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Chemistry > Bonding and Structure > Flashcards

Bonding and Structure Flashcards

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1
Q

Give the formula for a carbonate ion.

A

CO3-

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2
Q

Give the formula for a nitrate ion.

A

NO3-

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3
Q

Give the formula for a sulphate ion.

A

SO4 2-

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4
Q

Give the formula for an ammonium ion.

A

NH4+

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5
Q

Give the formula for a phosphate ion.

A

PO4 3-

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6
Q

Define ionic bonding.

A

An electrostatic attraction between positive and negatively charged ions.

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7
Q

What is the usual electronic configuration of the positive ion in a compound?

A

The same as that of the previous noble gas on the periodic table.

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8
Q

What is the usual electronic configuration of the negative ion in a compound?

A

The same as that of the next noble gas on the periodic table.

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9
Q

Define covalent bonding.

A

The overlap of atomic orbitals, each containing 1 electron, to give a shared electron pair.

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10
Q

Define a molecule.

A

A small group of atoms bonded together by covalent bonds.

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11
Q

What is the ‘octet rule’?

A

The general rule that atoms ideally contain 8 electrons on their outer shells when bonding covalently.

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12
Q

What is a shared pair of electrons called?

A

A bonding pair.

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13
Q

What is an electron pair that is not involved in bonding called?

A

A lone pair.

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14
Q

What is a dative covalent bond?

A

A covalent bond in which both shared electrons come from the same atom.

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15
Q

What kind of structure do ionic compounds form?

A

Giant ionic lattices.

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16
Q

What is a simple covalent structure?

A

A small, covalently bonded compound that forms discrete molecules.

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17
Q

What is a giant covalent structure?

A

An infinite lattice structure of atoms, where each atom is bonded covalently to many neighbouring atoms.

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18
Q

What kind of structures do metals have?

A

Giant lattice structures.

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19
Q

What property of metal atoms allows delocalised electrons?

A

A relatively low ionisation energy.

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20
Q

Describe metallic bonding.

A

A regular lattice structure of positively charged metal ions, held together by a sea of delocalised electrons.

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21
Q

Why do metals have high melting and boiling points?

A

Strong electrostatic forces between positively charged ions and electrons require a lot of energy to overcome, and therefore high temperatures are needed.

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22
Q

Why do metals conduct electricity?

A

The delocalised electrons in metals are free to move and therefore carry charge, allowing current to flow.

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23
Q

Why do metals have poor solubility?

A

Strong electrostatic attractions between positive ions and delocalised electrons require a lot of energy to break in order to dissolve. Metals usually react, rather than dissolve in water.

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24
Q

Why do some metals have higher melting points than others?

A

The bigger the charge on a metal atom, the stronger the electrostatic forces in the metal are, and the more energy required to overcome them. Different metals carry different charges.

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25
Q

Why are some metals better conductors than others?

A

Metals with different charges release different numbers of electrons into the sea of electrons. The higher the density of delocalized electrons, the more conductive a material will be.

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26
Q

Define malleable.

A

Easily hammered into shape.

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27
Q

Define ductile.

A

Easily drawn into a wire.

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28
Q

What elements are present in stainless steel?

A

Iron, Chromium and Carbon.

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29
Q

What elements are present in Brass?

A

Copper and Zinc.

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30
Q

What are alloys?

A

Metals containing other atoms of different sizes, interrupting the regular structure and making them less malleable.

31
Q

Why are metals malleable and ductile?

A

They form regular lattice structures with layers that can easily move past each other.

32
Q

Why do giant ionic lattices have high melting and boiling points?

A

Strong electrostatic attraction between positive and negative ions needs a lot of energy to overcome.

33
Q

Why do giant ionic lattices have poor conductivity when solid?

A

The ions are fixed in a lattice so cannot move and carry charge, and therefore current cannot flow.

34
Q

Why do giant ionic lattices have good conductivity when molten or aqueous?

A

The ions are no longer fixed in the lattice so are able to move and carry charge, allowing current to flow.

35
Q

Why do giant ionic lattices have good solubility in water?

A

The oxygen atom in the water molecules is attracted to and surrounds the positive ions in the lattice, and the hydrogen atoms in the water molecules are attracted to and surround the negative ions. This breaks down the ionic lattice and dissolves the compound.

36
Q

Why are giant ionic lattices brittle?

A

The slight movement of a layer of ions brings the positive ions next to other positive ions, and the negative ions next to other negative ions. This causes repulsion and the structure breaks apart.

37
Q

What is charge density in ionic compounds?

A

The charge on an ion compared to that ions mass. The higher charge density is, the stronger the electrostatic forces.

38
Q

Why do simple covalent structures have low melting points?

A

Weak intermolecular forces take little energy to overcome.

39
Q

Why do giant covalent structures have high melting points?

A

Every atom is bonded covalently to multiple other atoms, and each of these bonds must be broken in order for the structure to melt, which requires a huge amount of energy.

40
Q

Why do covalent structures have poor conductivity?

A

No free mobile electrons (or other charged particles) to carry charge

41
Q

Why do simple covalent structures have good solubility in non-polar solvents?

A

Solvent molecules form similar intermolecular bonds with molecules; weak intermolecular forces do not need much energy to overcome

42
Q

Why do giant covalent structures have poor solubility?

A

Strong covalent bonds between atoms require too much energy to overcome

43
Q

Describe the bonding in diamond.

A

A giant lattice of carbon atoms, where each atom is covalently bonded to 4 other surrounding atoms.

44
Q

Describe the physical properties of diamond.

A

Extremely hard, extremely high melting and boiling points, poor conductivity and low solubility.

45
Q

Describe the bonding in graphite.

A

Giant hexagonal lattices of carbon atoms each bonded to 3 other atoms in flat layers. Each carbon provides one electron to form a delocalized layer of electrons in a sea between planes of C atoms.

46
Q

Describe the physical properties of graphite.

A

Soft and malleable, good conductivity, low solubility.

47
Q

What is Valence Shell Electron Pair Repulsion (VSEPR) theory?

A

Electron pairs repel each other.
They will adopt positions where they are separated by the largest possible angle.
Lone pairs repel more strongly than bonded pairs so the repulsion between lone pair-lone pair > lone pair-bonded pair > bonded pair-bonded pair

48
Q

Describe the configuration and bond angle in linear molecules.

A

Two equally repulsive groups attached to the central atom and no lone electron pairs. 180 degree bond angle.

49
Q

Describe the configuration and bond angle in tetrahedral molecules.

A

Four equally repulsive groups attached to the central atom and no lone electron pairs. 109.5 degree bond angle.

50
Q

Describe the configuration and bond angle in pyramidal molecules.

A

Three equally repulsive groups attached to the central atom and repelled by a single lone electron pair. 107 degree bond angle.

51
Q

Describe the configuration and bond angle in trigonal planar molecules.

A

Three bonding electron pairs and no lone pairs attached to the central atom. All groups exist in a single, flat plane with a 120 degree bond angle.

52
Q

Describe the configuration and bond angle in octahedral molecules.

A

Six bonding electron pairs and no lone pairs attached to the central atoms. 90 degree bond angle.

53
Q

By what angle does a lone electron pair repel bonding electron pairs?

A

2.5 degrees.

54
Q

Describe the configuration and bond angle in non-linear molecules.

A

Two bonding pairs and two lone pairs of electron on the central atom. 104.5 degree bond angle.

55
Q

What are the two types of Van der Waals forces?

A

London forces and permanent dipole-dipoles.

56
Q

What is a dipole?

A

An imbalance in the distributions of electrons in a compound which causes random polarisation.

57
Q

What are London forces?

A

Intermolecular forces in simple covalent molecules that are caused by electrostatic attractions between randomly induced temporary dipoles.

58
Q

What affects the strength of London forces?

A

The number of electrons in a molecule which determines the strength of the dipoles and how branched a molecule is which determines the degree of contact between molecules.

59
Q

Describe the trend in London forces down a group.

A

London force strength increases down the group, as the increase in number of electrons allows stronger temporary dipoles to form, and hence stronger electrostatic attractions.

60
Q

Define electronegativity.

A

The measure of the attraction of a bonded atom with the pair of electrons in a covalent bond.

61
Q

What is a permanent dipole?

A

A small charge difference that does not change across a bond, with partial charges on the bonded atoms (the result of the bonded atoms having different electronegativities).

62
Q

What are the four most electronegative elements?

A

Fluorene, oxygen, nitrogen, chlorine.

63
Q

What are permanent dipole forces?

A

Permanent electrostatic forces between partially charged dipoles in small covalent molecules.

64
Q

What is a polar covalent bond?

A

A bond with a permanent dipole, having positive and negative partial charges on the bonded atoms

65
Q

What is a polar molecule?

A

A molecule with an overall dipole (ie one positive end and one negative end), having taken into account any dipoles across bonds and the shape of the molecule

66
Q

What is a hydrogen bond?

A

A hydrogen bond is a strong dipole-dipole attraction between an electron deficient hydrogen atom in NH, OH or HF on one molecule, and a lone pair of electrons on a highly electronegative atom (N, O or F) on a different molecule.

67
Q

What atoms can form hydrogen bonds?

A

Nitrogen, Oxygen and Fluorene.

68
Q

How do you draw a hydrogen bond?

A

A series of parallel lines between the lone pair on a delta- dipole on one molecule and a delta+ hydrogen atom on another molecule.

69
Q

What are three anomalous properties of water?

A

Unexpectedly high melting and boiling points, the density of ice compared to water, and unusually high surface tension.

70
Q

Why is the boiling point of water anomalous?

A

It is unusually high compared to other group 6 hydrides due to hydrogen bonding in water which causes much stronger intermolecular forces.

71
Q

Why is the density of ice anomalous?

A

Unlike most solids compared to their respective liquids, ice floats on water.

72
Q

Explain the relative densities of ice and water.

A

When water freezes, each molecule forms four hydrogen bonds with surrounding molecules (one from each hydrogen atom and both lone pairs on oxygen) forming a crystalline structure which forces molecules apart so that there are fewer molecules per unit volume.

73
Q

Explain the anomalously high surface tension of water.

A

At the surface of water, molecules are surrounded by other molecules on only one side. Therefore they form many hydrogen bonds with their neighbouring molecules at the surface, creating a layer of hydrogen bonds which leads to high surface tension.

Chemistry (13 decks)

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