University “POLITEHNICA” of Timisoara

Faculty of Industrial Chemistry and

Environmental Engineering

GENERAL CHEMISTRY

Assoc.Prof. dr.eng. Andrea Kellenberger

Lecture: 2h / week

Laboratory: 2h / week

Evaluation form: Examination

Nr. of credits: 5

GENERAL CHEMISTRY COURSE

Chapter 1

INTRODUCTION

Universe

Space

Matter

Substance

Energy

E – energy, in J

m – mass, in kg

c – light speed, in m s-1

E = mc2

Einstein’s equation:

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, copper, hydrochloric acid, sodium

hydroxide.

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.

Quantity Unit SymbolLength meter m

Mass kilogram kg

Time second s

Thermodynamic temperature

Kelvin K

Electric current Ampere A

Luminous intensity candela cd

Amount of substance mole mol

Table 1. Sİ Fundamental Units

Quantity Quantity equation Name of the unit Symbol

Area S = l2

l – lengthsquare meter m2

Volume V = l3 cubic meter m3

Speed v = l / tt – time

meter per second m s-1

Acceleration a = v / t meter per second squared m s-2

Molar concentration

cM = n / V

n- amount of substanceV-volume

mole per cubic meter mol m-3

Force F = m·am – massa – acceleration

kilogram· meter per second squared (Newton)

kg· m s-2( N )

Pressure P = F / SS – 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

Table 2. Sİ Derived Units

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:

6.36100 63.6%

10 Fe

and

3.64 100 36.4%10

S

Another sample of 25g FeS contains 15.91g Fe and

9.09g S. The composition of the second sample is the same :

15.91100 63.6%

25 Fe

and

9.09 100 36.4%25

S

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.

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-

Properties of the main fundamental particles

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 :

AZ Emass number

atomic number

symbol of the element

Atoms that have the same atomic number Z, but different

mass numbers A are called isotopes.

Hydrogen has as well another two isotopes :

Deuterium (heavy hydrogen): or D.21 H

Tritium (super-heavy hydrogen): or T.31 H

Natural abundance of hydrogen isotopes is :

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 : .11H

11H - 99.985%

21 H - 0.015%

31 H - insignificant percent

Abundance of the elementsWhat 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.

ElementClark Parameter [%]

Nr.crt.

ElementClark Parameter [%]

1 Oxygen 49,4 16 Samarium 510-4

2 Silicon 25,75 17 Gadolinium 510-4

3 Aluminum 7,51 18 Dysprosium 510-4

4 Iron 4,7 19 Ytterbium 510-4

5 Calcium 3,39 20 Erbium 410-4

6 Sodium 2,64 21 Argon 3,610-4

7 Potassium 2,40 22 Praseodymium 3,510-4

8 Magnesium 1,94 23 Lutetium 110-4

9 Hydrogen 0,88 24 Germanium 110-4

10 Titanium 0,58 25 Selenium 810-5

11 Chlorine 0,19 26 Cesium 710-5

12 Phosphor 0,12 27 Terbium 710-5

13 Carbon 0,087 28 Holmium 710-5

14 Manganese 0,085 29 Thulium 710-5

15 Sulfur 0,048 30 Niobium 410-5

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:

168 O - 99.759%

178 O - 0.037%

188 O - 0.204%

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 C126

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:

Mg2412

Mg2512

Mg2612

(78,70 %), exact atomic mass: 23,98504

(10,13 %), exact atomic mass: 24,98384

(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.

The first main shell is the nearest level to the nucleus and it

has a minimum energy.

Energetic level K L M N O P QPrincipal quantum number n

n = 1 n = 2 n = 3 n = 4 n = 5 n = 6 n = 7

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 m = 0

l = 1 m = -1; 0; +1

l = 2 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 m = 0

l = 1 m = -1; 0; +1

l = 2 m = -2; -1; 0; +1; +2

l = 3 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

H: 1s1

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.

1s

2s

Li: 1s22s1

The next element, beryllium (Z = 4), has 4 electrons

which occupy the 1s and 2s orbitals with opposite spins :

1s 2sBe: 1s22s2

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 : 1s 2s

2p

B: 1s22s22p1

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:

1s 2s

2p

C: 1s22s22p2

The configuration of nitrogen (Z = 7) is :

1s 2s 2p

N: 1s22s22p3

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 2s 2p

O: 1s22s22p4

The electron configuration for fluorine (Z = 9) and neon

(Z = 10) are : 1s 2s 2p

F: 1s22s22p5

Ne: 1s22s22p6

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].

1s 2s 2p 3s

Na [Ne]3s1

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.

K [Ar] 4s1

Ca [Ar] 4s2

For example, the configuration of potassium and calcium are:

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:

The S atom becomes S2- anion. Similarly, the Cl atom becomes Cl- anion:

S + 2e- S2-

[Ne] 3s23p4 [Ar]

Cl + e- Cl-

[Ne] 3s23p6 [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.

The sizes of atoms vary as shown in the following

figure.

Atomic radii [in pm]

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 size

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 sizesof 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).

Na+ Mg2+ Al3+

No. of protons 11 12 13

No. of electrons 10 10 10

Ionic radius [Å] 0,95 0,65 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 energy IonizationM (g) M+(g) + e-

energyIn 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 18

H2,20

He

Li0,98

Be1,57

B2,04

C2,55

N3,04

O3,44

F3,98

Ne

Na0,93

Mg1,31

Al1,61

Si1,90

P2,19

S2,58

Cl3,16

Ar

K0,82

Ca1,00

Sc1,36

Ti1,54

V1,63

Cr1,66

Mn1,55

Fe1,83

Co1,88

Ni1,91

Cu1,90

Zn1,65

Ga1,81

Ge2,01

As2,18

Se2,55

Br2,96

Kr3,0

Rb0,82

Sr0,95

Y1,22

Zr1,33

Nb1,6

Mo2,16

Tc1,9

Ru2,2

Rh2,28

Pd2,2

Ag1,93

Cd1,69

In1,78

Sn1,96

Sb2,05

Te2,1

I2,66

Xe2,6

Cs0,79

Ba0,89

La1,27

Hf1,3

Ta1,5

W2,36

Re1,9

Os2,2

Ir2,20

Pt2,28

Au2,54

Hg2,0

Tl1,62

Pb2,33

Bi2.02

Po2,0

At2,2

Rn

Fr0,7

Ra0,9

Ac1,10

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.