1. PERIODIC PROPERTIES
Chapter 6

2. Atomic structure and periodic trends
The atomic number (number of protons) defines the element and is now the order in which elements are arranged
The columns are called groups and were originally arranged because elements in the same group had similar properties, they are also called famiiles.
The rows are called periods ,
the elements on the left side of the period are more metallic and those on the right nonmetallic

3. Electron arrangement and periodic table
Electrons are the most important to chemists because they are the particles involved in chemical reactions.
The number of electrons = number of protons in a neutral atom.
As the atomic number increases, the number of electrons increases. For the first 20 elements, the period on the periodic table matches Bohr’s energy levels in the atom

4. Electron structure
In period 1 there are two elements : Hydrogen (Z=1) and Helium (Z=2), the maximum number of electrons in the first energy level is 2
In period 2 there are eight more elements ending at Neon, this period can hold 8 electrons
In period 3 there are eight more elements ending at Argon, this period can hold 8 electrons

5. Subdivisions of the periodic table and electron arrangement
The Bohr energy level can be divided into sublevels: s, p, d, and f.
Each sublevel differs slightly in energy
These match the sections of the periodic table

6. Valence electrons and the periodic table
Electrons fill the lowest energy locations in an atom first. For the first 20 elements the last electrons in the atom are placed in the highest energy level . These electrons are called the valence electrons.
The valence electrons are the ones most likely involved in chemical reactions
In the first two columns the valence electrons are in the s-sublevel.
In the last six columns the valence electrons are in the p-sublevel.

7. Valence electrons continued
There are only 8 valence electrons in any atom
By knowing the location of an element you can determine the location of the valence electrons in an element

8. Periodic trends in physical properties
Atomic Radius is the term used to describe the size of an atom
For elements on the right side of the periodic table (nonmetals) it is defined as half the distance between two covalently bonded atoms

9. Trends in atomic radii
Atomic radii is a major factor in explaining many periodic trends
Trend: Atomic radius increases top to bottom in a group and decreases left to right in a period

10. Cause of trends in Atomic radius
There are 3 primary factors:
1) the energy level of the valence electrons. The higher the energy level, the greater distance of the valence electron from the nucleus
2) Number of protons in the nucleus. The more protons the greater the attraction of the nucleus for the negatively charged electrons
3) The shielding of the valence electrons from the nucleus due to the repulsion between the inner energy shell electrons

11. Causes of trends continued
An effective way to consider the effect of the nucleus on a valence electron is to introduce the idea of the effective nuclear charge ( Z effective)
This is calculated by adding the total charge of the protons to the charge of the core electrons
For example the Z effective for Na is ((+11)+ (-10)) = +1

12. Trend in atomic radius in a group
As you go down a group the atomic radius increases.
Why?
1) Each element’s valence electrons are in a higher energy shell so they are farther from the nucleus
2) The Z effective is about the same because the atomic number increases but the number of core electrons also increases
This causes the force of attraction between the nucleus and the valence electron to weaken, allowing the cloud of electrons to expand.

13. Physical Properties
Define the terms first ionization energy and electron affinity
The first ionization energy is the minimum energy needed to remove one mole of electrons from one mole of gaseous ions at 25 C and 1 atm ( standard thermodynamic conditions)
This is the energy involved in forming cations ( positive ions) from neutral atoms and is part of the process of forming ionic bonds.
Electron affinity is the energy released when an atom attracts an electron forming an anion ( a negative ion).

14. First Ionization Energy trends
Ionization energy decreases down a group because:
1) The valence electrons are in higher energy levels, which are farther from the nucleus, weakening the attraction of the nucleus
2) The effective nuclear charge is essentially the same for each element in the group because another energy shell of core electrons is added.
3) The electron-electron repulsion increases with additional electrons.
Since the force between nucleus and electrons weakens with distance( which is increasing) and increases with effective nuclear charge( which is constant) The energy to remove an electron decreases

15. Trend in atomic radius in a period
3.2.3 Describe and explain the trends in atomic radius, first ionization energy and electronegativity for elements across period 3
As you go left to right in a period the atomic radius decreases.
Why?
1) Each element’s valence electrons are in the same energy shell so they start at the same distance from the nucleus
2) The Z effective increases because the atomic number increases but the number of core electrons stays the same
This causes the force of attraction between the nucleus and the valence electron to increase, causing the cloud of electrons to contract.

16. Trends in first ionization energy
The first energy increases moving to the right in the period.
What are the three factors that cause this??

17. Trends in electron affinity
3.2.4 Compare the relative electron affinity values of two or more elements based on their positions in the periodic table
Electron affinity values follow the same trend as ionization energy
What are the reasons???