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LC Chemistry > chemical bonding > Flashcards

chemical bonding Flashcards

1
Q

compound

A

substance that is made up of two or more DIFFERENT elements combined together chemically

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2
Q

octet rule

A

when bonding occurs, atoms tend to reach an electron arrangement with eight electrons in the outermost shell

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3
Q

an ion

A

a charged atom or group of atoms

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4
Q

an ionic bond

A

the force of attraction between oppositely charged ions in a compound

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5
Q

transition metal

A

one that forms at least one ion with a partially filled sublevel

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6
Q

molecule

A

group of atoms joined together. It is the smaller particle of an element or compound that can exist independently

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7
Q

valency of an element

A

defined as the number of atoms of hydrogen or any other monovalent element with which each atom of the element combines

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8
Q

electronegativity

A

a measure of the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond

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9
Q

electronegativity differences

A
  • difference > 1.7 indicates ionic bonding in a compound

- difference ≤ 1.7 indicates covalent bonding in a compound

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10
Q

the value of electronegativity

A
  • decreases down groups in Periodic Table. (Bc of increasing atomic radius, screening effect of inner electrons)
  • increases across table (bc of increasing nuclear charge, decreasing atomic radius)
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11
Q

the most electronegative element

A

F

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12
Q

halogens

A

decrease in reducing power down group due to drop in electronegativity values

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13
Q

noble gases

A
  • have 8 electrons on their outer shell + are quite stable

- generally unreactive

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14
Q

uses of noble gases

A

helium

argon

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15
Q

helium

A

used in airships as it is lighter than air

not as light as hydrogen (twice as heavy per volume) but does not burn

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16
Q

argon

A

most common noble gas

used to fill normal light bulbs to stop them imploding

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17
Q

limitations of the octet rule

A

hydrogen
lithium
transition metals

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18
Q

valency

A

the number of bonds an atom makes when it reacts

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19
Q

how to get valency

A
  • can be worked out calculating the number of electrons an atom needs to lose/gain to have 8e⁻ on outer shell
  • can be predicted from periodic table
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20
Q

examples of valency numbers

A

ammonia (contains only nitrogen and hydrogen)
methane (contains only carbon and hydrogen)
calcium bromide
silicon fluroide

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21
Q

transition metals (info)

A
  • transition elements have variable valency
  • form coloured compounds
  • used as catalysts
  • exceptions: zinc (Zn), scandium (Sc)
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22
Q

ionic bonding

A

-transfer of electrons
-ions are formed
-ionic bonding usually between Groups I and II (metals)
and groups VI and VII (non-metals)

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23
Q

cation

A

positive (+)

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24
Q

anion

A

negative (-)

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25
Q

crystal lattice

A

ionic substances usually form a structure called a crystal lattice

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26
Q

eg of crystal lattice

A

NaCl

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27
Q

why ionic substances form crystals

A

ionic substances form crystals bc positive ions attract negative ions in all directions

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28
Q

what shape do crystals form?

A

a lattice structure

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29
Q

characteristics of ionic substances

A
  • strong forces between ions -> v. hard to break up the lattice structure
  • high melting + boiling points
  • usually solid at room temp
  • cannot conduct electricity when solid (ions not free -> move and carry electricity)
  • most dissolve in water (when dissolved, ions can conduct electricity)
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30
Q

examples of everyday ionic substances

A
  • table salt (sodium chloride NaCl)

- fluoridation in water to prevent tooth decay (sodium fluoride NaF)

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31
Q

complex ions table

A

in hardback

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32
Q

covalent bonding

A

formed when atoms share electrons

eg. hydrogen molecule, chlorine molecule

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33
Q

types of covalent bonds

A

single

double triple

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34
Q

covalents bonds - single

A

1 pair of electron shared

eg. H₂

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35
Q

covalents bonds - double

A

2 pairs of electrons shared

eg. O₂

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36
Q

covalents bonds - triple

A

3 pairs of electrons shared

eg. N₂

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37
Q

bonding pairs

A

shared electron pairs that form covalent bonds

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38
Q

lone pairs

A

electron pairs not involved in bonding

39
Q

oribitals in a covalent bond

A

in a covalent bond, orbitals overlap

40
Q

sigma bond

A

formed by the head-on overlap of two orbitals (can be s-orbitals or p-orbitals)

41
Q

pi bond

A

formed by the sideways overlap of p-orbitals

42
Q

which is stronger between sigma and pi bonds?

A

sigma bonds are stronger than pi bonds as there is more overlap between orbitals

43
Q

sigma and pi bonds - double bonds

A

one sigma, one pi

44
Q

sigma and pi bonds - triple bonds

A

one sigma, two pi

45
Q

ionic bonds (facts)

A
  • transfer of electrons
  • ions formed
  • high melting point + boiling points
  • usually solid at room temp
  • conduct electricity when dissolved in water or molten (ions are free to move in these states)
46
Q

covalent bonds (facts)

A
  • sharing electrons
  • covalent bonds formed
  • usually low melting + boiling points
  • usually liquid or gaseous at room temp
  • do not conduct electricity (no ions present)
47
Q

sigma bonds (facts)

A
  • stronger
  • more overlap
  • both s and p orbitals
  • head on overlap
48
Q

pi bonds (facts)

A
  • weaker
  • less overlap
  • only p overlap
  • sideways overlap
49
Q

test for anions experiment

A

in hback

50
Q

electronegativity differences - polarity

A

0 - 0.4 -> non-polar

  1. 4 - 1.7 -> polar
  2. 7 - 4 -> ionic
51
Q

explain, in terms of atomic orbitals, how single bonds are formed

A

sigma, head-on overlap of atomic orbitals

52
Q

explain, in terms of atomic orbitals, how double bonds are formed

A

sigma: head-on overlap of p orbitals
pi: sideways overlap of p orabitals

53
Q

distinguish between sigma and pi covalent bonding

A
  • sigma bond formed by head on overlap of atomic orbitals

- pi bonds formed by sideways overlap of atomic orbitals

54
Q

identify in ammonia the type of (i) intramolecular (ii) intermolecular forces, present

A

intramolecular: polar (covalent) bonds within ammonia molecules
intermolecular: hydrogen bonds between ammonia molecules

55
Q

account for the difference in the shapes of the BF3 (boron trifluoride) and NH3 (ammonia) molecule

A

Boron trifluoride has three bond pairs of electrons and no lone pair

Ammonia has three bond pairs of electrons and one lone pair

56
Q

Ammonia (NH3) and Silane (SiH4) - account for difference in bond angles between them

A
  • Ammonia has 3 bond pairs and a lone pair whereas silane has 4 bond pairs (no lone pair)
  • Lone pair of electrons has greater repelling power than a bond pair of electrons
57
Q

When hydrogen bonding occurs

A

when hydrogen bonded to nitrogen, oxygen, or fluorine

58
Q

reasons why can a molecule with polar bonds be non-polar

A
  • centres of positive and negative charge coincide
  • dipole moments cancel
  • symmetrical distribution of bonds in 3-dimensional space around central atom / cancels due to symmetry of molecules
59
Q

In terms of bonding/forces, why is the boiling point of hydrogen lower than that of oxygen?

A

hydrogen has fewer electrons, (therefore) weaker intermolecular forces

60
Q

bonding/forces: why does iodine have a very low solubility in water?

A
  • iodine pure covalent (non-polar)
  • water is a polar solvent
  • intermolecular forces between iodine and water very waek
61
Q

bonding/forces: explain why: when a charged rod is held close to a thin stream of water flowing, the stream of water is deflected

A

-charge on rod attracts opposite charge on polar (dipole of) water molecule

62
Q

Ionic bonding definition

A

-Bond/Force of attraction betwen oppositely-charged ions
/
-Bond involving transfer (loss and gain) of electrons

63
Q

Covalent bonding definition

A

-Unequal sharing of bonding electrons
/
-Bond has slight positive (δ⁺) and slight negative (δ⁻) ends

64
Q

Why do ionic substances conduct electriity when molten or dissolved in water but not in the solid state?

A
  • Molten/dissolved: ions free to move

- Solid: ions not free to move/ions locked in position

65
Q

ammonia is polar covalent and is water-soluble.

Show that the molecule has polar covalent bonding

A
  • there is an electronegativity difference between N and H, showing unequal sharing
  • show electroneg diff.
66
Q

describe the processes involved when ammonia dissolves in water

A
  • hydrogen bonds between the slightly negative O of water and H of ammonia
  • and between the slightly H of water and N of ammonia

OR

  • breaking of hydrogen bonds in water
  • forming of hydrogen bonds between ammonia and water
67
Q

why boiling point of one molecule with hydrogen bonding can be smaller than boiling point of another with hydrogen bonding

A
  • weaker hydrogen bonding in one and stronger in another

- smaller electronegativity difference for (eg. NH) bond, while bigger electronegativity difference for (eg. OH) bond

68
Q

use electron pair repulsion theory to determine shape of the ammonia molecule and why bond angle is 107°

A
  • three bonding and one lone pair
  • giving bond arrangement to be pyramidal

-greater repelling power of lone pair pushes bonds closer

69
Q

illustrate hydrogen bonding in ammonia

A

know how to illustrate
(occurs between H of one ammonia molecule and the N of another)

-2009 qs has diagram

70
Q

why does methane have a low boiling point?

A

has weak intermolecular forces

71
Q

why does water deflect towards a positively charged rod whereas cyclohexane doesn’t?

what would happen if it was a negatively charged rod instead?

A
  • polarity of water causes attraction to charged rod
  • non-polarity of cyclohexane means it is not affected by charged rod

-stream of water still attracted to rod as molecules/dipoles arrange themselves with positive pole towards rod

72
Q

water - polar or non polar

A

polar

73
Q

when accounting for difference in boiling points

A

make sure to explain for both molecules.

eg. x has hydrogen bonds, while y has weak intermolecular forces

74
Q

why would something be non-soluble in water?

A
  • non-polar

- does not form hydrogen bonds with water

75
Q

conditions ionic compounds can conduct electricity

A
  • in solution/in water

- in the molten state

76
Q

label clearly any tetrahedrally bonded carbon atom in the molecule

A

just label which C atoms have a tetrahedral shape

77
Q

what type of intermolecular forces would you expect to find in nitrogen gas?

A

molecule is non-polar

78
Q

would you expect GeH₄ to be water soluble?

A

no, as GeH₄ is non-polar

79
Q

why the bonding in a molecule could be different from bonding in other molecules listed

A
  • might be non-polar
  • might be polar
  • might have hydrogen bonding
80
Q

properties affected by presence of hydrogen bonding

A
  • boiling point
  • melting point
  • solubility in water
  • surface tension
  • specific heat
81
Q

state how PH3 differs from other hydrides in terms of bonding (2012 qs)

A

virtually non-polar, due to tiny electronegativity difference between P and H in PH3, bigger electronegativity value differences in other hydrides

82
Q

“identify the molecule or molecules where there is hydrogen bonding”

A

list ALL the molecules with hydrogen bonding, not just one of them

83
Q

BCl3:
polar or non polar?: + why
B-Cl bond
BCl3 molecule -

A

B-Cl: polar, due to unequal sharing of electrons

BCl3 molecule: non-polar, cancels due to symmetry of molecule

84
Q

why boiling point of ammonia lower than bp of water

A

weaker h bonding in ammonia / smaller electroneg diff for NH bond

85
Q

account for diff in bond angle between methane (109.5) and water (104.5)

A

lone pairs in water have greater repelling power/push bonds closer

86
Q

why bp of water is higher than methane

A

stronger h bonds between water molecules

87
Q

diagram of bonds in magnesium chloride

A

ionic bond

88
Q

how to know if bond is ionic

A

check electronegativity

89
Q

drawing ionic bonds

A

know how to draw ionic bond diagrams

90
Q

when asked to draw dot and cross diagram

A

-check electronegativity first to check if bond is ionic or covalent

91
Q

hydrogen sulfide chemical formula

A

H2S

92
Q

why H2S bp is lower than water

A
  • h bonds in water

- weak intermolecular forces in H2S

93
Q

what type of intermolecular forces would you expect in nitrogen molecules? why

A
  • van der waals

- bc molecule is non polar

LC Chemistry (26 decks)

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