The Periodic Law Notes (Chapter 5)

I. History of the Periodic Table

About 70 elements were known by 1850 (no noble gases) but there didn’t appear to be a good way of arranging or relating them to study.

A. Mendeleev and Chemical Periodicity

Mendeleev placed known information of elements on cards (atomic mass, density, etc…). He arranged them in order of increasing atomic masses, certain similarities in their chemical properties appeared at regular intervals. Such a repeating pattern is referred to as periodic.

Dmitri Mendeleev 1869Dmitri Mendeleev 1869

http://www.chem.msu.su/eng/misc/mendeleev/welcome.html

1. Mendeleev’s table was published in 1869.2. He left blanks in his periodic table for undiscovered elements and he predicted their properties. Later elements were discovered with properties he predicted!3. Problems with his table – a few elements did not fit – the atomic mass arrangement did not match with other similar properties.4. Recognition – Mendeleev never received the Nobel Prize – the importance of the Periodic Table was not realized in his lifetime!

B. Moseley and the Periodic Law

1. Moseley, with Rutherford, in 1911 discovered a new pattern. The positive charge of the nucleus increased by one unit from one element to the next when elements are arranged as they are now on the Periodic Table2. This discovery led to definition of atomic number and the reorganization of the Periodic Table based on atomic number not atomic mass.3. The Periodic Law – the physical and chemical properties of the elements are periodic functions of their atomic numbers.4. Moseley died at the age of 28 – victim of WWI

C. The Modern Periodic Table

1. An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column (or group).Groups: vertical columns (#1-18)Periods: horizontal rows (# 1-7)2. Periodicity – the similarities of the elements in the same group is explained by the arrangement of the electrons around the nucleus.

Periodic Table Expanded ViewPeriodic Table Expanded View

The way the periodic The way the periodic table is usually seen is a table is usually seen is a compressed view, placing compressed view, placing the Lanthanides and the Lanthanides and Actinides at the bottom of Actinides at the bottom of the table.the table.

Layout of the Periodic Table: Metals vs. Nonmetals

1IA

18VIIIA

12

IIA13

IIIA14

IVA15VA

16VIA

17VIIA

2

33

IIIB4

IVB5

VB6

VIB7

VIIB8 9

VIIIB10 11

IB12IIB

4

5

6

7

MetalsMetalsNonmetalsNonmetals

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II. Electron Configurations and the Periodic Table

A. The s-block Elements: Groups 1 and 2

1. Group 1: Alkali metals -soft silvery metals

-most reactive of all metals, never found free in nature-reacts with water to form alkaline or basic solutions – store under kerosene-whenever you mix Li, Na, K, Rb, Cs, or Fr with water it will explode and produce an alkaline solution-ns1 (ending of all electron configurations for this group) http://www.teachertube.com/members/viewVideo.php?title=Group_1_Alkali_metals_and_water&video_id=51068

2. Group 2: Alkaline earth metals

-less reactive than Alkali, but still react in water to produce an alkaline solution

-never found free in nature

-harder, denser, stronger than alkali

- ns2 (ending of all electron configurations for this group), because they have 2 electrons in the s sublevel, this makes them a little less reactive then the Alkali metals in group 1

B. The d-Block Elements: Groups 3-12

-are all metals with metallic properties (malleability, luster, good conductors, etc…); are referred to as the Transition Metals

-Harder and denser than alkali or alkaline-Less reactive than alkali or alkaline-For the most part their outermost electrons are

in a d sublevel-Exceptions to the electron configuration are

found in these groups (Ex: Ni, Pd, Pt)

Irregular Electron configurations

– sometimes the electron configuration is NOT what we would predict it to be. Sometimes electrons are moved because (l) it will result in greater stability for that atom or (2) for some unknown reason?? It is very important to define “stable” here. STABLE means:all degenerate (equal energy) orbital’s are FULLall degenerate orbital’s are half-fullall degenerate orbital’s are totally empty.

Examples – draw the orbital notation (lines and arrows) for the predicted electron configuration for Cr #24:

However, the real E. C. is [Ar]4s13d5. The 4s1 electron has been moved to achieve greater stability. The elements with “irregular” electron configurations are Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, La, Pt, Au, Ac, Gd, Th, Pa, U, Np, and Cm. We will also use the actual electron configuration (found on the periodic table at the back of your book) instead of the predicted configuration for these elements. You do NOT have to memorize these, they will be highlighted or marked on your periodic table.

C. The p – Block Elements: Groups 13 – 18

-Contain metals and nonmetals

-Metalloids, along zigzag line, have characteristics of both metals and nonmetals (many are good conductors but are brittle). The metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium.

1. Group 17 - Halogens – most reactive nonmetals

-7 electrons in outermost (s and p) energy levels (that is why so reactive – only need one electron to have 8)

-called the salt formers (they react vigorously with metals to form salts) A salt is a metal and a nonmetal bonded together.

-most are gases

2. Group 18 - Noble gases – unreactive

- 8 electrons in outermost s and p energy levels

- all are gases

The s and p blocks are called the main group or representative elements!

D. The f-Block Elements: Inner Transition Metals

-final electrons fill an f sublevel

 

1. Lanthanides – shiny reactive metals; Ce-Lu (fill the 4f sublevel)

2. Actinides – unstable and radioactive; Th-Lr (fill the 5f sublevel)

Hydrogen and Helium - Oddballs

-Hydrogen is NOT an Alkali metal, it is a very reactive gas. It is placed with the Alkali metals because 1s1 is its electron configuration.

-Helium is a Noble gas, it is unreactive, but it does not have 8 electrons in outermost energy level, because it only has 2 total electrons!

III. Electron Configuration and Periodic Properties

A. Valence Electrons – electrons in the outermost s and p orbital’s; electrons available to be lost, gained, or shard in the formation of chemical compounds

When all of the valence orbital’s are full (have 8 electrons), the atom tends to be unreactive (like the Noble Gases)

# of Valence Electrons Group # Ending Configuration1 1 ns1 very reactive2 2 ns2 3 13 np1

4 14 np2

5 15 np3

6 16 np4

7 17 np5 very reactive8 18 np6 very unreactive

B. Atomic Radii

1. The size of an atomic radius cannot be measured exactly because it does not have a sharply defined boundary. However the atomic radius can be thought of as ½ the distance between the nuclei of identical atoms joined in a molecule

  

 

Atomic Radii Trend

2. Period trend – atomic radii decrease as you move across a period.

As you move across a period, from left to right, the size of an atom decreases because the nucleus is getting larger and more positive, but it is still pulling on the same number of energy levels. Which atom is larger?

(A) Zr or Sn (B) Li or Cs

3. Group trend - atomic radii decrease as you move up a group (or increase as you move down a group).

Shielding effect - an invisible barrier made of core electrons serve to decrease the pull of the nucleus on the outer (valence) electrons.

Shielding increases as you go down a group because there are more core electrons.

Shielding is considered to be constant as you move across a period because the number of energy levels is staying the same.

Which atom has more shielding? (A) K or Ca (B) Na or KWhich atom is smaller? (A) N or P (B) Li or K

C. Ionization Energy (IE)

– a measure of the ease with which an electron can be removed from a gaseous atom or ion (sometimes called ionization potential).

1. First ionization energy – the energy required to remove one outermost electron from a gaseous atom. Second ionization energy – the additional amount of energy needed to remove an outermost electron from the gaseous +1 ion.

IE Trend

2. Period trend – ionization energy increases as you move across the period.As you move across a period the ionization energy increases because the atoms get smaller. Another way to think of it: the number of valence electrons increases (the amount of energy needed to remove one electron is less then what is needed to remove 7 or 8 electrons).

3. Group trend – ionization energy increases as you move up a group (or decreases as you move down a group).In general, as you do down a group the ionization energy decreases because the size of the atom is increasing and the outermost electrons are further from the nucleus.

1. Which atom has the higher first ionization energy? (A) Hf or Pt (B) Cl or Ar

D. Electron Affinity (EA)

– the energy change that accompanies the addition of an electron to a gaseous atom. The desire of an atom to have another electron. Chlorine for example would like to have another electron and therefore has a very high electron affinity.

EA Trend

1. Period trend – electron affinity increases as you move across a period because atoms become smaller and the nuclear charge increases. This means there is a greater pull

from the nucleus. 2. Group trend – electron affinity increases as you move

up a group (or decreases as you move down a group) because the size of the atom increases and shielding increases. This means there is less pull from the nucleus.

Example: Which element has the greater electron affinity? Pb or Sn

E. Ionic Radii

1. Period trend – The size of ions decreases as you move across a period because you have more protons pulling on the same number of energy levels.

2. Group trend – the size of the ions decrease as you move up a group (or increase as you move down a group) because the number of energy levels increases.

Example: Which would be larger? K+1 or Ge +4

Which is larger? P ion or Cl ion

Ionic Radii Trend

3. Metals – metal ions (cations) always have a smaller radii than its corresponding atom because:

a. it loses its outer energy level electrons (valence electrons)b. the proton to electron ratio is greater in the ion than in the atom

sodium atom 11p+/11e- = 1.0sodium ion 11p+/10e- = 1.1

The value of p+ ro e- ratio varies inversely to the size of the ion. The sodium ion is smaller because it has a larger proton to electron ratio.

Which would be larger? Potassium atom or potassium ion

4. Nonmetals - nonmetal ions are always larger than their corresponding atoms because:

a. repulsion between electronsb. the p+ to e- ratio is less in the ion

chlorine atom 17p+/17e- = 1.0chlorine ion 17p+/18e- = 0.94

The chlorine ion is larger because it has a smaller proton to electron ratio.

Which would be larger? Sulfur atom or sulfur ion

F. Electronegativity (EN)

– the tendency of an atom to attract electrons to itself when it is chemically combined (bonded) to another element.

Electronegativity Trend

In general, metals have low EN and nonmetals have high EN. The actual amount of EN an atom has is indicated by a number of the Pauling Electronegativity Scale that goes from 0 to 4. Dr. Linus Pauling set up this scale and gave the element having the greatest EN an arbitrary number of 4, and he assigned numbers to the others relative to this element.Flourine is the most electronegative element at 4. (3.98) and Francium is the least electronegative at 0.7.

1. Period trend - EN increases as you go across a period (excluding the noble gases) because size decreases.

2. Group trend - EN increases as you go up a group (or decreases as you go down a group) because there is less pull from the nucleus as the electrons get further away.

Which would have the greater EN? (A) Ca or Se (B) Be or O

Electronegativity enables us to predict what type of bond will be formed when two elements combine.