TOPIC 1: Electron, Bonding and Structure Flashcards
(51 cards)
1
Q
Explain the different boiling points of NH3, F2 and BR2. (5)
A
- NH3 has hydrogen bonding b/w molecules
- F2 & Br2 have induced d.p-d.p forces
- intermolecular forces in Br2 stronger than in F2
^ bcs Br2 has more electrons
- Hydrogen bonding is stronger than the induced d.p-d.p forces in F2
2
Q
State and explain two anomalous properties of ice caused by hydrogen bonding. (4)
A
- ice is less dense than water
^ ice has an open lattice - ice has a relatively high melting point
^ hydrogen bonds are relatively strong
3
Q
Describe and explain the electrical conductivity of sodium oxide and sodium in their solid and molten states. (5)
A
- sodium conducts in solid and molten state
^ sodium has delocalised electrons in both states - sodium oxide conducts in molten state but not solid
^ molten sodium oxide has mobile ions
^ solid sodium oxide has immobile ions
4
Q
Why is Na3PO4 described as a salt of H3PO4. (1)
A
- the hydrogen ions are replaced by sodium ions
5
Q
Suggest why PH3 has a lower boiling point than NH3. (1)
A
- the intermolecular forces in PH3 are weaker
6
Q
Define: a dative covalent bond
A
- both electrons have been donated by one atom
7
Q
Explain why the student might expect the H-N-H bond angle to be larger in H3NBF3 than in NH3. (3)
A
- NH3 has 3 bonding pairs & 1 lone pair
of electrons - H3NBF3 has 4 bonding pairs of
electrons - lone pair on Nitrogen now becomes a
bonding pair
8
Q
Explain why there is a difference in the melting points of K, KBr and H2O. (6)
A
- In K, there’s electrostatic attraction
b/w cations & electrons - In KBr, there’s electrostatic attraction
b/w oppositely charged ions - KBr has ionic bonding
- In H2O, there’s hydrogen bonding b/w
molecules - strength of forces: (bonding)
ionic > metallic > hydrogen
9
Q
Define: electronegativity
A
- the ability of an atom to attract electrons
- in a covalent bond
10
Q
Molecules of BF3 contain polar bonds, but the molecules are non polar.
Explain why. (2)
A
- BF3 is symmetrical
- the dipoles cancel out
11
Q
Explain why a CH2Cl2 molecule is polar. (1)
A
- the dipoles do not cancel out
12
Q
SbCl3 molecules are polar.
Explain. (2)
A
- there’s a difference in
electronegativity - molecules are not symmetrical and
dipoles do not cancel out
13
Q
Describe how Van der Waals’ forces arise. (3)
A
- uneven distribution of electrons
- creates a temporary dipole
- causing induced dipoles in
neighbouring molecules
14
Q
Suggest why there are no other intermolecular forces in solid sulfur. (1)
A
- no polar bonds
15
Q
Explain why a molecule of SF6 has an octahedral shape. (2)
A
- sulfur has six bonded pairs
- electron pairs repel
16
Q
Explain why a solution of copper(II) nitrate conducts electricity. (1)
A
- ions are mobile
17
Q
State and explain the trend in the boiling points of chlorine, bromine and iodine. (3)
A
- b.p. increases down the group
- stronger intermolecular forces
- more energy needed to break these
intermolecular forces
18
Q
Why would astatine be expected to react similarly to other halogens. (1)
A
- same number of outermost electrons
19
Q
Explain why a solution of Copper(II) nitrate conducts electricity. (1)
A
- ions are mobile
20
Q
Explain how ionic bonding holds the particles together in an ionic compound. (1)
A
- strong electrostatic attraction between positively charged ions & negatively charged ions
21
Q
Describe the bonding process that forms lithium halides, with respect to electrons. (3)
A
- Li donates one electron
- The halogen accepts one electron
- Li forms a 1+ ion
& Cl forms a 1- ion
22
Q
Describe the type of bond that forms b/w the boron triflouride and ammonia molecules. (2)
A
- a dative covalent bond
- the N provides both electrons for the bond
23
Q
Describe the bonding in an (S2)2- ion. (1)
A
- 2 sulfurs attached by a covalent bond
^ each with a negative charge
24
Q
Explain why silver chloride has such a high melting point. (3)
A
- giant ionic structure
- strong electrostatic attraction b/w oppositely charged ions
^ which requires large amounts of energy to overcome
25
Explain why Al2Cl6 doesn't conduct when molten but Al2O3 does. (4)
Al2Cl6:
- covalent bonding
- therefore no free charges when molten, so not an electrical conductor
Al2O3:
- has an ionic lattic
- therefore there's mobile ions in the molten state
26
Describe the bonding in calcium carbide, CaC2. (4)
- (C2)2- is a diatomic ion
- triple bond b/w carbons
- Ca2+ ion
- ionic bond b/w Ca2+ and (C2)2-
27
Explain why ionic compounds have high melting and boiling points (2)
- Strong electrostatic attractions between oppositely charged ions.
- High temperature needed to provide sufficient energy to overcome the attractions.
28
Explain why ionic compounds dissolve in water (2)
- Polar water molecules are attracted towards ions on the surface of the ionic lattice. Water molecules bond to the ions, weakening and breaking them.
- Ions become surrounded by water molecules and break free from the lattice
29
Define: covalent bonding
The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
30
Define: orbital
A region around the nucleus that can hold up to two electrons with opposite spins
31
Define: isoelectronic
The same number of electrons
32
State the maximum number of electrons that can be held in each of the first three shells of an atom. (3)
1st = 2
2nd = 8
3rd = 18
33
State how many orbitals there are in a p-subshell and how the electrons are arranged if the subshell is full (3)
- 3 atomic orbitals
- each orbital has 2 electrons
- with opposite spins
34
Explain why both G1 & G2 are known as s-block elements (1)
The outer electrons are s-subshell electrons for all
35
Explain the strength of the ionic bond in sodium chloride (2)
- strong electrostatic attraction
- b/w Na+ and Cl- ions
36
Explain why it’s easier to use aqueous MgF2 in a lab setting than the molten version (2)
- High melting point
- as strong electrostatic attraction b/w oppositely charged ions in all directions
37
Compare the binding in phosphorus trichloride and ammonia (4)
- both undergo covalent bonding
- phosphorus atom and nitrogen atom are both central atoms
- each have 1 lone pair
- 3 shared pairs of electrons
38
Describe how the shape of the periodic table is linked to the electronic structure (6)
- elements in same period have same number of filled shells
- elements in same group have same number of valence electrons
- s-block includes G1 & G2 as it takes only 2 electrons to completely fill s-subshell
- p-block includes G3-0 as it takes 6 electrons to completely fill the p-subshell
- d-Block includes transition metals as it takes 10 electrons to completely fill d-subshell
- d-block begins on period 4 as d-orbitals have higher energy than 4s orbitals
39
Suggest why there are three possible p-subshells but only one possible s-subshell in an atom (3)
- s-orbitals are spherical so multiple subshells not possible
- p-orbitals are propeller shaped so 3 p-orbitals would not overlap significantly
40
Justify why hydrogen is positioned in the middle of the periodic table and not apart of G1 (3)
- very diff physical properties to G1 metals
- and chemical properties
- despite outer shell containing 1 s-subshell electron only
41
Explain why giant ionic structures have high melting points (2)
- strong electrostatic attraction between oppositely charged ions
- large amounts of energy required to overcome them
42
Explain why sodium bromide has a higher melting point than sodium and sodium iodide (6)
Stage 1:
- Na has metallic bonding & a giant structure
- there’s attraction b/w positive nucleus and delocalised electrons in Na
Stage 2:
- ionic bonding in NaBr & giant structure
- there’s attraction b/w + and - ions in NaBr
Stage 3:
- ionic bonds are stronger than metallic bonds
- stronger attraction b/w the + and - ions in NaBr than NaI
- since Br- ion is smaller than I- ion
43
Why is an arrow used to represent a
N ➡️ H bond? (1)
To show both electrons come from nitrogen
44
Suggest how methanol and methanethiol could be separated (1)
(Fractional) distillation
45
Define: ionic lattice (2)
- Repeating pattern
- of oppositely charged ions
46
State whether the following conduct electricity when solid or molten: (5)
aluminium
aluminium fluoride
boron tribromide
Aluminium:
- conducts in solid and molten stated
- has delocalised electrons
Aluminium flouride:
- conducts when molten because it has mobile ions
- doesn’t conduct when solid (ions fixed in position in an ionic lattice)
Boron tribromide:
- does not conduct in solid and molten states (no mobile ions)
47
Explain how the structure and bonding in bromine account for its relatively low melting point (3)
- forces b/w molecules
- which are induced dipole-dipole forces
- are weak so overcome easily by increased kinetic energy
48
Predict the type of structure and bonding of SO2 and MgO & explain the difference in their melting points (4)
- MgO = giant ionic
- SO2 = simple molecule
- ionic bonds in MgO much stronger than intermolecular bonds in SO2
- ionic bonds in MgO require more energy to overcome
49
Describe the relative energies of the 2s orbitals and each of the three 2p orbitals in a nitrogen atom (2)
- p-orbitals have greater energy than
s-orbitals
- three p-orbitals have equal energy
50
What are the 3 rules for filling electron orbitals? (3)
- Hund’s rule (each orbital must be singlely occupied before being paired up)
- Each orbital can hold up to 2 electrons with opposite spins
- Lowest energy orbitals must be filled first
51
Why is hyrogen H+
fills 1s orbital
