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IB CHEM HL > Topic 3 + 13 > Flashcards

Topic 3 + 13 Flashcards

1
Q

lanthanoids

A
  • 1st row of f-block
  • metal
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2
Q

actinoids

A
  • 2nd row of f-block
  • metal
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3
Q

metalloids

A
  • have characteristics of both metals & non-metals
  • similar physical properties & appearance to metals
  • similar chemical properties to non-metals
  • B, Si, Ge, As, Sb, Te, Po
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4
Q

liquid at SATP (298K, 100kPa)

A
  • Br (Bromine), Hg (Mercury)
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5
Q

gas at SATP (298K, 100kPa)

A
  • H, He, N, O, F, Cl, He, NE, Ar, Kr, Xe, Rn
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6
Q

effective nuclear charge (Zeff)

A
  • attractive positive charge of nuclear protons acting on valence electrons
  • Zeff experienced by outer e- < full nuclear charge
  • Zeff = (# of total e-) - (# of inner e-)
  • increases left to right, bottom to top
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7
Q

atomic radius

A
  • not intuitive
  • increases down a group bc # of occupied electron shells increase
  • decreases across a period bc attraction btwn nucleus & electrons increases as nuclear charge increases
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8
Q

ionic radius

A
  • positive ions have smaller ionic radius than parent atoms due to loss of outer shell
  • negative ions have bigger ionic r than parent atoms due to increased electron repulsion in outer principal E level
  • decreases from groups 1-14 for positive ions due to increase in nuclear charge (increased attraction btwn nucleus & electrons pulls outer shell closer to nucleus)
  • decreases from group 14-17 for negative ions due to increase in nuclear charge
  • increases down a group as # of electron energy levels increase
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9
Q

first ionization energy (I1)

A
  • energy required to remove 1 mole of electrons from ground state of 1 mole of gaseous atom
  • [ H(g) -> H+(g) + e- ]
  • always endothermic & positive value
  • measure of attraction btwn nucleus & outer electrons
  • increases across a period bc increase In Zeff causes increase in attraction btwn outer electrons & nucleus, making electrons harder to remove
  • decreases down a group bc increased distance btwn electron and nucleus due to # of electron energy levels increasing reduces attraction btwn e- and nucleus
  • opposite trend of atomic radii
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10
Q

small departures from first ionization trend

A
  • provide evidence for division of energy levels into sub-levels
  • group 2 [ns2] has greater I1 than group 13 [ns2p1]: p orbitals have higher energy than s orbitals
  • group 15 [ns2npx1py1pz1] has greater I1 than group 16 [ns2npx2py1pz1]: electron removed from group 16 is taken from doubly occupied p-orbital (unlike group 15). It is easier to remove as it’s repelled by its partner
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11
Q

first electron affinity

A
  • change in energy when 1 mole of electrons is added to 1 mole of gaseous atoms to form 1 mole of gaseous ions
  • [ X(g) + e- -> X-(g) ]
  • exothermic & negative
  • negative of I1 of anion
  • group 17: attract e- the most bc incomplete outer E level & high Zeff (≈+7)
  • group 1: attract e- the least bc lowest Zeff (≈+1)
  • group 2&15: maximum EA bc
    • group 2 [ns2]: added e- must be placed into p-orbital which is further from nucleus and experiences reduced electrostatic attraction due to shielding from electrons in s-orbital
    • group 15 [ns2npx1py1pz1]: added electron must occupy p-orbital that’s already singly occupied. so attraction btwn electron & atom is less than expected due to increased electron repulsion
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12
Q

electron affinity

A
  • not available for noble gases as they generally don’t form negatively charged ios
  • 2nd & 3rd EA: endothermic & positive bc added e- is repelled by the negative charge
  • e.g. 2nd EA for oxygen [ O-(g) + e- -> O2-(g) ]
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13
Q

electronegativity

A
  • measure of ability of its atom to attract electrons in covalent bond
  • measure of attraction btwn nucleus and outer electrons that are bonding
  • element with higher EN have higher electron pulling power
  • element with lower EN have lower electron pulling power
  • increases left to right across period bc increase in Zeff leads to increase in attraction btwn nucleus & bond electrons
  • decreases down a group bc bonding electrons are further from nucleus and more shielded, reducing attraction
  • doesn’t apply to group 18 bc they don’t form covalent bonds
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14
Q

compare/contrast electronegativity and ionization energy

A
  • same general trend as ionization energy
  • ionization energy is directly measured and property of gaseous atoms (only)
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15
Q

compare/contrast electronegativity and electron affinity

A
  • elements with higher EN have most exothermic EA
  • EA: property of isolated gaseous atoms
  • EN: property of atom in molecule
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16
Q

metals vs. non-metals

A
  • non-metals have higher ionization energy & electronegativity
  • metals can conduct electricity bc their valence electrons can move away from nucleus. valence electrons form a sea of electrons around fixed metal cations
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17
Q

melting point

A
  • very complex since it depends on bonding & structure of elements
  • simple for group 1 & 17
  • group 1: melting point decreases down the group bc its metallic structure is held together by attractive forces btwn delocalized outer electrons & positively charged ions and this attraction decreases with distance
  • group 17: melting point increases down the group bc its molecular structure is held together by weak intermolecular forces (London dispersion forces) and this attraction increases with number of electrons in molecule
  • generally rises across a period, max at group 14, fall back to reach minimum at group 18
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18
Q

group 3 element that is not solid at room temperature

A

chlorine and argon

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19
Q

chemical properties

A
  • detemrined by electron configuration of its atoms
  • elements of same group have similar chemical properties as they have same number of valence electrons
  • group 18: colorless gas, monoatomic, very unreactive bc of inability to lose/gain electrons
    • don’t form cations bc highest IE
    • don’t form anions as extra e- would need to be added to an empty outer energy level shell where e- would experience negligible Zeff
    • stable octet except Helium
  • all other atoms: reactive bc unstable incomplete electron energy levels. lose/gain electrons to achieve electron configuration of nearest noble gas
    • group 1, 2, 13: lose electrons, generally metals
    • group 15, 16, 17: gain electrons, generally non-metals
  • metalloids: intermediate properties
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20
Q

group 1: alkali metals

A

silvery, too reactive to be found in nature (stored in oil to prevent contact with water/air)
- reactivity increases down a group as elements with higher atomic number have lower IE
- react with water to produce hydrogen & metal hydroxide
- lithium: floats & reacts slowly, releases hydrogen but keeps its shape
- sodium: reacts vigorously with release of hydrogen, heat produced is used to melt metal, forming small ball that moves around water surface
- potassium: reacts more vigorously to produce sufficient heat to ignite the hydrogen produced
- e.g. [ 2K(s) + 2H2O(l) -> 2K+(aq) + 2OH-(aq) + H2(g) ]

21
Q

physical properties of alkali metal

A
  • good conductors of electricity & heat
  • low density
  • grey shiny surfaces when freshly cut with knife
22
Q

chemical properties of alkali metal

A
  • very reactive
  • form ionic compounds with non-metals
23
Q

group 17: halogens

A
  • exist as diatomic molecules (X2)
  • accept electrons bc nuclei have high Zeff, exerting strong pull on any electrons from other atoms
  • reactivity decreases down a group as atomic radius increases & attraction for outer electrons decrease
24
Q

physical properties of halogens

A
  • colored
  • show gradual change from gases (F2, Cl2) to liquid (Br2) and solids (I2, At2)
25
Q

chemical properties of halogens

A
  • very reactive non-metals
  • reactivity decreases down the group
  • form ionic compounds with metals
  • form covalent compounds with other non-metals
26
Q

group 17 & 1 reactions

A
  • halogens reacting with alkali metals form ionic halides: halogen atom gains one electron from group 1 element and forms halide ion X-, resulting in both ions having stable octet
    • e.g. [ 2Na(s) + Cl2 (g) -> 2NaCl(s) ]
      • electrostatic force of attraction btwn oppositely charged Na+ & Cl- ions bonds ions together
  • most vigorous reaction occurs btwn alkali metal at bottom of group 1 & halogen at top of group 17
27
Q

displacement reaction

A
  • seeing relative reactivity of elements by placing them in direct competition for extra electron
  • e.g. [ 2Br-(aq) + Cl2(aq) -> 2Cl-(aq) + Br2(aq) ]
    • Cl nucleus has stronger attraction for electron than Br nucleus due to smaller atomic radius, so Cl takes electron from Br ion to form Cl- while Br- forms Br
28
Q

identifying halide ion in solution

A
  • halogens form insoluble salts with silver
  • [ Ag+(aq) + X-(aq) -> AgX(s) ]
  • adding solution containing a halide to solution containing silver ions produces precipitate which is useful for identifying halide ion in solution (based on color)
29
Q

bonding of period 3 oxides

A
  • all emmebts in peroid 3 can bond with oxygen except Ar (noble gas, doesn’t interact w/ oxygen)
  • ionic compounds: formed btwn metal & non-metal so oxides of Na, Mg, Al have giant ionic structures (crystal)
  • covalent compounds: formed btwn non-metals so oxides of P, S, Cl are molecular covalent
  • oxide of silicon (metalloid): giant covalent
  • ionic character of compound depends on difference in EN btwn elements
    • ionic character of oxides decrease from left to right across a period as EN increases toward 3.4 (=EN O), increase down group as EN decreases
30
Q

sodium oxide

A
  • Na2O(s)
  • +1 oxidation number
  • high electrical conductivity in molten state
  • giant ionic
  • base
31
Q

magnesium oxide

A
  • MgO(s)
  • +2 oxidation number
  • high electrical conductivity in molten state
  • giant ionic
  • base
32
Q

aluminum oxide

A
  • Al2O3 (s)
  • +3 oxidation number
  • high electrical conductivity in molten state
  • giant ionic
  • amphoteric
33
Q

silicon dioxide

A
  • SiO2 (s)
  • +4 oxidation number
  • very low electrical conductivity in molten state
  • giant covalent
  • acidic
34
Q

phosphorus forming oxide

A
  • P4O10 (s): phosphorus (V) oxide
  • P4O6 (s): phosphorus (III) oxide
  • no electrical conductivity in molten state
  • molecular covalent
  • acidic
35
Q

sulfur forming oxide

A
  • SO3 (l): sulfur trioxide, +6
  • SO2 (g): sulfur dioxide, +4
  • no electrical conductivity in molten state
  • molecular covalent
  • acidic
36
Q

chlorine forming oxide

A
  • Cl2O7 (l): dichloride heptoxide, +7
  • Cl2O (g): dichloride monoxide, +1
  • no electrical conductivity in molten state
  • molecular covalent
  • acidic
37
Q

electron configuration & atomic radii of first row d-block

A
  • simillar, relatively small range in atomic radii
  • relatively small decrease in atomic radii across the period due to correspondingly small increase in Zeff experienced by outer 4s electrons
    • increase in nuclear charge is largely offset by addition of electrons in 3d sub-level
  • similarity in atomic radii explains how transition metals can form alloys: atoms of one d-block can be replaced w/ atoms of another w/o too much disruption of solid structure
  • small increase in Zeff accounts for small range in I1
38
Q

alloy

A
  • substance that combines more than one metal or mixes metal with other non-metallic elements
  • e.g. brass: copper + zinc
39
Q

physical properties of 1st row d-block elements

A
  • high electrical & thermal conductivity
  • high melting point
  • malleable
  • high tensile strength
  • ductile
  • iron, cobalt, nickel: ferromagnetic (forms permanent magnet)
  • bc strong metallic bonding in elements
    • 3d & 4s electrons are close in energy and all are involved in bonding & forms parts of delocalized sea of electrons that holds metal lattice together. this large # of delocalized e- accounts for the strength of metallic bond & high electrical conductivity
40
Q

chemical properties of 1st row d-block elements

A
  • form compounds with more than one oxidation number
  • form variety of complex ions
  • form coloured compounds
  • act as catalyst when either elemtns or compounds
41
Q

zinc

A
  • not transitional metal
  • doesn’t form colored solutions
  • only +2 oxidation state in its compounds
    • both Zn atom and Zn2+ atom have complete d sub-level
  • Sc3+ is colorless in aqueous solution (no d electrons) but still a transition metal bc its atom has incomplete d sub-shell
42
Q

oxidation state of 1st row d-block

A
  • all transitional metals show both +2 & +3 oxidation state (M3+ ions for Sc-Cr, M2+ more common for later ones bc hard to remove 3rd e-)
  • maximum oxidation state at manganese (use of both 3d & 4s electrons)
  • oxidation states above 3+ generally show covalent character
43
Q

complex ions

A
  • formed when transition metal ions in solution have high charge density and attract water molecules which form coordinate bonds with positive ions
  • central ion is surrounded by ligands possessing lone pair of electrons
44
Q

coordination number

A

number of coordinate bonds from ligands to central ion

45
Q

EDTA 4-

A
  • polydentate ligand, six atoms (4 O, 2 N) with lone pairs available to form coordinate bonds
  • hexadentate ligand
  • occupy all octahedral sites & grip the central ion in chelate (six-pronged claw, important in food)
46
Q

polydentate ligand

A
  • ligand with more than one lone pair available to form coordinate bond with central transition ion
47
Q

splitting

A
  • transitional metals absorb light bc d-orbitals split into 2 sub-levels
  • octahedral: 3 (lower E) & 2 (higher E) splitting pattern
  • change in energy depends on
    • nuclear charge & identify of central metal ion
    • charge density of ligand (higher E -> higher charge density)
    • geometry of complex ion
    • # of d electrons present (oxidation #)
48
Q

how to find the visible color when replacing ligand

A

1) replacing with ligand with higher E -> increase change in energy -> shorter wavelength -> absorbs color +1 to the right on color wheel -> sees the color on the opposite side

IB CHEM HL (5 decks)

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