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hl chemistry > unit 2 > Flashcards

unit 2 Flashcards

(75 cards)

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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 (preferably 8 valence electrons)

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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
large molecules are more/less volatile
more
26
what allows substance to conduct electricity
- must contain charged particles that can move - ionic compounds must have cations and anions
27
shape of water molecule surrounding positive ion (solubility)
. V > O < ^
28
shape of water molecule surrounding negative ion (solubility)
. ^ < O > V
29
lattice enthalpy meaning
measure of the strength of an ionic bond/energy required to build lattice structure
30
factors affecting lattice enthalpy
ionic radius, ionic charge
31
solubility meaning
maximum quantity of solute that can dissolve in certain amount of solvent/solution and certain temp
32
main factors of solubility
1. nature of solute and solvent 2. molecular size 3. temperature
33
how does the nature of solute and solvent affect solubility
less soluble when: solute-solute strength > solute- solvent attraction more soluble when: solute-solute strength
34
what molecular size is most soluble
- generally, smaller molecules are more soluble than larger ones
35
how does higher temperature affect solubility of solid solute
- an increase in temperature of solution increases solubility - more energy to break bonds
36
how does temp affect solubility of liquid solute
- not greatly affected - bonds aren't as strong as solid, extra energy is not needed
37
how does temp affect solubility of gas solute
- 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
what is a solubility curve
graph experimentally determined to determine solubility of any solute at any temp range
39
how does pressure affect solubility of solids and liquids
almost no effect
40
how does pressure affect solubility of gases
- as pressure increases, solubility increases - collision frequency b/w gas and liquid molecules increase
41
what are factors of rate of solution
1. size of particles 2. stirring 3. amount of solute already dissolved 4. temperature
42
creating a bond b/w water and an ion is exothermic/endothermic
exothermic
43
breaking a bond b/w ions and lattice is exothermic/endothermic
endothermic
44
guidelines for substance to be soluble
more than or equal to 1g solute/100cm^3 H2O can dissolve
45
what makes a substance insoluble
less than or equal to 0.1g solute/100cm^3 H2O can dissolve
46
what makes a substance slightly soluble
0.1 - 1.0g solute/100cm^3 H2O is soluble
47
why are compounds of ions with small charges more soluble than ones with large charges?
- the greater the charge, the greater the forces that hold them together
48
what does the dissociation equation represent
- soluble ionic compounds dissociate/separate into existing ions in water - represents separation of ions
49
supersaturate solution meaning
- more solute dissolved than able at certain temp - occurs when solid solute is put in high temp solvent then cooled down
50
how is covalent bond formed
by electrostatic attraction b/w a bonding pair of electrons and positively charged nucleus
51
a covalent bond is a region of relatively low/high electron density
high
52
how is distance of covalent bond determined
- 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
bond types vs bond lengths
single bonds = longest bond length triple bonds = shortest bond length
54
bond types vs bond strengths
single bond = weakest bond triple bond = strongest bond
55
what did nevil sidgwick discover about lewis theory of bonding
- one atom could contribute both electrons shared - octet is desirable but not necessary, especially in low d orbitals
56
nomenclature meaning
specific name given to molecule
57
molecular formula vs. empirical formula
includes actual number of atoms of each element in a molecule vs. simplest whole number ratio of atoms molecule contains (good for ionic)
58
what is empirical formula derived from
percentage composition data from combustion analysis
59
what is the full structural formula
shows every bond + atom - shown using 180, 90, 120 degree angles
60
what is the condensed structural formula
- omits bonds where can be assumed, groups atoms together - brackets are used to indicate locations of branched chains
61
what is the skeletal formula
- 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
62
what is stereochemical formula
- shows relative positions of atoms in 3d
63
what are resonance structures
- set of lewis structures that describe bonding of atoms when one is insufficient
64
____ formal charge represents most stable structure
low
65
what is formal charge used for
compare stability of non-equivalent lewis structures to find most stable
66
if it's unable to make a formal charge of a molecule 0, which atom should get the charge?
most electronegative atoms should get charges if there are any
67
guidelines for determining solubility: #1 (common salts + ammonium...)
all common salts of group 1 elements and ammonium are soluble
68
guidelines for determining solubility: #2 (acetates+nitrates...)
all common acetates + nitrates are soluble
69
guidelines for determining solubility: #3 (group 17 + metals...)
all binary compounds of group 17 (except F) w/ metals are soluble (except silver, mercury, lead)
70
guidelines for determining solubility: #4 (sulfates...)
all sulfates are soluble except barium, strontium, lead, calcium, silver mercury
71
guidelines for determining solubility: #5 (carbonates, oxides, hydroxides, phosphates...)
except for ones in rule 1, carbonates, hydroxides, oxides, phosphates are insoluble
72
what makes a compound have high conductivity in solid state
- if electrons can move around freely, carrying energy along material - ex. delocalized electrons of metals
73
why would a solution feel cold when dissolving a solute in solvent?
- 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
74
non polar covalent bonds definition
- when electrons are shared more evenly compared to polar covalent - occurs b/w atoms of same element (pure covalent) or symmetrical molecules
75
why do products of metallic bond have high luster
- electrons get excited when they absorb energy, as they return to original state, release energy emitting light

hl chemistry (7 decks)

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