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Ismail awcabdi Atomic structure Chapter one: Atomic structure Atomic number The number of protons in an atom is known as its atomic number. The atomic number of an element shows: • The number of protons in the nucleus of an atom of that element. • The number of electrons in a neutral atom of that element. • The position of the element in the periodic table

Mass number The mass number is the sum of the number of protons and the number of neutrons in the nucleus of an atom.

Summary table Particle name Relative mass Relative charge electron Almost no mass -1 Proton 1 +1 neutron 1 none

Isotopes Atoms of the same element with same number of protons but different number of neutrons. Isotopes: • have the same atomic number • have different mass number • have same chemical properties the word isotope mean (equal place) i.e occupying the same place in the periodic table and having the same atomic number.

                                       Electrons in an atom Electrons hold the key to almost the whole chemistry. Protons and neutrons give atoms their mass. Only electrons determines (are involved in) changes that happen during chemical reactions 
                                        
                                       Ionization energy

When an atom loses an electron it becomes a positive charge ion. So it has been ionized and the energy needed to remove one electron is called ionization energy.

The first ionization energy. The first ionization energy of an element is the amount of energy needed to remove one electron from each atom in a mole of atom of an element in the gaseous state.

The general symbol for the ionization energy ∆Hi and the first ionization energy is ∆Hi1. For example:

Ca(g)------------------------> Ca+ + e- ∆Hi1 = +590kj mol-1

The second ionization energy The energy needed to remove a second electron from each ion is the second ionization energy. For example calcium:

Ca+(g) ------------- Ca2+ (g) + e- ∆Hi2 = +1150

                                       The successive ionization energy

We can continue removing electrons until only the nucleus of an atom is left. The sequence of the first, second, third, fourth etc ionization energies is called successive energy. for the first eleven elements in the periodic table are shown this table

Removed electrons 1 2 3 4 5 6 7 8 9 10 11 1 H 1310 2 He 2370 5250 3 Li 520 7300 11800 4 Be 900 1760 14850 2100 5 B 800 2420 3660 25000 32800 6 C 1090 2350 4620 6220 37800 47300 7 N 1400 2860 4580 7480 9450 53300 8 O 1310 3390 5320 7450 11000 13300 71300 84100 9 F 1680 3470 6040 8410 11000 15200 17900 92000 106000 10 Ne 2080 3950 6120 9370 12200 15200 20000 23070 115380 131400 11 Na 510 4560 6940 9540 13400 16600 20100 25500 28900 141000 158700

• The ionization energy increase as each electron is removed from an atom, the remaining ion becomes more positively charged. Moving the next electron away from the increased positive charge is more difficult and the next ionization energy is even larger. • There are one or more particularly large rises within the set of ionization energies of each element (except hydrogen and helium)

Factors influencing the ionization energies  The size of the positive nuclear charge: this charge affects all the electrons in an atom. The increase in nuclear charge with atomic number will tend to cause an increase in an ionization energy  The distance of the electron from the nucleus: this distance effects means that all force attraction decrease rapidly as the distance between the attracted bodies increase and the lower ionization energy.

The shielding effect by electrons in filled inner shells: all electrons are negatively charged and repel each other. Electrons in the inner shells repel electrons in the outer shells and reduce the effect of the positive nuclear charge. This called the shield

 ing effect. the greater the shielding effect upon an electron the lower is the energy required it.

Evidence for the existence of electron shells. In the successive ionization energies of lithium, we see a low first ionization energy, followed by much larger second and third ionization energies. This is confirms that lithium has one electron in its outer shell n = 2, which is easier to remove than of the two electrons in the inner shell n = 1. The larger increase in ionization energies indicates where there is a change from shell n = 2 to n = 1.

Subshells and orbitals  The energy levels (shells) of principle quantum numbers n =1, 2, 3, 4, etc. do not have precise energy values. Instead, they each consist of a set of Subshells, which contain orbitals with different energy values.  The subshells are labeled s, p, d, and f. an s subshell contains one orbital; a p subshell contains three orbitals; a d subshell contains five orbitals; an f subshell contains seven orbitals.  An electron orbital represents a region of space around the nucleus of an atom, within which there is a high chance of finding that particular electron.

The shapes of s and p orbitals Each orbital has its own, three dimensional shape. It is not possible to draw shape orbitals precisely. They do not have exact boundaries but are fussy, like clouds; they are often called charge clouds. Note that there is only one type of s orbital but three different p orbitals ( px, py, pz). Approximate representation of orbitals are shown below:

S Orbital

ppppx orbital py orbital pz orbital P orbitals

Order of filling shells and orbitals

Any individual orbital can hold one or two but not more. In each successive of the periodic table, the order of filling the shells and orbitals is the order of their relative energy. The lowest energy orbitals are filled first. The order of filling is:

First 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p…

An expected order is followed up to the 3p subshells, but there is a variation, as the 4s is filled before the 3d. this variation and other variations further along in the order are caused by the increasingly complex influences of nuclear attractions and electron repulsions upon individual electrons.

Electronic configurations The most common way of representing the electronic configuration is shown below. For example, hydrogen, has one electron in an s orbital in the shell with the principle quantum number n = 1. We show this as:

                                                    	Number of electrons
                                                                               

          1s1
                    Principle quantum number
                                                Orbitals

Principle quantum number

Electronic configuration for the first 36 elements in the periodic table are shown here

Num. Symbol K L M N O P Q 1. Period 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 1 H 1 2 He 2 2. Period 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 3 Li 2 1 4 Be 2 2 5 B 2 2 1 6 C 2 2 2 7 N 2 2 3 8 O 2 2 4 9 F 2 2 5 10 Ne 2 2 6 3. Period 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 11 Na 2 2 6 1 12 Mg 2 2 6 2 13 Al 2 2 6 2 1 14 Si 2 2 6 2 2 15 P 2 2 6 2 3 16 S 2 2 6 2 4 17 Cl 2 2 6 2 5 18 Ar 2 2 6 2 6 4. Period 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 19 K 2 2 6 2 6 .. 1 20 Ca 2 2 6 2 6 .. 2 21 Sc 2 2 6 2 6 1 2 22 Ti 2 2 6 2 6 2 2 23 V 2 2 6 2 6 3 2 24 Cr 2 2 6 2 6 5 1 25 Mn 2 2 6 2 6 5 2 26 Fe 2 2 6 2 6 6 2 27 Co 2 2 6 2 6 7 2 28 Ni 2 2 6 2 6 8 2 29 Cu 2 2 6 2 6 10 1 30 Zn 2 2 6 2 6 10 2 31 Ga 2 2 6 2 6 10 2 1 32 Ge 2 2 6 2 6 10 2 2 33 As 2 2 6 2 6 10 2 3 34 Se 2 2 6 2 6 10 2 4 35 Br 2 2 6 2 6 10 2 5 36 Kr 2 2 6 2 6 10 2 6

The following points should be noted:  When the 4s orbital is filled, the next electron goes into 3d orbital (see scandium).this begins a pattern of filling up 3d subshells which finishes at zinc. The elements that add electrons to the d subshells are called the-d block elements (transition elements)  There are variations in the pattern of filling the d subshell at elements 24(chromium) and 29(copper). These elements have only one electron in there 4s orbitals. Chromium has 5d electrons rather than four, copper has ten d electrons rather than nine. This is the outcome of the complex interaction of attraction and repulsion in their atoms.  From elements 31(gallium) to 36(krypton) the electrons add to the 4p subshells. This is similar to the pattern of filling the 3p subshell from elements 13(aluminum to 18(argon) in period 3.

Electronic configuration of ions The number of electrons in an ion is found from the atomic number of the element and the charge of the ion.

Sodium atom Sodium ion Fluorine atom Fluoride ion Symbol Na Na+ F F- Atomic number 11 11 9 9 Electrons 11 10 9 10 configuration 1s2 2s22p63s1 1s2 2s22p6 1s2 2s22p5 1s2 2s22p6 Note that both the sodium ion Na+ and Fluoride ion F- have the same electronic configuration as the noble gas neon.

Electronic configuration of boxes Another useful way of representing electronic configuration is in box form. We can show the electrons as arrows with their clockwise or anticlockwise spin as or The electronic configuration of some elements represented in this way: