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chem bonding Flashcards

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1
Q

what is electronegativity?

A

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

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2
Q

how does electronegativity change across the period?

A

electronegativity increases across the period.

across the period, nuclear charge increases + shielding effect remains relatively constant -> effective nuclear charge increases -> electronegativity increases

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3
Q

how does electronegativity change down the group?

A

electronegativity decreases down the group.

in the same grp, nuclear charge and shielding effect increases, so elements have roughly the same effective nuclear charge. BUT no. of quantum shells increases -> atomic radius increases -> electronegativity decreases

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4
Q

what are the 3 most electronegative atoms

A

F, O, N

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5
Q

describe metallic bonding.

A

metals have a giant metallic lattice structure, which is held together by strong electrostatic forces of attraction between metal cations and the sea of delocalised electrons

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6
Q

**

what are the factors that affect the strength of metallic bonding?

A

number of valence electrons contributed per metal atom AND charge and radius of metal cation

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7
Q

how does the number of valence of electrons contributed per metal atom affect strength of metallic bonding?

A

number of valence electrons contributed per metal atom increases, metallic bonding stronger

larger number of valence electrons contributed per atom -> greater number of delocalised electrons -> stronger metallic bonding

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8
Q

how does charge and radius of metal cation affect the strength of metallic bonding?

A

increased charge and decreased atomic radius leads to stronger bonding

charge density = ionic charge/ionic radius

higher charge + smaller radius = higher charge density -> stronger metallic bonding

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9
Q

why do metals have high melting and boiling points?

A

a large amount of energy is required to overcome the strong electrostatic forces of attraction between cations and the sea of delocalised electrons

stronger metallic bonding -> more energy required to overcome metallic bonding -> higher melting & boiling point

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10
Q

why are metals good electrical & thermal conductors?

A

metals have delocalised electrons that act as mobile charge carriers

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11
Q

why are metals malleable and ductile?

A

when a large stress is applied to a piece of solid metal, the layers of ions will slide over one another into new positions. the overall shape of the metal changes but doesn’t break apart as the sea of delocalised electrons prevent repulsion

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12
Q

describe ionic bonding

A

ionic compounds have a giant ionic lattice structure which is a lattice of cations and anions held together by strong electrostatic forces of attraction between oppositely charged ions

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13
Q

what is lattice energy?

A

lattice energy is the heat evolved when 1 mole of pure ionic solid is formed from its constituent gaseous ions

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14
Q

why do ionic compounds have high melting and boiling points?

A

a large amount of energy is needed to overcome the strong electrostatic attractions between oppositely charged ions

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15
Q

can ionic compounds conduct electricity?

A

ionic compounds can conduct electricity in molten or aqueous state but NOT in solid state

in solid state, ions are in fixed positions and are not free to move. in molten/aqueous state, ionic compounds are good electrical conductors as ions can act as mobile charger carriers. the higher the concentration of the compound, the better its conductivity

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16
Q

why are ionic compounds hard?

A

oppositely charged ions are held in fixed positions by strong ionic bonding. large amounts of energy are required to move the ions out of position is ionic lattices are quite hard

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17
Q

why are ionic compounds brittle?

A

repulsion between ions of like charges causes the lattice to shatter

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18
Q

what is a covalent bond?

A

a covalent bond is formed when two atoms share a pair of electrons which are found in the valence orbitals of the atom

19
Q

what does bond energy (strength) depend on?

A

bond length and bond order

  1. shorter bonds (usually formed by smaller atoms) are stronger as the overlap between orbitals is more effective
  2. multiple bonds are shorter and stronger than single bonds. when more electrons are being shared, the attraction between the 2 positively charged nuclei is stronger
20
Q

what elements can have more than 8 electrons in their valence electron shell?

A

elements from period 3 onwards !!!!!

21
Q

what is VSEPR theory?

A

valence shell electron pair repulsion theory

22
Q

what are the 2 main principles of the VSEPR theory?

A
  1. electron pairs around they central atom of a molecule arrange themselves as far apart as possible to minimise their mutual repulsion
  2. the repulsion between 2 lone pairs > lone pair & bond pair > 2 bond pairs
23
Q

why is the repulsion between 2 lone pairs > lone pair & bond pair > 2 bond pairs

A

a lone pair is attracted by only 1 nucleus while a bond pair is attracted by two nuclei -> electron cloud of a lone pair is less ‘elongated’ than a bond pair so a lone pair of a central atom takes up more space around the atom -> it is closer to the central atom so it repels other electron pairs around it more than a bond pair would

24
Q

how does the percentage of s character in hybrid orbitals affect bond strength?

A

the higher the percentage of s character, the stronger the hybrid orbital is.

since the p orbital is more elongated compared to a spherical s orbital, a hybrid orbital that has more p character tends to form a longer and weaker bond

25
Q

how does difference in electronegativity affect the type of bond formed?

A

small difference in electronegativity -> covalent/nonpolar bond
large difference in electronegativity -> ionic/polar bond

26
Q

what are the 3 types of intermolecular forces of attraction?

A

instantaneous dipole-induced dipole attractions (dispersion forces), permanent dipole-permanent dipole attractions and hydrogen bonding

27
Q

where do dispersion forces occur?

A

between all molecules or between noble gas atoms

28
Q

what factors affect the strength of dispersion forces?

A

number of electrons and surface area of contact between adjacent molecules

29
Q

how do the number of electrons affect the strength of dispersion forces?

A

larger number of electrons = more polarisable = stronger dispersion forces

30
Q

how does the surface area of contact between adjacent molecules affect the strength of dispersion forces?

A

larger surface area of contact = dipoles are more easily induced = stronger dispersion forces

31
Q

where do permanent dipole-permanent dipole attractions occur?

A

between polar molecules that do not have hydrogen bonding. the attraction is between the oppositely charged ends of 2 polar molecules

32
Q

where does hydrogen bonding occur?

A

between polar molecules that contain a hydrogen atom bonded to F/O/N and an atom (F/O/N) with a lone pair of electrons

33
Q

how does hydrogen bonding occur?

A

the hydrogen atom is highly electron deficient so the lone pair on F/O/N is able to attract the electron deficient H atom strongly

34
Q

what factors affect the number of hydrogen bonds that can be formed?

A

number of hydrogen atoms and the number of lone pairs available

35
Q

what factor affects the strength of hydrogen bonding?

A

the electronegativity of the atom that H is bonded to. the more electronegative the atom is (eg F), the stronger the hydrogen bonding

36
Q

how does hydrogen bonding affect boiling point?

A
  1. more electronegative = larger electronegativity difference = stronger hydrogen bonding = more energy required to overcome = higher boiling point
  2. molecule can form more hydrogen bonds = more extensive hydrogen bonding = more energy required to overcome = higher boiling point
37
Q

how does intramolecular hydrogen bonding affect boiling point?

A

more intramolecular hydrogen bonding = less intermolecular hydrogen bonding = less energy required to overcome = lower boiling point

38
Q

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

A

the simple molecular lattices are held together by weak intermolecular forces of attraction between molecules and little thermal energy is required to overcome them

39
Q

why is diamond unable to conduct electricity?

A

it has no mobile charger carriers as all valence electrons are held between the atoms are not free to move

40
Q

why is graphite a good conductor of electricity?

A

each carbon atom in graphite is covalently bonded to 3 other carbon atoms so there are delocalised electrons that are free to move along the layers

41
Q

why is graphite soft and slippery?

A

layers of graphite are held together by weak dispersion forces so they can slide over each other easily

42
Q

rank solute-solvent interactions in terms of how strong they are

A
  1. ion-dipole interaction (strongest)
  2. hydrogen bonding
  3. pd-pd interaction (between 2 polar molecules that cannot form hydrogen bonding)
  4. pd-id interaction (between 1 polar and 1 non-polar molecule)
  5. dispersion forces (weakest)
43
Q

what interactions can be formed between water and a solute?

A

hydrogen bonding, pd-id interactions and ion-dipole interactions

chem (11 decks)

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