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Higher Chemistry > Bonding and Structure > Flashcards

Bonding and Structure Flashcards

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1
Q

covalent bonding definition

A

electrostatic attraction between positive nuclei and the shared pair of negative electrons.

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2
Q

list of metallic elements

A

Li, Be, Na, Mg, Al, K, Ca

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3
Q

list of covalent molecular (diatomic) elements

A

H2, N2, O2, F2, Cl2

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4
Q

list of discrete covalent molecular elements

A

P4, S8, C60 (fullerene)

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5
Q

list of covalent network elements

A

B, C, Si

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6
Q

list of monatomic elements

A

He, Ne, Ar… (noble gases)

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7
Q

metallic bonding definition

A

electrostatic attraction between positive metal ions and a sea of delocalised negative outer electrons

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8
Q

London dispersion forces definition

A

the weak forces between all atoms and molecules including monatomic elements. (weakest form of intermolecular bonding)

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9
Q

strength of LDF

A

the strength of LDF is related to the number of electrons within an atom or molecule. The more electrons the stronger the LDF will be.

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10
Q

difference between intermolecular and intramolecular forces

A

intramolecular- within molecules

intermolecular- between molecules

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11
Q

metallic bonding properties

A
  • conduct electricity when solid or liquid due to delocalised electrons.
  • high density due to closely packed lattice structure
  • usually high bp (lots of bonds to break)
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12
Q

mp and bp of metallic elements trend

A
  • increase across period because more electrons in outer shell, the stronger the metallic bond.
  • decrease down a group because the additional filled electron shells weaken electrostatic forces between nuclei and delocalised electrons.
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13
Q

monatomic elements properties

A
  • exist as single atoms
  • form no bonds due to full outer shell (no intramolecular bonding)
  • do not conduct electricity because there are no free electrons.
  • contain LDF
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14
Q

covalent diatomic properties

A

-low mp/bp because it is the weak intermolecular forces which break between the molecules.

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15
Q

covalent diatomic mp/bp trend

A

down the group, mp/bp increase because the strength of the LDF is increasing due to larger atoms with more electrons.

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16
Q

Fullerene

A
  • one of the three forms of carbon.
  • smallest is C60, known as buckminsterfullerene
  • does not conduct electricity as no free electrons.
  • NOT A COVALENT NETWORK
  • large molecule=strong LDF
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17
Q

covalent network properties

A
  • every atom linked to another by strong covalent bonds

- high mp/bp because lots of energy needed to break strong covalent bonds

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18
Q

covalent network forms of carbon

A

DIAMOND
-tetrahedral structure
-does not conduct electricity (no free electrons)
-hardest natural substance (used for drills/cutting tools)
GRAPHITE
-layered structure with LDF between layers
-very soft (layers break away due to weak LDF)
-conducts electricity (delocalised electrons)

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19
Q

allotrope

A

different form of same element

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20
Q

discrete covalent molecular mp/bp

A

low because it is the weak intermolecular forces which break between molecules not covalent bonds within molecule

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21
Q

noble gases mp trend

A

increases down the group as larger atoms with more electrons mean stronger LDF.

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22
Q

3 types of Intramolecular bonding

A
  1. covalent molecular (polar or non-polar)
  2. Metallic
  3. Ionic
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23
Q

covalent bonding definition

A
  • shared pair of electrons
  • both nuclei try to pull electrons towards themselves
  • creates a strong bond between the two atoms.
24
Q

pure/non-polar covalent bonding definition

A

when both atoms have an equal ‘pull’ on the shared electrons because they have the same electronegativity.
the electrons sit in the middle of the two nuclei.
Eneg difference 0.0-0.3

25
Q

polar covalent bond definition

A
  • a bond formed when the shared pair of electrons in a covalent bond are not shared equally.
  • this is due to different elements having different electronegativities.
  • Eneg difference 0.4-2.5ish
26
Q

dipole moment

A

shows direction electrons are being pulled in

27
Q

the bonding continuum

A

a spectrum which starts at pure/non-polar covalent and ends at ionic

28
Q

how to work out if a molecule is polar or non-polar

A
  1. is there a polar bond in the molecule (E difference>0.3) –> if not, non-polar
  2. is the molecule symmetrical? yes=non-polar no=polar
29
Q

testing polarity

A

non-polar: charged rod does no attract or repel a stream of water.
polar: charged rod attracts/repels stream of water.

30
Q

Ionic bond

A
  • usually between metal and non-metal with a large E difference
  • forms a lattice of + and - ions
  • electrostatic force of attraction between positively and negatively charged ions
  • non metal (high eneg) gains electron to form negative ion
  • element with low eneg loses electron to form positive ion.
  • formula tells you ratio of ions
31
Q

increasing ionic character/polarity

A

increasing electronegativity difference

32
Q

3 types of intermolecular/ Van der Waals’ forces

A
  1. London Dispersion Forces
  2. Permanent dipole-permanent dipole interactions
  3. Hydrogen bonds
33
Q

what do intermolecular bonds dictate

A

melting/boiling points of molecules

34
Q

are intra or inter stronger?

A

intramolecular bonds are stronger (what holds atoms together)

35
Q

permanent dipole-permanent dipole

A
  • occurs between polar molecules
  • the permanent dipole in one molecule is attracted to the permanent dipole of another
  • higher mp/bp than just LDF because pdp-pdp are stronger and take more energy to break.
36
Q

what is needed for a fair comparison between the mp/bp of a polar and non-polar molecule

A

similar formula mass

37
Q

polar or non-polar if only has LDF

A

non-polar

38
Q

polar or non-polar if molecule has LDF and pdp-pdp

A

polar

39
Q

hydrogen bonding

A
  • strongest intermolecular bonding
  • occurs between molecules where there is an atom of H joined to an atom of N,O,F.
  • higher mp/bp as they are held most tightly together.
  • extreme case of pdp-pdp due to very large difference in electronegativity
40
Q

properties related to intermolecular forces: melting and boiling points

A

stronger bonds (pdp-pdp/hydrogen) require more energy to break as molecules are held more tightly together.

41
Q

viscosity

A

how thick/gloopy a liquid is

42
Q

increasing viscosity trend

A

more OH groups=more H bonding=molecules are held more tightly=more viscous

43
Q

miscibility

A

how well liquids mix together

miscible liquids mix thoroughly without any visible boundary

44
Q

why are some liquids immiscible

A

if one is polar, other is non-polar; unable to bond to one another

45
Q

what molecules will mix well with water

A

polar molecules

46
Q

why does ‘like dissolve like’

A

due to attraction between the charges in each compound

polar substance+polar solvent–>soluble

47
Q

density

A

how tightly packed the particles are

48
Q

density of ice

A

ice is less dense than liquid water because hydrogen bonding between H20 molecules in ice gives a very open and expanded structure.

49
Q

what decides bonding type

A

difference in Eneg can decide bonding type

ultimately properties of a substance which dictate bonding type

50
Q

ionic bonding definition

A

electrostatic forces of attraction between positive and negative ions, usually between metal and non-metal.

51
Q

delocalised definition

A

delocalised electrons, in metallic bonding, are free from attachment to any one metal ion and are shared amongst the entire structure.

52
Q

dipole definition

A

an atom or molecule in which a concentration of positive charges is separated from a concentration of negative charges.

53
Q

isoelectronic

A

means having the same arrangement of electrons

Eg: Neon, Na+ and Mg2+

54
Q

lattice

A

a lattice is a regular 3D arrangement of particles in space. The term is applied to metal ions in a solid, and to positive and negative ions in an ionic solid.

55
Q

lone pairs

A

pairs of electrons in the outer shell of an atom which take no part in bonding.

Higher Chemistry (16 decks)

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