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Chemistry A-level > Structure - Topic 2.1 - Chemistry > Flashcards

Structure - Topic 2.1 - Chemistry Flashcards

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1
Q

what is an ionic bond

A

the strong electrostatic attraction between two oppositely charged ions

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2
Q

state 2 factors which affect the strength of an ionic bond

A

ionic radii and ionic charge

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3
Q

how does a stronger ionic charge affect the strength of an ionic bond

A

the greater the charge on an ion, the closer the ions are held together, so stronger ionic bond. This means more energy is needed to break the electrostatic forces, so a higher melting point

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4
Q

how does a smaller ionic radii affect the strength of an ionic bond

A

smaller ions can pack closer together than larger ions, as electrostatic attraction is larger at a small distance. This means ionic compounds with small ions has a higher melting point

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5
Q

why does ionic radius increase as you go down a group

A

as you go down the group, the atomic number increases, so extra electrons are added to the shells.

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6
Q

what are isoelectronic ions

A

ions of different atoms with the same number of electrons

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7
Q

why does the ionic radius of a set of isoelectronic ions decrease as the atomic number increases

A

as the atomic number increases, the number of electrons stays the same, but the number of protons increases. This means that the electrons are attracted to the nucleus more strongly

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8
Q

what are ionic crystals

A

giant lattices of ions formed when each ion is electrostatically attracted in all directions to ions of the opposite charge

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9
Q

what are the 4 physical properties of ionic compounds

A
  • high melting point
  • soluble in water, but not in non-polar solvents
  • don’t conduct electricity when solid
  • brittle
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10
Q

what is a molecule

A

formed when 2 or more atoms bond together by covalent bonds

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11
Q

what is a covalent bond

A

the strong electrostatic attraction between the two positive nuclei and the shared electrons

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12
Q

between which types of molecules does covalent bonding occur

A

between non-metals

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13
Q

between which types of molecules does ionic bonding occur

A

between a metal and a non-metal

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14
Q

where does attraction occur in a covalent molecule

A

the positive nuclei are attracted to the area of electron density between the two nuclei

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15
Q

where does repulsion occur in a covalent molecule

A

the two positively charged nuclei and repel each other, as do the electrons

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16
Q

what is bond length

A

the distance between the 2 nuclei or the distance where the attractive and repulsive forces balance each other out

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17
Q

how does a higher electron density cause a higher bond enthalpy

A

the higher the electron density, the stronger the attraction between the atoms, so the higher the bond enthalpy

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18
Q

how many electrons are shared in a carbon-carbon triple bond

A

6 electrons

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19
Q

how many electrons are shared in a carbon-carbon double bond

A

4 electrons

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20
Q

how many electrons are shared in a carbon-carbon single bond

A

2 electrons

21
Q

what is bond enthalpy

A

how much energy is needed to break a particular chemical bond. The higher the bond enthalpy, the stronger the bond

22
Q

what is a dative covalent bond

A

where one atom donates both electrons to a bond

23
Q

how does type of electron pair affect how much it repels other electron pairs

A

lone pairs repel more than bonding pairs, meaning bond angles between bonding pairs are reduced, because they are pushed together by lone pair repulsion

24
Q

what are the 2 factors affecting the shape of the molecule

A

type of electron pairs surrounding the atom and the number

25
Q

by how many degrees is the bond angle reduced by when 1 lone pair is added to the molecule

A

-2.5

26
Q

state 2 similarities between carbon and silicon

A
  • both form giant covalent structures
  • both can form 4 strong, covalent bonds
27
Q

molecular formula of nitrogen

A

N2

28
Q

molecular formula of oxygen

A

O2

29
Q

molecular formula of fluorine

A

F2

30
Q

molecular formula of Phosphorus

A

P4

31
Q

molecular formula of sulfur

A

S8

32
Q

molecular formula of chlorine

A

CL2

33
Q

describe the structure of diamond

A
  • giant covalent structure
  • each carbon atom is bonded to 4 neighbours
  • has a tetrahedral arrangement
34
Q

describe the structure of Silicon (IV) dioxide

A
  • giant covalent structure
  • each silicon atom is bonded to 4 neighbours, with oxygen atoms in between
  • has a tetrahedral arrangement
35
Q

state the 5 properties of giant covalent structures

A
  • high melting point
  • hard
  • good thermal conductors
  • insoluble
  • can’t conduct electricity
36
Q

how can graphite conduct electricity

A

delocalised electrons are free to move along the sheet, thus can carry a charge

37
Q

what is metallic bonding

A

when the positive metal ions are electrostatically attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons

38
Q

why do giant metallic lattices have a high melting point

A

because of the strong metallic bonding. Also, the more electrons, the stronger the bonding, so the higher the melting point

39
Q

why are giant metallic lattices malleable and ductile

A

the layers of metal atoms can slide over each other without breaking the metallic bonds.

40
Q

why are giant metallic lattices good thermal conductors

A

the delocalised electrons can pass on the kinetic energy

41
Q

why are giant metallic lattices good electrical conductors

A

the delocalised electrons are free to move and can carry a charge

42
Q

why are giant metallic lattices insoluble

A

because of the strength of the metallic bonding

43
Q

what is electronegativity

A

the ability of an atom to attract the bonding electrons in a covalent bond

44
Q

why does electronegativity increase across a period and up the groups

A

the more electronegative elements have higher nuclear charges and smaller atomic radii

45
Q

when is a covalent bond non-polar

A

when both atoms have similar or identical electronegativities, so the electrons will sit midway between the two nuclei and the bond

46
Q

why are covalent bonds in homonuclear, diatomic gases (H2, CL2) non-polar

A

because the atoms have equal electronegativities, so the electrons are equally attracted to both nuclei

47
Q

when is a covalent bond polar

A

If the bond is between two electrons with different electronegativities, the bonding electrons will be pulled more towards the more electronegative atom. This causes the electrons to be spear unevenly, causing a dipole

48
Q

what is a dipole

A

a difference in charge between the two atoms caused by a shift in electron density in the bond

49
Q

when is a bond purely covalent

A

between atoms of a single element, like diatomic gases because the electronegativity difference is zero, so the bonding electrons are arranged evenly

Chemistry A-level (4 decks)

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