charge

Explain why ionic compounds are brittle and shatter easily

when a force / stress is applied ions of the same charge are brought side by side and so repel each other as like charges repel

What is the difference in strength of the ionic bond and the melting point between magnesium oxide and sodium chloride

Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges.

• E.g. MgO has a higher melting point than NaCl as the ions involved (Mg2+& O2- are smaller and have higher charges than those in NaCl , Na+& Cl- )

Draw the dot and cross diagram for sodium cyanide (NaCN)

Draw the dot and cross diagram for lithium hydroxide (LiOH)

Write down the charge of the following ions

• hydrogen

• lithium

• sodium

• potassium

• silver

• magnesium

• calcium

• zinc

• lead

• copper(II)

• iron (II)

• iron (III)

• aluminium

• fluoride

• chloride

• oxide

• sulfide

• nitride

• hydroxide

• nitrite

• nitrate

• hydrogencarbonate

• hydrogen sulfate

• cyanide

• sulfate

• sulfite

• carbonate

• phosphate

• ammonium

• hydronium

What is a covalent bond and what is a dative covalent bond? give and draw an example

covalent bond - a shared pair of electrons

dative /co-ordiante bond - contains a shared pair of electrons with both electrons donated by one atom. e.g the ammonium ion

The ion H₂F⁺ is formed when hydrogen fluoride gains a proton as shown in the equation

HF + H⁺ → H₂F+

Name the type of bond formed when HF reacts with H⁺ Explain how this bond is formed.

dative covalent bond

• the fluorine atom in HF donates a pair of electrons to the electron deficient H⁺ ion

What are alloptropes of carbon, give examples?

allotrope - a different structural form of the same element in the same physical shape

• diamond

• graphite

• graphene

State and and explain the structure and properties of diamond

• giant covalent structure (macromolecular)

• each carbon atom is covalently bonded to 4 other carbon atoms —→ has strong covalent bonds which require lots of energy to overcome hence has a high melting point

• has many strong covalent bonds and so is hard

• does not conduct electrcity —→ no delocalised electrons present to move through the structure and carry electric charge

• tetrahedral structure —→ the bond angle about each atom is 109.5 degrees

State and and explain the structure and properties of graphite

• giant covalent structure (macromolecular)

• each carbon atom is covalently bonded to 3 other carbon atoms —→ has strong covalent bonds that require lots of energy to overcome hence has a high m.p. point

• can conduct electricity —→ due to the presence of one delocalised electron per carbon atom that can move through the structure and carry electric charge

• soft and slippery, used as a lubricant —→ arranged in layers that have weak van Der Waal forces between the layers and so can slide past one another

• trigonal planar structure - the bond angle about each atom is 120 degrees

Describe the structure and bonding in graphite and explain why the melting point in graphite is so high

• a gaint covalent structure

• carbon atoms are arranged in layers

• which are connected by covalent bonds (within each layer)

• van der waal forces are present between layers

• there are many covalent bonds that need to be broken

State and explain the structure and properties of silicon dioxide

• giant covalent structure (macromolecular)

• each silicon atom is covalently bonded to 4 oxygen atoms, and each oxygen atom is covalently bonded to 2 silicon atoms

• has a high m.p. —→ due to the presence of many covalent bonds

What is a metallic bonding? and state the block in which metals are found

found in metals due to the strong electrostatic force of attraction between the positive ions and the negative sea of delocalised electrons

• metals are found in the d-block

Explain why nickel has a high melting point, refer to structure and bonding

• a giant metallic structure

• whereby there is a strong electrostatic force of attraction between the positive metal ions and the negative sea of delocalised electrons

• hence requires a lot of energy to overcome

Explain why metals are ductile (can be stretched into wires) or malleable

positive metal are identical and so ions are arranged in layers that can slide past each other

Explain the difference in melting point between Na ions and Mg ions

• Magnesium ions has a higher charge of +2 compared to Na of +1

• Magnesium has more electrons

• thus magnesium has a greater electrostatic force of attraction between the positive ion and electrons

• hence requires more energy to overcome

State the four crystal structures and give examples of each

• ionic - sodium chloride / sodium nitride

• giant covalent (macromolecular) - diamond / graphite

• metallic - aluminium / gallium

• molecular - iodine

State the forces or bonds found in each crystal structure

• ionic - electrostatic force of attraction between oppositely charged ions

• giant covalent - covalent bonds (in diamond and graphite) / van der waal forces between the layers in graphite

• metallic - electrostatic force of attraction between the positive metal ions and the negative sea of electrons

• molecular - weak van der waal forces between molecules so are easily broken

Define the term electronegativity

the relative ability for an atom to attract the bonding pair of electrons in a covalent bond

Describe and explain the trend of electronegativity in the periodic table as you go across a period

As you go across a period (from left to right), electronegativity increases

• this is because the nuclear charge (number of protons) increases

• but the shielding remains constant, yet the number of electrons on the outer shell increases

• therefore a smaller atomic radius = so bonding electrons are closer to the nucleus

• smaller atomic radius have less electron shielding

• less shielding = stronger attraction between nucleus and bonding pair of electrons

• therefore smaller atoms have a higher elctronegativity

Describe and explain the trend of electronegativity in the periodic table as you go down a group

• as you go down a group the atoms become bigger due to more electrons and so have fewer bonding electrons that are closer to the nucleus

• larger atoms have greater electron shielding

• greater electron shielding = weaker attraction between nucleus and bonding pair of electrons

• therefore bigger atoms have a lower electronegativity

How is a polar covalent bond produced?

The distribution of electron density in a covalent bond between elements with different electronegativities will be unsymmetrical —→ this causes a polar covalent bond

Explain why the C–Cl bond is polar. [2]

• Chlorine has a higher electronegativity (than carbon)

• so chlorine becomes δ- and carbon becomes δ+

Explain why the O-H bond in a methanol molecule is polar

oxygen is delta negative whilst hydrogen is delta positive

this causes electron density arround oxygen resulting in an unsymmetrical molecule

State the difference between a polar and non-polar molecule

polar molecules:

• Polar molecules have a permanent dipole.

• Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.

non-polar molecules:

• A symmetrical molecule will not be polar even if individual bonds within the molecular are polar.

• The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule.

Give the shape and angle of molecules with:

• 2 bonding pairs

• no lone electron pairs

give examples

• a linear shape

• bond angle is 180^o

examples: carbon dioxide CO₂, berrylium chloride BeCl₂

Give the shape, angle and diagram of molecules with:

• 3 bonding pairs

• no lone electron pairs

give examples

• trigonal planar shape

• bond angle is 120^o

examples: BF₃, SO₃

Give the shape, angle and diagram of the molecule with:

• 4 bonding pairs

• no lone electron pairs

give examples

• tetrahedral shape

• bond angle is 109.5^o

examples - methane CH₄

Give the shape, angle and diagram of the molecule with:

• 5 bonding pairs

• no lone electron pairs

give examples

• trigonal bypyramidal shape

• bond angle can be 90^o or 120^o

examples: PCl₅

Give the shape, angle and diagram of molecules with:

• 6 bonding pairs

• no lone pairs of electrons

give examples

• octahedral shape

• the bond angle is 90^o

examples: SF₆

Give the shape, angle and diagram of molecules with:

• 2 bonding pairs of electrons

• 2 lone pairs of electrons

give examples

• bent/ v-shape

• bond angle is 104.5^o

examples: H₂O

Give the shape, angle and diagram of molecules with:

• three bonding pairs

• 1 lone pair of electron

give examples

• trigonal pyramidal shape

• bond angle is 107^o

examples: NH₃

Give the shape, angle and diagram of molecules with:

• 4 bonding pairs

• 1 lone pair of electron

give examples

• see-saw shape

• bond angle can be <90^o and <120^o

examples: SF₄

Give the shape, angle and diagram of molecules with:

• 3 bonding pairs

• 2 lone pairs of electrons

give examples

• T-shape

• the bond angle is 90^o

examples: CIF₃

Give the shape, angle and diagram of molecules with:

• 4 bonding pairs

• 2 lone pairs of electrons

give examples

• square planar

• bond angle is 90^o

examples: XeF₄

Top Tip: how to explain shape in an exam question

State the three types of intermolecular forces and state which one is the strongest and which one is the weakest

• van der waal forces (also known as induced dipole-dipole or London or dispersion forces)

• hydrogen bonding

• permanent dipole-dipole

strongest - hydrogen bonding

weakest - van der waal forces

There are van der Waals forces between non-polar molecules.

Explain what causes these forces.

• there is a random movement of electrons in one molecule which creates a dipole

• this induces a dipole in another molecule

• Thus there is a temporary attraction between δ+ and δ- in different molecules

Describe and explain hydrogen bonding

• a type of permanent dipole-dipole force

• which is the attraction between the lone pair of electrons on either a O,N,F atom and δ+ H of OH on another molecule

what is meant by the term permanent dipole?

polar molecules have permanent dipoles

• when a molecule has a δ+ region and a δ- region, this creates a polar molecule

• the δ+ end of a molecule is attracted to the δ- of another molecule

Suggest, in terms of the intermolecular forces for each compound, why CBr₄ has a higher boiling point than CHBr₃ (3)

• CBr₄ has van Der Waal forces between it’s molecules

• CHBr₃ has van Der Waal forces and permanent dipole-dipole forces

• yet the van Der Waal forces in CBr₄ are stronger than the combined van Der Waal and permanent dipole dipole forces in CHBr₃

• As it has more electrons

Explain why chlorine has a lower boiling point than bromine (2)

Bromine has a greater number of electrons than chlorine

therefore bromine has greater and stronger van Der Waals forces between bromine molecules than chlorine molecules hence require more energy to overcome

Which would be a more suitable solution to dissolve iodine crystals, I₂ in, water H₂O or carbontetrachloride CCI₄. Explain your answer

• iodine would need to be dissolved in carbon tetrachloride CCI₄

• as iodine is a non-polar molecule therefore can only be dissolved in non-polar molecules, H₂O is a polar molecule whereas CCI₄ is a non-polar molecule

Explain the anomalous nature of ice

generally the solid state of a substance is denser than it’s liquid state, however ice is less dense than liquid water

• this is because ice is an open crystalline structure whereby there are gaps in the structure of ice because water molecules are held further apart