2.2.2 Bonding and structure

Question 1

Define ionic bonding

What does a dot and cross diagram show ?

Answer 1

the electrostatic attraction between positive and negative ions

where the electrons in a bond come from

Question 2

Define metallic bonding

Answer 2

the** strong electrostatic attraction **between a **lattice of positive ions **and a sea of delocalised electrons

Question 3

What is a giant ionic lattice ?

Answer 3

regular structure of oppositely charged ions that are strongly attracted in all directions

Question 4

Explain how the structure of ionic compounds impacts melting and boiling points ?

Answer 4

High melting/boiling points
* giant ionic lattice is held my many electrostatic forces of attraction
* these need a lot of energy to be overcome
* so high melting/boiling points

Ionic compounds are typically solid at room temp
Strength depends on size and charge
**-Smaller ion/higher charge = stronger ionic bonds
**

Question 5

Explain how the structure of ionic compounds impacts electrical conductivity

Answer 5

• Do not conduct when solid as ions are held in fixed positions so no charge is conducted
• Conduct when molten or dissolved as ions are **mobile **and can carry a charge

Question 6

Explain how the structure of ionic compounds impacts solubility

What is the exception to this rule ?

Answer 6

Soluble in polar solvents such as water
* The slight negative oxygen will be attracted to positive ions
* The slight positive hydrogens will be attracted to the negative ions
Ions are attracted and surrounded.
This causes the ionic lattice to break down and the ionic compound dissolves

If the ions have large charges the ionic attraction may be too strong for the water to break down the lattice structure (solubility decreases as ionic charge icreases)

Question 7

Define covalent bonding

Answer 7

strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

It is the overlap of atomic orbitals - The attraction is localised, so only goes in two directions

Question 8

What does the dot and cross diagram show in covalent bonding ?

Answer 8

shows the electrons transferred

Question 9

Why is BF^3 an exception to the covalent bonding rules ?

Answer 9

Boron is **electron deficient **as it only has 6 electrons in its outer shell
This is allowed because:
* as long as the central atom maximises the covalent bonds it can form it will be stable
* Boron is in group 3
* So it forms a maximum of 3 covalent bonds

Question 10

Why is SF^6 (sulfur hexofluoride) an exception to the covalent bonding rules ?

Answer 10

It is has an expanded octet
* outer shell has 12 electrons
* It is in period 3 so has access to the third shell
* This can hold 18 electrons, and the 6 unpaired are available for bonding

Question 11

What are the three types of covalent bond ?

Answer 11

• single
• Double/triple
• Dative (coordinate)

Question 12

Define dative bond

Answer 12

a covalent bond in which the shared pair of electrons comes from only one of the bonding atoms
-The shared pair was originally a LONE PAIR
- Movement of the lone pair is shown by an arrow

Question 13

Define average bond enthalpy

Answer 13

measurement of covalent bond strength
* larger value = stronger covalent bond

Question 14

What are the 4 rules in electron pair repulsion theory ?

Answer 14

• **electron pairs **in the valence shell repel as far as possible
• The greater the number of electron pairs = the smaller the bond angles
• multiple bonds have the same repulsion to single bonds
• Lone pairs repel more strongly than bonded pairs (which reduced the bond angles)

Question 15

What do the 3 wedges represent in molecule shape diagrams ?

Answer 15

• solid line = bond on the plane
• solid wedge = comes out of the plane
• dotted wedge = goes into the plane

Question 16

Give the bonded/lone pairs and the angle in linear molecules

Answer 16

2 bonded pairs
No lone pairs
180 degrees

eg Ab^2

Question 17

Give the bonded/lone pairs and angle in trigonal planar molecules

Answer 17

3 bonded pairs
No lone pairs
120 degrees

eg AB^3

Question 18

Give the bonded/lone pairs and angle in a tetrahedral molecule

Answer 18

4 bonded pairs
No lone pairs
109.5 degrees

eg AB^4
Shape is now 3D as one bond goes into the plane and one goes out

Question 19

Give the bonded/lone pairs and angle in a trigonal bipyramidal molecule

Answer 19

5 bonded pairs
No lone pairs
There are now two bond angles
* 3 bonds on the same plane = 90 degree
* 2 bonds going in and out = 120 degree

AB^5

Question 20

Give the bonded/lone pairs and angles in an octahedral molecule

Answer 20

6 bonded pairs
No lone pairs
90 degree angles

Question 21

Why do lone pairs have greater repulsion ?
How much do they reduce the bond angle by ?

Show the relative repulsion of the different pairs

Answer 21

slightly closer to the central atom and occupies more space
-reduce bond angle by 2.5 degrees

Bonded-bonded < bonded-lone < lone-lone

Question 22

Give the bonding/lone pairs and angle of a pyramidal molecule

Answer 22

4 bonding pairs
1 lone pair
-lone pairs repel more than the bonding pairs, so reduce the bond angles by 2.5
Angle = 109.5 - 2.5 = 107

examples = CH^4 (methane), NH^4 (ammonia)

Question 23

Give the bonding/lone pairs and angle of a non-linear molecule

Answer 23

2 bonding pairs
2 lone pairs which reduces the bond angles by 5 degrees
Angle = 109.5 - 5 = 104.5 degrees

Example = H^2O (water)

Question 24

How should you explain the shape of a molecule in an exam question ?

Answer 24

1. state the number of bonding/lone pairs
2. state that electron pairs repel
3. if no lone pairs = electrons repel equally
4. if lone pairs = lone pairs repel more than bonding pairs
5. state the actual shape and bond angle

Question 25

Define electronegativity

Answer 25

ability of an atom to **attract** the **bonding electrons** in a covalent bond

Question 26

What does a high **pauling value** show ?

Answer 26

a high electronegativity (increases up and across the periodic table) * fluorine, oxygen and nitrogen are the most electronegative elements * Group one metals are the least electronegative

Question 27

What is the electronagativity difference of an ionic bond ?

Answer 27

high as both of the ions have full charges ## Footnote - one bonded atom will have a much** greater attraction f**or the shared pair than the other - so will **attract** these electrons causing them to be **transfered **(and so forms an ionic bond)

Question 28

What is the electronegativity difference of a polar covalent bond ?

Answer 28

moderate as both the atoms have a partial charge

Question 29

What is the electronegativity difference of a non-polar covalent bond ?

Answer 29

low as the atoms are uncharged

Question 30

How does electronegativity change across a period ?

Answer 30

increases * as the number of protons increases * and the atomic radius decreases * so the attraction of the bonding atoms will be greater

Question 31

How does electronegativity change down a group ?

Answer 31

decreases * distance between the nucleus and outer electrons increases * shielding of inner electrons increases * so the attraction of the bonding atoms will be reduced

Question 32

Define non-polar bond

Answer 32

bonded electron pair is **shared equally between the bonded atoms** * occurs when bonded atoms are THE SAME or have the **SAME ELECTRONEGATIVITY** ## Footnote eg H-H or Cl-Cl

Question 33

Define polar bond

Answer 33

bonded electron pair is **shared unequally **between the bonded atoms * occurs when bonded atoms have **DIFFERENT electronegativity **values ## Footnote eg H-Cl - chlorine is more electronegative - so has a greater attraction for the bonded pair - causing polar bond (there is a slight negative and slight positive charge - polarised) - shown by a delta sign

Question 34

Define dipole

Answer 34

the separation of opposite charges

Question 35

Define permanent dipole

Answer 35

the difference in electronegativity of two atoms in a **polar bond** (dipole is unchanging)

Question 36

What type of dipole is produced when atoms are arranged: * symmetrically * unevenly

Answer 36

Symmetrical - dipoles cancel - molecule has NO overall dipole so is NON-POLAR Uneven -dipoles don't fully cancel - molecule has an overall dipole - so is POLAR

Question 37

Define intermolecular forces

Answer 37

weak attractive forces between dipoles of DIFFERENT molecules (less energy needed to break then compared to ionic or covalent) ## Footnote Molecule = a small group of covalently bonded atoms

Question 38

What is an induced dipole-dipole interaction ? ## Footnote aka London forces weakest intermolecular force

Answer 38

attraction between induced dipoles - exist between ALL MOLECULES

Question 39

How do induced dipole-dipole interactions arise ?

Answer 39

* electrons are **constantly moving** - electron density shifts * at any moment** electron density** could be** uneven** * forming an **instantaneous dipole** * Electron density on one end **repels,** causing an **induced dipole **on the neighbour molecule * This causes a chain reaction (induced dipole induces further dipoles) ## Footnote They are TEMPORARY

Question 40

How does the number of electrons in an induced dipole-dipole interaction impact is properties ?

Answer 40

* more electrons * bigger induced dipoles (as there is a greater difference in electronegativity) * greater dipole-dipole interactions * **stronger attractive forces** between molecules so.... * more energy needed to overcome the intermolecular forces, increasing the boiling points ## Footnote Larger surface area also increases forces as there is a more exposed electron cloud

Question 41

What is a permanent dipole-dipole interaction ?

Answer 41

weak electrostatic forces of attraction between molecules with a **permanent dipole **(significant electronegativity difference - polar molecules) ## Footnote These happen in addition to induced dipole-dipole forces - Extra energy is needed to break this force, so molecules with both induced AND permanent interactions have a higher boiling point

Question 42

What is hydrogen bonding ?

Answer 42

specific type of permanent dipole-dipole interaction between molecules containing: * an electronegative element with a long pair of electrons **(oxygen, nitrogen, fluorine)** * a **hydrogen **bonded to the en. element

Question 43

How does a hydrogen bond form ? ## Footnote Strongest intermolecular force

Answer 43

* a **dipole forms **between the electronegative element and the hydrogen *This bond is so polorised that a HYDROGEN BOND forms between the hydrogen and the lone pair ## Footnote molecules with hydrogen bonding usually have -OH or -NH groups

Question 44

Why does water have a higher boiling point than expected ?

Answer 44

* large number of **hydrogen bonds** * a lot more energy is needed to break these bonds * increasing the melting/boiling point

Question 45

Why does water have a high surface tension ?

Answer 45

* many hydrogen bonds * give water a tight structure * this creates surface tension

Question 46

Why is ice less dense than water?

Answer 46

* hydrogen bonds in ice hold the molecules in the **open lattice** further apart * hydrogen bonds are long, so volume increases and **density decreases** * so the i**ce floats** ## Footnote when ice melts, the hydrogen bonds are broken and the ice lattice collapses - molecules move closer together so water is more dense

Question 47

Define simple molecular lattice

Answer 47

**covalently bonded molecules** attracted by **weak intermolecular forces** ## Footnote The atoms WITHIN each molecule are held by **strong covalent bonds**

Question 48

Why do simple molecules have low melting and boiling points ?

Answer 48

* intermolecular forces are weak * minimal energy to break * low melting and boiling points (liquids and gases at room temp) ## Footnote larger molecule = greater intermolecular forces

Question 49

Why do simple molecules not conduct electricity ?

Answer 49

* no charged particles * so cannot carry charge

Question 50

How does solubility vary in simple molecules?

Answer 50

In polar solvents... Non-polar molecule * only form weak induced dipole-dipole interactions * the intermolecular forces in the polar solvent are too strong to be broken * so molecule is insoluble in polar solvents Polar molecule * can form stronger hydrogen bonds * these interact with the dipoles in the water * so can dissolve ## Footnote Polar molecules = soluble in polar solvents (water, ethanol) Non-polar molecules = soluble in non-polar solvents (alkanes and aromantics)