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chemistry > 2.2.2 Bonding and structure > Flashcards

2.2.2 Bonding and structure Flashcards

1
Q

Define ionic bonding.

A

Electrostatic attraction between oppositely charged ions.

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2
Q

Why do ionic compounds form giant ionic lattices?

A

Because oppositely charged ions are strongly attracted in all directions.

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3
Q

Explain why NaCl forms a giant ionic lattice.

A

Na+ and Cl- ions attract each other strongly in all directions

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4
Q

Explain the high melting point of ionic compounds.

A

Strong electrostatic forces between ions require large amounts of energy to break.

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5
Q

Explain the solubility of ionic compounds in water.

A

Ionic compounds are soluble as polar water molecules attract and surround ions.

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6
Q

Explain why ionic compounds conduct electricity when molten or aqueous.

A

Ions are free to move and carry charge.

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7
Q

Define covalent bonding.

A

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

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8
Q

What is a single covalent bond?

A

A shared pair of electrons between two atoms.

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9
Q

What is a multiple covalent bond?

A

Two or more shared pairs of electrons between atoms (e.g.

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10
Q

What is a dative covalent (coordinate) bond?

A

A shared pair of electrons donated by one atom.

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11
Q

What does average bond enthalpy measure?

A

Strength of a covalent bond.

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12
Q

State the relationship between bond enthalpy and bond strength.

A

Larger average bond enthalpy = stronger covalent bond.

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13
Q

What is electron pair repulsion theory?

A

Electron pairs around a central atom repel to positions of minimum repulsion.

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14
Q

Explain why lone pairs repel more than bonded pairs.

A

Lone pairs are closer to the nucleus and occupy more space.

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15
Q

State the bond angle in CH4.

A

109.5° (tetrahedral).

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16
Q

State the bond angle in NH3.

A

107° (pyramidal).

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17
Q

State the bond angle in H2O.

A

104.5° (non-linear).

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18
Q

State the shape and bond angle of a molecule with 2 bonding pairs.

A

Linear

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19
Q

State the shape and bond angle of a molecule with 3 bonding pairs.

A

Trigonal planar

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20
Q

State the shape and bond angle of a molecule with 4 bonding pairs.

A

Tetrahedral

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21
Q

State the shape and bond angle of a molecule with 3 bonding pairs and 1 lone pair.

22
Q

State the shape and bond angle of a molecule with 2 bonding pairs and 2 lone pairs.

A

Non-linear

23
Q

State the shape and bond angle of a molecule with 6 bonding pairs.

A

Octahedral

24
Q

Define electronegativity.

A

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

25
State the trend in electronegativity across a period.
Electronegativity increases towards fluorine.
26
What is a polar bond?
A covalent bond with an unequal distribution of electrons due to different electronegativities.
27
Define permanent dipole.
Separation of charge within a molecule due to polar bonds.
28
When is a molecule polar?
When it has polar bonds and the dipoles do not cancel due to shape.
29
Why is CO2 non-polar despite having polar bonds?
Dipoles cancel due to the linear shape.
30
Why is H2O polar?
Dipoles do not cancel due to bent shape.
31
Name two types of van der Waals' forces.
Permanent dipole–dipole and induced dipole–dipole interactions.
32
Another name for induced dipole–dipole interactions?
London (dispersion) forces.
33
When does hydrogen bonding occur?
Between a hydrogen atom bonded to N
34
Explain why ice is less dense than water.
Hydrogen bonds hold molecules in an open lattice in ice.
35
Explain why water has high melting and boiling points.
Hydrogen bonding requires more energy to overcome.
36
Describe a simple molecular lattice.
Covalently bonded molecules held together by intermolecular forces.
37
Explain why simple molecular compounds have low melting points.
Weak intermolecular forces require little energy to overcome.
38
Explain why simple molecular compounds do not conduct electricity.
No free ions or electrons to carry charge.
39
Describe what is meant by the term ionic lattice.
Repeating pattern of oppositely charged ions ✓
40
Explain what is meant by the term electronegativity.
The ability of an atom to attract electrons (in a covalent bond) ✓
41
Explain what is meant by ionic bonding.
Electrostatic attraction between oppositely charged ions ✓
42
Describe the bonding in magnesium and silicon.
Answer: Magnesium: metallic bonding (delocalised electrons and cations) ✓ Silicon: covalent bonding (shared pairs between atoms) ✓
43
Name the structure of solid chlorine and bromine.
Simple molecular lattice ✓
44
Predict the shape of SbCl₃.
Trigonal pyramidal ✓ (Sb has) three bonding pairs and one lone par of electrons ✓ Pairs of electrons repel ✓
45
What is the shape of COCl₂?
Trigonal planar ✓
46
Explain the boiling points of NH₃, F₂, and Br₂.
NH₃: hydrogen bonding ✓ F₂ and Br₂: van der Waals’ forces (Br₂ has more electrons) ✓
47
Why does H₂S have a lower boiling point than H₂O?
H₂O has hydrogen bonding; H₂S has weaker London forces because H2S has fewer electrons ✓ Hydrodgen bonding is stronger so requires more energy to break forces.
48
Describe the relative energies of the 2s orbital and each of the three 2p orbitals in a nitrogen atom.
Energy Difference: "p-orbitals have greater energy than the s-orbital in the same shell" ✓ (ALLOW reverse argument: "2s is lower energy than 2p") Degeneracy of p-orbitals: "(All three) p-orbitals have equal energy" ✓ (ALLOW: "pₓ, pᵧ, p_z are degenerate")
49
SbCl3 molecules are polar. Explain why.
There is a difference in electronegativities (between Sb and Cl) OR (Sb-Cl) bonds are polar OR have a dipole OR Dipoles seen on the diagram ✔ The molecule is not symmetrical AND dipoles do not cancel ✔
50
State and explain two anomalous properties of ice caused by hydrogen bonding.
Property 1 Ice is less dense than water ✔ Explanation 1 The molecules in ice are held apart by hydrogen bonds ✔ OR ice has an open lattice OR structure Property 2 Ice has a relatively high melting point ✔ Explanation 2 Hydrogen bonds are relatively strong OR Hydrogen bonds are stronger (than other intermolecular attractions or forces) OR More energy is needed to overcome hydrogen bonding. ✔
51
Describe and explain the electrical conductivity of sodium oxide, Na2O, and sodium in their solid and molten states.
M1: Sodium conducts in the solid and molten states ✔ Reason for conductivity of Na mark M2: Sodium has delocalised electrons (in both solid and liquid state) ✔ Conductivity of Na2O mark M3: Na2O conducts when molten and not when solid ✔ Reason for conductivity of Na2O marks M4: Molten Na2O has ions which are mobile ✔ M5: Solid Na2O has ions which are fixed (in position) OR ions are held (in position) OR ions are not mobile AND in an (ionic) lattice OR structure ✔

chemistry (14 decks)

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