LITERATURE Blackboard test PREVENTIVE CONSERVATION:
http://www.cci-icc.gc.ca/resources-ressources/agentsofdeteriorationagentsdedeterioration/index-eng.aspx
Please make sure you read the links underneath the caption:
Agents of Deterioration
The agents of deterioration identify 10 primary threats specific to heritage environments and encourage their
prevention at the collections level first, by avoiding, blocking and detecting possible damage, and then by
responding and treating damage. This integrated approach to conservation has also led to developments in risk
assessment.
Literature and Material for Selection CR 2016-2017
Page 1
LITERATURE for Blackboard test Principles of Conservation and Restoration:
Ed. Nicholas Stanley Price, M. Kirby Talley Jr. and Alessandro Melucco Vaccaro, Readings in
Conservation: Historical and Philosophical Issues in the Conservation of Cultural Heritage, The
Getty Conservation Institute, Los Angeles, The J. Paul Getty Trust 1996.
Part II The Original Intent of the Artist
• Introduction, pp. 162-175
• Ernst van de Wetering, Reading 19, pp. 193-199
Part III The Emergence of Modern Conservation Theory
• Introduction, pp. 202-211
• Cesare Brandi, Reading 22, pp. 230-235
• Albert France-Lanord, Reading 24, pp. 244-247
Part IV Historical Perspective
• Introduction, pp. 262-267
• Paul Philippot, Reading 26, pp. 268-274
• R. H. Marijnissen, Reading 27, pp. 275-280
Part V Restoration and Anti-Restoration
• Introduction, pp. 308-313
• Eugène-Emmanuel Viollet-le-Duc, Reading 30, pp. 314-318
• William Morris, Reading 31, pp. 319-321
• John Ruskin, Reading 32, pp. 322-323
Part VI Reintegration of Losses
• Introduction, pp. 326-331
• Max J. Friedlander, Reading 33, pp. 332-334
• Albert Philippot and Paul Philippot, Reading 34, pp. 335-338
• Cesare Brandi, Reading 35, pp. 339-342
• Paul Philippot, Reading 38, pp. 358-363
Part VII The Idea of Patina
• Introduction, pp. 366-371
• Paul Philippot, Reading 39, pp. 372-376
• Ernst van de Wetering, Reading 43, pp. 415-421
Part VIII The Role of Science and Technology
• Introduction, pp. 424-431
• Paul Coremans, Reading 44, pp. 432-438
• Giorgio Torraca, Reading 45, pp. 439-444
Literature and Material for Selection CR 2016-2017
Page 2
Literature for Blackboard testing Chemistry for Conservation 2016-2017
Fundamental (A, B, etc) and Chapter/ paragraph numbering:
Peter Atkins, Loretta Jones, Leroy Laverman – Chemical principles. The quest for insight
6th edition (W.H. Freeman and Company,New York)
Elements and atoms, Matter, energy, radiation and the quantum-mechanic model of atoms
A.3 Energy B.1 Atoms B.2 The Nuclear Model B3 Isotopes B4 The Organization of the Elements
2.1* The principal quantum number 2.2* Atomic Orbitals 2.3 Electron Spin 2.4 The Electronic
Structure of Hydrogen
Exercises: B3, B5, B9, 2.25, 2.31
*Note: 2.1 The principal quantum number and 2.2 Atomic orbitals are difficult to understand without
knowledge of mathematics and physics at university level. However, for Chemistry for Conservation
this mathematics and physics is not needed.
A short summary of the most important concepts in 2.1 and 2.2: the atomic orbital, in particular the
spatial ‘shape’ (boundary surface) of the s- and p-orbitals. The principal quantum number
n (= 1, 2, 3....) refers to an energy level (shell) relative to the nucleus and the average volume of the
orbital(s) at this level. The notation of the orbitals: the n followed by the type of orbital
(s,p,d,f), e.g.1s,2p, 3p. Which type and how many (energy-)equivalent orbitals exist for each value
of n ( = 1, 2,3,4, 5…). In each shell (each n) there is only one s orbital (1s, 2s, 3s, 4s, 5s) ; for
n = 2,3, 4,5, …there are three p orbitals (2px, 2py, 2pz , 3px, 3py, 3pz, etc) ;
for n = 3,4,5,.. there are five d orbitals; for n = 4, 5, … there are seven f orbitals.
Electronic structure of many-electron atoms, periodic table
2.5 Orbital Energies 2.6 The Building-Up Principle 2.7 Electronic Structure and the Periodic Table
Exercises: 2.37-2.40, 2.49-2.52
Ionic compounds, molecular compounds, covalent bonds
C.1 What are Compounds? C.2 Molecules and Molecular Compounds C.3 Ions and Ionic Compounds
Exercises: C7-C9, C13-C14
Covalent bonds, valence bond theory
3.5 Lewis Structures 3.6 Lewis Structures of Polyatomic Species 3.7 Resonance 3.8 Formal Charge
3.9 Radicals and Biradicals 3.10 Expanded Valence Shells
Exercises see: Exercises_Small molecules and ions relevant to Chemistry for Conservation
4.4 Sigma and pi bonds 4.5 Electron promotion and hybridization of orbitals 4.6 Other common types
of hybridization 4.7 Characteristics of double bonds
Molecular shape, electronegativity and polarity of molecules,
3.12 Correcting the Covalent Model: Electronegativity
4.1 The Basic VSEPR Model 4.2 Molecules with Lone Pairs on the Central atom 4.3 Polar molecules
Exercises see: Exercises_Small molecules and ions relevant to Chemistry for Conservation
Note: only the electron pair arrangements: linear, trigonal planar and tetrahedral
Note: only the molecular shapes: linear, bent (=angular), trigonal planar, trigonal pyramidal and
tetrahedral
Intermolecular forces
3.13 Correcting the Ionic Model: Polarizability
6.1 The origin of intermolecular forces 6.2 Ion-Dipole Forces 6.3 Dipole-Dipole Forces 6.4 London
Forces 6.5 Hydrogen Bonding 10.8 The limits of solubility 10.9 The Like-Dissolves-Like Rule
Exercises: see Exercises_Small molecules and ions relevant to Chemistry for Conservation
Literature and Material for Selection CR 2016-2017
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Acids and bases
J.1 Acids and Bases in Aqueous Solution J.2 Strong and Weak Acids and Bases J.3 Neutralization
12.1 Brønsted-Lowry Acids and Bases 12.4 Proton Exchange Between Water Molecules
12.5 The pH Scale 12.6 The pOH of Solutions 12.7 Acidity and Basicity Constants 12.8 The Conjugate
Seesaw 12.10 The Strengths of Oxoacids and Carboxylic Acids 12.13 The pH of salt solutions (till
Example 12.10, no calculations)
Exercises: 12.3acd, 12.4acd, 12.21, 12.43,12.44, 1247, 12.48 , 12.51bcde
Organic compounds and functional groups
20.1 Haloalkanes 20.2 Alcohols 20.3 Ethers 20.4 Phenols 20.5 Aldehydes and Ketones 20.6 Carboxylic
Acids 20.7 Esters 20.8 Amines, Amino Acids, and Amides
In 20.1 NOT the nucleophilic substitution
In 20.3 NOT the crown ethers
Exercises: 20.62a, 20.67, 20.68
Exercises small molecules and ions relevant to Chemistry for Conservation
a) Draw for each particle its Lewis structure, or its Lewis structures if equivalent
resonance structures exist
b) Determine for each particle the electron-pair geometry and molecular shape
at the central atom. Consider in particles 22-25 all non-hydrogen atoms as
central atoms.
c) Argue for each neutral particle whether it is expected to be polar or apolar.
d) For the neutral particles, indicate the intermolecular forces that are present
when the particles are in the liquid state
1) CO2 (g) carbon dioxide
2) NO (g) nitrogen monoxide
3) NO2 (g) nitrogen dioxide
4) CN- cyanide anion
5) SO2 (g) sulfur dioxide
6) SO3 (g) sulfur trioxide
7) H2SO4 (ℓ) sulfuric acid
8) HNO3 (ℓ) nitric acid
9) CO32- carbonate anion
10) NO3- nitrate anion
11) HNO2 (ℓ) nitrous acid
12) HClO (ℓ) hypochlorous acid
13) H3PO4 (ℓ) phosphoric acid
14) PO43- phosphate anion
15) HClO2 (ℓ) chlorous acid
16) O3 (g) ozone
17) CH3CH2OH (ℓ) ethanol
18) CH3COOH (ℓ) acetic acid (= ethanoic acid)
19) CH3COCH3 (ℓ) acetone (= dimethylketone)
20) CH3CH2OCH2CH3 (ℓ) (= diethylether)
Literature and Material for Selection CR 2016-2017
Page 4
Summary of Chemistry for the selection of potential students Master
Conservation and Restoration
A In order to pass the chemistry test successfully you should:
-
-
-
-
Know how to calculate the amount of protons, neutrons and electrons of atoms and
ions of the elements with Z = 1-57 and 72-89 using the Periodic Table (1A4 sheet,
will be supplied)
Know how to calculate the total amount of valence electrons of the atoms and ions of
the elements with Z = 1-57 and 72-89 using the supplied Periodic Table
Be able to establish whether a molecule or a polyatomic ion is a radical (odd amount
of valence electrons) or not (even amount of valence electrons)
Be able to recognize in structural formulas common anions and cations (listed on a
2A4 sheet, will be supplied)
Be familiar with the concept of atomic orbitals (regions in space around the nucleus
that have a high probability for an electron to be present)
Know that various types (s, p, d, f) of atomic orbitals exist, that s-orbitals have a
spherical shape and p-orbitals are dumb-bell shaped
Know that only valence (=outer-shell) electrons are involved in chemical bonds
Know two types of chemical bonds, the ionic bond and the covalent bond
Know that a single covalent bond (σ bond, sigma bond) is formed by an end-to-end
overlap of (atomic) orbitals
Know that for some elements double and triple bonds are possible, involving one and
two pi (π) bonds respectively
Know that a pi bond is formed by side-side overlap of parallel p-orbitals
Know that valence electrons occur in pairs, either as bonding pair (sigma bond, pi
bond) between two atoms, or as lone pair at an atom
Know that the Lewis structure of a covalently bonded molecule or a polyatomic ion
depicts all the elements and their valence electrons (in pairs), how to construct
this Lewis structure, and that in this construction the octet rule (max. eight valence
electron near a nucleus) strictly applies to the elements C, N, O and F
Know elements that from periods 3-5 can accommodate more than eight valence
electrons near the nucleus
Know that for some molecules and polyatomic ions several (equivalent) resonance
structures are needed that together describe the Lewis structure.
Know that the occurrence of resonance structures implies delocalization of pi-bond
electrons and lone-pair electrons
Know that the electron-pair geometry (=electron-pair arrangement, EG) is the spatial
arrangement of all single (covalent) bonds and lone pairs around an atom in a
molecule or polyatomic ion, and how to establish this EG (only the EG: linear,
trigonal planar, tetrahedral).
Know how to establish from the EG the molecular shape (MS) at a chosen atom in a
molecule or polyatomic ion. Only the MS: linear, bent=angular, trigonal planar,
tetrahedral and trigonal pyramid.
Know that the absolute difference (so leaving out any minus sign) in Pauling
electronegativity (Δχ) between two bonded atoms is a measure for the type of bond.
For 0 < Δχ < 0.5 the bond is non-polar covalent. For increasing Δχ (> 0.5 till 1.5) the
covalent bond becomes increasingly polar, and eventually if Δχ > 2.0 the bond is
ionic.
Literature and Material for Selection CR 2016-2017
Page 5
-
-
-
Know how to establish whether a part of a molecule at a chosen central atom is polar
or non-polar (= apolar) using the MS at this central atom and the (absolute) difference
in Pauling electronegativity between the central atom and each of the atoms bonded to
this central atom.
Know that parts of molecules that contain the elements O, N, F or Cl bonded to a C
atom are usually polar, unless the MS around this C atom is completely symmetric
Know the three possible types of intermolecular interactions: London (dispersion)
forces, dipole-dipole interactions and hydrogen bonds. When these interactions occur
and how they affect the boiling point of fluids and the miscibility of fluids.
Know how to interpret a line structure of a small organic molecule in terms of the
complete Lewis structure
Be able to recognize in a picture of a complete Lewis structure, an incomplete Lewis
structure (lone pairs left out) and a line structure of an organic molecule the following
functional groups:
alcohol, aldehyde, ketone, ether, carboxylic acid, phenol, ester, amine, amide
Know that a primary alcohol can be oxidized to an aldehyde, and an aldehyde to a
carboxylic acid
Know that a secondary alcohol can be oxidized to a ketone
Know that an ester can be hydrolyzed into an alcohol and a carboxylic acid, and vice
versa that the condensation reaction of the latter two compounds gives an ester and
water
Know that an amide can be hydrolyzed into an amine and a carboxylic acid, and vice
versa that the condensation reaction of the latter two compounds gives an amide and
water
B No questions will be asked about the following subjects:
-
-
The elements with Z = 58-71 and Z = 90-111
f-orbitals
d-orbitals, except for the total amount of valence electrons present (can be read of the
supplied Periodic Table)
diradicals
mathematical and/or physical formulas
excited states
The EG’s : trigonal bipyramidal, octahedral
The MS’s : T-shape, See-saw, trigonal bipyramidal
The MS’s: Square planar, Square pyramidal, Octahedral
The concept of hybridization
Specific names of compounds (only the functional group classes alcohol, aldehyde,
ketone, ether, carboxylic acid, phenol, ester, amine, amide)
Mathematical calculations, except addition (total amount of valence electrons) and
subtraction (calculation of Δχ)
Supplied material (have print-outs of 1-3 ready when answering the questions)
1 A 1A4 sheet Periodic Table (Wikimedia)
2 A 2A4 sheet with common cations and anions (ho1-anions-cations)
3 A 1A4 sheet with two figures from Atkins and Jones (6th edition)
Figure 3.12 Pauling Electronegativity of main group elements
Figure 4.1 Possible EG and MS at central atoms (see under A which are
applicable)
4 A 1A4 summary of the rules how to construct Lewis structures
Literature and Material for Selection CR 2016-2017
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Literature and Material for Selection CR 2016-2017
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Symbols and Charges for Monoatomic Ions
Symbol Name
Symbol Name
+
H¯ hydride
H hydrogen ion
+
Li lithium ion
F¯ fluoride
+
Na sodium ion
Cl¯ chloride
+
Br¯ bromide
K potassium ion
+
Rb rubidium ion
I ¯ iodide
+
2
Cs cesium ion
O ¯ oxide
2+
2
Be beryllium ion
S ¯ sulfide
2+
2
Mg magnesium ion
Se ¯ selenide
2+
2
Te ¯ telluride
Ca calcium ion
2+
Sr strontium ion
2+
+
3
Ba barium ion
Ag silver ion
N ¯ nitride
2+
2+
3
Ra radium ion
Ni nickel ion
P ¯ phosphide
2+
3+
3
Zn zinc ion
Al aluminum ion
As ¯ arsenide
Systematic name
(Stock system)
copper(I)
copper(II)
iron(II)
iron(III)
tin(II)
tin(IV)
chromium(II)
chromium(III)
manganese(II)
Symbol
+
Cu
2+
Cu
2+
Fe
3+
Fe
2+
Sn
4+
Sn
2+
Cr
3+
Cr
2+
Mn
Common
name
Symbol
2+
cuprous
Hg2
2+
cupric
Hg
2+
ferrous
Pb
4+
ferric
Pb
2+
stannous
Co
3+
stannic
Co
+
chromous Au
3+
chromic
Au
3+
manganous Mn
Systematic name
(Stock system)
mercury(I)
mercury(II)
lead(II)
lead(IV)
cobalt(II)
cobalt(III)
gold(I)
gold(III)
manganese(III)
Common
name
mercurous
mercuric
plumbous
plumbic
cobaltous
cobaltic
aurous
auric
manganic
Formula Name
NO3¯ nitrate
Symbols and Charges for Polyatomic Ions
Formula Name
Formula Name
Formula Name
ClO4¯ perchlorate BrO4¯ perbromate IO4¯ periodate
NO2¯ nitrite
ClO3¯ chlorate
BrO3¯ bromate
IO3¯
iodate
2¯
CrO4 chromate
ClO2¯ chlorite
BrO2¯ bromite
IO2¯ iodite
2¯
Cr2O7 dichromate ClO¯ hypochlorite BrO¯ hypobromite IO¯ hypoiodite
CN¯ cyanide
OH¯ hydroxide
MnO4¯ permanganate NH2¯ amide
2¯
SO4 sulfate
2
C2O4 ¯ oxalate
3
PO4 ¯ phosphate
2
O2 ¯ peroxide
2¯
CO3 carbonate
O2¯
superoxide
HCO3¯ hydrogen carbonate (bicarbonate)
2
HSO4¯ hydrogen sulfate (bisulfate) SO3 ¯ sulfite HSO3¯ hydrogen sulfite (bisulfite)
HC2O4¯ hydrogen oxalate (binoxalate) HCOO- formate CH3COO- acetate
2
HPO4 ¯ hydrogen phosphate H2PO4¯ dihydrogen phosphate
2
HPO3 ¯ phosphonate (phosphite) H2PO3¯ hydrogen phosphonate
Literature and Material for Selection CR 2016-2017
Page 8
More Symbols and Charges for Polyatomic ions
2
S2O3 ¯ thiosulfate
HS¯ hydrogen sulfide
CrO42- chromate
3
3
AsO4 ¯ arsenate
BO3 ¯ borate
MnO42- manganate
2¯
2¯
SeO4 selenate
B4O7 tetraborate
MnO4- permanganate
2¯
2¯
SiO3 silicate
SiF6 hexafluorosilicate
2¯
C4H4O6 tartrate
SCN¯ thiocyanate
+
The most common positively charged polyatomic ion is NH4 , the ammonium ion.
Prefixes Used to Indicate Number in a Name Involving Two Non-Metals
mono– 1 hexa–
6
di–
2 hepta–
7
tri–
3 octa–
8
tetra–
4 nona–
9
penta– 5 deca–
10
These prefixes are used in naming binary compounds involving two non–metals, e.g. P2O5, Cl2O, NO, N2O, NO2,
N2O5, PCl3, PCl5, SO2, SO3, SiO2. Sometimes metal ions are involved in a Greek prefix name, but these are less
common. Examples include UF6, SbCl3, SbCl5, OsO4, BiCl3.
There is a preferred order of the non-metals when writing them in a formula. It is: Rn, Xe, Kr, B, Si,
C, Sb, As, P, N, H, Te, Se, S, I, Br, Cl, O, F.
CO is carbon monoxide, NOT carbon monooxide. As4O6 is tetrarsenic hexoxide, NOT tetraarsenic
hexaoxide.
Acid Names – add the word acid to each name when saying or writing.
Non–oxygen containing
Oxygen containing (oxyacids)
Name when disName when a pure
Formula
solved in water
compound
Formula Name
nitric acid
HF
hydrofluoric acid
hydrogen fluoride
HNO3
nitrous acid
HCl
hydrochloric acid
hydrogen chloride
HNO2
HBr
hydrobromic acid hydrogen bromide
H2SO4
sulfuric acid
HI
hydroiodic acid
hydrogen iodide
H2SO3
sulfurous acid
phosphoric acid
HCN
hydrocyanic acid
hydrogen cyanide
H3PO4
H2S
hydrosulfuric acid hydrogen sulfide
H3PO3
phosphorous acid
H2CO3
carbonic acid
HC2H3O2 acetic acid (also written CH3COOH)
HCOOH formic acid
H2C2O4 oxalic acid
Halogen and oxygen containing (halogen oxyacids)
HOF hypofluorous acid
*
*
*
HClO hypochlorous acid HClO2 chlorous acid HClO3 chloric acid HClO4 perchloric acid
HBrO hypobromic acid HBrO2 bromous acid* HBrO3 bromic acid HBrO4 perbromic acid
HIO hypoiodic acid
*
HIO3 iodic acid
HIO4 periodic acid
* does not exist or instable
Literature and Material for Selection CR 2016-2017
Page 9
Rules for drawing Lewis structures , the (VSEPR) electron-pair geometry
and the molecular shape at central atoms in covalently bonded particles
Lewis structure ‘rules’
- The elements C, N, O and F (all period 2) strive for 8 valence electrons in the vicinity
of the nucleus (octet rule) by sharing valence electrons, but they have NEVER more than 8.
- A Lewis structure in which a C, N, O or F has much less than 8 valence electrons (e.g
only 6) is usually NOT the Lewis structure that approach the true structure the most closely.
- Elements from periods 3, 4 and higher have d-orbitals and can accommodate more than 8
valence electrons in the vicinity of the nucleus (expanded valence shells) (e.g. the S in H2SO4).
- Always calculate the formal charges of the atoms and indicate them in the Lewis structure if
they are non-zero (+1 or -1)
- If an atom has a very high (+2) or low (-2) formal charge, the Lewis structure is almost never
the best possible Lewis structure (use expanded valence shells if the element is from period 3
or higher)
- The sum of all formal charges in a particle must equal the charge of the particle ( = 0 if the
particle is a neutral molecule
- The amount of electron pairs (bond and lone pairs) present in a Lewis structure, including a
single electron in case of a radical, must equal the total amount of valence electrons in the
particle.
- Always draw a set of two electrons (bond and lone pairs) as a dash (−), NOT as a set of two
dots (this suggests a diradical).
- Pi-bond electron pairs and lone pairs can often be delocalized: always check if more equivalent (
so the same distribution of formal charges but at different atoms) resonance structures exist and:
- Always check that all resonance structures comply with the above-mentioned rules (esp. the
octet rule in case of the elements C, N, O en F!).
Chemical ‘rules’ regarding Lewis structures
- In oxoacids (acids containing O atoms), e.g. H2SO4, HNO3 etc. the H atom(s) are almost
always bonded to the O atom(s), not to the central atom (S, N). Exception: in H3PO3 is one H
bonded to the P
- 'Triangles' of covalently bonded atoms do exist (e.g cyclopropane, and epoxides) but without
C they are very instable and usually Lewis structures can be drawn that approach the true
structure more closely.
- Long sequences of single bonded N and/or O atoms are very instable (e.g. -N-N-Nand -O-O-O- do not exist in chemicals. Also a single bonded -O-O- is instable and, except for
in peroxides ( e.g. H2O2), usually Lewis structures can be drawn that approach the true
structure more closely.
- Unless a halogen atom (F, Cl, B, I) is the central atom in a particle, it always has a single bond
to another non-metal element. A double bond would imply a positive formal charge, highly
unlikely for the electronegative halogens. I3- is lineair, without double bonds.
- If a halogen atom (Cl, Br or I; not F) is the central atom in a particle (e.g. in oxoacids like
HClO4), double bonds do exist.
Electron-pair geometry and molecular shape
- When deducing the molecular shape at a central atom, always give the electron-pair geometry
(= electron-pair arrangement) at this central atom: a correct electron-pair geometry but incorrect
name of the molecular shape yields more points than just the incorrect name of the molecular
shape.
Literature and Material for Selection CR 2016-2017
Page 10
PERIODIC TABLE OF THE ELEMENTS
GROUP
IA
1
1
1.0079
1
1s
1
H
HYDROGEN
3
2
6.941(2)
1
[He] 2s
Li
LITHIUM
11
3
22.990
1
[Ne] 3s
19
4
39.098
1
[Ar] 4s
K
POTASSIUM
37
85.468
1
[Kr] 5s
5
Rb
RUBIDIUM
55
132.905
1
[Xe] 6s
6
FAMILY
2
4
(223)
1
[Rn] 7s
ATOMIC NUMBER
2
[He] 2s
78
ELECTRON
CONFIGURATION (3)
12
24.305
2
[Ne] 3s
MAGNESIUM
20
40.078
2
[Ar] 4s
Ca
CALCIUM
38
87.62(1)
2
[Kr] 5s
Sr
STRONTIUM
56
137.327
2
[Xe] 6s
Ba
BARIUM
88
(226)
2
[Rn] 7s
Fr
Ra
FRANCIUM
RADIUM
3
21
IIIB
44.956
1
2
[Ar] 3d 4s
Sc
SCANDIUM
39
88.906
1
2
[Kr] 4d 5s
Y
YTTRIUM
57-71
Lanthanides
89-103
Actinides
IVB
4
22
47.867
2
2
[Ar] 3d 4s
2
[Kr] 4d 5s
Zr
178.49(2)
14
VB
5
23
50.942
3
2
2
41
2
Hf
HAFNIUM
104
(261)
Rf
RUTHERFORDIUM
138.905
1
2
[Xe] 5d 6s
La
LANTHANUM
1
f
1
Halogens
Transition metals
Noble gases
Physical State (25°C. 1 atm)
Actinides
Ne
Hg - liquid
5
2
42
95.94(2)
5
1
43
8
26
55.845
6
2
[Ar] 3d 4s
(98)
6
VIIIB
1
3
2
74
183.84(1)
14
4
2
75
14
5
2
76
1
6
2
2
7
10
195.084
14
13
9
1
1
2
48
2
1
80
Ta
W
105
(262)
DUBNIUM
(227)
1
2
[Rn] 6d 7s
10
1
2
Ce
CERIUM
Ac
ACTINIUM
(264)
Bh
108
(277)
HASSIUM
(268)
110
(281)
Au
2
111
MERCURY
(272)
112
49
PHOSPHORUS
2
2
33
74.922
10
16
2
3
2
4
10
17
5
18
2
4
35
2
6
Ar
ARGON
79.904
10
39.948
[Ne] 3s 3p
CHLORINE
78.96(3)
6
NEON
Cl
SULFUR
2
Ne
35.453
2
20.180
[He] 2s 2p
[Ne] 3s 3p
S
10
5
F
32.065
34
2
HELIUM
FLUORINE
[Ne] 3s 3p
P
72.64(1)
10
3
18.998
[He] 2s 2p
O
30.974
2
4
OXYGEN
[Ne] 3s 3p
GERMANIUM
114.818
10
2
1
In
Hg
GOLD
32
15
2
[He] 2s 2p
VIIA
2
5
36
83.798
10
2
6
50
118.710
10
2
2
Se
As
ARSENIC
51
SELENIUM
121.760
10
Br
2
3
BROMINE
52 127.60(3) 53
10
2
Kr
KRYPTON
126.904
4
10
2
5
54
131.293
10
2
6
Sn
(285)
81
204.383
14
10
2
1
[Xe] 4f 5d 6s 6p
Tl
THALLIUM
113
Sb
(284)
82
207.2(1)
14
10
2
2
[Xe] 4f 5d 6s 6p
Pb
LEAD
114
Te
ANTIMONY
TIN
83
TELLURIUM
208.980
14
10
2
3
[Xe] 4f 5d 6s 6p
Bi
BISMUTH
(289)
115
(288)
84
(209)
14
10
2
4
[Xe] 4f 5d 6s 6p
Po
POLONIUM
116
(292)
I
Xe
IODINE
85
XENON
(210)
14
10
2
5
[Xe] 4f 5d 6s 6p
At
86
10
2
6
Rn
ASTATINE
117
(222)
14
[Xe] 4f 5d 6s 6p
RADON
118
(294)
Ds Rg Cn Uut Uuq Uup Uuh Uus* Uuo
MEITNERIUM DARMSTADTIUM ROENTGENIUM COPERNICIUM UNUNTRIUM UNUNQUADIUM UNUNPENTIUM UNUNHEXIUM UNUNSEPTIUM UNUNOCTIUM
59
140.908
3
[Xe] 4f 6s
2
60
144.242
4
[Xe] 4f 6s
2
61
(145)
5
[Xe] 4f 6s
2
[Xe] 4f 6s
231.036
2
1
2
92
238.029
3
1
2
93
(237)
4
1
2
[Rn] 5f 6d 7s [Rn] 5f 6d 7s [Rn] 5f 6d 7s
THORIUM
PROTACTINIUM
Pa
2
151.964
7
[Xe] 4f 6s
2
64 157.25(3) 65
7
1
2
U
URANIUM
SAMARIUM
94
(244)
6
2
[Rn] 5f 7s
EUROPIUM
95
(243)
7
2
158.925
9
[Xe] 4f 6s
GADOLINIUM
TERBIUM
96
(247)
7
1
2
[Rn] 5f 7s
[Rn] 5f 6d 7s
AMERICIUM
CURIUM
Tb
97
(247)
9
2
[Rn] 5f 7s
Np Pu Am Cm Bk
NEPTUNIUM PLUTONIUM
2
[Xe] 4f 5d 6s
Nd Pm Sm Eu Gd
Pr
91
62 150.36(2) 63
6
2
[Rn] 6d 7s
Th
109
PLATINUM
Mt
Hs
BOHRIUM
Pt
IRIDIUM
PRASEODYMIUM NEODYMIUM PROMETHIUM
232.038
2
107
Ir
OSMIUM
2
15.999
17
9
(1) Pure & Applied Chemistry, Vol. 78,
No. 11, pp. 2051–2066 (2006)
http://www.iupac.org/publications/pac/2006/pdf/7811x2051.pdf
[Xe] 4f 5d 6s
90
(266)
SEABORGIUM
140.116
1
106
Os
RHENIUM
Sg
Db
58
Re
TUNGSTEN
TANTALUM
1
INDIUM
200.59(2)
14
2
VIA
16
8
[Kr] 4d 5s 5p [Kr] 4d 5s 5p [Kr] 4d 5s 5p [Kr] 4d 5s 5p [Kr] 4d 5s 5p [Kr] 4d 5s 5p
Cd
10
28.086
SILICON
2
3
N
Si
69.723
2
NITROGEN
[Ne] 3s 3p
GALLIUM
112.411
10
14
14.007
[He] 2s 2p
Ga Ge
CADMIUM
196.967
1
ALUMINUM
10
2
C
AI
31
2
VA
15
7
[Ar] 3d 4s 4p [Ar] 3d 4s 4p [Ar] 3d 4s 4p [Ar] 3d 4s 4p [Ar] 3d 4s 4p [Ar] 3d 4s 4p
[Kr] 4d 5s
Ag
14
10
ZINC
SILVER
79
65.409
Zn
107.868
10
IIB
12
30
12.011
CARBON
26.982
2
IVA
14
6
[He] 2s 2p
[Ne] 3s 3p
[Ar] 3d 4s
[Kr] 4d 5s
Pd
78
1
COPPER
46 106.42(1) 47
PALLADIUM
2
63.546
Cu
[Kr] 4d
192.217
IB
11
29
[Ar] 3d 4s
10
1
Rh
14
8
NICKEL
RHODIUM
77
58.693
Ni
102.906
8
10
28
1
B
Fe
Tc - Man-made
[Ar] 3d 4s
[Kr] 4d 5s
190.23(3)
14
2
COBALT
IRON
44 101.07(2) 45
Ru
186.207
7
Co
[Kr] 4d 5s
Tc
58.933
[Ar] 3d 4s
Fe
7
[Kr] 4d 5s
9
27
2
BORON
- solid
- gas
10.811
[He] 2s 2p
Lanthanides
MOLYBDENUM TECHNETIUM RUTHENIUM
180.947
14
54.938
Mn
Nb Mo
NIOBIUM
73
VIIB
[Ar] 3d 4s
[Kr] 4d 5s
ACTINIDES
d
Alkaline earth metals
IIIA
13
5
[Xe] 4f 5d 6s [Xe] 4f 5d 6s [Xe] 4f 5d 6s [Xe] 4f 5d 6s [Xe] 4f 5d 6s [Xe] 4f 5d 6s [Xe] 4f 5d 6s [Xe] 4f 5d 6s [Xe] 4f 5d 6s
57
p
5
7
25
CHROMIUM MANGANESE
LANTHANIDES
s
51.996
Cr
92.906
4
VIB
6
24
[Ar] 3d 4s
[Kr] 4d 5s
ZIRCONIUM
72
ELEMENT NAME
He
Non-metal
Chalcogens
PLATINUM
VANADIUM
91.224
1
Metalloids
Alkaline metals
Pt
V
TITANIUM
2
9
[Ar] 3d 4s
Ti
40
195.084
14
2
1s
Metal
(1)(2)
[Xe] 4f 5d 6s
ATOMIC
SYMBOL
89
ORBITAL
FILLING
RELATIVE
ATOMIC MASS
(g.mol-1)
BERYLLIUM
Cs
87
ELECTRONEGATIVITY
9.0122
Be
CESIUM
7
IIA
Na Mg
SODIUM
18 VIIIA
2 4.0026
66
162.500
10
[Xe] 4f 6s
2
Dy
DYSPROSIUM
98
(251)
10
2
[Rn] 5f 7s
Cf
67
164.930
11
[Xe] 4f 6s
2
Ho
HOLMIUM
99
(252)
11
2
[Rn] 5f 7s
68
167.259
12
[Xe] 4f 6s
2
Er
13
2
THULIUM
(257)
12
168.934
[Xe] 4f 6s
70 173.04(3) 71
14
[Xe] 4f 6s
2
[Rn] 5f 7s
101
(258)
13
2
[Rn] 5f 7s
YTTERBIUM
102
(259)
14
FERMIUM
2
[Rn] 5f 7s
Es Fm Md No
BERKELIUM CALIFORNIUM EINSTEINIUM
2
Tm Yb
ERBIUM
100
69
MENDELEVIUM NOBELIUM
(3) The electronic configurations for which there is doubt are not given.
14
1
2
Lu
LUTETIUM
103
(262)
Lr
LAWRENCIUM
(2) The relative atomic mass is given with five significant digits. For items that do not have a stable radionuclide, the value in parentheses indicates the mass number of the isotope of the element with the longest half-life.
However, the three elements Th, Pa and Pu which have a characteristic terrestrial isotopic composition, an atomic weight is indicated.
Literature and Material for Selection CR 2016-2017
174.967
[Xe] 4f 5d 6s
Page 11