1
Q
lanthanoids
A
- 1st row of f-block
- metal
2
Q
actinoids
A
- 2nd row of f-block
- metal
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
4
Q
liquid at SATP (298K, 100kPa)
A
- Br (Bromine), Hg (Mercury)
5
Q
gas at SATP (298K, 100kPa)
A
- H, He, N, O, F, Cl, He, NE, Ar, Kr, Xe, Rn
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
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
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
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
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
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
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) ]
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
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)
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
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
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
18
Q
group 3 element that is not solid at room temperature
A
chlorine and argon
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
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
chemical properties of halogens
- very reactive non-metals
- reactivity decreases down the group
- form ionic compounds with metals
- form covalent compounds with other non-metals
26
group 17 & 1 reactions
- 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
displacement reaction
- 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
identifying halide ion in solution
- 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
bonding of period 3 oxides
- 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
sodium oxide
- Na2O(s)
- +1 oxidation number
- high electrical conductivity in molten state
- giant ionic
- base
31
magnesium oxide
- MgO(s)
- +2 oxidation number
- high electrical conductivity in molten state
- giant ionic
- base
32
aluminum oxide
- Al2O3 (s)
- +3 oxidation number
- high electrical conductivity in molten state
- giant ionic
- amphoteric
33
silicon dioxide
- SiO2 (s)
- +4 oxidation number
- very low electrical conductivity in molten state
- giant covalent
- acidic
34
phosphorus forming oxide
- P4O10 (s): phosphorus (V) oxide
- P4O6 (s): phosphorus (III) oxide
- no electrical conductivity in molten state
- molecular covalent
- acidic
35
sulfur forming oxide
- SO3 (l): sulfur trioxide, +6
- SO2 (g): sulfur dioxide, +4
- no electrical conductivity in molten state
- molecular covalent
- acidic
36
chlorine forming oxide
- Cl2O7 (l): dichloride heptoxide, +7
- Cl2O (g): dichloride monoxide, +1
- no electrical conductivity in molten state
- molecular covalent
- acidic
37
electron configuration & atomic radii of first row d-block
- 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
alloy
- substance that combines more than one metal or mixes metal with other non-metallic elements
- e.g. brass: copper + zinc
39
physical properties of 1st row d-block elements
- 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
chemical properties of 1st row d-block elements
- 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
zinc
- 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
oxidation state of 1st row d-block
- 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
complex ions
- 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
coordination number
number of coordinate bonds from ligands to central ion
45
EDTA 4-
- 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
polydentate ligand
- ligand with more than one lone pair available to form coordinate bond with central transition ion
47
splitting
- 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
how to find the visible color when replacing ligand
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
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