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S3.1 Flashcards

periodic table

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1
Q

arrangment of elements in the periodic table

A
  • elements are arranged according to their atomic number (Z)
  • vertically arranged into groups (1-18) ⇒ elements with the same number of valence electrons, hence, chemical and physical properties
  • horizontally arranged into periods (1-7) ⇒ number of principal energy levels
  • elements are arranged into 4 blocks associated with the 4 sublevels (s, p, d, f)
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2
Q

groups of the periodic table

A

group 1 = alkali metals
group 2 = alkaline earth metals
groups 3-14 = transition metals
group 15 = pnictogens
group 16 = chalcogens
group 17 = halogens
group 18 = noble gases

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3
Q

group 1

A
  • alkali metals
  • extremely reactive (reactivity increases down the group)
  • soft
  • Li, Na, K have lower density than water
  • low melting and boiling points (decrease down the group)
  • react with water ⇒ create hydroxides ⇒ stored in oil to prevent reactions with oxygen and moisture in the air
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4
Q

group 17

A
  • halogens
  • reactive non-metals
  • melting and boiling points increase down the group
  • diatomic molecules
  • reactivity decreases down the group
  • a halogen will react with a solution of halide ions, formed by a less reactive halogen (more reactive halogen molecule + less reactive ion) ⇒ the halogen displaces the ion
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5
Q

properties of metals

A

good conductors of electricity and heat
malleable, ductile, lustre
act as reducing agents
most often form cations

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6
Q

properties of non-metals

A

do not conduct electricity
act as oxidising agents
most often form anions

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7
Q

properites of metalloids

A

metallic and non-metallic properties
semiconductors ⇒ their conductivity relies on temperature

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8
Q

“tightness of hold” on electrons - trend of the periodic trend

A

Across the period (to the right) ⇒ higher effective nuclear change with the same number of principal energy levels → electrons are held tighter
Down the group ⇒ growing number of principal energy levels → larger orbitals, electrons held looser, shielding effect

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9
Q

effective nuclear charge

A

the net positive charge experienced by outer valence shell electrons from the force of attraction between protons in the nucleus and the force of repulsion between inner core electrons

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10
Q

atomic radius

A
  • = the distance from an atom’s nucleus to the outermost orbital of one of its electrons
  • usually measured as half of the distance between the nuclei of two of the same atoms bonded together (metals - half the distance between two atoms adjacent to each other in the crystalline lattice, noble gases - half the distance between two nuclei of atoms in a frozen gas)
  • the electron cloud has no definite edge
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11
Q

atomic radius is influenced by

A
  • energy level ⇒ higher energy levels are further away
  • charge on nucleus ⇒ more charge pulls electrons closer to nucleus
  • shielding effect ⇒ core electrons in inner non-valence levels of the atom reduce positive nuclear charge experienced by valence electrons
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12
Q

net charge experienced by an electron

A

Zeff = Z - S
Z … atomic number, actual charge
S … shielding constant

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13
Q

trends in atomic radius across the periodic table

A

down a group ⇒ each atom has an additional energy level → atoms get bigger
across a period ⇒ electrons are in the same energy level, larger effective nuclear charge → outermost electrons are closer → smaller radius

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14
Q

ionic radius of cations and anions

A

Cations ⇒ smaller than the original atom (as a result of losing electrons)
Anions ⇒ larger than the original atom (as a result of gaining electrons)

less electrons => smaller

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15
Q

ionic radius of ions of transition elements

A

ions of non-transition metals have noble gas configurations (isoelectronic with them)
down a group ⇒ ionic radius of ions increases
across a period ⇒ nuclear charge increases → ionic radius decreases, even with energy level differences

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16
Q

ionic radius of isoelectronic ions

A

positive ions with the most protons are smaller

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17
Q

losing electrons is an endo or exothermic reaction?

A

endothermic

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18
Q

do filled and half-filled orbirals have higher or lower energy? how does this affect IE?

A

filled and half-filled orbitals have lower energy ⇒ lower IE

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19
Q

IE down a group, across a period

A

down a group ⇒ electron is further away from the nucleus, more shielding → decreased first IE
across the period ⇒ atoms have the same energy level and shielding, nuclear charge is increasing ⇒ increased first IE

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20
Q

electron affinity

def, reaction, exo/endothermic, unit

A

= energy change associated with the addition of a mole of gaseous electrons [kJ/mol]
X (g) + e– → X– (g) ⇒ exothermic reaction

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21
Q

trend in electron affinity across the period and down the group

A

across the period ⇒ greater nuclear charge → atoms are smaller → increased electron affinity
down the group ⇒ electrons are in higher energy levels → atoms are larger → decreased electron affinity

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22
Q

what is the electron affinity of nobel gases?

A

positive

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23
Q

electronegativity

A

= tendency of an atom to attract electrons to itself when chemically combined with another element (i.e. how strong electrons are attracted to the nucleus)
higher electronegativity ⇒ pulls electrons towards itself
atoms with higher electronegativity also have negative electron affinity

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24
Q

trends across a period and down a group with electronegativity

A

down a group ⇒ decreased electronegativity
across a period ⇒ increased electronegativity

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25
Q

trend in metallic character across the period, down the group?

A

across the period ⇒ decreases
down the group ⇒ decreases

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26
Q

lewis acids and baces

A

Lewis acid ⇒ accepts electron pair
Lewis base ⇒ donates electron pair

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27
Q

trend of increasing acidic character across a period

A

increasing acidic character across a period due to greater nuclear charge, pulling electrons closer (less likely to give them away)

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28
Q

amphoteric substance

A

able to react as both an acid and base (f.i. Al2O3)

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29
Q

classifications of acids + naming

A
  1. binary acids ⇒ hydro- + non-metal atom + -ic acid
  2. oxyacids ⇒ anion:-ate, acid: -ic; anion: -ite, acid: -ous OR -ate/-ic form with oxidation number
  3. organic acids
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30
Q

hydrofluoric acid

A

HF

binary acid

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31
Q

fluoride

A

F-

anion

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32
Q

hydrochloric acid

A

HCl

binary acid

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33
Q

chloride

A

Cl-

anion

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34
Q

hydroiodic acid

A

HI

binary acid

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35
Q

iodide

A

I-

anion

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36
Q

iodic acid

A

HIO3

oxyacid

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37
Q

iodate

A

IO3-

anion

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38
Q

hydrosulfuric acid

A

H2S

binary acid

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39
Q

sulfide

A

S(2-)

anion

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40
Q

nitrous acid

A

HNO2

oxyacid

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41
Q

nitrite

A

NO2-

anion

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42
Q

nitric acid

A

HNO3

oxyacid

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43
Q

nitrate

A

NO3-

anion

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44
Q

acetate

A

HC2H3O2

organic acid

45
Q

acetate

A

C2H3COO–

anion

46
Q

methanoic/formic acid
methonate, formate

A

HCOOH
HCOO–

47
Q

ethanoic/acetic acid
ethanoate, acetate

A

CH3COOH
CH3COO–

48
Q

sulfuric acid

A

H2SO4

49
Q

sulfate

A

SO4(2-)

50
Q

sulfurous acid

A

H2SO3

51
Q

sulfite

A

SO3(2-)

52
Q

carbonic acid

A

H2CO3

53
Q

carbonate

A

CO3(2-)

54
Q

phosphoric acid

A

H3PO4

55
Q

phosphate

A

H3PO4

56
Q

hypochlorous acid

A

HClO

57
Q

hypochlorite

A

ClO-

58
Q

chlorous acid

A

HClO2

59
Q

chlorite

A

ClO2-

60
Q

chloric acid

A

HClO3

61
Q

chlorate

A

ClO3-

62
Q

perchloric acid

A

HClO4

63
Q

perchlorate

A

ClO4-

64
Q

hydrocyanic acid

A

HCN

65
Q

cyanide

A

CN-

66
Q

oxidation state

A

= the number of electrons shared or transferred when forming a bond

67
Q

oxidation state of hydrogen

A

hydrogen is combined with more electronegative elements (most non-metals) ⇒ +1
hydrogen is combined with less electronegative elements (metals) ⇒ -1

68
Q

transition metals

characteristics, ions, property similarity, role in reactions

A
  • have an incomplete d-subshell
  • able to form more ions since they can lose ions from both s and d orbitals
  • properties of transition metals are similar to each other, different from the properties of the main group elements due to similarities in valence electron configuration
  • act as surface catalysts
69
Q

surface catalyst

A

= reactant molecules are adsorbed onto a solid surface before they react with the catalyst to form the product

70
Q

which elements of the d-block are not transition metals and why

A

Zn, Cd, Hg, Cn have full d orbitals ⇒ are not transition metals

71
Q

properties of transition metals

A
  • high melting points
  • high density
  • hard
  • good electrical conductors
72
Q

magnetism of transition elements

A

paramagnetic
ferromagnetic
diamagnetic

73
Q

paramagnetism of transition elements

A

half-filled orbitals ⇒ weak attraction to external magnetic fields
half-filled orbitals have a net magnetic field

74
Q

ferromagnetism of transition elements

A

only Fe, Ni, Co
have half-filled orbitals ⇒ strongly attracted to external magnetic fields and retain the direction in which these external fields act

75
Q

diamagnetism of transition elements

A

have only filled orbitals ⇒ weakly repelled by external magnetic fields

76
Q

electron configuration of transition elements

A
  • first and second series ⇒ ns2 (n-1)dX
  • third and fourth series ⇒ ns2 (n-2)f14 (n-1)dX
  • f is filled before d - Ac and La are the exceptions – they have their d orbital filled before f sublevel, only after them is the f-sublevel filled
77
Q

formation of ions of transition metals

A

lose the ns electrons first, then the (n-1)d electrons
(n - 1)d5 and (n - 1)d10 are the most stable configurations

78
Q

what are the most stable electron configurations

A
  • all orbitals are completely filled
  • all orbitals are half-filled
  • orbitals are empty
79
Q

electron configuration of Cr

A

[Ar] 4s1 3d5

80
Q

electron configuration of Cu

A

[Ar] 4s1 3d10

81
Q

electron configuration of Mo

A

[Kr] 5s1 4d5

82
Q

electron configuration of Ru

A

[Kr] 5s1 4d7

83
Q

electron configuration of Pd

A

[Kr] 5s0 4d10

84
Q

atomic radii of transition metals

A

very similar
small increase down a group
lanthanide contraction ⇒ third transition series atoms are about the same size as the second

85
Q

lanthanide contraction

A

= decrease in expected atomic size for the transition series atoms that come after lanthanides
electrons in f orbitals have a smaller shielding ability than in other orbitals ⇒ a stronger pull on the valence electrons due to greater effective nuclear charge

86
Q

ionization energy trends of transition metals

A
  • across a period ⇒ increased first IE
  • first IE of the third series is higher than those of the first and second series due to lanthanide contraction and poor shielding
87
Q

electronegativity trend of transition metals

A
  • across a period ⇒ increases, except for the last element in the series (full d-shell)
  • electronegativity slightly increases between the first and second series, the third is about the same as the second
88
Q

oxidation states of transition metals

A

most commonly exhibit multiple oxidation states that vary by 1
3B-7B ⇒ highest oxidation state = number of the group the element is in

89
Q

complex ion

def, type of reaction, type of bond

A

monatomic cation combines with multiple monatomic anions or neutral molecules
complex ion formation is a type of Lewis acid-base reaction (metal is the base, ligands are the acid)
formed with coordinate covalent bonds

90
Q

ligand

A

= attached monatomic anions and neutral molecules to a complex ion

91
Q

coordination compound, number

A

complex ion + counterions = coordination compound (complex ion made neutral)
coordination number = number of atoms donating lone pairs to the central atom

92
Q

primary and secondary valence

A

primary valence = oxidation number of the central transition metal
secondary valence = number of ligands bonded to the metal (coordination number)

93
Q

coordinate covalent bond

A

a bond formed when the pair of electrons is donated by one atom

94
Q

chelate

A

= complex ion containing a multidentate ligand
the multidentate ligand in a complex ion is called a chelating agent (EDTA)

95
Q

types of ligands (dentates)

A
  1. unidentate ligand ⇒ has one lone pair of electrons, creating only one coordinate covalent bond with the central atom
  2. bidentate ⇒ two lone pairs of electrons, can create 2 coordinate covalent bonds with the metal
  3. multidentate ligand ⇒ 2+ lone pairs, same number of possible bonds with the metal
96
Q

geometry of complex ions

A

2 ligands ⇒ linear
4 ligands ⇒ square planar or tetrahedral
6 ligands ⇒ octahedral

97
Q

what does bonding of transition metal complexes involve?

A

the d orbitals of the central metal atom ⇒ a set of 5 degenerate orbitals

98
Q

how does the crystal field splitting energy occur?

A

bonds in coordination compounds form due to the attraction of the electrons on the ligand for the + charge on the metal cation → electrons on the ligands repel electrons in the d orbitals of the metal ion ⇒ energies of the d orbitals are split
the difference in energy of the orbitals ⇒ crystal field splitting energy … Δo = h

99
Q

how is the energy difference in Δo created?

A

after electrons on the ligands repel electrons of the d orbitals of the metal ion, the energies split
where ligands overlap with orbital lobes ⇒ strong repulsions ⇒ higher energy
ligands come in between orbital lobes ⇒ weak repulsions ⇒ lower energy

100
Q

size of Δo depends on

A
  1. the identity of the metal ion ⇒ down a group of metals in the same oxidation state → increased energy gap
  2. the types of ligands
  3. oxidation state of the metal ion ⇒ charge of the metal cation increases → increased energy gap
101
Q

types of ligands and effect on Δo

A

strong ligands ⇒ Δo is larger
weak-field ligands ⇒ Δo is smaller

102
Q

electron jumps between d orbitals in transition metal complexes

A

electrons can jump from a lower energy level (the orbitals between the x, y and z axes) to a higher energy level (lie along the cartesian plane) if the higher energy level orbitals are either empty, half-filled or one is half-filled
transition metal ions show many intense colours in crystals or solutions due to the energy, caused by the transitions of electrons between the split d sublevel orbitals

103
Q

application of coordinate compounds

A
  • extraction from metal ores
  • use of chelating agents in response to heavy metal poisoning
  • qualitative analysis of metal ions
  • commercial coloring agents, inks, cosmetics, paints, …
  • biomolecules (porphyrin rings that entrap the metal ions - hemoglobin Fe, chlorophyll Mg, drugs - cisplatin)
104
Q

transition elements with unstandard configurations

A
  1. Cr
  2. Cu
  3. Mo
  4. Ru
  5. Pd
105
Q

why do melting/boiling points of alkali metals decrease down the group?

A

greater size and mass => weaker metallic bonds => less energy required to break them

106
Q

why does the reactivity of halogens decrease down the group?

A

atoms become larger, more shielding => incoming electrons are not able to “add on” to the orbitals as well => smaller tendency to gain electrons, react

107
Q

why does reactivity of alkali metals increase down a group?

A

atoms become larger, more shielding => lower IE => higher tendency to give electrons, react

108
Q

why do the melting/boiling points of halogens increase down the group?

A

greater mass => stronger intermolecular (dispersion, London) forces => more energy to break them

chemistry hl (8 decks)

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