PERIODIC TABLE

VOCABULARY:

Isoelectronic Elements: are elements that have the same electron configuration.

Periodic Law: The properties of the elements are a periodic function of their atomic numbers.

Periodic Table: A pictorial arrangement of the elements based upon their atomic numbers and electron configurations.

Transition Element: An element whose highest energy electron is in the d orbital.

Lanthanoid Series: Fourteen elements beginning with lanthanum in which the highest energy electrons to be in the 4f sublevel.

Actinoid Series: Fourteen elements beginning with actinium in which the highest energy electrons to be in the 5f sublevel.

Period: A horizontal row on the periodic table.

Group: The elements of a vertical column in the periodic table.

Octet Rule: An especially stable arrangement of four pairs of electrons in the outer energy level of an atom.

Family: The elements composing a vertical column of the periodic table.

Metal: An element that tends to lose electrons in chemical reactions.

Nonmetal: An element that tends to gain electrons in chemical reactions.

Metalloid: An element that has properties characteristic of a metal and a nonmetal.


DEVELOPING THE PERIODIC TABLE:

Elements were first grouped by:

- ___________________ / ___________________ Properties

- Metallic / Non-Metals

- Increasing ___________________

- How they react with other ___________________

* * However, predicting physical and chemical properties was more difficult without an organizational scheme.


ORGANIZING BY PROPERTIES:

Mendeleev (1869, Russian Chemist)

Developed the first ___________________ :

- Horizontal Rows (Periods) by increasing Atomic ___________________

- Vertical Columns (Groups) by similar chemical ___________________ (how the elements reacted with oxygen).

* Example: Group 1A, Li2O, Group 2A, MgO

Mendeleev's Periodic Table had some missing spots because those elements had not been discovered yet.

Mendeleev not only predicted that the missing elements would be discovered, he also predicted the ___________________ of some of those elements.

Moseley (1913, British Physicist)

Moseley arranged the periodic table by increasing ___________________ , and the discrepancies disappeared.


PERIODIC LAW:

Periodic Law states that the properties of elements repeat ___________________ when the elements are arranged in ___________________ order by their atomic numbers.


CHEMICAL REACTIVITY:

From the above table, one can see the repeating patterns in Ratios of how the elements react with Hydrogen and Fluorine.

Ratios 1, 2, 3, 4, 3, 2, 1


IONIZATION ENERGIES:

First Ionization Energy is the energy needed to remove One mole of ___________________ from One mole of Atoms ( in the gaseous phase ).

Ionization Energy increases from Left to Right along a Period.


THE PERIODIC TABLE - TODAY:

In the modern periodic table, each vertical column identifies a Group of elements, or a Chemical ___________________ . The Groups are identified by numbers across the top of the periodic table. The Horizontal Rows of elements are called ___________________ .


CHEMICAL FAMILIES:

ALKALI METALS: GROUP 1A

All alkali metals react with H2O to form an alkaline (Basic) Solution.

ALKALINE EARTH METALS: GROUP 2A

All alkali earth metals also react with H2O to form a basic solution.

HALOGENS: GROUP 7A

Salt Formers

NOBLE GASES: GROUP 8A

Noble Gases have all their orbitals Filled.

The other groups are identified by the element at the top of the column. Example, Group 4A is the called the Carbon Family (or Group).


REPRESENTATIVE ELEMENTS:

The elements in Groups 1A - 8A, are known as the ___________________ or the Representative Elements.

TRANSITION METALS: The "B" Group Elements between 2A and 3A. (d-orbitals area).

INNER-TRANSITION METALS: Lathanide and Actinide Elements (are separated to make it easier to fit on a single page).


PERIODS:

Periods have different numbers of elements:

PERIOD 1 2 ELEMENTS
PERIOD 2 & 3 8 ELEMENTS
PERIOD 4 & 5 16 ELEMENTS
PERIOD 6 32 ELEMENTS
PERIOD 7 NOT COMPLETED - Less than half are Natural Elements (Synthetic- Radioactive)


PATTERNS IN ELECTRON CONFIGURATION:

Mendeleev's organization of the elements was based on the Macroscopic chemical and physical properties. Once Subatomic particles were discovered, chemists sought models of the atom that would relate the subatomic structure of the atom to its chemical properties.

ELECTRON CONFIGURATION OF NOBLE GASES:

For Noble Gases, each has a total of 8 electrons in the ___________________ s and p orbitals. The outermost orbitals are the orbitals with the highest principal quantum number that are occupied by electrons in the ground state electron configuration.

Electrons in the Outermost s and p Orbitals are often involved in ___________________ (Valence Electrons).

The remaining electrons are called the ___________________ Electrons.

Remember that the maximum number of electrons in an orbital is ___________________ and that there are one s orbital and three p orbitals for a given principal quantum number. Therefore, the maximum number of valence electrons is ___________________ .


ELECTRON CONFIGURATION OF ALKALI METALS:

Each element in the Alkali metal group has ___________________ valence electron in the Outermost s orbital.

ELECTRON CONFIGURATION OF ALKALINE EARTH METALS:

Each element in the Alkaline Earth metal group has ___________________ valence electrons in its Outermost s orbital.

ELECTRON CONFIGURATION OF HALOGENS:

Each element in the Halogen group has ___________________ valence electrons in its Outermost s and p orbitals. (2 in s, and 5 in p)

DETERMINING THE NUMBER OF VALENCE ELECTRONS:

Generally speaking, The number of Valence Electrons is ___________________ to the Group number ( For Groups 1A - 8A ).


PATTERNS FOR REPRESENTATIVE ELEMENT IONS:

When the Valence Orbitals are ___________________ , the Atom tends to be Un-Reactive. (example: Noble Gases)

ALKALI METALS:

One electron is in the "s" orbital, typically alkali metals lose ___________________ electron in a chemical reaction to form a plus one charge ( example: Na+1 ion).

Example: Electron Configuration for sodium:

Na: 1s22s22p63s1

Na: [ Ne ] 3s1 (where [ Ne ] is representing the Core Electrons)

When an electron is donated by sodium, the electron leaves the 3s orbital.

The new electron configuration is now:

Na+: 1s22s22p6

Na+: [ Ne ]

ISOELECTRONIC ELEMENTS: are elements that have the same Electron Configuration. ( Na+ and Ne )

ALKALINE EARTH METALS:

Lose ___________________ electrons, and has a +2 charge as an ion. (loses two valence electrons)

HALOGENS:

Seven Valence Electrons - Easier to gain one electron, therefore has a -1 charge.

NOBLE GAS ELEMENTS:

Are Un-Reactive - ___________________ Valence Orbitals


HOMEWORK PROBLEMS:

1a. State the "periodic trend" for the periodic table in your own words.

1b. From the representative elements (Groups 1A - 8A, no Group B's), name one group (column) that contains only metallic elements and two groups that contain only nonmetallic elements.

1c. What is the name and symbol for the element in each of the following positions on the periodic table?
i) Period 3, Group 7A.
ii) Period 2, Group 6A
iii) Period 5, Group 1B
iv) Period 6, Group 4A

1d. From the following list of elements: K, Ca, Cl, U, La, Sr, Kr
i) Which elements are inner-transition metals (if any)?
ii) Which are alkaline earth metals from the following (if any)?

1e. Explain isoelectronic, and give one example of two isoelectronic atoms (hint: be sure to show the correct charges)

Answers:


FORMING IONS:

In General, Metals ___________________ electrons and Non-Metals ___________________ electrons when they React together to form Compounds.

Ions of Representative Elements are formed from the loss and gain of electrons often have electron configurations that are identical to the electron configurations of the ___________________ (Group 8A).


PATTERNS FOR TRANSITION ELEMENTS AND IONS:

FORTH PERIOD TRANSITION ELEMENTS:

The actual electron configuration can be determined ONLY by experiments.

Some unexpected configurations occur (Chromium, Cr, & Copper, Cu for example) because a special ___________________ occurs when sets of d-orbitals are exactly half filled OR completely filled (more stable) AND s-orbitals are half filled OR completely EMPTY - as a result of the d-orbitals becoming half filled or completely filled with electrons.

Cr: [ Ar ] 4s13d5

Cu: [ Ar ] 4s13d10

IONS FORMATION:

One of the characteristics of the transition metals that is most different from the Representative Elements is that the transition metals have a tendency to form more than one ion. They do so because ___________________ can be lost from both the s orbitals and the d orbitals. However, each of these "multiple" charges are NOT the result of special stability.

Example:

Fe+2: [Ar] 4s1 3d5 (1 electron is lost from the 4s orbital and one electron is lost from the 3d orbital)

Fe+3: [Ar] 4s0 3d5 (2 electrons are lost from the 4s orbital and one electron is lost from the 3d orbital)


HOMEWORK PROBLEMS:

2a. For each of the following elements, name the chemical group or family. Then determine the number of valence electrons for each element:
i) carbon
ii) S
iii) nickel
iv) Xe.
v) lead
vi) Ba

2b. Use the position on the periodic table to determine the number of valence electrons in the elements:
i) Be
ii) Ga
iii) As
iv) F

2c. Use the position on the periodic table to determine how many electrons are in the 3d orbitals for nickel.

2d. Write the neutral electron configurations (both the long form and the short-hand notation) for:
i) Ca
ii) P
iii) B
iv) Cr

2e. Write the correct electron configurations for the element copper when it is a +1 charge.

2f. The electron configuration for silver is [Kr]5s14d10. What is the electron configuration of Ag+1 (remember "orbital stability).

Answers:


PERIODIC TRENDS:

So far you have seen that the periodic table was organized to illustrate the repetitive nature of the properties of the elements. Knowing the position of an element in the periodic table allows you to make predictions about its chemical behavior. Some additional properties of elements that can be related to the element's position the periodic table - Such patterns are referred to as Periodic Trends. An awareness of certain periodic trends is necessary for an understanding of chemical bonding.

ATOMIC AND IONIC RADII TRENDS:

Atomic Radius is 1/2 the distance from the canter of the two nuclei in a Solid Crystal.

Size (Radius) ___________________ : as you go across a Period (From Left (Group 1A) to Right (Group 7A)).

More of a ___________________ Charge (Protons) causes the force of attraction on the electrons to ___________________ .

Size (Radius) Increases: From Top to Bottom of the Periodic Table.

[ Diameter scale is equal to 1.0 picometer (1x10-12m) ]

The cause for increase down the table is due to the valence electrons in different energy levels (Higher Energy Levels and More Orbitals).

Also "___________________" accounts for the larger size. The core electrons shield the valence electrons from the full charge from the nucleus. Also the addition of more energy levels causes the atom's diameter to increase.

___________________ is associated with shielding. For example, Calcium has a total of 20 protons, and from the trends we have already learned, calcium wants to lose only 2 electrons. The remaining CORE electrons "shield" the twenty "positive charges" that the protons want to "emit" and 18 of those 20 positive charges are neutralized. Therefore, Calcium will exhibit a +2 charge.


ION SIZE vs. ATOMIC SIZE:

Positive Ion: Is generally ___________________ in size than its Neutral atom Radius (loses electrons).

Negative Ion: Is generally ___________________ in size than its Neutral atom Radius (gains electrons).


IONIZATION ENERGY:

Another important property that shows periodic trends is ionization energy. Recall, that Ionization Energy is the energy needed to remove a mole of electrons from a mole of neutral gaseous atoms. This process can be represented as:

Element (g) + Ionization Energy --> Ion+ (g) + e-

First Ionization Energy removes the most ___________________ held electron ( +1 charge )

Low Ionization Energy forms positive ions more ___________________ than an element with a High Ionization Energy.

High Ionization Energy = Negative Ions

Low Ionization Energy = Positive Ions

Ionization Energy across a Period ( Left to Right ) ___________________ .

Ionization Energy Down a Family (Group) (Top to Bottom) Decreases.

DEVIATIONS IN THE TREND:

If you closely examine the graphic (above) you will find two decreases in the generally increasing ionization energies going across the Period.

The ___________________ of an electron is easier from a p orbital than from an s orbital. Recall that p orbitals are farther away from the nucleus than the s orbitals (within the same energy level).

What accounts for the Second Drop (example: from Nitrogen to Oxygen)? Once the p orbitals have one electron in each, it is relatively easy to remove the electron from that p orbital that has the two electrons because of ___________________ within an orbital.


IONIZATION ENERGY WITHIN A GROUP:

As you move Down a Group, the Ionization Energy Decreases (Why?)

- An ___________________ in the Atomic Radius (causes a Decrease in Ionization Energy)

- More ___________________, Shield the Outer Electrons from the charge associated with the nucleus. This shielding interferes with the protons' ability to pull on the valence electrons thus causing the atoms with many core electrons to have a larger atomic radius.


SUCCESSIVE IONIZATION ENERGIES:

Ionization Energies, IE, increase as you remove electrons from an atom.

[ The Energies Above are in kJ ]

Why is Sodium's (Na) Second Ionization Energy Nine Times greater than its First Ionization Energy?

IE1 ___________________ an electron from the 3s1 location (easy)

IE2 ___________________ an electron from the 2p6 location (difficult)

Therefore, it is easier to Remove Electrons with a greater Principal Quantum Number, but the Core Electrons are More difficult to remove.


SUMMARY OF TRENDS IN THE PERIODIC TABLE:


HOMEWORK PROBLEMS:

3a. Explain why the size of atoms increases as one moves down a group on the periodic table and decreases as one move across a period.

3b. Why is a negatively charged ion larger than its corresponding neutral atom?

3c. Predict the charge on the most common ion and predict whether each ion would be smaller or larger than its neutral atom:
i) Mg
ii) Cl
iii) Al
iv) S
v) Cs
vi) I
vii) O

3d. Describe ionization energy?

3e. Explain why the second ionization energy of barium is relatively "low" while the third ionization energy is very "high".

Answers:


CLICK HERE TO VIEW THE SHORT-HAND ORBITAL SEQUENCING FOR ALL "KNOWN" ELEMENTS