Bonding Flashcards

1
Q

What is ionic bonding?

A

the electrostatic force of attraction
between oppositely charged ions formed by electron transfer.

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2
Q

What structure do ionic compounds have?

A

giant ionian lattice of alternating + and - ions

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3
Q

When is ionic bond stronger and the meltingpoints higher?

A

when the ions are smaller or have higher charges

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4
Q

What happens to the size of the ionic radi as you go down a group?

A

it is increases because there are more shells of electrons

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5
Q

When do ions form?

A

when electrons are transferred between elements that have a large difference in electronegativity (metals and non metals)

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6
Q

Are positive or negative ions smaller?

A

positive because they have lost electrons

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7
Q

What forces hold the ionic lattice together?

A

very strong electrostatic forces

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8
Q

What are the properties of ionic bonding?

A

-very high melting point
-soluble in h20
-electrical insulator when solid
-electrical conductors whenmolten/dissolved
-brittle

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9
Q

Why do ionic bonds have a very high melting point?

A

strong electrostatic forces between oppositely charged ions

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10
Q

Why are ionic bonds soluble in h20?

A

h20 is very polar so can disrupt forces between ions and break them up

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11
Q

Why are ionic bonds brittle?

A

when moved ions are no longer alternately arranged so they repel each other and the lattice breaks

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12
Q

What is covalent bonding?

A
  • the sharing of outer electrons to achieve a full shell
  • with electrostatic attraction between the shared electrons and the positive nucleus.
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13
Q

What types of covalent bonds exist?

A

Single, double, and triple bonds,

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14
Q

What is a dative covalent (coordinate) bond?

A
  • A bond where one atom donates a pair of electrons to another atom
  • For example, ammonia (NH₃) donates a lone pair to H⁺.
  • formed when an electron deficient atom accepts a lone pair of electrons from an atom with a lone pair of electrons
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15
Q

How is a dative covalent bond represented?

A
  • With an arrow showing the direction of electron donation
  • from the donor to the acceptor
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16
Q

What is the structure of graphite?

A
  • consists of hexagonal layers
  • each carbon bonded three times, and the fourth electron delocalized.
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17
Q

Why can graphite conduct electricity?

A

Delocalized electrons can carry charge, enabling conductivity

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18
Q

Why does graphite have a high melting point?

A

Strong covalent bonds require a lot of energy to break.

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19
Q

Why is graphite slippery and low density?

A

Weak intermolecular forces between layers allow them to slide, and the layers are far apart compared to covalent bond lengths.

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20
Q

What is the structure of diamond?

A

Each carbon atom is bonded four times in a tightly packed lattice.

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21
Q

Why is diamond non-conductive?

A

no free electrons to charge

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22
Q

What are the similarities between diamond and graphite?

A
  • Both have high melting points due to strong covalent bonds
  • both insoluble in water because the bonds are too strong to break.
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23
Q

What are the names and bond angles of molecule with no lone pairs?

A
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24
Q

What are the names and bond angles of molecules WITH lone pairs?

A
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25
Q

What happens to bond angles with 4 bond pairs and 2 lone pairs?

A

bond angle remains at 90 as Lone pairs repel equally, so the bond angles remain unchanged.

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26
Q

What is electronegativity?

A

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

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27
Q

What makes a covalent bond polar?

A

shared pair of electrons are not shared equally

28
Q

Why is chlorine (Cl₂) nonpolar?

A

Both atoms have equal electronegativity, so electrons are shared equally, so molecules are symmetrical

29
Q

Why are hydrocarbons nonpolar?

A

Electrons are shared equally in C–H bonds, and the molecules are symmetrical

30
Q

What makes water a polar molecule?

A

Oxygen is more electronegative than hydrogen, creating uneven charge distribution

31
Q

Can a molecule with polar bonds be nonpolar?

A

YES
if the molecule is symmetrical (e.g., CO₂), the dipoles cancel out.

32
Q

What are van der Waals forces?

A

Weak intermolecular forces due to temporary dipoles.

33
Q

How do van der Waals forces form?

A
  • Temporary dipoles arise when electrons unevenly distribute.
  • These induce dipoles in neighboring molecules, causing weak attraction.
34
Q

How does molecular size affect van der Waals forces?

A

Larger molecules with more electrons have stronger van der Waals forces

35
Q

What are dipole-dipole forces?

A

Forces between molecules with permanent dipoles.

36
Q

How do dipole-dipole interactions compare to van der Waals?

A

They are STRONGER

37
Q

What is an instantaneous dipole?

A
  • A temporary dipole created by the random movement of electrons in an atom or molecule
  • causes uneven charge distribution at any given moment.
  • happens in an instant
38
Q

What is an induced dipole?

A

A dipole created when an instantaneous dipole in one molecule induces a dipole in a nearby molecule by attracting or repelling its electrons

39
Q

What is a Dipole?

A

a dipole occurs when there is a separation of charge within a molecule due to differences in electronegativity

40
Q

What is hydrogen bonding?

A

A strong type of dipole-dipole interaction between hydrogen and highly electronegative atoms like N, O, and F

41
Q

Which atoms must be present for hydrogen bonding to occur?

A

Nitrogen (N)
Oxygen (O)
Fluorine (F)
- bonds will form between lone pairs of N, F OR O and a Hd+

42
Q

Why is ice less dense than liquid water?

A
  • In liquid water, hydrogen bonds constantly break and reform as molecules move about
  • In ice, the hydrogen bonds hold the molecules in fixed positions; this makes them slightly further apart than in liquid water
43
Q

Why does HF have a higher boiling point than HCl?

A

HF exhibits hydrogen bonding, which is stronger than dipole-dipole forces in HCl.

44
Q

Do molecules with hydrogen bonding also have other intermolecular forces?

A

Yes, they also have van der Waals and dipole-dipole forces

45
Q

What is the importance of hydrogen bonding?

A
  • ice is less dense than water
  • water has a much high melting and boiling point than would be expected
  • protein folding
  • DNA base pairing
    Enzyme reactions
46
Q

What is metallic bonding?

A

Electrostatic attraction between positive metal ions and delocalized electrons in a metallic lattice

47
Q

How does the number of delocalized electrons affect metallic bonding?

A

More delocalized electrons result in stronger bonding and higher melting points

48
Q

Why are metals good conductors of electricity?

A

The delocalized electrons can move freely and carry an electrical charge

49
Q

Why are solid metals insoluble?

A

The metallic bonds are too strong to be broken by interaction with solvents

50
Q

Do ionic compounds conduct electricity? why?

A

Yes
When molten as the ions are free to move and carry charge

51
Q

what is simple molecular covalent bonding?

A
  • strong covalent bonds between atoms
  • weak van der waals forces of attraction between molecules
52
Q

Are there any lone electrons in simple covalent bonding?

53
Q

Can simple molecular covalent molecules conduct electricity? why?

A

NO
all electrons used in boding and aren’t free to move

54
Q

Do simple molecular substances have a high/low melting point? why?

A

Low - weak van der waals forces of attraction between molecules that don’t take much energy to overcome

55
Q

Describe macromolecular covalent bonding

A

Lattice of many atoms held together by strong covalent bonds

56
Q

Do metallic compounds conduct electricity? why?

A

Yes as delocalised electrons can move throughout the metal to carry charge

57
Q

How does the strength of metallic bonds change across the periodic table?
Why?

A
  • Increases → higher Melting and boiling points, stronger Higher charge on metal ions
  • More delocalised electrons per ion
    Stronger force of attraction between them
58
Q

What affects electronegativity

A
  • nuclear charge
  • atomic radius
  • electron shielding
59
Q

What is the strongest type of inter-molecular force?

A

Hydrogen bonding

60
Q

Are van der waals forces stronger in smaller or larger molecules

A

Larger - more electrons

61
Q

describe permanent dipole-dipole attraction

A
  • some molecules with polar bonds have permanent dipoles (forces of attraction between those dipoles and those of neighboring molecules)
62
Q

What does the electron pair repulsion theory state?

A

the electron pairs will take up positions as far away from each other as possible to minimise the repulsive forces between them

63
Q

Which experience the most repulsion?

A

LP-LP repulsion strongest
LP-BP repulsion middle
BP-BP repulsion weakest

64
Q

What are the similar structural and bonding features of diamond and graphite

A
  • Both consist of entirely carbon atoms
  • Both have covalent bonding
  • Both have giant covalent (crystal lattices) structures
65
Q

What is a giant ionic lattice?

A

Repeating/ alternating pattern of oppositely charged ions

66
Q

Why are metals malleable?

A

Layers can slide over each other

67
Q

Why is graphite soft?

A

Weak van der waals forces between its layers