Upcoming SlideShare
×

# Periodic Relationships

1,889 views

Published on

Published in: Technology, Economy & Finance
1 Like
Statistics
Notes
• Full Name
Comment goes here.

Are you sure you want to Yes No
Your message goes here
• Be the first to comment

Views
Total views
1,889
On SlideShare
0
From Embeds
0
Number of Embeds
2
Actions
Shares
0
39
0
Likes
1
Embeds 0
No embeds

No notes for slide

### Periodic Relationships

1. 1. Periodic Relationships T- 1-855-694-8886 Email- info@iTutor.com By iTutor.com
2. 2. PERIODIC TRENDS IN PROPERTIES OF ELEMENTS  There are many observable patterns in the physical and chemical properties of elements as we descend in a group or move across a period in the Periodic Table.  We will rationalize observed trends in  Sizes of atoms and ions.  Ionization energy.  Electron affinity. The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. The Size of an Atom
3. 3. Trend in Atomic Radius  Different methods for measuring the radius of an atom, and they give slightly different trends  Van der Waals radius = Nonbonding  Covalent radius = Bonding radius  Atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds.
4. 4.  Valence shell farther from nucleus  Effective nuclear charge fairly close  Atomic Radius Decreases across period (left to right) The number of energy levels increases as you move down a group as the number of electrons increases. Each subsequent energy level is further from the nucleus than the last. Therefore, the atomic radius increases as the group and energy levels increase.  Atomic Radius Increases down group Atomic Radius adding electrons to same valence shell effective nuclear charge increases valence shell held closer - As you go across a period, electrons are added to the same energy level. At the same time, protons are being added to the nucleus. The concentration of more protons in the nucleus creates a "higher effective nuclear charge." In other words, there is a stronger force of attraction pulling the electrons closer to the nucleus resulting in a smaller atomic radius.
5. 5. Atomic Radius
6. 6. Sizes of Ions • Ionic size depends upon – The nuclear charge. – The number of electrons. – The orbital in which electrons reside.
7. 7. Sizes of Ions • Cat ions are smaller than their parent atoms: – The outermost electron is removed and repulsions between electrons are reduced. • Anions are larger than their parent atoms – Electrons are added and repulsions between electrons are increased. • In an Isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge.
8. 8. Ionization Energy The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. – The first ionization energy is that energy required to remove the first electron. – The second ionization energy is that energy required to remove the second electron, etc. • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap.
9. 9. Ionization Energy
10. 10. First Ionization Energies  Larger the effective nuclear charge on the electron, the more energy it takes to remove it  The farther the most probable distance the electron is from the nucleus, the less energy it takes to remove it  1st IE decreases down the group  valence electron farther from nucleus  1st IE generally increases across the period  effective nuclear charge increases
11. 11. First Ionization Energies
12. 12. First Ionization Energies
13. 13. Irregularities in the Trend  Ionization Energy generally increases from left to right across a Period  except from 2A to 3A, 5A to 6A Be 1s 2s 2p B 1s 2s 2p N 1s 2s 2p O 1s 2s 2p Which is easier to remove an electron from B or Be? Why? Which is easier to remove an electron from N or O? Why?
14. 14. Irregularities in the First Ionization Energy Trends Be 1s 2s 2p B 1s 2s 2p Be+ 1s 2s 2p To ionize Be you must break up a full sublevel, cost extra energy B+ 1s 2s 2p When you ionize B you get a full sublevel, costs less energy
15. 15. Irregularities in the First Ionization Energy Trends To ionize N you must break up a half-full sublevel, cost extra energy N+ 1s 2s 2p O 1s 2s 2p N 1s 2s 2p O+ 1s 2s 2p When you ionize O you get a half-full sublevel, costs less energy
16. 16. Trends in Successive Ionization Energies  Removal of each successive electron costs more energy – shrinkage in size due to having more protons than electrons – outer electrons closer to the nucleus, therefore harder to remove  Regular increase in energy for each successive valence electron  Rarge increase in energy when start removing core electrons
17. 17. Ionization Energies (kJ/mol)
18. 18. Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl + e− Cl− 1) As you move down a group, electron affinity decreases. 2) As you move across a period, electron affinity increases.
19. 19. Electron Affinity  The first occurs between Groups IA and IIA. – The added electron must go in a p orbital, not an s orbital. – The electron is farther from the nucleus and feels repulsion from the s electrons.  The second discontinuity occurs between Groups IVA and VA. – Group VA has no empty orbitals. – The extra electron must go into an already occupied orbital, creating repulsion.
20. 20. Oxidation States  A way of keeping track of the electrons.  Not necessarily true of what is in nature, but it works.  need the rules for assigning . The oxidation state of elements in their standard states is zero. Oxidation state for monatomic ions are the same as their charge. Oxygen is assigned an oxidation state of -2 in its covalent compounds except as a peroxide. In compounds with nonmetals hydrogen is assigned the oxidation state +1. In its compounds fluorine is always –1. The sum of the oxidation states must be zero in compounds or equal the charge of the ion.
21. 21. Oxidation States 1. The oxidation state of any element such as Fe, H2, O2, P4, S8 is zero (0). 2. The oxidation state of oxygen in its compounds is -2, except for peroxides like H2O2, and Na2O2, in which the oxidation state for O is -1. 3. The oxidation state of hydrogen is +1 in its compounds, except for metal hydrides, such as NaH, LiH, etc., in which the oxidation state for H is -1. 4. The oxidation states of other elements are then assigned to make the algebraic sum of the oxidation states equal to the net charge on the molecule or ion. 5. The following elements usually have the same oxidation states in their compounds:+1 for alkali metals - Li, Na, K, Rb, Cs; 6. +2 for alkaline earth metals - Be, Mg, Ca, Sr, Ba; 7. -1 for halogens except when they form compounds with oxygen or one another;
22. 22. Element Oxidation state Compound or ion Fe +2 Fe2+ Fe = Fe2+ + 2 e- +3 Fe3+ Fe2+ = Fe3++ e- Zn 0 Zn Zn is reducing agent +2 Zn2+ O -1 H2O2 H2O2 = O2 + H2O 0 O2 -2 H2O Oxidation States
23. 23. The End For more information call us 1-855-694-8886 Visit www.iTutor.com