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hl chemistry > unit 2 > Flashcards

unit 2 Flashcards

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1
Q

electronegativity meaning

A

relative attraction that an atom of an element has for the shared pair of electrons (likeliness of participating in covalent bond)

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2
Q

what is used to determine the type of bond b/w atoms

A

electronegativity value difference

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3
Q

electronegativity value difference + electron distribution of pure covalent bond

A

△Xp = 0, equal distribution

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4
Q

electronegativity value difference + electron distribution of polar covalent bond

A

0 <△Xp <= 1.8, unequal distribution

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5
Q

electronegativity value difference + electron distribution of ionic bond

A

△Xp > 1.8, ionized (no sharing)

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6
Q

what is it called when an atom is more likely to form covalent bond

A

greater covalent character

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7
Q

what is it called when an atom is more likely to form metallic bond

A

greater metallic character

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8
Q

what bond is formed b/w metal (low IE, low EN) and metal (low IE, low EN)

A

metallic

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9
Q

what bond is formed b/w metal (low IE, low EN) and non-metal (high IE, high EN)

A

ionic

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10
Q

what bond is formed b/w non-metal (high IE, high EN) and non-metal (high IE, high EN)

A

covalent

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11
Q

what is another name for the triangular bonding diagram

A

van Arkel - ketelaar triangle (in data booklet)

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12
Q

octet rule

A

atoms losing and gaining electrons to get full orbital

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13
Q

what results in formation of ionic bond

A

electrostatic force of attraction

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14
Q

ways to tell multivalent transition metals appart

A
  • visually, solutions will have diff colours
  • naming using roman numerals
  • old: latin system
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15
Q

cations are positively charged and have been ________

A

oxidized

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16
Q

anions are negatively charged and have been _______

A

reduced

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17
Q

what is the giant ionic lattice

A
  • 3D crystalline structure of anions and cations
  • fixed arrangement of ions based on a repeating unit
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18
Q

what is the coordinate number

A

expresses number of ions that surround given ion in lattice

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19
Q

what is lattice energy

A

measure of strength of attraction b/w ions within lattice

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20
Q

what ions have a higher lattice energy

A
  • small and highly charged ions since they have larger charge density
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21
Q

number of ions in lattice

A
  • very large number and can grow indefinitely
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22
Q

what does melting point depend on, what gives high melting point?

A

ionic radius and ionic charge of ions
- large charge, small ionic radius

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23
Q

volatility meaning

A

how easily a substance evaporates

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24
Q

ionic compounds have very low/high volatility. why?

A

low - strong forces of attraction b/w ions in lattice

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25
Q

large molecules are more/less volatile

A

more

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26
Q

what allows substance to conduct electricity

A
  • must contain charged particles that can move
  • ionic compounds must have cations and anions
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27
Q

shape of water molecule surrounding positive ion (solubility)

A

. V
> O <
^

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28
Q

shape of water molecule surrounding negative ion (solubility)

A

. ^
< O >
V

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29
Q

lattice enthalpy meaning

A

measure of the strength of an ionic bond/energy required to build lattice structure

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30
Q

factors affecting lattice enthalpy

A

ionic radius, ionic charge

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31
Q

solubility meaning

A

maximum quantity of solute that can dissolve in certain amount of solvent/solution and certain temp

32
Q

main factors of solubility

A
  1. nature of solute and solvent
  2. molecular size
  3. temperature
33
Q

how does the nature of solute and solvent affect solubility

A

less soluble when: solute-solute strength > solute- solvent attraction
more soluble when: solute-solute strength<solute-solvent attraction

34
Q

what molecular size is most soluble

A
  • generally, smaller molecules are more soluble than larger ones
35
Q

how does higher temperature affect solubility of solid solute

A
  • an increase in temperature of solution increases solubility
  • more energy to break bonds
36
Q

how does temp affect solubility of liquid solute

A
  • not greatly affected
  • bonds aren’t as strong as solid, extra energy is not needed
37
Q

how does temp affect solubility of gas solute

A
  • the higher the temperature, the lower the solubility of gas
  • kinetic energy in gas is so great, liquid molecules cannot hold them (heat breaks attractive force b/w liquid and gas)
38
Q

what is a solubility curve

A

graph experimentally determined to determine solubility of any solute at any temp range

39
Q

how does pressure affect solubility of solids and liquids

A

almost no effect

40
Q

how does pressure affect solubility of gases

A
  • as pressure increases, solubility increases
  • collision frequency b/w gas and liquid molecules increase
41
Q

what are factors of rate of solution

A
  1. size of particles
  2. stirring
  3. amount of solute already dissolved
  4. temperature
42
Q

creating a bond b/w water and an ion is exothermic/endothermic

A

exothermic

43
Q

breaking a bond b/w ions and lattice is exothermic/endothermic

A

endothermic

44
Q

what makes a substance soluble

A

more than or equal to 1g solute/100cm^3 H2O can dissolve

45
Q

what makes a substance insoluble

A

less than or equal to 0.1g solute/100cm^3 H2O can dissolve

46
Q

what makes a substance slightly soluble

A

0.1 - 1.0g solute/100cm^3 H2O is soluble

47
Q

why are compounds of ions with small charges more soluble than ones with large charges?

A
  • the greater the charge, the greater the forces that hold them together
48
Q

what does the dissociation equation represent

A
  • soluble ionic compounds dissociate/separate into existing ions in water
  • represents separation of ions
49
Q

supersaturate solution meaning

A
  • more solute dissolved than able at certain temp
  • occurs when solid solute is put in high temp solvent then cooled down
50
Q

how is covalent bond formed

A

by electrostatic attraction b/w a bonding pair of electrons and positively charged nucleus

51
Q

a covalent bond is a region of relatively low/high electron density

A

high

52
Q

how is distance of covalent bond determined

A
  • too close: repulsion
  • too far: no interaction
  • state of balanced attraction and repulsion requires least amount of energy to achieve, more likely to occur
53
Q

bond types vs bond lengths

A

single bonds = longest bond length
triple bonds = shortest bond length

54
Q

bond types vs bond strengths

A

single bond = weakest bond
triple bond = strongest bond

55
Q

what did nevil sidgwick discover about lewis theory of bonding

A
  • one atom could contribute both electrons shared
  • octet is desirable but not necessary, especially in low d orbitals
56
Q

nomenclature meaning

A

specific name given to molecule

57
Q

what is molecular formula

A

includes actual number of atoms of each element in a molecule

58
Q

what is empirical formula

A

simplest whole number ratio of atoms molecule contains, does not tell actual number of atoms in molecule (good for ionic)

59
Q

what is empirical formula derived from

A

percentage composition data from combustion analysis

60
Q

what is the full structural formula

A

shows every bond + atom
- shown using 180, 90, 120 degree angles

61
Q

what is the condensed structural formula

A
  • omits bonds where can be assumed, groups atoms together
  • brackets are used to indicate locations of branched chains
62
Q

what is the skeletal formula

A
  • bonds represented by lines, carbon atoms assumed where 2 lines meet/where line begins+ends
  • hydrogen atoms assumed
  • takes into account angles b/w carbon atoms
63
Q

what is stereochemical formula

A
  • shows relative positions of atoms in 3d
64
Q

what are resonance structures

A
  • set of lewis structures that describe bonding of atoms when one is insufficient
65
Q

____ formal charge represents most stable structure

A

low

66
Q

what is formal charge used for

A

compare stability of non-equivalent lewis structures to find most stable

67
Q

if it’s unable to make a formal charge of a molecule 0, which atom should get the charge?

A

most electronegative atoms should get charges if there are any

68
Q

guidelines for determining solubility: #1 (common salts + ammonium…)

A

all common salts of group 1 elements and ammonium are soluble

69
Q

guidelines for determining solubility: #2 (acetates+nitrates…)

A

all common acetates + nitrates are soluble

70
Q

guidelines for determining solubility: #3 (group 17 + metals…)

A

all binary compounds of group 17 (except F) w/ metals are soluble (except silver, mercury, lead)

71
Q

guidelines for determining solubility: #4 (sulfates…)

A

all sulfates are soluble except barium, strontium, lead, calcium, silver mercury

72
Q

guidelines for determining solubility: #5 (carbonates, oxides, hydroxides, phosphates…)

A

except for ones in rule 1, carbonates, hydroxides, oxides, phosphates are insoluble

73
Q

what makes a compound have high conductivity in solid state

A
  • if electrons can move around freely, carrying energy along material
  • ex. delocalized electrons of metals
74
Q

why would a solution feel cold when dissolving a solute in solvent?

A
  • when the bond breaking energy (endo) is greater than the bond making energy (exo)
    -ex. salt dissolving in water –> energy salt ion bonds breaking>bonds b/w salt ions and water molecules
75
Q

non polar covalent bonds definition

A
  • when electrons are shared more evenly compared to polar covalent
  • occurs b/w atoms of same element (pure covalent) or symmetrical molecules
76
Q

why do products of metallic bond have high luster

A
  • electrons get excited when they absorb energy, as they return to original state, release energy emitting light
77
Q
A

hl chemistry (4 decks)

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