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Chemistry > Module 2 + 4 organic > Flashcards

Module 2 + 4 organic Flashcards

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1
Q

Define a homologous series

A

Is a series of organic compounds with the same functional group but each successive member differing by CH2.

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2
Q

Define a functional group

A

Is the part of an organic molecule responsible for its chemical reactions.

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3
Q

Define the term aliphatic

A

An aliphatic molecule contains carbon atoms joined together in straight or branched chains. It does not contain a benzene ring.

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4
Q

Define the term alicyclic

A

Is an aliphatic molecule that contains carbon atoms joined together in a ring that is not a benzene ring.

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5
Q

Define the term aromatic

A

An aromatic molecule contains at least 1 benzene ring.

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6
Q

Define the term benzene ring

A

The hexagonal unsaturated ring of six carbon atoms present in benzene and many other aromatic molecules.

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7
Q

What does a curly arrow represent?

A

The movement of a pair of electrons.

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8
Q

What is meant by heterolytic bond fission?

A

The breaking of a covalent bond where one atom gets both bonding atoms.

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9
Q

What is a lone pair?

A

Is a pair of electrons not involved in a bond.

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10
Q

What is a covalent bond?

A

Is a pair of electrons shared between 2 atoms.

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11
Q

What is a sigma bond?

A

Is a shared pair of electrons, directly between the 2 bonding atoms.

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12
Q

What is a Pi bond?

A

Is a shared pair of electrons, above and below the 2 bonding atoms.

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13
Q

What is a single bond?

A

Is a sigma bond. It contains one shared pair of electrons.

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14
Q

What is a double bond?

A

Is a sigma bond and a pi bond. It contains 2 shared pairs of electrons.

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15
Q

Where does a curly always start?

A

At a pair of electrons.

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16
Q

How/why do chemical reactions happen?

A

All because electrons are moving to make bonds or to break bonds.

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17
Q

What does a mechanism show?

A

The movement of pairs of electrons using curly arrows.

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18
Q

What structure does metallic bonding have?

A

Giant metallic lattice

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19
Q

What structure does ionic bonding have?

A

Giant ionic lattice

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20
Q

What structure does covalent bonding have?

A

Simple molecular structure

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21
Q

What type of particles does metallic bonding contain?

A

Metals. Positive ions and delocalised electrons.

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22
Q

What type of particles does ionic bonding contain?

A

Metal and non-metal. Positive and negative ions.

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23
Q

What type of particles does covalent bonding contain?

A

Non-metal atoms.

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24
Q

Does metallic bonding have high or low melting and boiling point?

A

High melting and boiling points due to the strong electrostatic forces of attractions.

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25
Q

Does ionic bonding have high or low melting and boiling point?

A

High melting and boiling points due to strong electrostatic forces of attraction between positive and negative ions.

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26
Q

Does covalent bonding have high or low melting and boiling point?

A

Low melting and boiling points due to weak intermolecular forces - less energy to overcome.

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27
Q

Does metallic bonding conduct electricity?

A

Yes as the delocalised electrons are free to move and carry a charge

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28
Q

Does ionic bonding conduct electricity?

A

When molten or dissolved ions (negative) are free to move and carry a charge. Don’t conduct when solid as ions are not free to move.

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29
Q

Does covalent bonding conduct electricity?

A

No as there are no delocalised electrons or negative ions to carry a charge.

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30
Q

Does metallic bonding dissolve in water?

A

No

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31
Q

Does metallic bonding dissolve in covalent solvents?

A

No

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32
Q

Does ionic bonding dissolve in water?

A

Some due to waters slightly charged but not very well.

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33
Q

Does ionic bonding dissolve in covalent solvents?

A

No.

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34
Q

Does covalent bonding dissolve in water?

A

No

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35
Q

Does covalent bonding dissolve in covalent solvents?

A

Yes

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36
Q

What is meant by the term electronegativity?

A

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

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37
Q

4 bond pairs and no lone pairs - what is the shape of bond angle?

A

Tetrahedral, 109.5 deg

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38
Q

3 bond pairs and no lone pairs - what is the shape and bond angle?

A

Trigonal planar, 120 deg

39
Q

3 bond pairs and 1 lone pair - what is the shape and bond angle?

A

Trigonal pyramidal, 107 deg

40
Q

6 bond pairs and no lone pairs - what is the shape and bond angle?

A

Octahedral, 90 deg

41
Q

2 bond pairs and 2 lone pairs - what is the shape and bond angle?

A

Non-linear, 104.5 deg

42
Q

When the bonding electrons (shared pair) are evenly distributed between the atoms that make up a bond; we say that the bond is what?

A

Symmetrical

43
Q

If the bonding electrons are unevenly distributed between the atoms that make up the bond; we say that the bond is what?

A

Non-symmetrical

44
Q

In a polar bond the pair of electrons are attracted to the atom which is significantly more what?

A

Electronegative

45
Q

What are the electronegativity values for H,O,F,C,N and Br?

A

H= 2.20 O=3.44 F=3.98 C=2.55 N=3.04 Br=2.96

46
Q

How does electronegativity change across the periodic table?

A

Electronegativity increases from bottom to top in groups, and increases from left to right across periods. With fluorine being the most electronegative element.

47
Q

What is a polar molecule?

A

A polar molecule is a molecule in which one end of the molecule is slightly positive, while the other end is slightly negative.

48
Q

When do polar bonds form?

A

Polar bonds form between atoms of the same element with different electronegativities so that the electrons are asymmetrically distributed between the atoms.

49
Q

How can a molecule be non-polar but still have polar bonds?

A

E.g.: tetrachloromethane, each C-Cl bond is polar, but a CCl4 molecule is symmetrical: tetrahedral, so the dipoles act in different directions and cancel each other out.

50
Q

What is an intermolecular force?

A

Intermolecular forces are the attractive and repulsive forces that arise between the molecules of a substance?

51
Q

What are the three types of intermolecular forces?

A

> Induced dipole-dipole attraction forces/interactions (London dispersion forces) - all molecules.
Permanent dipole-dipole forces/interactions ( only for polar molecules).
Hydrogen bonding - molecules contains atom with lone pair of e- and H-O, H-N or H-F.

52
Q

How do induced dipole-dipole attraction forces/interactions (London dispersion forces) arise?

A

They arise because of the constant movement of electrons which results in the uneven distribution of charge, an instantaneous dipole. These instantaneous dipoles induce temporary dipoles on neighbouring atoms: The attraction between the induced dipoles is what constitutes induced dipole-dipole attraction forces and is much weaker than attractions between permanent dipoles.

53
Q

What are induced dipole-dipole attraction forces/interactions (London dispersion forces) and where do they occur?

A

These occur between all atoms and simple covalent molecules and are the weakest type of intermolecular force.

54
Q

How does the strength of induced dipole-dipole attraction forces/interactions (London dispersion forces) change?

A

The strength of the induced dipole-dipole attraction forces increases as the number of electrons increases ( boiling point increases down group VII as more energy is needed to overcome stronger intermolecular forces ) and as the surface area of contact increases between molecules.

55
Q

What are hydrogen bonds?

A

Hydrogen bonds occur between an electron deficient hydrogen on one molecule and the lone pair of electrons on a highly electronegative atom usually on a different molecule. These are the strongest type of intermolecular force. Hydrogen must be attached to F, N or O (the most electronegative elements) for hydrogen bonding to oocur?

56
Q

What does the magnitude of induced dipole-dipole interactions depend on?

A

The number of electrons in the molecule and the number of points of contact between molecules.

57
Q

How does the magnitude of the force of attraction between molecules affect the boiling point?

A

As the strength of intermolecular forces increases, the more energy is needed to overcome the intermolecular forces and seperate molecules from each other -> boiling point increases.

58
Q

Why does water have a much higher boiling point than expected?

A

Because when water molecules are in liquid state they form hydrogen bonds between molecules (between an electron deficient hydrogen atom and a lone pair on an oxygen atom of another water molecule). Hydrogen bonds are the strongest of all IM so a large amount of energy is needed to break these interactions before it can begin to boil.

59
Q

Why is solid water(ice) less dense than liquid water?

A

Because when ice forms, hydrogen bonds form between water molecules. These hydrogen bonds hold the water molecules apart in an open lattice structure meaning that the water molecules are futher apart than they are in liquid water making ice more dense than water. Also in ice there are fewer molecules per unit volume.

60
Q

Outline the solubility rules for simple molecular substances.

A

Simple molecular substances are non-polar, intermolecular forces form between molecules and solvent. IM forces in lattice break and solid dissolves. Molecules are soluble in non-polar solvents.

61
Q

Do simple molecular substances conduct electricity? give reasons

A

No they do not conduct electricity, even when liquid because their molecules are uncharged, so there are no free charged particles to carry the current.

62
Q

Why do simple molecular substances have low melting and boiling points?

A

Simple molecular substances have very weak intermolecular forces, so very little energy is needed to overcome.

63
Q

What does a metal + acid produce?

A

Salt + hydrogen

64
Q

What does a base + acid produce?

A

salt + water

65
Q

What does an acid + alkali produce?

A

salt + water

66
Q

What does metal carbonate + acid produce?

A

salt + water + carbon dioxide

67
Q

What does ammonia + acid produce?

A

Ammonium salt

68
Q

What is a salt?

A

When the H+ ions in the acid are replaced by a metal ion or ammonium ion.

69
Q

What is a base?

A

A proton (H+) acceptor

70
Q

What is an alkali?

A

A soluble base that donates OH- ions in aqueous solution.

71
Q

What is an acid?

A

A proton (H+) donator

72
Q

What is a subshell?

A

A sub-shell is agroup of the same type of atomic orbitals within a shell. These can be s sub-shells made from one single s- orbital, p sub-shells - made up from three p- orbitals, d sub-shells - made up from five d- orbitals or f sub-shells - made from seven f- orbitals.

73
Q

What is an atomic orbital?

A

A region within an atom that can hold up to two electrons with opposite spins.

74
Q

What is the shape of an s orbital and what’s the max number of electrons it can hold?

A

Sphere shape
2 electrons

75
Q

What is the shape of an p orbital and what’s the max number of electrons it can hold?

A

Dumbbell shape
2 electrons

76
Q

Explain an orbitals function

A

An orbital can be empty, contain one electron or contain two electrons. An orbital is shown as a box. If there are two electrons in an orbital, they must have opposite spins, shown by the up and down arrows.

77
Q

What’s the number of orbitals and the maximum number of electrons that each sub-shell can hold?

A

s sub-shell - one orbital, two electrons.
p sub-shell - three orbitals, six electrons.
d sub-shell - five orbitals, ten electrons.
f sub-shell - 7 orbitals, 14 electrons.

78
Q

What sub-shells do the shells contain?

A

shell 1 - 1s sub shell
shell 2 - 2s and 2p sub shells
shell 3 - 3s, 3p and 3d sub shell
shell 4 - 4s, 4p, 4d and 4f sub shells.

79
Q

What is the Aufbau principle?

A

Electrons fill the subshells in order of increasing energy. The order is 1s 2s 2p 3s 3p 4s 3d 4p for the elements hydrogen to krypton.

80
Q

What is Hund’s rule?

A

Within a sub-shell, the orbitals are first occupied singly by unpaired electrons. The electrons only pair up when there are no empty orbitals left in the sub-shell.

81
Q

What is the exception to the rule for chromium?

A

According to Aufbau principl the electron configurstaion of Cr should be [Ar] 4s2 3d4 however the actual electronic configuration is [Ar] 4s1 3d5 - this is because half-filled subshells (d5) are more stable than partially filled ones (d4).

82
Q

What is the exception to the rule for copper?

A

Following the typical rules copper would be [Ar] 3d9 4s2, but since 3d10 has less energy/is more stable, the actual electron configuration would be [Ar] 3d10 4s1.

83
Q

Which elements are so stable that they exist as single atoms?

A

Noble gases

84
Q

What can we say about the outer shell of noble gases?

A

They have full outer shells (8e-)

85
Q

What will oppositely charged ions do?

A

Attract and share a pair of electrons.

86
Q

What is an ionic bond?

A

An electrostatic force of attraction between oppositely charged ions.

87
Q

What two factors need to be considered to work out the charge density of the ions?

A

Charge of an ion and the size of an ion.

88
Q

Many ionic compounds dissolve in polar solvents. What happens when an ionic compound dissolves?

A

Polar molecules break down the lattice and surround in solution.

89
Q

When considering solubility what two processes do we have to consider?

A

The strength of attraction between oppositely charged ions in the lattice. Attraction between ion and water molecule.

90
Q

Why do ionic compounds not conduct when solid, but do conduct when molten?

A

When solid, ions are not free to move as they are in a fixed position so they cant carry a charge (electrons), whereas in a liquid they are free to move so can carry a charge.

91
Q

What is a covalent bond?

A

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

92
Q

Sometimes, instead of each atom contributing one electron to the shared pair, one atom contributed both electrons - the other atom must have an empty orbital to except this pair. What is this bond called?

A

A dative covalent bond or a coordinate bond.

93
Q

What are the two sentences that always need to be included in shape questions?

A

> lone pairs of electrons repel each other more than bonding pairs.
bonding pairs repel each other as far/as much as possible.

Chemistry (10 decks)

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