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Done By Lecturer: Amal Abu- Mostafa.  OBJECTIVES: ◦ Describe periodic trends for:  A) Atomic and Ionic sizes.  B) Ionization energy.  C) Electron.

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Presentation on theme: "Done By Lecturer: Amal Abu- Mostafa.  OBJECTIVES: ◦ Describe periodic trends for:  A) Atomic and Ionic sizes.  B) Ionization energy.  C) Electron."— Presentation transcript:

1 Done By Lecturer: Amal Abu- Mostafa

2  OBJECTIVES: ◦ Describe periodic trends for:  A) Atomic and Ionic sizes.  B) Ionization energy.  C) Electron affinity.  D) And Electronegativity.

3  The wave nature of the electron makes it difficult to define exactly what we mean by the size of an atom or ion.  They get around this by measuring more than 1 atom at a time.

4  Is half the distance between the two nuclei of a diatomic molecule.

5  One factor is the principal quantum number n of the orbital; the larger n is, the larger the size of the orbital, so the atomic radius for the atom will be larger.  The other factor is the effective nuclear charge is the positive charge that an electron experiences from the nucleus, equal to the nuclear charge but reduced by any shielding or screening from any intervening electron distribution.

6  Atomic Size - Group trends  Within each group (vertical column), the atomic radius tends to increase with increasing the period number(energy level). H Li Na K Rb

7  Going from left to right across a period, the size gets smaller.  Electrons are in the same energy level.  But, there is more nuclear charge.  So Outer electrons are pulled closer. Na MgAlSiP S ClAr


9  Energy levels and Shielding have an effect on the GROUP  Nuclear charge has an effect on a PERIOD.

10  When electrons are added to an atom, the mutual repulsions between them increase. This causes the electrons to push apart and occupy a larger volume.  So, negative ions are 1.5 to 2 times larger than the atoms from which they are formed.

11  When electrons are removed from an atom, the electron - electron repulsions decrease, which allows the remaining electrons to be pulled closer together around the nucleus.  Therefore, positive ions are always smaller than the atoms from which they are formed

12  The ionization energy (IE) is the energy required to remove an electron from an isolated, gaseous atom or ion in its ground state.  The energy required to remove only the first electron is called the first ionization energy.  The second ionization energy is the energy required to remove the second electron. ◦ Always greater than first IE.

13  The third IE is the energy required to remove a third electron. ◦ Greater than 1st and 2nd IE


15  From the table, Lithium, for example, has three ionization energies because it has three electrons.  Removing the outer 1s 2 2s 1 Electron from the isolated Lithium atoms to give lithium ions Li + requires 520 kJ/mol, so the first IE = 520 kJ/mol  The second IE of lithium is = 7297 kJ/mol  Li + (g) Li 2+ (g) + e –  Removal of an electron from the now-exposed 1s 2 core of lithium requires more than thirteen times the energy used to remove the first electron


17  The greater the nuclear charge, the greater IE.  Greater distance from nucleus decreases IE  Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE.  Shielding effect.

18  The electron on the outermost energy level has to look through all the other energy levels to see the nucleus.  Electrons have the same shielding, if they are in the same period

19  1)Ionization Energy - Group trends  As you go down a group, the first IE decreases because... ◦ The electron is further away from the attraction of the nucleus, and ◦ There is more shielding. ◦ So, it will be easier to remove the electron.

20  2) Ionization Energy - Period trends  All the atoms in the same period have the same energy level.  Same shielding.  But, increasing nuclear charge.  So IE generally increases from left to right.  Exceptions at full and 1/2 full orbitals.

21  Full Energy Levels require lots of energy to remove their electrons. ◦ Noble Gases have full orbitals.  Atoms behave in ways to try and achieve a noble gas configuration.

22  The electron affinity (EA) is the potential energy change associated with the addition of an electron to a gaseous atom or ion in its ground state.  If the negative ion is stable, the energy change for its formation is a negative number (EA= - value).  When a second electron must be added, work must be done to force the electron into an already negative ion. this is an endothermic process, (EA= + value).

23  From up to down in the group the EA decreases.  2) Electron affinity –Period Trend  The general trend from left to right in any period is toward more negative electron affinities. The EA increases.  Note especially that the Group VIA and Group VIIA elements have the largest negative electron affinities of any of the main-group elements.

24  Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element.  They share the electron, but how equally do they share it?  An element with a big electronegativity means it pulls the electron towards itself strongly.

25  1)Electronegativity - Group Trend  The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has.  Thus, more willing to share.  Low electronegativity.

26  2) Electronegativity - Period Trend  Metals at the left of the table, let their electrons go easily  Thus, low electronegativity  At the right end of the table, are the nonmetals.  They want more electrons.  Try to take them away from others  So, high electronegativity.

27  The arrows indicate the trend: Ionization energy and Electronegativity INCREASE in these directions.



30  Thank you

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