C5-6. Atoms, Structure And Bonding

Question 1

What is the difference between relative atomic mass and relative isotopic mass?

Answer 1

• Relative atomic mass is the weighted mean mass of an atom.
• Relative isotopic mass relates to the weighted mean mass of an isotope of the element in question.
• Both of these are given in comparison to 1/12th the mass of a carbon-12 atom.

Question 2

What characterises ionic bonding?

Answer 2

• The transfer of electrons between the highest energy levels
• Production of ions with opposite charges and stable electronic configurations isoelectric to noble gases
• The ionic bond itself is the electrostatic attraction between cations and anions

Question 3

What characterises covalent bonding?

Answer 3

• Atoms share electrons from their highest energy levels
• Covalent bonds are the electrostatic attractions between nuclei and shared electrons
• Multiple pairs of electrons can be shared
• Sometimes involves expanding the octet

Question 4

How do we define the number of electrons a shell holds?

Answer 4

Max no. electrons in a given shell = 2n^2, where n is the principal quantum number.

Question 5

How are shells divided into sub shells?

Answer 5

• Subshells are labelled s, p and d in relation to the orbitals they contain.
• The order of energy of electrons increases s<p<d.
• The order of subshell energies overlaps as energy increases, hence explaining why 4s precedes 3d.

Question 6

How do we define orbitals?

Answer 6

These are areas about nuclei is which up to two electrons can be present, with opposite spin arrangements.

Question 7

What are the common shapes of orbitals?

Answer 7

• S-orbitals have a spherical shape.
• P-orbitals have a dumbbell-like shape.

Question 8

How do orbitals fill, by Hund’s rule?

Answer 8

• Electrons undergo spin-pair repulsion. When they spin, a magnetic field is induced causing repulsion.
• To prevent this electrons initially occupy separate orbitals as they are added.
• They then begin to pair up.

Question 9

Using the Aufbau principle, what happens to electronic configuration during ion formation?

Answer 9

• Generally, electrons are gained and lost to/from the highest-energy outermost subshell.
• The transition metals have an exception to the rule - the 4s subshell is filled AND emptied first.
• Also note that Cr and Cu have unusual configurations - 3d5 and 3d10 are more stable.

Question 10

How do we write out electronic configurations of atoms?

Answer 10

We write from the lowest-energy subshell upwards, with a superscript number describing the number of electrons which fill the sub shell. This is based on the number of orbitals within an s, p or d sub shell - 2, 3 and 5 respectively.

Question 11

How does dative covalent bonding work?

Answer 11

• Molecules with lone electron pairs donate these to form a bond with an electron-deficient ion.
• So both electrons in the covalent bond come from one ion.

Question 12

How do we describe metallic bonding?

Answer 12

• The outer shell electrons of a metal atom are donated to form a pool of delocalised electrons.
• This leads to strong electrostatic attractions between the oppositely-charged ions.
• The donated electrons are now delocalised so can move freely through the lattice structure, carrying a charge.

Question 13

What principle do we use to determine molecule shapes?

Answer 13

• We use valence shell electron pair repulsion theory. This essentially means that, pairs of electrons repel to move as far away as possible.
• Lone pairs have a greater repulsive force than bonded pairs.

Question 14

What is electronegativity?

Answer 14

• The ability of an atom to attract a bonded pair of electrons in a covalent bond.
• This exhibits an increase across the period and a decrease down the group.

Question 15

Explain the difference between polar and non-polar bonds.

Answer 15

• Polar bonds are formed with atoms of varying electronegativities. Where one atom is more electronegative, electrons are attracted more - inducing a dipole.
• Non-polar bonds form when atoms of equal electronegativity interact, so the electron pair is equally distributed.

Question 16

Are molecules of the same element - such as H2, N2, etc. - polar or non-polar?

Answer 16

These are non-polar, because the electronegativities of the atoms involved are the same. That means that the electrons in their overlapping orbitals are shared evenly and equally between them; so the bond between them is not polarised and and therefore has no dipole.

Question 17

Explain permanent-dipole, permanent-dipole forces.

Answer 17

• Two molecules with permanent dipoles attract one another where their constituent atoms with opposite charges attract one another.
• Therefore an intermolecular force occurs - weakly, since no electrons are actually transferred.

Question 18

Explain how London forces occur and work.

Answer 18

Random electron motion in all molecules causes the formation of an instantaneous dipole moment. This induces a dipole moment in another neighbouring molecule, causing a slight attraction between the molecules.

Question 19

Explain permanent-dipole, induced-dipole forces.

Answer 19

• Molecules with a permanent dipole (polar molecules) can induce a dipole moment in another neighbouring molecule supplementary to the pre-existing London forces.
• This causes a weak induced dipole, and therefore a weak attraction between the two molecules.
• The resulting induced dipole then induces dipoles on neighbouring molecules - hence the attraction spreads throughout the system.

Question 20

Explain the relationship between number of electrons and strength of London forces.

Answer 20

• As the number of electrons increases, the greater the fluctuation in the movements of these electrons.
• This causes the difference in the electronegative and electropositive ends to increase.
• Therefore the London forces become stronger.

Question 21

What is hydrogen bonding?

Answer 21

• This occurs between a highly electronegative atom with a lone pair of electrons and the the lone pair of a hydrogen atom in another molecule.
• This is the strongest type of intermolecular bonding.

Question 22

Describe the effect of hydrogen bonding on the physical properties of water.

Answer 22

• Solid ice is less dense than liquid water. This occurs because the hydrogen bonds arrange the water molecules in a more open lattice structure, so the water molecules in ice are further apart than in water.
• Water has a high melting/boiling point. This is because hydrogen bonds require extra forces to overcome as they are stronger.

Question 23

Describe a giant covalent structure.

Answer 23

• Bonds between atoms continues indefinitely, there are no molecules.
• Strong covalent bond between all atoms. This gives rise to a high melting point and boiling point.
• For example, the allotropes of carbon.

Question 24

Describe a simple molecular structure.

Answer 24

• Within molecules there are strong covalent bonds.
• Between these molecules we find much weaker intermolecular bonds, which require less energy to overcome.
• This gives rise to a much lower melting/boiling point.

Question 25

How many orbitals are there in each type of subshell?

Answer 25

• s: 1 orbital
• p: 3 orbitals
• d: 5 orbitals
• f: 7 orbitals

Question 26

What are the bond angles of the following structures?
- Linear
- Trigonal planar
- Tetrahedral
- Octahedral

Answer 26

• 180 degrees
• 120 degrees
• 109.5 degrees
• 90 degrees

Question 27

Why does electronegativity decrease down the group?

Answer 27

• Atomic radius increases so the outermost electrons are further away from the nucleus
• Shielding increases, increasing outwards repulsion
• So the overall attraction between the nucleus and bonded electrons decreases

Question 28

Why does electronegativity increase across the period?

Answer 28

• More protons in the nucleus, so nuclear charge increases
• Reduced atomic radius due to increase in no. of electrons
• So the overall attraction between nucleus and bonded electrons increases

Question 29

How do we determine the strength of metallic bonding?

Answer 29

• A larger ion has reduced charge density, so the electrostatic attraction to the pool of electrons decreases.
• A greater charge on the metal ion leads to increased attraction.
• A greater number of delocalised electrons causes increased attraction.

Question 30

Why does boiling point increase down group 7?

Answer 30

• As we go down the group, the number of electrons in an atom increases.
• This causes an increase in London forces, so the difference in the delta-positive and delta-negative ends of a molecule increases.
• Therefore an instantaneous dipole occurs, inducing a dipole in another molecule and leading to attraction.

Question 31

Why do water, hydrogen fluoride and ammonia all have higher boiling points than expected?

Answer 31

• They all use hydrogen bonding, which is much stronger than other forms of intermolecular bond.
• Therefore more thermal energy is required to overcome these forces of attraction, than if only London forces were present.

Question 32

What do we mean if we say something is isoelectric to another?

Answer 32

• They have the same electronic configuration.
• This happens for example during the formation of ions; the electronic configuration of an ion is isoelectric to that of the precedent noble gas.

Question 33

What is the difference between mass number and relative atomic mass?

Answer 33

• Mass number refers to the sum of protons and neutrons contained within a specific isotope of an element.
• Relative atomic mass is a weighted mean mass of all of the isotopes of a given element, in comparison to 1/12th the mass of carbon-12.

Question 34

Describe and explain the structure and bonding in a giant ionic lattice.

Answer 34

• Giant structures of oppositely-charged ions with an indefinite number of constituent atoms.
• Held together by strong electrostatic forces of attraction due to the transfer of electrons between highest energy-levels.
• This gives them a high melting/boiling point.

Question 35

Explain the following properties of ionic structures…
- Melting/boiling point
- Conductivity

Answer 35

• Melting/boiling point is high, due to the strong electrostatic forces of attraction between oppositely charged ions
• They are only conductive in the liquid/molten state; ions are free to move when molten, but cannot do so in solid state

Question 36

Explain the following properties of covalent structures…
- Melting/boiling point
- Conductivity

Answer 36

• Melting/boiling point is low, due to the weak intermolecular forces between molecules
• They are unconductive, because there are no mobile charge carriers available to carry a current

Question 37

Explain the following properties of metallic structures…
- Melting/boiling point
- Conductivity

Answer 37

• Melting/boiling point is high, due to the strong electrostatic forces of attraction between +ve metal cations and the -ve electron pool.
• Conductivity is high, due to the pool of delocalised electrons

Question 38

Why does the 4s sub-shell fill before the 3d sub-shell?

Answer 38

This occurs because the energy levels of sub-shells begin to overlap as energy level increases, to the point of which the 3d sub-shell is actually higher in en energy than the 4s sub shell. Lower-energy sub shells fill before higher-energy ones, so 4s fills prior to 3d.

Question 39

Why are non-polar molecules soluble in non-polar solvents?

Answer 39

Intermolecular forces form between the non-polar molecules and solvent, which weakens the intermolecular forces within the simple molecular lattice of the non-polar molecule. These then break down, causing the compound to dissolve.

Question 40

Why are polar molecules soluble in polar solvents?

Answer 40

This process is very simple, and resembles the process of dissolving an ionic lattice. The polar molecule and polar solvent both have permanent dipoles, which means that the two can attract one another.

Question 41

Why are non-polar molecules insoluble in polar solvents, and vice versa?

Answer 41

There is little interaction between the non-polar molecules in the lattice and the polar solvent molecules. Therefore the intermolecular bonding within the lattice is too strong to be overcome by these considerably weaker forces of attraction.

Question 42

How and why do ionic compounds dissolve in polar solvents?

Answer 42

• The polar water molecules have an overall dipole - they have electropositive and an electronegative sides.
• These are attracted to charged particles (anions and cations) within the ionic lattice, therefore overcoming the electrostatic forces of attraction within the lattice itself.
• This leads to each ion being surrounded by the solution.