Atomic Structure KeyPoints

Alex Teoh / Update 14 July 2007 / 20 Feb 2005 / 16 Feb 2002


Note: points 2 through 9 below are learning and examination assessment objectives

 

1. Historical Consideration

 

John Dalton (Britain)

developed the Atomic Theory of matter

J.J. Thompson (Britain)

discovered electrons

Henry Moseley (Britain)

atomic numbers

Niels Bohr (Denmark)

Bohr's model of the atom; electron shells

Ernest Rutherford (New Zealand)

nucleus is + charged; discovered protons

James Chadwick (Britain)

discovered neutrons

 

 

2. The 3 Sub-Atomic Particles & Their Properties

 

 

Proton

Electron

Neutron

Relative mass

1

1/2000

1

Relative charge

+1

-1

0

 

 

3. The Structure of an Atom

 

a.The structure of an atom is analogous to our solar system. Planets (electrons) revolve round the Sun ( nucleus) in fixed orbits ( electron shells ).

Thus, an atom consists of a central nucleus surrounded by moving electrons.

b.The electrons move round the nucleus in fixed paths known as electron shells.

c.The nucleus contains only protons and neutrons. Thus, the nucleus of an atom is

positively charged. However, the whole atom is neutral in charge.

Thus, in an atom, the number of protons must be equal to the number of electrons.

 

 

4. Definitions of Mass Number and Atomic Number

 

The Periodic Table is a classification of elements (refer to the Periodic Table in your textbook) Any element in the Periodic Table can be represented by its chemical symbol. Moreover, a chemical symbol consists of two numbers - one at the top left corner and one at the bottom left corner. The number at the top left corner is the Mass Number or Nucleon Number. The number at the bottom left corner is the Atomic or Proton Number. Let’s look at the element Sodium. The symbol of sodium is Na. Its mass number is 23 and its atomic number is 11. Let’s look at another element – Fluorine. Its chemical symbol is F and its mass number and atomic number are 19 and 9 respectively.

 

 

 

By definition, for an atom,

Atomic Number ( or Proton number ) = number of protons = number of electrons

Mass Number ( or Nucleon number ) = number of protons + number of neutrons

Egs: Find the number of protons, neutrons and electrons present in an atom of
a) Sodium Na
The atomic number of sodium is 11; its mass number is 23 (refer to periodic table). Thus by definition, it has 11 protons as well as 11 electrons. From the definition, we know that the Mass number = protons + neutrons. Thus, the number of neutrons = mass number – protons = 23 – 11 = 12.
b) Fluorine F
The atomic number of fluorine is 9; its mass number is 19 (refer to periodic table). Thus by definition, it has 9 protons as well as 9 electrons. From the definition, we know that the Mass number = protons + neutrons. Thus, the number of neutrons = mass number – protons = 19 – 9 = 10.

 

 

 

5.  Electronic Structure of an Atom

 

The atomic number of sodium is 11. Thus, sodium has 11 electrons revolving round its nucleus. How do we arrange the 11 electrons in their electron shells ? Can we place 5 in the 1st shell and 6 in the 2nd shell ? Or put 2 in the 1st , 4 in the 2nd and 5 in the 3rd shell ? Obviously NOT.

a) The electronic structure of an atom shows/tells us how the electrons in an atom

 are arranged in their electron shells. The rule below applies

The 1st shell can contain a maximum of 2 electrons.
The 1st shell is the shell closest to the nucleus

The 2nd shell can contain a maximum of 8 electrons.

The 3rd shell can contain a maximum of 8 electrons.
In reality, the 3rd shell can contain a maximum of 18 electrons.

 

b) The shell furthest away from the nucleus is called the outermost shell (OMS). Electrons contained in the OMS are called the outermost shell electrons or valence electrons. In all chemical reactions, it is the valence electrons that are involved in chemical reactions via chemical bonding.

c) Drawing the electronic structure of an atom

Example-let’s draw the electronic structure of a FLUORINE atom

Step 1: Refer to a Periodic Table. Look for Fluorine – find its mass and atomic numbers.
They are 19 and 9 respectively.
Step 2: The atomic number tells us that there are 9 protons as well as 9 electrons in the nucleus. The number of neutrons is computed as 19-9= 10. Draw the nucleus and state the number of protons (use the letter P) and neutrons (use the letter N) it contains.
Step 3: Draw 2 electron shells; place 2 electrons in the 1st shell and 7 electrons in the 2nd shell (draw crosses or dots for electrons)  See diagram below.

 

NOTE: If you are asked to STATE the electronic structure, you need not draw. Write the answer as 2.7

According to the examination syllabus requirements, you are expected to draw the electronic structures of the first 20 elements, ie. from Hydrogen, atomic number 1 to Calcium, atomic number 20 )



6. Electronic Structure and The Periodic Table

 

ai) Each element in the periodic table is characterized by its atomic number. No two elements in the Periodic Table can have the same atomic number !

ii) An element is composed of atoms. All the atoms in one element are identical and similar. For instance, if I were to draw a triangle to represent an atom of Iron, then all the atoms present in Iron must look like triangles. However, the atoms present in one element are different from those in another element.
iii) All elements in the periodic table are arranged in increasing atomic number

iv) Vertical columns in the PT are called Groups. Horizontal rows are called Periods. There are 8 groups (Groups I-VIII) and 7 periods.
v) Elements in Groups I-III are metals (except for Boron) while and Groups V-VIII consist mainly of non-metals (there are exceptions to the rule). Elements in Group IV contain both metals (Ge, Sn, Pb) and non-metals (Si, C).


b) All elements in the same GROUP have the same number of valence electrons eg. The electronic structures of lithium, sodium and potassium in Group I are 2.1; 2.8.1 and 2.8.8.1 respectively. Their outermost shells contain only ONE valence electron. Thus, elements in the same group have similar CHEMICAL properties However, their physical properties are different (eg. All Gp I elements react with water to form alkalis and hydrogen BUT their melting points are all different) Group Number = Number of Valence electron [eqn 1]

c) All elements in the same PERIOD have the same number of electron shells eg. The electronic structures of sodium, magnesium and aluminium m in Group I are 2.8.1; 2.8.2 and 2.8.3 respectively. Each element has THREE electron shells. Period Number = Number of Electron shells [eqn 2]

 

Note
Applying Eqn 1, any element in Group V has 5 valence electrons and any element that has 2 valence electrons must belong to Group II in the periodic table.   
Applying Eqn 2, any element in the 4th Period must contain 4 electron shells and any element that has 3 electron shells must be in the 3rd period in the periodic table
.   

 

OK girls, it’s 12 midnite now. Will continue tomorrow. Goodnite !


[ alexteoh / 15 July 2007 / Sun  / 0630h ]


 
7. Isotopes

 

a) Isotopes are atoms of the SAME element whose ATOMIC numbers are the SAME but having DIFFERENT Mass numbers Eg: Chlorine has many isotopes.



The 2 important ones are Chlorine-35 (Cl-35) and Chlorine-37 (Cl-37). Observe that both Cl-35 and Cl-37 have different mass numbers, ie. 35 and 37 respectively BUT they have the same atomic number of 17. Thus, both are isotopes.

 

Isotopes

Mass Number

neutrons

Atomic Number

protons

electrons

Electronic
structure

Valence
electrons

chlorine 35

35

18

17

17

17

2.8.7

7

chlorine 37

37

20

17

17

17

2.8.7

7

 

If isotopes of the same element have different mass numbers but the same atomic number, this implies that they have the:
- SAME number of protons
- SAME number of electrons
- SAME electronic structure
- SAME number of electron shells
- SAME number of VALENCE electrons
- SAME CHEMICAL properties.
 
but
- DIFFERENT number of neutrons
- DIFFEENT PHYSICAL properties (eg. melting, boiling point etc)

 

b) Some isotopes may be radioactive. We call them Radioisotopes or Radioactive isotopes. Radioactive isotopes are isotopes whose nuclei are unstable and dissipate their excess energy by spontaneously emitting radiation (alpha, beta or gamma rays)

Names of some Radioactive Isotopes and their Uses

* Carbon-14 – used for archaeological dating
* Cobalt-60 – used for food irradiation (food treatment- destroys bacteria in food)
* Iodine-131 – used for the treatment of thyroid cancer
* Americium 241- present in ionization chambers of smoke detectors
* Iridium 192 – locate weakness spots/areas in metal pipes or aircraft parts

 

c) Isotopic Abundance – the mass number of an element is also known as its Relative Atomic Mass. Refer to the periodic table - the relative atomic mass of chlorine is 35.5; why is it not a whole number? The two main isotopes of chlorine are present in different amounts (75% of Cl-35 and 25% of Cl-37) in our environment. Thus the average mass Or relative atomic mass of chlorine is 0.75 X 35 + 0.25 X 37 = 35.5

 

 

 

8. NOBLE GASES and Stability

 

The Group VIII elements are also known as Noble gases or Inert gases. Noble gases have STABLE electronic structures (except for Helium, all noble gases have 8 electrons in their outermost shells. This arrangement of 8 valence electrons makes noble gases very stable). Thus, they are chemically unreactive and do NOT form bonds with other atoms.


 

9. Ions & Properties

 

a) First, how are ions formed?
It is the aim of every atom to achieve a Noble gas configuration or electronic structure - this makes the atom very stable. So how does an atom attain an electronic structure that is similar to that of a Noble gas? It does this either by gaining electrons or losing electrons when forming bonds with other atoms. The electronic structures of ions resemble those of noble gases.

 

b) All ions are charged particles. They are either positively or negatively charged. Positively charged ions also called Cations; Negatively charged ones are called Anions

 

c) A POSITIVE ion(s) is/are formed when an atom LOSES electron(s)

Atoms of METALS lose electrons to form POSITIVE charged ions
Note: The electronic structure of a magnesium atom is 2.8.2 When the magnesium atom loses 2 electrons to form the magnesium ion, Mg2+ , the ion now has an electronic structure of 2.8  which resembles the electronic structure of the noble gas, Neon. That is: Mg - 2 e è Mg2+. Thus, the magnesium ion has attained stability while the magnesium atom has not.

 

d) A NEGATIVE ion(s) is/are formed when an atom GAINS electron(s)

Atoms of NON metals gain electrons to form NEGATIVE charged ions

Note: The electronic structure of a chlorine atom is 2.8.7 When the chlorine atom gains 1 electron to form the chloride ion, Cl - , the ion now has an electronic structure of 2.8.8  which resembles the electronic structure of the noble gas, Argon. That is: Cl + e è Cl- Thus, the chloride ion has attained stability while the chlorine atom has not.

 

d) IMPORTANT An atom and an ion of the same element will have the SAME number of protons and neutrons BUT different number of electrons. That means, the nucleus of an atom and the nucleus of an ion - of the same element - will always contain the same number of protons and neutrons
See the examples below.

 

 

Sodium Atom

Sodium Ion

Fluorine Atom

Fluorine Ion

Protons

11

11

9

9

Neutrons

12

12

10

10

Electrons

11

10

9

10

Electronic Structure

2.8.1

2.8

2.7

2.8

Note that the electronic structures of ions resemble those of noble gases.

Thus, the outermost shell of an ion will contain only either 2 or 8 electrons.

 


Jesus, thank you for giving me the “strength” and your Grace in completing this write-up for my students.
alexteoh / 15 July 2007 / 1317 h



Are there particles smaller in size than protons, electrons and neutrons ?
Further Optional Reading- click on the link
 The Particle Adventure