Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

CHEMISTRY Physical > Bonding > Flashcards

Bonding Flashcards

You may prefer our related Brainscape-certified flashcards:
AP® Chemistry flashcards Periodic Table flashcards Chemistry 101 flashcards Organic Chemistry 101 flashcards
1
Q

What is ionic bonding? (2)

A

⇒ An ionic bond is the electrostatic force of attraction between oppositely charged ions (1)
⇒ Formed by electron transfer. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

State the two conditions for ionic bonding, explain your answer. (3)

A

⇒ Only occurs between metals and non-metals (1)
⇒ Electrons are transferred from metal atoms to non-metal atoms. (1)
⇒ Metals with 1,2 or 3 electrons in their shells find it easier to lose electrons as it requires energy than gaining electrons (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

What are some common formulas to remember? (5)

A 
⇒ SO₄²⁻ (1) (sulfate)
⇒ OH⁻ (1) (Hydroxide)
⇒ NO₃⁻ (1) (Nitrate)
⇒ CO₃²⁻ (1) (Carbonate)
⇒ NH₄⁺ (1) (Ammonium)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

State the type of structure ionic bonding forms, give an example. (2)

A

⇒ Giant Ionic Lattice (1)

⇒ Na⁺+ Cl⁻ (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What is the general trend of melting and boiling points in an ionic bond? (3)

A

⇒ High (1)
⇒ The giant lattice contains strong electrostatic forces between oppositely charged ions (1)
⇒ requires a lot of energy to break the electrostatic bonds (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What are the general trends of conductivity for ionic compounds in Solids and molten/aqueous? (2)

A

⇒ When solid; conductivity is poor as the ions cannot move in the fixed lattice, thus no charge is being carried (1)
⇒ When molten/aqueous; conductivity is good, ions can move around and carry a charge (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Explain why ionic compounds are brittle? (2)

A

⇒ A blow in a direction may move the ions and produce a contact between ions with like charges. (1)
⇒ The ions will repel, thus breaking the compound apart (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Define covalent bonding (1)

A

⇒ A covalent bond is a shared pair of electrons. (occurs between non-metals) (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

What does a single covalent bond contain? (1)

A

⇒ A shared pair of electrons. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

What do multiple covalent bonds contain? (1)

A

⇒ Multiple shared pairs of electrons. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

What is the general trend of B.Ps and M.Ps of simple covalent compounds, hence or otherwise stating why this is the case? (3)

A

⇒ Generally low melting and boiling points (1)
⇒ There is a weak attraction between the molecules (1)
⇒ Therefore, great amounts of energy are not required to overcome/break the bonds. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Explain why simple covalent molecules are bad conductors of electricity. Hence, or otherwise, is this the same case when molten? (2)

A

⇒ The molecules are neutral, ergo there are no charged particles to carry the current. (1)
⇒ Even when molten, the solution will not conduct electricity as there are no charged particles. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

What are crystals? (2)

A

⇒ Solids that have a regular arrangement and are held together by forces of attraction (1)
⇒ There are four basic crystals: Ionic, Metallic, molecular, and macromolecular (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Give an example of each type of crystal, briefly explain each crystal’s bonding structure. (4)

A

⇒ Ionic crystals: NaCl (regular lattice of +ve/-ve ions. (1)
⇒ Metallic crystals: any metal (metals exist as a lattice of positive metal ions embedded in a sea of electrons) (1)
⇒ Molecular crystals: I₂ (Strong covalent bonds held in a regular array of intermolecular forces) (1)
⇒ macromolecular: Diamond and graphite (Covalent bonds extend throughout the system) (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What are giant covalent structures? (1)

A

⇒ A structure that has a huge network of covalently bonded atoms. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What are two examples of macromolecular structures? (2)

A

⇒ Diamond (1)

⇒ Graphite (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Describe the structure of diamond, and the shape of its molecule. (2)

A

⇒ Each carbon atom is covalently bonded to 4 other carbon atoms (1)
⇒ Tetrahedral shape (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

State the 5 properties of diamond? (5)

A 
⇒ very hard (1)
⇒ high melting points (1)
⇒ does not conduct electricity (1)
⇒ does not dissolve in solvents. (1)
⇒ good thermal conductors as vibrations travel easily (1)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Describe the structure of graphite. (3)

A

⇒ Carbon atoms are covalently bonded to 3 other carbon atoms (1)
⇒ 4th outer electron of each C atom is delocalised (1)
⇒ sheets of hexagons (graphene) is bonded by weak van der Waal forces (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Describe the properties of graphite. (4)

A

⇒ slippery and soft (1)
⇒ electrical conductor (1)
⇒ insoluble in any solvent (1)
⇒ high melting points. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Why does graphite conduct electricity? (2)

A

⇒ Only three of carbon’s valence electrons are used in covalent bonding, the fourth electron becomes delocalised. (1)
⇒ The delocalised electrons can move and carry a charge. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Why is graphite slippery and soft? (2)

A

⇒ There are weak van der Waal forces between the layers of graphite (1)
⇒ Which can be easily broken so the sheets can slide past each other. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Why do diamond and graphite have high melting points? (2)

A

⇒ They have many strong covalent bonds (1)

⇒ Which requires a lot of energy to break (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

Describe the structure of molecular crystals. (2)

A

⇒ molecules that are held together in a regular arrangement (1)
⇒ By weak intermolecular forces. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

What are two examples of molecular crystals? (2)

A

⇒ Iodine (1)

⇒ Ice (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

State the properties of molecular crystals and explain your answer. (4)

A

⇒ Low melting point: It does not take much energy to break the weakly induced dipole-dipole (weak intermolecular) forces between the molecules (1)
⇒ Non-conductor of electricity: All valence electrons are used in bonding and are not free to move and carry the charge. (1)
⇒ Insoluble in water: They are non-polar and do not interact with water (1)
⇒ Soft: atoms are held weakly by intermolecular bond
and so can be removed easily by force (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

What is co-ordinate (dative covalent) bonding? (1)

A

Shared pairs of electrons where both electrons (a lone pair) are supplied by one atom. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

State two conditions for co-ordinate bonding to take place. (2)

A

⇒ The atom that accepts the electron pair does not have a full outer shell (e⁻ deficit) (1)
⇒ The atom that donates the e⁻’s has a pair of electrons not in use (a lone pair) (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

How do we represent a co-ordinate bond? (1)

A

⇒ Using an arrow. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

What is metallic bonding? (2)

A

⇒The attraction between a sea of delocalised electrons and positive ions (1)
⇒ Arranged in a lattice. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

Why are metals good conductors of electricity and heat? (2)

A

⇒ Good thermal conductivity: Delocalised electrons pass kinetic energy to each other. (1)
⇒ Good electric conductor: Delocalised electrons can move and carry charge through the metal. (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

Why do metals have high melting points? (2)

A

⇒ Strong attractions between positive ions and negative electrons form a strong electrostatic bond (1)
⇒ Requires a lot of energy to overcome (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

Why are metals malleable and ductile? (2)

A

⇒ Layers of positive ions can slide over each other (1)

⇒ Without disrupting the bonding (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

What are 3 factors affecting the strength of metallic bonding? (6)

A

⇒ Number of electrons/Charge of ion: The greater the charge of the ion, the greater number of e⁻, the stronger the electrostatic attraction (2)
⇒ number of protons: The stronger the electrostatic attraction (2)
⇒ Size of the ion: The smaller the ion, the closer the electrons are to the +ve nucleus→stronger bond. (2)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

Explain why an Mg ion has a stronger metallic bonding and higher melting point than an Na ion? (4)

A

⇒ There are more electrons in the outer shell of the Mg ion (1)
⇒ it has more protons than Na ion (1)
⇒ It is a smaller ion than a Na ion (smaller atomic radii) (1)
⇒ so there is stronger attraction and higher energy required to break bonds (1)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

State and explain the trend in melting points of the Group II elements Ca-Ba. (3)

A

⇒ Trend: Decreases (1)
⇒ Increase in size of ion/atom (1)
⇒ A weaker attraction for delocalised/free/sea of electrons / weaker metallic bonding (1)

37
Q

What is electronegativity? (2)

A

⇒ The power of an atom to attract the pair of electrons/electron density to itself (1)
⇒ In a covalent bond (1)

38
Q

What is electron density? (1)

A

⇒ How negative charge is distributed in a molecule (1)

39
Q

How is the electronegativity of atoms measured? (1)

A

⇒ 0-4 on the Pauling scale. (1)

40
Q

Which element is the most electronegative? (1)

A

⇒ Fluorine with 4.0 (1)

41
Q

What does electronegativity depend on? (3)

A

⇒ Nuclear charge of an atom (1)
⇒ The distance between the nucleus and the valence electrons (1)
⇒ The shielding of the nuclear charge by electrons in the inner shells (1)

42
Q

Describe the trend of electronegativity as you go across a period: (4)

A

⇒ EN increases (1)
⇒ nuclear charge increases/proton number increases (1)
⇒ number of shells remain the same (1)
⇒ so atomic radii decreases as there is a strong attraction between nucleus and electrons in shells (1)

43
Q

Describe the trend of electronegativity as you go down a group, explain your answer. (3)

A

⇒ EN decreases (1)
⇒ distance between the nucleus and outer electrons increases (1)
⇒ shielding of inner shells increases. (1)

44
Q

Define polarity of covalent bonds (polar molecules)? (2)

A

⇒ The unequal sharing of electrons between atoms (1)
⇒ That is bonded covalently (1)
e.g HF (The difference in electronegativity causes, causes the bonding electrons to be pulled more towards F.)
The greater the difference in electronegativity between the two atoms, the more polar the bond is.

45
Q

What are non-polar molecules? (1)

A

⇒ When the electron sharing is equal in a bond. (1)

e.g: F₂ (The molecule is completely non-polar, so the bonding electrons are equidistant to each atom)

46
Q

How do we represent partial charges? (1)

A

⇒ δ⁺ and δ⁻ (1)

47
Q

How would an electron cloud look like between an HF molecule? (1)

A

⇒ The electron cloud will be distorted towards the fluorine atom. (1)

48
Q

What is a dipole? (1)

A

⇒ A dipole is a difference in charge between two electronegative atoms (1)

49
Q

Molecules with polar covalent bonds will have what force? (1)

A

⇒ permanent dipole force (1)

50
Q

What are the 3 types of intermolecular forces? (In order from the weakest-strongest) (3)

A

⇒ Van der Waals (Induced dipole-dipole forces) (1)
⇒ Dipole-Dipole forces (1)
⇒ Hydrogen bonding (1)

51
Q

What is the difference between INTER and INTRA molecular forces?

A

⇒ INTER = are between molecules in a substance and weaker. (e.g: Hydrogen bonding) (1)
⇒ INTRA = are between the atoms in a molecule and stronger. (e.g metallic bonding) (1)

52
Q

Where do van der Waal forces arise? (1)

How do these Van der Waal forces arise? (3)

A

⇒ Occur in all molecules. (1)
⇒ Spontaneous movement of electrons forms a dipole in one molecule (1)
⇒ Induces a dipole in another molecule (1)
⇒ causes an induced temporary attraction (1)

53
Q

How is the strength of intermolecular forces different in different molecules? (2)

A

⇒ Larger molecules have larger electron clouds, thus a greater number of electrons, thus a stronger van der Waals force (1)
⇒ The closer together the two molecules, the stronger the VdW force (1).

54
Q

Suggest why methaneselenol (CH₃SeH) has a higher boiling point than methanethiol (CH₃SH). (2)

A

⇒ (Methaneselenol is a) bigger molecule (1)

⇒ With stronger/more VdW forces between molecules (1)

55
Q

What are permanent dipole-dipole forces? (1)

A

⇒ Weak electrostatic attraction between polar molecules. (1)

56
Q

When do permanent dipole forces arise in molecules? (2)

A

⇒ Unsymmetrical molecule: (length of bonds are different, dipoles don’t cancel out e.g: dichloromethane) shape causes a difference in electronegativity between two atoms in the molecule forming a dipole moment, ergo a permanent dipole. (1)
⇒ A lone pair (1)

57
Q

Why does SiH₄ not have permanent dipole? (3)

A

⇒ SiH₄ has a completely symmetrical shape (1)
⇒ with all hydrogen bonds being pushed at the same distance (1)
⇒ thus canceling out the partial charges (1)

58
Q

Why does PH₃ have a permanent dipole? (1)

A

⇒ Contains a lone pair (1)

59
Q

Explain how permanent dipole-dipole forces arise between hydrogen chloride molecules? (2)

A

⇒ The difference in electronegativity leads to bond polarity (1)
⇒ Dipoles don’t cancel out (1)

60
Q

How does hydrogen bonding arise? (3)

A

⇒ A hydrogen atom is “sandwiched” between two very electronegative atoms with lone pairs (1)
⇒ forms δ⁺ and δ⁻ on hydrogen and the other atom. (1)
⇒ A lone pair will be attracted to the δ⁺H forming the hydrogen bond. (1)

61
Q

What is hydrogen bonding? (3)

A

⇒ Hydrogen bonds are permanent dipole-dipole attractions (1)
⇒ Between a ᵟ⁺H atom (that is covalently bonded to ᵟ⁻O, ᵟ⁻N or ᵟ⁻F) in one molecule (1)
⇒ And a ᵟ⁻O, ᵟ⁻N or ᵟ⁻F atom in another molecule. (“sandwiched” between two molecules that have either ᵟ⁻O, ᵟ⁻N or ᵟ⁻F) (1)

62
Q

What are two conditions needed for hydrogen bonding? (2)

A

⇒ A highly electronegative atom (1)

⇒ With lone pairs (1)

63
Q

State the three molecules that have hydrogen bonding. (from highest boiling to lowest) (3)

A

⇒ H2O (1)
⇒ HF (1)
⇒ NH3(DNA) (1)

64
Q

Why do H2O, HF, NH3 not fit the trend in the hydrides graph of boiling points? (2)

A

⇒ higher boiling points than expected for their molecular mass (1)
⇒ As they have an extra intermolecular force: hydrogen bonding thus more energy is needed to overcome their bonds. (1)

65
Q

How are hydrogen bonds represented? (1)

A

⇒ Continuous 90 degrees parallel lines between the two molecules
(| | | | |) (1)

66
Q

How are permanent dipole-dipole forces represented? (1)

A

⇒ With a dotted line between the two molecules. (1)

67
Q

Explain why the O-H bond in methanol is polar (2)

A

⇒ There is a large difference in EN between the O and H atom (1)
⇒ So O will attract higher e- density around it to form δ⁺and δ⁻ molecule. (1)

68
Q

What’s the difference between saturated and unsaturated fats in terms of intermolecular forces? (2)

A

Saturated fats have regular chains/branches so stronger IMFs so higher M.P. (1)
- Unsaturated fats have random branched chains so weaker IMFs so lower M.P. (1)

69
Q

Why is hydrogen bonding important? (2)

A

⇒ Ice is less dense than water because the particles are further apart/less closely packed. (1)
⇒ they enable organisms to thrive in ponds as they would reflect and absorb sunlight to keep the bottom of the pond warm. (1)

70
Q

What does Electron Pair Repulsion Theory state? (1)

A

⇒ Electron pairs repel each other equally if there are all bond pairs to get as far away as possible to minimise repulsive forces between them. (1)

71
Q

What are the steps in determining a molecule’s structure? (5)

A

⇒ S1: Determine number of electrons of the central atom (e.g: PCl₆⁻, central atom = p, p=group 5 = 5e⁻)
⇒ S2: Add one electron onto S1 for each bond formed. (e.g: PCl₆⁻, Cl has 6 bonds; 5(S1) + 6(Cl₆⁻)= 11 electrons)
⇒ S3: (For ions): Allow for any charge of central atom, If
-ve add an electron, if +ve, remove an electron. (e.g: PCl₆⁻, (-ve) add electron, So, Σe⁻ = 12)
⇒ Divide previous Σe⁻ by 2 to give the number of electron pairs/lone pairs: (12/2=6, 6 bonding pairs)
⇒ Determine the shape (6 bonding pairs = octahedral)
⇒ Determining lone/bonding pairs:
→ e.g: NH₃, if we do the steps for NH₃ we get ΣBonds = 4 (But there are only 3 hydrogen atoms to bond, so one bond is left in excess so it’s a lone pair.
→ e.g: But, in PCl₆⁻, ΣBonds = 6, and there are 6 Cl atoms to bond so there are no lone pairs.

72
Q

What is the order of repulsion for electron pairs? (1)

A

⇒ lone pair - lone pair> lone pair to bond pair> bond pair-bond pair. (1)

73
Q

What is the angle degree rule for lone pairs? (1)

A

1 lone pair = 2.5°
2 lone pairs = 5
e.g: NH₃ is a tetrahedral (Normal tetrahedral has bond angles of 109.5 degrees, but due to the lone pair, the angles are reduced by 2.5 =109.5-2.5 = 107 degrees between the bonds.

74
Q

What is the angle in a linear shape? State its number of Bonding/Lone pairs.

A

180°
2 bonding pairs
0 Lone pairs

75
Q

What is the angle in a trigonal planar shape? State its number of Bonding/Lone pairs.

A

120°
3 bonding pairs
0 Lone pairs

76
Q

What is the angle in a bent shape with 2 BP and 1 LP?

A

118°

77
Q

What is the angle in a tetrahedral shape? State its number of Bonding/Lone pairs.

A

109.5°
4 Bonding pairs
0 Lone pairs

78
Q

What is the angle in a trigonal pyramidal shape with 3 BP and 1 LP?

A

107°

79
Q

What is the angle in a bent shape with 2 BP and 2 LP

A

104.5°

80
Q

What are the 2 angles in a trigonal bipyramidal shape? State its number of Bonding/Lone pairs.

A

120° and 90°
5 bonding pairs
0 Lone pairs

81
Q

What are the 2 angles in trigonal pyramidal with 4 BP and 1 LP?

A

119° and 89°

82
Q

What is the angle(s) in a t-shape?

A

120° 89°

83
Q

What is the angle in an octahedral? State its number of Bonding/Lone pairs.

A

90°
6 Bonding pairs
0 Lone pairs

84
Q

What is the angle in a square pyramidal? State its number of Bonding/Lone pairs.

A

89°
5 bonding
1 Lone

85
Q

What is the angle in a square planar? State its number of Bonding/Lone pairs.

A

90°
4 Bonding pairs
2 Lone pairs

86
Q

Explain why CF₄ has a bond angle of 109.5°. (2)

A

⇒ Around carbon there are 4 bonding pairs of electrons (and no lone pairs) (1)
⇒ Therefore, these repel equally and spread as far apart as possible (1)

87
Q

Identify one molecule with the same number of atoms, the same number of electrons, and the same shape as H3O⁺ (1)

A

⇒ NH3. (1)

88
Q

Deduce the molecule of a compound that has the same number of atoms and electrons and the same shape as AlCl₄⁻ (1)

A

⇒ SiCl₄ (1)

CHEMISTRY Physical (4 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions