Chemistry 1(Vet) – CHEM1405
Welcome: Bachelor of Veterinary Science
This Week: Atomic Structure and the Periodic Table




Overview of CHEM1405, administrative matters
Resources and study in chemistry
Atomic structure
Periodic Table
1
Assistance in Chemistry
Dr Adrian V George
 Room 224
 [email protected]
Assistance - administrative
 First Year Enquiry Office (10 am - 3.15 pm)
 E-mail: [email protected]
Assistance - Course Work
 Duty Tutor Room (Monday– Friday, 1-2, back of Lab D)
 Chem1405 discussion board
2
Assistance in Chemistry
Formal:
Informal:
Lectures
Study groups
Laboratories
Duty tutor
On-line:
eLearning
ChemCAL
Self-help
problems
Off-line:
Problem sets
Textbooks
Lecture notes
Plan your study time from the start – now!
3
1
Assistance in Chemistry
Information and Resources





First Year Chemistry web site (http://firstyear.chem.usyd.edu.au/)
eLearning (learn-on-line.usyd.edu.au ) (NB change of address from notes)
Resource Page (http://firstyear.chem.usyd.edu.au/chem1405/
(htt //fi t
h
d d
/ h 1405/ )
ChemCAL (http://chemcal.chem.usyd.edu.au/)
Text book: Blackman et al. Chemistry
4
Course information
Laboratory Work
 Start in week 2, attendance a requirement of the course
Assessment
15 % laboratory assessment


15 % tutorial quizzes (Week 6, 9, 12)

10 % organic spectroscopy assignment

60 % 3 hour exam at the end of semester
Staff-Student Liaison Committee
5
Recap: Atomic Structure

1803

1897
1909

1909

1920s

J Dalton
provided evidence for fundamental
indivisible particles - atoms
J J Thomson studied cathode rays - electrons
R A Millikan measured the charge of an electron
(1.6 x 10-19 C)
E Rutherford proposed an atom be composed of a
small positive nucleus surrounded by
a lot of space occupied by the
electrons
N Bohr
electrons occupy orbits of defined
energy
6
2
Atomic nomenclature
Atomic Number - The number of
protons.
Mass Number - The number of
protons + neutrons.
Isotopes – different no of neutrons,
same no of protons eg 12C, 13C.
7
Relative Atomic Mass
Based on a standard that the mass of




12C
is exactly 12.
A mole is the same number as the number of 12C
atoms in exactly 12 grams of 12C atoms.
The number "12 grams" is chosen to coincide with
the defined atomic mass of one 12C atom, 12 amu.
1 mole = 6.022 x 1023.
The relative mass of a single atom can be measured
by a mass spectrometer.
8
Mass Spectrometer
accelerating plates
70 eV
–
+
+
+
+
+
+
+
+
+ +
+
+
+
+
ion beam
+
9
3
Mass Spectrometer
magnetic field
ion beam
+ + + ++ +
+ + + +
+
+
+
10
Mass Spectrometer
The mass spectrometer measures relative mass of a single
atom related to the mass of a single atom of 12C.


20 elements occur in nature as single isotopes:
Be F
Be,
F, Na
Na, Al
Al, P
P, Sc
Sc, Mn
Mn, Co
Co, As
As, Y
Y, Nb
Nb, Rh
Rh, II, Cs
Cs, Pr
Pr, Tb
Tb,
Ho, Tm, Au, Bi.
Their atomic masses are shown on the periodic table.
Thus, the atomic mass of a sodium atom (23Na) is
22.9898 amu.
The remaining elements are each mixtures of several
isotopes. The atomic mass being a weighted average of
the naturally occurring atomic masses.
11
E.g. Chromium
The atomic weight of chromium is obtained by
multiplying each isotopic mass by its fractional
abundance and then summing:
49.9461 x
51.9405 x
52.9407 x
53.9389 x
Total
4.35%
83.79%
9.50%
2.36%
=
=
=
=
2.17
43.52
5.03
1.27
= 51.99
The atomic mass of chromium is 51.99.
12
4
Molar mass
The relative molar mass of a substance is the sum of its
constituent atoms or ions
e.g. sodium chloride, NaCl
relative molar mass = 22.99 + 35.45 = 58.34
e.g. glucose, C6H12O6
relative molar mass = (6 x 12.01) + (12 x 1.008)
+ (6 x 16.00) = 180.16
e.g. sodium lactate, C3H5O3Na
relative molar mass = (3 x 12.01) + (5 x 1.008)
+ (3 x 16.00) + 22.99 = 112.06
13
Atomic Theory
Light of different colour
has a different
wavelength
Wavelength:
short  long
Colour:
blue  red
Energy:
high  low
400
 -Ray
500
X-Ray
UV
600
700 nm
Infrared Microwave and
Radio frequency
14
Atomic Spectrum of Hydrogen

Light emitted from a hydrogen arc lamp is composed of only
a few lines:
 Only light of certain energy is emitted
 The pattern of lines is unique to hydrogen
Atomic spectrum
Continuous spectrum

Suggests the process emitting light in the atom is
quantised

The electron in the atom may possess only certain
energies
15
5
Bohr Atom
16
Other Elements
Atomic emission spectrum is a characteristic of the element
 Bohr model of the atom works well for H but not for
other elements (see http://onsager.bd.psu.edu/~jircitano/periodic4.html)
 Quantum mechanics gives a better description
17
Quantum Mechanical Model


Light has a dual nature and the de Broglie
equation relates wavelength to momentum
 = h/mv
Schrödinger Equation
 = E

E

This can only be solved if various boundary conditions are
applied. That is, the waves must be standing waves that are
 continuous
 single valued
 multiples of a whole number of half wavelengths
18
6
Quantum Mechanical Model



There are then discrete solutions that represent the energy
of each electron orbital.
A point in 3-D space may be described by three coordinates;
an electron orbital is described by four coordinates.
coordinates
The coordinates of the orbital are given by quantum
numbers.
19
Principal Quantum Number, n

n = 1, 2, 3 …
 Describes the size and extent of the orbital.
 The larger the value of n, the bigger & the higher
energy the orbital.
n=1
n=2
n=3
20
Angular Momentum Quantum No, l
 l = 0, 1, 2…(n -1)
 Describes the shape of the orbital
 e.g. if n = 2; l = 0 or 1
l=0
l=1
l=2
l=3
l=0
"s"
"p"
"d"
"f"
l=1
21
7
Magnetic Quantum Number, ml



ml = -l, -(l -1) … 0 … (l -1), l
Describes the orientation of the orbital
e.g. if l = 0; ml = 0
(s orbital)
if l = 1; ml = -1, 0, +1
(px, py, pz orbitals)
if l = 2; ml = -2, -1, 0, +1, +2
(dxy, dyz, dxz, dx -y , dz )
2
2
2
22
Spin Quantum Number, ms



ms = + 1/2 , - 1/2
Describes the spin of the electron.
Each orbital, uniquely described by n, l and ml may contain a
maximum of two electrons, one spin +1/2, the other spin -1/2 .
Why is this important?
•Relates to a thorough understanding of the periodic table
•Size of atom/ion related to metal toxicity
•Relates to type of bonds (σ or ) formed in compounds
•Shape – essential for design of selective drugs
23
Question:
Complete the table
n=
1
2
l=
0
0, 1
Description
s
s, p
Maximum no. of
2
2, 6
electrons in sub-shell
Total electrons
2
8
Shell,,
Sub-shell,
3
4
24
8
Polyelectronic Atoms




When determining the ground state
electron configuration, there are
three rules:
Pauli exclusion principle - no two
electrons can have an identical set of
four quantum numbers. i.e. there are a
maximum of 2 electrons in any one
orbital.
Aufban principle - fill up low energy
orbitals before high energy ones.
Hund’s rule - orbitals with the same
energy (i.e. the same sub-shell) have
the maximum number of unpaired
electrons.
25
Polyelectronic Atoms

The orbital energy of a polyelectronic atom
increases:
1s < 2s < 2p < 3s < 3p < 4s < 3d <
4p < 5s < 4d < 5p < 6s <4f <5d …
26
Question: complete the table
Element
H
He
Li
B
C
O
Ne
Al
Ca
Sc
Cr
Fe
Cu
No of electrons Electron configuration
1
1s1
2
1s2
3
1s2 2s1
5
1s2 2s22p1
6
8
10
13
20
[Ar] 4s2
21
24
[Ar] 4s1 3d5
26
29
[Ar] 4s1 3d10
27
9
Periodic Table
The vertical columns are called Groups.
The horizontal rows are called Periods.
The Group number = The number of Valence electrons
(electrons in the outer shell).
The Period number = The number of occupied energy shells. 28
Periodic Table & Electrons
The orbital energy of a polyelectronic atom increases
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s <4f <5d …
29
Periodic Table & Quantum No
Angular Momentum Quantum No
Principal
Quantum No
l = 0 s-block; l = 1 p-block; l = 2 d-block
n=1
n=2
n=3
n=4
n=5
n=6
n=7
Magnetic Quantum No:
1 s-orbital; 3 p-orbitals; 5 d-orbitals
Spin Quantum No:
2 electrons in each orbital
30
10
Summary
You should now be able to
 Recognise how relative atomic masses are derived.
 Calculate relative molar mass for any substance.
 Understand the difference between a Bohr model and
quantum mechanical model of an atom.
 Understand the relationship between the four quantum
numbers and electron configuration.
 Determine the electron configuration of an element
from its position in the Periodic Table.
 Recognise whether an element is a metal, non-metal or
semi-metal from its position in the Periodic Table.
31
Biological Periodic Table
This Lecture: Periodic Table
 Overview of the ‘Biological’ Periodic Table
 Ions
 Inorganic nomenclature
32
Periodic Table & Properties
Metals
Semi-metals
or metalloids
Non-metals
33
11
Metals and non-metals
Metals show:
Non-metals are typically:
•malleability and ductility.
•good conduction of electricity
and heat.
•luminous surface appearance.
•form
f
positively
i i l charged
h
d ions
i
called cations.
•brittle solids or gases (there is one
liquid, bromine).
•have poor thermal and electrical
conductivity.
•form
f
negatively
ti l charged
h
d ions
i
called
ll d
anions.
34
Essential and Toxic Metals
All elements, just like all
substances, are toxic at
sufficiently high doses.

Some metallic elements
are essential trace
elements needed by the
body to maintain good
health, others are of no
benefit to a healthy body
and are toxic at even very
low concentrations.

The difference between these types
of elements is demonstrated in the
following curves showing the
relationship between concentration
and the health of an organism.
Well
Healthy
Dead
Concentration
Concentration
Essential Trace Element
Toxic Metal
35
Essential and Toxic Metals
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
1
18
2
HYDROGEN
HELIUM
H
He
1.008
4.003
3
4
5
6
7
8
9
10
LITHIUM
BERYLLIUM
BORON
CARBON
NITROGEN
OXYGEN
FLUORINE
NEON
Li
6.941
Be
9.012
B
C
10.81
12.01
11
12
13
SODIUM
MAGNESIUM
ALUMINIUM
N
O
14.01
16.00
14
15
SILICON
PHOSPHORUS
F
Ne
19.00
20.18
16
17
SULFUR
CHLORINE
18
ARGON
Na
Mg
Al
Si
P
S
Cl
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
20
21
22
23
24
25
26
27
28
29
30
31
32
33
34
35
36
POTASSIUM
CALCIUM
SCANDIUM
TITANIUM
VANADIUM
CHROMIUM
MANGANESE
IRON
COBALT
NICKEL
COPPER
ZINC
GALLIUM
GERMANIUM
ARSENIC
SELENIUM
BROMINE
KRYPTON
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
Kr
40.08
44.96
47.88
50.94
58.69
63.55
65.39
78.96
79.90
83.80
39.10
69.72
72.59
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
RUBIDIUM
STRONTIUM
YTTRIUM
ZIRCONIUM
NIOBIUM
MOLYBDENUM
52.00
TECHNETIUM
54.94
RUTHENIUM
55.85
RHODIUM
58.93
PALLADIUM
SILVER
CADMIUM
INDIUM
TIN
ANTIMONY
74.92
TELLURIUM
IODINE
XENON
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
Xe
85.47
87.62
88.91
91.22
92.91
95.94
[98.91]
101.07
102.91
106.4
107.87
112.40
114.82
118.69
121.75
127.60
126.90
131.30
55
56
57-71
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
CAESIUM
BARIUM
HAFNIUM
TANTALUM
TUNGSTEN
RHENIUM
OSMIUM
IRIDIUM
PLATINUM
GOLD
MERCURY
THALLIUM
LEAD
BISMUTH
POLONIUM
ASTATINE
RADON
Cs
Ba
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
Rn
132.91
137.34
178.49
180.95
183.85
186.2
190.2
192.22
195.09
196.97
200.59
204.37
207.2
208.98
[210.0]
[210.0]
[222.0]
87
88
FRANCIUM
RADIUM
89-103
104
105
106
107
108
109
RUTHERFORDIUM
DUBNIUM
SEABORGIUM
BOHRIUM
HASSIUM
MEITNERIUM
Fr
Ra
Rf
Db
Sg
Bh
Hs
Mt
[223.0]
[226.0]
[261]
[262]
[266]
[262]
[265]
[266]
Essential
Toxic
Medicinal
Medicinal
Archaea
Nearly all of the trace essential elements are from the first row of
the d block.
Most of the highly toxic elements are from the late d and early p
blocks of the fifth and sixth periods.
36
12
Reasons?
•Nature has made use of the more abundant
elements.
•Toxic elements from inert bonds (in general) and
the toxic elements don’t bind well to the O and N
donor ligands
g
preferred
p
byy nature.
An important aspect of abundance is availability. For
example, Al is one of the most abundant elements but is
toxic. It is not “bioavailable” because of the insolubility
of the forms it is found in nature.
Increased use of Al is making it much more bioavailable
but our bodies have evolved only limited ability to deal
with it.
37
Biological Roles for Metals




Trace Element: A chemical element required by an
organism in only a trace amount. Typically require
less than 3 mg/day intake.
Bulk elements - Essential elements with typical
intakes of greater than 100 mg/day.
Enzyme: A protein specialised to catalyse a specific
metabolic reaction.
Cofactor: A small-molecular-weight substance
required for the action of an enzyme.
38
Bulk Elements
Bulk Elements
Biochemical Function
Calcium
Bone and Teeth ca 1kg
Chlorine
Electrolyte
Magnesium
Bone, some enzymes ca 25g
Phosphorous
Bone, Nucleic acids
Sodium
Extracellular cation - water
electrolyte balance
Intracellular cation
Potassium
39
13
Trace Elements
Trace Elements
Biochemical Function
Fluorine
Bone and Teeth
Iodine
Chromium
Vanadium
Manganese
Thyroid hormone production
Utilisation of blood glucose
Co-factor of nitrate reductase
Co-factor for enzymes
Iron
Iron proteins, such as heme, ferritins
Cobalt
Nickel
Vitamin B12
Cofactor of urease
Copper
Zinc
cytochrome oxidase
Enzymes eg alcohol dehydrogenase
Molybdenum
Selenium
aldehyde oxidase
Glutathione peroxidase
Arsenic
Silicon
Tin
Not known
Connective tissue and Bone
Formation of bone
40
Toxicity & Health Effects Of Chromium
Claims:
•May reduce weight and increase muscle.
•May help diabetes and lower cholesterol.
•Regulates blood sugar.




Cr(VI) is classified as "Carcinogenic to Humans".
Cr(VI) compounds are soluble in water thus may have a
harmful effect on the environment.
Cr(VI) is readily reduced by Fe2+ and dissolved sulfides.
Cr(III) is considered to be an essential nutrient.
41
Arsenic






As(III) combines with -SH groups and interferes with the function of
a number of enzymes.
As(V) is considerably less toxic than As(III). Forms H2AsO4- which is
chemically similar to phosphate and can interfere with phosphate
metabolism.
The lethal dose to humans is estimated at 1 to 4 mg of arsenic per
k off b
kg
body
d weight.
i ht
Long-term exposure to low concentrations of arsenic has been
reported to cause skin cancer and it may carry risk of various
internal organ cancers.
Other effects of high exposure levels include nausea, vomiting and
diarrhoea; decreased production of red and white blood cells;
abnormal heart rhythms; blood vessel damage; and a "pins" and
"needles" sensation in hands and feet.
Arsenic is carcinogenic. Breathing it increases the risk of lung
cancer. Ingesting it increases the risk of skin cancer and tumours of 42
the bladder, kidney, liver and lung.
14
Ions and Oxidation States
H
Li
Na
K
Rb
Cs
Be
Mg
Ca
Sr
Ba
F
Fr
R
Ra
Sc
Y
5771
89
89105
Ti
Zr
Hf
V Cr Mn
Nb Mo Tc
Ta W Re
Group 2: M2+ cations
Fe Co Ni
Ru Rh Pd
Os Ir Pt
Cu
Ag
Au
Zn
Cd
Hg
Eg Ca [Ar] 4s2
d-block: variable
oxidation state
Ca2+ [Ar]
Eg Fe [Ar] 4s2 3d6
Group 1: M+ cations
Fe2+ [Ar] 3d6
Eg Na 1s2 2s2 2p6 3s1
Fe3+
[Ar]
3d5
Na+ 1s2 2s2 2p6
B
Al
Ga
In
Tl
C
Si
Ge
Sn
Pb
N
P
As
Sb
Bi
O
S
Se
Te
Po
F
Cl
Br
I
At
He
Ne
Ar
Kr
Xe
Rn
Down a group,
more metallic,
O.No two units
apart. Eg As(III)
or As(V)
Group 17: X- anions
Eg Cl 1s2 2s2 2p6 3s2 3p5
Cl- 1s2 2s2 2p6 3s2 3p6
43
Size
 Atomic size related to
electron configuration
Cations are always
smaller than the atoms
from which they formed
Anions are always larger
than the atoms from
which they formed
44
Ionisation Energy
 Ionisation energy is
always positive
M(g)
M+(g) + e-
It depends
d
d on the
h
strength with which the
outermost electron is held
by the nucleus
45
15
Naming Ionic Compounds
Ions can be either monatomic or polyatomic.
Ionic compounds contain cations and anions in a ratio that
maintains electrical neutrality.
cation
anion
Formula
Name
Ba2+
NO3-
Ba(NO3)2
barium nitrate
K+
PO43-
Ag
+
Cd2+
O
2-
potassium phosphate
silver oxide
S2-
cadmium sulfide
potassium permanganate
K+
MnO4-
Na+
HCO3-
sodium hydrogencarbonate
Fe2+
SO42-
iron(II) sulfate
Fe3+
SO42-
iron(III) sulfate
Complete
the
formula
46
Questions
1. What is the formula and name of the compound formed
between barium and phosphate (PO43-)?
2. Order
d the
h following
f ll
atoms in terms off increasing radius:
d
Ar, Li, Na, P, Sb
3. Order the following species in terms of increasing
radius: S2-, Cl-, Ar, K+, Ca2+
47
16