University “POLITEHNICA” of Timisoara
Faculty of Industrial Chemistry and
Environmental Engineering
GENERAL CHEMISTRY
Assoc.Prof. dr.eng. Andrea Kellenberger
GENERAL CHEMISTRY COURSE
Lecture: 2h / week
Laboratory: 2h / week
Evaluation form: Examination
Nr. of credits: 5
Chapter 1
INTRODUCTION
Space
Universe
Substance
Matter
Energy
Einstein’s equation:
E = mc2
E – energy, in J
m – mass, in kg
c – light speed, in m s-1
Substance is what all things consists of. One of the
most important properties of the substance is the mass.
Other characteristics of the substances are homogeneity
and constant composition.
Homogeneity = same characteristics in all the volume.
Constant composition = in any given portion of a
substance there are the same particles which interact in the
same way.
Examples of substances: water, sugar, oxygen,
sodium
chloride,
hydroxide.
copper,
hydrochloric
acid,
sodium
The great majority of things consist of a mixture of
substances. The air, for example, is not a substance,
because it can be separated through distillation in oxygen,
nitrogen, argon and other gases. The petrol is not a
substance as well – through distillation it can be separated
in hydrocarbons.
Chemistry
is
the
science
of
substances.
Chemistry studies the structure, properties and changes
of substances.
International System of Units ( Sİ units)
(from French: Le Systeme İnternational d’Unites).
Making observations is a key part of the scientific process.
Sometimes observations are qualitative, for example: the
substance is a green gas, and sometimes they are
quantitative (the mass is 10 grams, the temperature is
25ºC or the pressure is 1015 mbars). A quantitative
observation is called a measurement. A measurement
always consists of two parts: a number and a unit.
Table 1. Sİ Fundamental Units
Quantity
Unit
Symbol
meter
m
Mass
kilogram
kg
Time
second
s
Thermodynamic
temperature
Kelvin
K
Electric current
Ampere
A
Luminous intensity
candela
cd
mole
mol
Length
Amount of substance
Table 2. Sİ Derived Units
Quantity
Quantity equation
Name of the unit
Symbol
Area
S = l2
l – length
square meter
m2
Volume
V = l3
cubic meter
m3
Speed
v=l/t
t – time
meter per second
m s-1
Acceleration
a=v/t
meter per second squared
m s-2
Molar
concentration
cM = n / V
mole per cubic meter
n- amount of substance
V-volume
Force
F = m·a
m – mass
a – acceleration
kilogram· meter per second kg· m s-2( N )
squared (Newton)
Pressure
P=F/S
S – area
kilogram· meter per square
second per square meter
kg· m s-2· m-2
N· m-2= Pa
Density
Ρ=m/V
kilogram per cubic meter
kg· m-3
mol m-3
Examples:
- for volume (liter, L: 1L= 1dm3), for pressure (bar:
1 bar = 105 Pa);
- for temperature, there are two non SI scales: Celsius
(ºC) and Fahrenheit (ºF). Temperature conversions can be
made using the equations shown below:
Kelvin from Celsius:
T (K)= t (ºC) + 273.15
Celsius from Fahrenheit:
t (ºC )= 5/9 [t(ºF) – 32]
CHAPTER 2
ATOMIC STRUCTURE
OF THE SUBSTANCES
One of the oldest scientific concepts is that all matter
can be broken down until finally the smallest possible
particles are reached; these particles cannot be further
subdivided. The Greek philosopher Democritus (460 – 370
B.C.) considered these particles to be in a constant motion,
but able to fit together into stable combination. The
characteristics of substances resulted from the different size,
shape and arrangements of the particles, named atoms. In
Greek, atom means indivisible.
Modern atomic theory was developed by John Dalton
(1776 – 1844), based on quantitative data, not only
qualitative observations or speculations. Two natural
laws serve as the basics of Dalton’s atomic theory:
a) Law of conservation of mass
b) Law of definite composition (proportions):
a. Law of conservation of mass:
The total mass of materials present after a
chemical reaction is the same as before the
reaction.
Example: the reaction between Mg and O2.
0.24g Mg react with 0.16g oxygen. After the reaction,
0.40g MgO (magnezium oxide) were obtained.
b. Law of definite composition (proportions):
All samples of a compound have the same
composition, the same proportions by mass of the
constituent elements.
To see how the law of constant composition works,
let’s consider the compound FeS (iron sulfide). A sample of
10g FeS contains 6.36g Fe and 3.64g S. That means:
6.36
 100  63.6% Fe
10
and
3.64
 100  36.4%S
10
Another sample of 25g FeS contains 15.91g Fe and
9.09g S. The composition of the second sample is the same :
15.91
 100  63.6%
25
Fe
and
9.09
 100  36.4% S
25
J. Dalton was aware of these observations and he
offered an explanation for them. This is known as Dalton’s
atomic theory. The main ideas of this theory can be stated
as follows :
 Chemical elements are made of small particles called
atoms.
 All atoms of a given element are identical.
 The atoms of a given element are different from those
of any other element.
 Atoms of one element can combine with atoms of other
elements to form compounds. A given compound
always has the same relative number and type of atoms.
 Atoms are indivisible in chemical processes. That is,
atoms are not created or destroyed in chemical reactions.
A chemical reaction simply changes the way the atoms are
grouped together.
In order to apply Dalton’s theory in predicting new
phenomena, it was necessary to assign characteristic masses
to atoms. These masses became known as atomic weights.
Since they are very small, it is impossible to isolate and weigh
individual atoms. Dalton tried to establish relative atomic
weights. If an atom of hydrogen, for example, is taken to have
the mass of 1 unit, the mass of oxygen atom is 16.
2.1. Atom structure
At the end of the 19th century, the English physicist J. J.
Thomson showed that the atoms of any element emit tiny
negative particles. Thus, he concluded that all types of atoms
must contain these negative particles called electrons.
Although atoms contain negative particles, the whole
atom is not negatively or positively charged. So, Thomson
concluded that the atom must also contain positive particles
that balance the negative charge given by the electrons.
He imagined the “plum pudding model” of the atom, in
which electrons are scattered like plums into the uniform
“pudding” of positive charges.
Rutherford’s experiment
Rutherford measured the deviation of alpha particles (helium
ions with a positive charge) directed normally onto a sheet of very
thin gold foil. Assuming the plum pudding model, the alpha particles
should all have been deviated by, at most, a few degrees. The results
of the experiment were very different from those Rutherford
anticipated. Most of the α particles passed straight through the foil,
some of them were deflected at large angles and some were even
reflected backward.
Rutherford’s experiment
Conclusions:
• the plum pudding model for the atom could not be correct;
• the atom must contain a very small (compared with the size
of the atom) positive charge which causes the large
deflections of the α particles;
• the atom is mostly empty space because most of the α
particles past directly through the foil.
These results could be explained only in terms of
nuclear atom: an atom with a dense center of positive
charge (nucleus) around which tiny electrons moved in
a
space otherwise empty.
He concluded that the nucleus has a positive charge
to balance the negative charge of the electrons and that it
must be small and dense. This picture of the atom was the
planetary model of the atom.
Planetary model of the atom
In 1919 Rutherford discovered that the nucleus
contains positive particles named protons. Furthermore,
in 1932, James Chadwick discovered that nucleus
contains also neutral particles – neutrons. Protons and
neutrons are known as nucleons.
Electrons, protons and neutrons are fundamental
particles of the matter. There are more than 30 other
fundamental particles. Properties of the main fundamental
particles are given in the next table.
Properties of the main fundamental particles
Electric charge
Mass
Symbol
[C]
Relative
charge
kg
amu
Proton
+1.602·10-19
+1
1.67262·10-27
1.0073
p; p+
Neutron
0
0
1.67493·10-27
1.0087
n; no
Electron
-1.602·10-19
-1
9.10939·10-31
0.0005486
e; e-
The attraction force between the positive charges
(protons) and the negative ones (electrons) keeps the atom
together. In this image the nucleus is like a sun and electrons
like planets. Several problems arise with this concept – the
electrons might be expected to slow down gradually and fall on
the nucleus?
To explain why this did not occur, Niels Bohr (1913)
postulated:
 The electrons can move around the nucleus only on
certain orbits (allowed orbits).
 The electrons can gain or lose energy only by jumping
from one allowed orbit to another. When an electron moves
towards the nucleus energy is radiated and if it moves away
from the nucleus energy is absorbed.
 For an electron to remain in its orbit the electrostatic
attraction force between the electron and the nucleus must
equal to the centrifugal force which tends to throw the
electron out of its orbit.
N. Bohr admitted that the orbits of the electrons are
circular.
Bohr’s model of the atom
A. Sommerfeld (1916), based on the atomic spectra of
hydrogen, suggested that the permissive orbits of the
electrons may be elliptic.
Bohr – Sommerfeld model of the atom
Atomic number, mass number and chemical element
The number of protons in an atom is called the atomic
number Z. In an atom, which must be electrically neutral, the
number of electrons are also equal to Z. The total number of
protons and neutrons in an atom is the mass number A.
Thus, the number of neutrons is A-Z.
The three subatomic particles considered, the electron,
proton and neutron, are the only ones involved in chemical
phenomena. A study of matter at its most fundamental level
must consider a lot of additional subatomic particles.
All atoms with the same number of protons signify
a chemical element. Each element has a name and a
distinctive symbol.
Chemical
symbols
are
one
or
two
letter
abbreviations of the elements name (usually the Latin
name). The first letter, but never the second is capitalized.
For example :
Hydrogenium –
H
Aurum -
Au
Nitrogenium –
N
Cuprum -
Cu
Carbonum –
C
Silicium -
Si
Oxygenium –
O
Ferrum –
Fe
Phosphorus –
P
Tellurium -
Te
Sulphur –
S
Natrium -
Na
Fluorum -
F
Aluminium -
Al
Hydrargirum –
Hg
Strontium -
Sr
Stibium –
Sb
Protactinium -
Pa
Platinum -
Pt
Plutonium –
Pu
To represent the composition of any particular atom,
we need to specify its number of protons, neutrons and
electrons. We can do this with the symbol :
mass number
atomic number
A
E
Z
symbol of the element
Atoms that have the same atomic number Z, but
different mass numbers A are called isotopes.
The most simple element is hydrogen. The nucleus of
hydrogen consists only of one proton. Consequently, the
hydrogen atom has just one electron. This is isotope called
light hydrogen : 1 H .
1
Hydrogen has as well another two isotopes :
Deuterium (heavy hydrogen): 12 H or D.
Tritium (super-heavy hydrogen): 3 H or T.
1
Natural abundance of hydrogen isotopes is :
1
- 99.985%
1H
2
- 0.015%
H
1
3
- insignificant percent
1H
Abundance of the elements
What is the most abundant element? This simple
question does not have a simple answer. If we consider the
entire Universe, hydrogen accounts for about 90% of all the
atoms and 75% of the mass, and helium accounts for most
of the rest. If we consider only the elements present on
Earth, iron is probably the most abundant element. However,
most of the iron is in Earth’s core. The currently accessible
elements are those present in Earth’s atmosphere, oceans
and solid continental crust up to 16 km depth. The relative
abundance in these parts of the Earth are called Clark
parameters.
Nr.
crt.
Element
Clark
Parameter [%]
Nr.
crt.
Element
Clark
Parameter [%]
1
Oxygen
49,4
16
Samarium
510-4
2
Silicon
25,75
17
Gadolinium
510-4
3
Aluminum
7,51
18
Dysprosium
510-4
4
Iron
4,7
19
Ytterbium
510-4
5
Calcium
3,39
20
Erbium
410-4
6
Sodium
2,64
21
Argon
3,610-4
7
Potassium
2,40
22
Praseodymium
3,510-4
8
Magnesium
1,94
23
Lutetium
110-4
9
Hydrogen
0,88
24
Germanium
110-4
10
Titanium
0,58
25
Selenium
810-5
11
Chlorine
0,19
26
Cesium
710-5
12
Phosphor
0,12
27
Terbium
710-5
13
Carbon
0,087
28
Holmium
710-5
14
Manganese
0,085
29
Thulium
710-5
15
Sulfur
0,048
30
Niobium
410-5
Not all the known elements exist in Earth’s crust.
There are only 88 natural elements. The rest of known
elements can be produced only artificially by nuclear
processes. Moreover, most of the elements do not occur
free in nature, that is, as uncombined element. Only about
20% of them do. The remaining elements occur in chemical
combinations with other elements.
We can see in the last table that oxygen is the most
abundant element in the Earth’s crust (49.4%).
There are 3 natural isotopes of oxygen:
16
- 99.759%
O
8
17
- 0.037%
O
8
18
- 0.204%
O
8
Large amount of oxygen exists in water and rocks as
well in free state like molecular oxygen (O2) and ozone (O3).
Molecular oxygen (O2) and ozone (O3) are allotropes of
the element oxygen. The second element in Clark’s table is
silicon (Si – 25.75%), but silicon occurs only in chemical
combinations.
Atomic mass unit
Because the fundamental particles and the atoms are
very tiny it is difficult to operate with the small values of
their masses. This is the reason why the atomic mass
unit (amu) was introduced.
1 amu is the 12th part of the mass of the isotope
12
6
C
1 amu = 1.66 · 10-27 kg
Generally, atomic masses of the elements are fractional
number because the natural elements are a mixture of
two or more isotopes.
For example, magnesium has 3 stable isotopes:
24
12
Mg (78,70 %), exact atomic mass: 23,98504
25
12
Mg (10,13 %), exact atomic mass: 24,98384
26
12
Mg (11,17 %), exact atomic mass : 25,98259
Knowing the abundance of the stable isotopes one may
calculate the atomic mass of magnesium:
AMg=0.787 x 23.98504 + 0.1013 x 24.98384 + 0.1117 x 25.98259 =
= 24.30934
Electronic configuration of the atoms
Louis de Broglie (France) and Werner Schrödinger
(Austria) in the mid 1920s, suggested that like a light, the
electron has both a wave and particle properties.
When
Schrödinger
carried
out
a
mathematical
analysis based on this idea, he obtained a new model for the
atom: wave model.
In this model the electron has not a well defined orbit.
The motion of the electron seems to be rather a vibration.
The three-dimensional region of space around the nucleus in
which we can find the electron is called orbital. In fact, it is a
region of probability where the electron is likely to be found.
Let us consider a multi electronic atom. We can
assume that each electron has a specific mean path from
nucleus. The electrons having a similar mean path form a
main
energetic
level
or
main
electronic
shell,
characterized by the principal quantum number n. The
main shells are denoted by letters.
Energetic level
Principal quantum
number n
K
L
M
N
O
P
Q
n=1
n=2
n=3
n=4
n=5
n=6
n=7
The first main shell is the nearest level to the nucleus and it
has a minimum energy.
K shell (n = 1) consists of one s sublevel, containing one
orbital with a spherical symmetry named s orbital:
s orbital
L shell (n = 2) has 2 sublevels:
one s sublevel containing one spherical shaped s orbital
and one p sublevel containing 3 p orbitals.
Each p orbital consists in two lobes distributed along
one of the three rectangular axes through the nucleus:
In order to characterize the shape of the orbital the
orbital quantum number or azimuthal number l has been
introduced. For s orbital l = 0 and for p orbital l = 1.
All orbitals having the same l value form a subshell
or sublevel.
The orientation of the orbitals is given by the
magnetic quantum number m. It may be 0; ±1; ±2; …; ±l.
For example, if l = 1 provided that m = -1; 0; +1, that
is there are 3 p orbitals: px, py and pz.
For the third electronic level M (n = 3), the values of the
orbital and magnetic numbers are the following:
n=3
l=0
l=1
l=2
m=0
m = -1; 0; +1
m = -2; -1; 0; +1; +2
Beside s and p orbital, on the third electronic level there
are 5 orbitals characterized by orbital number l = 2,
named d orbital. There are 5 different d orbitals:
The d orbitals
Furthermore, for the 4th electronic level N, the quantum
numbers are:
n=4
l=0
l=1
l=2
l=3
m=0
m = -1; 0; +1
m = -2; -1; 0; +1; +2
m = -3; -2; -1; 0; +1; +2; +3
In this case, the orbitals with l = 3 are called f orbitals.
According to the values for m, there are 7 f orbitals:
The f orbitals
The energy of the levels increases as the values of
the n increase.
An orbital can be empty or it can contain one or
maximum two electrons. If two electrons occupy the same
orbital, they must have opposite spins, associated with
spin quantum number s, which may be ±1/2.
The order of filling orbitals
The electron configuration is the arrangement of electrons
on shells, subshells and orbitals. Electrons fill low energy orbitals,
closer to the nucleus, before they fill higher energy ones.
The order of energy levels is not identical to the
principal quantum number, due to the interaction between
electrons and nucleus. The attractive force of the nucleus
for a given electron increases as the nuclear charge
increases.
Therefore, the orbital energy should decrease with
increasing the atomic number.
The energy for the principal quantum levels are
given in the next figure.
The order in which electrons occupy orbitals is:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
In order to establish the electronic configuration of an
atom a lot of algorithms have been proposed. One of these
algorithms is known as a “minimum (n + l) rule”. According
to this rule, the electronic levels and sublevels will be filled in
the increasing order of the sum of the principal quantum
number and the orbital one. For the same sum n + l, the low
energy corresponds to the orbital with the lower n.
Minimum n + l rule can be illustrated by the following
table:
n
1
2
3
4
5
6
l
0
0
1
0
1
2
0
1
2
3
0
1
2
3
4
0
1
2
3
4
5
s
s
p
s
p
d
s
p
d
f
s
p
d
f
g
s
p
d
f
g
h
n+l 1
2
3
3
4
5
4
5
6
7
5
6
7
8
9
6
7
8
9
10
11
A suggestive image of the “n + l rule” has been given by
Goldansky’s chessboard:
The Goldansky’s chessboard
Pauli’s exclusion principle:
Two electrons in an atom cannot have all four
quantum numbers alike.
If two electrons exist in the same orbital, that is they
have identical principal number, orbital and magnetic ones,
these electrons must have opposite spins (different spin
numbers). Maximum number of electrons in an electronic
shell with principal number n is 2.
Hund’s rule
When orbital of identical energy are available,
electrons occupy these singly rather than in pairs. As a
result, an atom tends to have as many unpaired electrons
as possible.
In its ground state, hydrogen has its electron in the 1s
orbital. This is commonly represented in two ways: the
electron configuration or orbital diagram (box diagram). A
small arrow indicates the electron.
1s

1s1
H:
Electron configuration
orbital (box) diagram
The next element is helium (He). It has two protons in
the nucleus and so it has two electrons. Both electrons are
placed in the 1s orbital, but they have opposite spins.
1s
He:
1s2

Lithium (Z = 3) has three electrons, two of which go
into the 1s orbital. That is, the 1s orbital is full. The third
electron must occupy an orbital with n = 2, in this case 2s
orbital.
Li:
1s22s1
1s
2s


The next element, beryllium (Z = 4), has 4 electrons
which occupy the 1s and 2s orbitals with opposite spins :
Be:
1s22s2
1s

2s

Boron (Z=5) has five electrons, four of which occupy the
1s and 2s orbitals. The fifth electron goes into the second type
of orbitals with principal quantum number 2, one of the 2p
orbitals :
B:
1s22s22p1
1s
2s 2p

 
Because all the 2p orbitals have the same energy, it does
not matter which 2p orbital the electron occupies.
Carbon (Z=6), the next element, has six electrons: two in
1s orbital, two in 2s orbital and two in 2p orbital. The last
electrons occupy different 2p orbitals:
C:
1s22s22p2
1s
2s
2p


 
The configuration of nitrogen (Z = 7) is :
N:
1s22s22p3
1s
2s
2p



Oxygen (Z = 8) has four electrons in the 2p orbitals, one
of the 2p orbitals is occupied by a pair of electrons with
opposite spins :
1s
O:
1s22s22p4

2s

2p
  
The electron configuration for fluorine (Z = 9) and neon
(Z = 10) are :
F:
1s22s22p5
Ne: 1s22s22p6
1s
2s
2p


  


  
With neon, the K (n = 1) and L (n = 2) shells are filled.
For sodium (Z = 11), the eleventh electron fill the first orbital of
the M (n = 3) shell, that is 3s orbital. Thus, the electron
configuration for sodium will be 1s22s22p63s1. To avoid writing
the inner-level electrons the configuration of neon 1s22s22p6 is
abbreviated with [Ne].
Na
[Ne]3s1
1s
2s


2p
  
3s

We have to note that the configuration of an element
differs from the previous only by an electron named
“differentiating electron”.
Periodic table of the elements
The chemical and physical properties are determined
by the number and arrangement of the electrons, that is by
the atomic number. If the elements are arranged in groups,
each group having a characteristic electronic structure, then
elements should show similarities in chemical and physical
properties.
A classification scheme of the elements, similar to that
used today, was discovered independently by Dmitri
Mendeleev and Luther Meyer in 1869. Their classifications
were based on an early version of the periodic law :
If the elements are arranged in order of
increasing atomic mass, certain sets of properties are
found to reappear periodically.
The arrangement of the elements based on the
periodic law is called periodic table. In Mendeleev’s
periodic table the elements were arranged in 12 horizontal
rows and 8 vertical columns or groups. The eight groups
were further divided in 16 sub-groups. Mendeleev’s
periodic table was a short form.
The modern periodic table is a long form.
The horizontal rows of the table are called periods.
The first period of the table consists of only two
elements: hydrogen (H) and helium (He).
The second and third periods have eight elements
each: from lithium (Li) to neon (Ne) and from sodium
(Na) to argon (Ar).
The fourth and fifth periods comprise 18 elements
each, ranging from potassium (K) to krypton (Kr) and
from rubidium (Rb) to xenon (Xe).
The sixth period is a long one with 32 members.
From this period 14 elements are extracted and
placed at the bottom of the table. This series of 14
elements, which fits between lanthanum (La, Z=57) and
hafnium (Hf, Z=72) is called the lanthanides or rare earth
series.
The seventh and final period is incomplete for the
moment, but is believed to be as long as the sixth one. A 14
member series, extracted from the seventh period and
placed at the bottom of the table is called the actinide
series.
All the atoms of the elements from group 1 possess a
single outer – shell electron in an s orbital. Elements of the
first group are called alkali-metals.
The atoms of the elements from group 2 have two
electrons in an outer shell. These elements are alkaline
earth metals. These two groups are known as the s block,
because their properties arise from the presence of s
electrons.
Elements of group 13 have three electrons in their
outer shell, two s electrons and one p electron. The p
electron is the differentiating electron.
Elements of group 14 have 4 electrons in the outer shell (s2p2)
Elements of group 15 have 5 electrons (s2p3),
Elements of group 16 have 6 electrons(s2p4),
Elements of group 17 have 7 electrons (s2p5)
Elements of group 18 have 8 electrons(s2p6).
Elements of 18 group have an outer shell full of electrons.
Because their properties are dependent on the presence of p
electrons, groups 13 – 18 are called p – block elements.
Elements of groups from 3 to 12 are called the d block
or transition elements and, in a similar way, lanthanoid and
actinoid elements are f block.
Periodic properties of the elements
The elements of the 18th group, rare gases, have the
configuration ns2 np6, except helium, whose configuration is
1s2. That means the outer shells of the atoms are full. These
prove to be very stable configurations and they can be
altered with great difficulty. As a result, rare gases have a
very low reactivity, they are also known as noble gases.
The electron configuration of the elements of groups 1
and 2 differ from these of noble gases by only one or two
electrons in the s orbital of a new shell.
For example, the configuration of potassium and calcium are:
K
Ca
[Ar] 4s1
[Ar] 4s2
Except hydrogen, the elements of groups 1 and 2 are
metals. The characteristic chemical properties of metallic
elements are based on the ease of removal of one or more
electrons from their atoms to produce positive ions:
K → K+ + e -
Ca → Ca2+ + 2e-
Some physical properties of metals – ability to conduct heat
and electricity, ductility, malleability – also arise from these
distinctive electron configurations.
Elements of the groups 16 and 17 have an electron
configuration with two or one electron less that those of
the
corresponding noble gas. Atoms of these elements can
realize the electronic configuration of a noble gas by gaining
the appropriate number of electrons. For example, the
electron configuration of S becomes that of Ar by gaining two
electrons:
S
+ 2e-  S2[Ne] 3s23p4
[Ar]
The S atom becomes S2- anion.
Similarly, the Cl atom becomes Cl- anion:
Cl
+ e[Ne] 3s23p6
 Cl[Ar]
These elements whose atoms can acquire a noble
gas configuration by a small number of electrons are non –
metals. Non – metals are H from group 1, C from group 14,
N and P from group 15, O, S and Se from group 16 and F,
Cl, Br and I from group 17.
B (13), Si, Ge (14), As, Sb (15), Tc, Po (16) and At
(17) are metalloids or semi – metals.
18th group is a special family of elements, but noble
gases may be considered non – metals. The rest of the
elements, including of course the lanthanides and actinides
are metals.
Atomic size
The sizes of atoms vary as shown in the following
figure.
Atoms get larger as we
go down in a group of
the periodic table and
they get smaller as we
go from left to right
along a period.
Atomic radii [in pm]
The atomic radius tends to increase on descending a
group due to the increment of the number of energetic
levels.
The outer shell electrons are further and further
from the nucleus, therefore less attracted by the positive
charge of the nucleus.
Along a period the charge of the nucleus increases
with the atomic number Z while the electrons are still filling
the same shell. The outer shell electrons are attracted more
strongly by the core and, as a result, the atomic radius
decreases from left to right through a period.
Ionic radius
When electrons are removed from a metal atom to form a
positive ion (cation), a significant reduction in size occurs.
Usually, the electrons are lost from the shell with the highest
principal quantum number.
The relative sizes of the cations are given in the figure:
The relative sizes
of the cations
The ionic radius for cations in the same period
decreases from left to right.
For example, in the series of cations Na+, Mg2+, Al3+
the number of electrons is the same (10), while the number
of protons increases together with the atomic number Z.
Al3+ is smaller that Mg2+ because the electrostatic
force between the 10 electrons and the nuclear charge of
Al (+13) is more powerful than that between the 10
electrons and the nuclear charge of Mg (+12).
No. of protons
No. of electrons
Ionic radius [Å]
Na+
11
10
0,95
Mg2+
12
10
0,65
Al3+
13
10
0,50
When a nonmetal atom gains one or more electrons to
form a negative ion (anion) the size increases
compared with the original atom.
The relative sizes of the anions
Along a period the ionic radius of anions decreases from
left to right.
Ionization energy
Is the energy required to remove one electron from
an individual atom in the gaseous phase. This is the first
Ionization
ionization energy
M (g)
M+(g) + eenergy
In case of metals, which have a small number of
electrons in the outer shell, a small amount of energy is
needed to remove an electron, that is metals have low
ionization energies.
Inside a group, the ionization energy tends to
decrease from top to bottom because the attraction force of
the nucleus decreases in the same way and the electron is
more easily removed.
Nonmetals have large ionization energies because
they have a large number of electrons in the outer shell.
Nonmetals tend to gain, not to lose electrons.
Ionization energies tend to increase from left to right along a
period of the periodic table.
In general, the elements that appear in the lower left
region of the periodic table have the lowest ionization
energies and are therefore the most chemically active
metals. On the other hand, the elements with the highest
ionization energies occur in the upper right hand region of
the periodic table.
The first ionization energy of the elements is a
function of atomic number Z:
The first ionization energy of the elements
The loss of the second electron occurs with greater
difficulty than the first. Therefore the second ionization
energy is higher than the first one.
A property used to describe the bond type that results
when atoms combine is electronegativity.
Electronegativity describes the ability of an atom to
attract electrons towards itself. The most widely used
electronegativity scale was devised by Linus Pauling.
Pauling’s
electronegativities
are
dimensionless
numbers ranging from about 1 for very active metals to 4.0
for fluorine, the most active nonmetal.
1
2
3
4
5
6
7
8
9
10
11
12
13
14
15
16
17
H
2,20
18
He
Li
0,98
Be
1,57
B
2,04
C
2,5
5
N
3,04
O
3,44
F
3,9
8
Ne
Na
0,93
Mg
1,31
Al
1,61
Si
1,9
0
P
2,19
S
2,58
Cl
3,1
6
Ar
K
0,82
Ca
1,00
Sc
1,36
Ti
1,54
V
1,63
Cr
1,66
Mn
1,55
Fe
1,83
Co
1,88
Ni
1,9
1
Cu
1,90
Zn
1,65
Ga
1,81
Ge
2,0
1
As
2,18
Se
2,55
Br
2,9
6
Kr
3,0
Rb
0,82
Sr
0,95
Y
1,22
Zr
1,33
Nb
1,6
Mo
2,16
Tc
1,9
Ru
2,2
Rh
2,28
Pd
2,2
Ag
1,93
Cd
1,69
In
1,78
Sn
1,9
6
Sb
2,05
Te
2,1
I
2,6
6
Xe
2,6
Cs
0,79
Ba
0,89
La
1,27
Hf
1,3
Ta
1,5
W
2,36
Re
1,9
Os
2,2
Ir
2,20
Pt
2,2
8
Au
2,54
Hg
2,0
Tl
1,62
Pb
2,3
3
Bi
2.02
Po
2,0
At
2,2
Rn
Fr
0,7
Ra
0,9
Ac
1,10
Pauling’s electronegativities of the elements
As a rough rule, most metals have electronegativities of about
1.7 or less; semi-metals about 2 and nonmetals greater than 2.