Chapter 13 - Chemical Periodicity

IONIZATION ENERGY WS

13.1 Development of the Periodic Table. 

  1. Dimitri Mendeleev, Russian chemist. Constructed 1st periodic table. Listed elements in vertical columns in order of increasing mass. Noticed regular (periodic) recurrence of physical and chemical properties. Elements with similar properties listed side by side.
  2. 1913 Henry Moseley, British physicist, determined nuclear charges, called atomic number. Arranged elements by atomic number as is today.

13.2 The Modern Periodic Table
  1. Modern table is 7 horizontal rows in order of increasing atomic number. 354-355 Rows are called PERIODS. (Correspond to principal electron levels)
  2. Vertical columns are called groups or families. Families are labeled by numbers and letters at top. (Correspond to increasing number of electrons).
  3. 1A-7A & 0 are representative elements.
  4. Group B are transition metals.
  5. Two rows at bottom are lanthanides and actinides.
  6. Elements in group have similar physical and chemical properties.
  7. Properties change from group to group. Change is a sequence and is same in all periods.
  8. Periodic law: When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

13.3 Electron Configurations and Periodicity:

Periodic Table1

Periodic table 2

Periodic table 3

 

1. Electron plays greatest part in determining properties of element. Elements classified into 4 categories by electron configuration.

Noble (group 0) gases are elements in which the outermost s and p sublevels are filled. Show He, Ne, Ar, Kr. (inert)

Representative elements are elements whose outermost s or p sublevels are only partially filled.

1 A elements are alkali metals

2 A elements are alkaline earth metals.

7 A nonmetalleic elements are halogens.

For any representative element the group number is equal to the # electrons in outermost energy level.

See examples page 357.

Transition metals are elements whose outermost s sublevel and d sublevel contain electrons. Called group B elements. They are characterized by having electrons in d orbitals. Energy level(d) is period number minus 1.

Inner transition metals are elements whose outermost s sublevel and nearby f sublevel generally contain electrons. Characterized by having electrons in f orbital. Energy level (f) is period number (6 or 7) minus 2.

2. Do example 1, 359

3. Sections or blocks of table correspond with levels s, p, d, f. See page 358.

4. s block is 1A, 2A, and He.

5. p block is 3A-0, except He.

6. d block is transition metals.

7. f block is inner transition metals.

8. Use periodic table to determine electron configuration. Read left to right, top to bottom until filled.

9. # Electrons in partially filled sublevel is determined by counting from left to element.

10. Example 1, problems 10, 11.

13.4 Periodic Trends in Atomic Size

Atomic and Ionic Radii

1. Radius of atoms cannot be measured directly. Use X-ray diffraction to estimate distance between nuclei in crystals.

2. In diatomic molecules we use half the distance between the nuclei

3. Atomic radius: Half the distance between the nuclei of two like atoms.

4. Group Trends. Size increase moving down a family due to increase in # of electrons and shielding of outer levels from attraction of nucleus, which increase in size.

5. Periodic Trends: Size generally decreases left to right. Remain in same principal energy level. Positive charge of nucleus increases (protons added), pulling level closer in and decreasing size.

6. Trend is less pronounced in atoms with more periods and electrons due to shielding.

7. Since in any period the # electrons between nucleus and period is the same, the shielding effect is constant within a period.

13. 5 Periodic Trends in Ionization energy

1. The energy that is required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom. This produces an ion.

2. Group 1A:easy to remove 1st outer electron. 2A easy to remove 1st and 2nd electron. 3A easy to remove all three. See table 13.1- 362 and figure 13.7-363.

3. Energy required to remove outmost electron, producing and ion+ is called 1st ionization energy. Next is 2nd ionization energy and so on.

4. Easy to predict removal of electrons. E.g., 1 A ionization energy jumps after the 1st electron is removed. See table 13.1.

Figure 13.7 First Ionization Energy vs Atomic Number

Group trends: Decrease as move down group. Atoms become larger and electrons farther from nucleus. Lower ionization energy is required.

Periodic trends: Representative elements increase ionization energy left to right. Nuclear charge becomes larger (shielding effect stays same in a period). Atoms become smaller. Electrons held more tightly. Higher ionization energy.

13.6 Periodic Trends in Ionic size:

1. Metallic elements have low ionization energies. Form positive ions easily.

2. Nonmetallic atoms form negative ions easily.

3. Cations (+) are always smaller than atoms. Loss of electrons causes greater attraction. E.g., Na+ is .095 nm, Na is .186 nm.

4. Anions (-) are always larger than their atoms. Gain of electron causes a lessening of attraction. E.g., Cl- is .181 nm, Cl is .099nm.

5. Group: Increases going down for both cations and anions.

6. Period: Generally decreases. See 365. Cations decrease from 1A to 3 A. Then Anions (larger to start with) decrease in size beginning with group

13.7 Periodic Trends in Electronegativity

1.Electronegativity: The tendency for the atoms of the element to attract electrons when they are chemically combined with another element.

2. Electronegativity is expressed in the Pauling electronegativity scale. See Table 13.2, 366

3. Noble gases omitted. Don't combine

4. Period: representative elements increases.

5. Group: decreases as move down a group.

6. Transition metals are not regular.

Figure 13.10