IONIC SIZE AND ELECTRONEGATIVITY

Electronegativity  

-   chemical property that describes the ability of an atom (or, more rarely, a functional group) to attract electrons (or electron density) towards itself

-   a measure of the tendency of an atom to attract a bonding pair of electrons.

-   affected by both its atomic number and the distance that its valence electrons reside from the charged nucleus. The higher the associated electronegativity number, the more an element or compound attracts electrons towards it 

-    increases on passing from left to right along a period, and decreases on descending a group.

 

Ionic Size

-           Ionic size is the size of the atom after it loses its outermost electron (after the first Ionization Energy). 

-           When atoms gain or lose electrons, the atom becomes an ion. When an atom gains an electron, it becomes a negatively charged ion that we call an anion. Anions are larger in size than their parent atoms because they have one or more additional electrons, but without an additional proton in the nucleus to help moderate the size 

-           When an atom loses an electron, it becomes a positively charged ion called a cation. Cations are smaller than their parent atoms because they have lost electrons (sometimes the entire outermost energy level) and the electrons that remain behind simply don't take up as much room.

-           Also note that when comparing cations and anions, the anions are larger.

 By
Michelle Cruz

Nina De los Reyes

Nicolle Groves

Lalique Lorenzo

Marga Villarosa

II-9

 

Metals, Non-Metals and Metalloids

METALS

• is a chemical element that is a good conductor of both electricity and heat and forms cations and ionic bonds with non-metals. 
• is an element,compound, or alloy characterized by high electrical conductivity. In a metal, atoms readily lose electrons to form positive ions (cations). Those ions are surrounded by delocalized electrons, which are responsible for the conductivity. The solid thus produced is held by electrostatic interactions between the ions and the electron cloud, which are called metallic bonds.
• Most elements are metals. 
• are elements that have atoms arranged in rows. The electrons are easily released from metal atoms so that layers of metal atoms exist in a 'sea' of electrons.

Location on the Periodic Table: 

• Metals are located on the left side and the middle of the periodic table. Group IA and IIA (the alkali metals) are the most active metals. The transition elements, groups IB to VII B, are also considered metals.

Physical Properties of Metals:

•  include shiny lustre, greyish - silver colour, hardness, good heat and electricity conductivity, high melting and boiling points, malleability (can be hammered into a sheet) and ductility (can be pulled into a wire).
• Some exceptions to these are the metals - sodium and calcium (very soft), gold and copper (yellowish colour), and mercury (low melting and boiling points).

Chemical Properties of Metals

• Some metals are more reactive than others. This is because very reactive metals lose electrons easily. Metals such as sodium are very reactive and are explosive in air. Metals such as gold are very unreactive, and therefore do not corrode or tarnish in air.

NON-METALS

• are located on the upper side of the periodic table. Non-metals are separated from metals by a line that cuts diagonally through the region of the periodic table containing elements with partially filled p orbitals. Technically, the halogens and noble gases are non-metals, but the non-metal element group usually is considered consist of hydrogen, carbon, nitrogen, oxygen, phosphorus, sulfur, and selenium. 

Properties

• Non-metals have high ionization energies and electronegativites. They are generally poor conductors of heat and electricity. Solid non-metals are generally brittle, with little or no metallic luster. Most non-metals have the ability to gain electrons easily. Non-metals display  a wide range of chemical properties and reactivities. 

METALLOIDS

Metalloids are the elements found along the stair-step line that distinguishes metals from non-metals. This line is drawn from between Boron and Aluminum to the border between Polonium and Astatine. The only exception to this is Aluminum, which is classified under "Other Metals". Metalloids have properties of both metals and non-metals. Some of the metalloids, such as silicon and germanium, are semi-conductors. This means that they can carry an electrical charge under special conditions.  Physical Properties Metalloids can be shiny like metals or dull like non-metals. They are ductile in nature and can be drawn in shapes of pipes. They are conductors of heat and electricity, but not as good as metals. Metalloids like boron, germanium, arsenic are used as dopants in glasses in semiconductor chips. Metalloids are usually brittle in nature and behave as electrical insulators at room temperature.  Chemical Properties Metalloids tend to have an intermediate property between metals and non-metals. They may look like metals, in case of arsenic and antimony that are crystalline solids. However, in chemical reactions, they may behave either as metals or non-metals. The metalloids are usually amphoteric oxides as metals are basically basic oxides and non-metals are generally acidic oxides. Some metalloids like boron, silicon and germanium behave as semiconductors. Their chemical reactivity depends on the substance they react with. Like boron acts as a metal when reacting with fluorine and behaves as non-metals when reacting with sodium. Many metalloids have different allotropes. For a given metalloid, one of its allotrope may react as a metal and the other allotrope may behave as a non-metal.

Members:

Romina Espuerta

Tina Silvestre

Maegan Zablan

Pamela Lopez

Jam Villanueva

Modern Periodic Table

During the nineteenth century, chemists though it logical to arrage the known elements in a table according to their atomic masses. Which such an arrangement, several scientists observed some trends in the chemical and physical properties of successive elements

1829

Green Chemst Johann W. Döbereiner showed that the atomic mass of the element strontium (Sr) lies midway between the atomic masses of calcium (Ca) and barium (Ba), forming what he called a triad. Some years later, he shows that other such triads existed such as halogens chlorine (Cl), bromine (Br), and iodine (I), and the alkali metals lithium (Li), sodium (Na), and potassium (K).

In each triad, the middle element had an atomic mass almost equal to the average atomic mass of the other two elements. However, there was no explainable to be offered to account for this observation at the time.

Other scientists expanded on Döbereiner's suggestions by showing that similar relationships among elements extended further than the triads. Florine (F) was added to the halogens, and magnesium (Mg) to the alkaline earth metals. On the other hand, the elements oxygen (O), sulfur (S), selenium (Se), and tellurium (Te) were classified as a family of elements.

1864

English chemist John Newlands observed that when elements are arranged in increasing atomic masses, there appeared to be a repetition of similar properties for every eight element. Newlands referred to this arrangement as the law of octaves. However, the law of octaves was found to inapplicable for elements after calcium and noble gases were unknown at this time. 

Newland's Law of Octaves

German chemist Julius Lothar Meyer devised a classification of elements into a table that accounted for the periodic variations in properties. His table included 56 elements. 

Russian chemist Dmitri Mendeleev observed that the elements were arranged in increasing atomic masses with similar physic and chemical properties were periodically prepared. By arranging the elements so that those with similar properties were in the same column, he constructed the periodic table. Mendeleev noted that in order for the elements with similar properties to fall under the same column, there had to be gaps in the middle. This led him to believe that the gaps represented new elements to be discovered in the future. Further more, Mendeleev predicted he properties of these yet to be discovered elements, as well as the properties of their compounds. 

Mendeleev VS Meyer

• Both Mendeleev and Meyer arranged elements in increasing atomic mass.
• Both left vacant spaces where unknown elements should fit.
• Mendeleev is called the "Father of the Modern Periodic Table" because he corrected the atomic masses of beryllium (Be), indium (In), and uranium (U).
• Mendeleev's table was based on the original periodic law, which stated that the physical and chemical properties of elements are period functions of their atomic masses. So if the elements were arranged in increasing atomic masses, those elements with similar properties would appear at regular intervals. His table contained 66 known elements. 

Problems arose later when more accurate masses were determined. The periodic law and its classificaltion to increased atomic masses cause several elements to be incorrectly placed in the table, which led to scientists thinking that another propery, aside from atomic mass, must be the basis for the similarity of the properrites of the elements that are close together in the periodic table. 

1913

English physicist Henry Moseley observed that the frequencies of X-rays emitted from atoms of elements were correlated with the sizes of their nuclear charges. Moseley called this nuclear charge of an atom atomic number (Z). He published the results of his work on 39 elements, arranging them in increasing atomic numbers. His work was perhaps the most fundamental single step in the development of the periodic table because e was able to solve Mendeleev's irregularities in his periodic table. 

1941

Glenn Seaborg co-discovered 10 elements with Edwin McMillan. He moved 14 elements out of the main body to their current locations. He was the only person to have an element named after him while he was still alive (Seaborgium, Sg)

The Modern Periodic Table

The modern version of the periodic table is arranged in increasing atomic number oder. Each square in the periodic table shows the symbol of each element, its atomic number, atomic mass, and other additional information. 

As of 2002, the modern periodic table has 114 elements. By 1940, all the natural-occuring elements, which are about 90, had already been discovered. More than 20 elements are artificially made or are naturally-occurring, but is short lived. 

The modern periodic law states that the chemical and physical properties of elements are periodic functions of their atomic numbers. Here, periodic function refers to the repetition of the properties of elements a regular intervals. 

Periodic Table Geography

Features of the Periodic Table

• Most useful to a chemist
• Organizes information about elements
• Elements are arranged in rows by atomic number

Groups and Periods 

• Horizontal rows are Periods. 
• Vertical columns are Groups or Families. Elements have similar chemical and physical properties and have the same valence shells (Found on the last energy level). 

Classifications 

• Metals
• Metalloids
• Noble gases
• Non-metals

The Modern Periodic Table of Elements

Group I

Michaela Panaguiton

Nadine Leano

Ella Canuel

Jenny Gonzales

Published by Michaela Panaguiton

Groups, Families, Periods and Valence of the Periodic Table

What are Groups? 

• A group is a vertical column in the Periodic Table of Elements. They are considered the most important method in classifying the different elements. 
• Groups are the elements having the same outer electron arrangement. The outer electrons are also called the valence electrons.

• Since they have the same number of valence electrons, the elements in a group shares the same chemical properties.
• The Roman numerals listed ABOVE each group are the usual number of valence electrons. 
• These are the groups: IA, IIA, IIIA, IVA, VA, VIA, VIIA, VIIIA

What are Families?

• Groups are also known as Families. The families are the names representing each group. 
• There are 9 different families, and these are the:
• Alkali Metals
• Alkaline Earth Metals
• Transition Metals
• Boron 
• Carbon
• Nitrogen
• Chalcogens
• Halogens
• Noble Gases

What are Periods? 

• Periods are the horizontal rows of elements found in the Periodic Table of Elements. There are 7 periods in the Periodic Table.
• The periods represents the energy level of an atom.

• The number of elements in a period increases as you move down because there are more sub levels per level as the energy level of the atom increases.
• According to the table below, not all the periods have the same number of elements.    

Period Number	Number of Elements
1	2
2	8
3	8
4	18
5	18
6	32
7	28

What are Valence?

• Valence is also known as valency or valency number, is a measure of the number of chemical bonds formed by the atoms of a given element.
• Valence is the number of electrons needed to fill the outermost shell of an atom. 

• For example, when you go across the table from carbon to nitrogen to oxygen, the number of valence electrons increases from 4 to 5 to 6. As we go from fluorine to neon to sodium, the number of valence electrons increases from 7 to 8 and then drops down to 1 when we start the new period.

How to get the element using the periods and groups!

• You can find an element even without the given atomic number just by using the groups and periods. Example: Look for the element that is located on Period 2, Group 4. The element is Carbon because its on the row of period 2 and it is under Group 4.

Try this! Find the element given the period and group number only.

• Period 4, Group 6
• Period 2, Group 2
• Period 5, under the Alkali Metals
• Halogens family, along period 5. 

________________________________________________________________

Made by: Group 3

Marj Mendoza

Roxy Trillanes

Isabella Meily

Eula Manibog

Djoseth Macomb

Ionization Energy and Electron Affinity

IONIZATION ENERGY:

• A certain amount of energy needed to knock off the electrons from a neutral gaseous atom to form a positive ion (cation)
• Related to Atomic Size. "The bigger the atom, the farther is the outermost electron from the nucleus." 

Note: Ionization energy decreases with increasing atomic size :D 

• Atoms going across a period from left to right generally increases. The rise is due to the increasing nuclear charge (Atomic No.), although the electronsbelong to the same principal energy level.
• There are certain exceptions to this rule. The elements in Group IIIA have lower first ionization energies than those in Group IIA. The same is true for elements in Group VA and VIA.

EXAMPLES: 

Going down a column, ionization energy decreases. The ionization energy of halogens are given below: 

Fluoride (F) ----------> 1879 kJ/mole

Chloride (Cl) ----------> 1254 kJ/mole

Bromine (Br) ----------> 1139 kJ/mole

Iodine (I) ----------> 1031 kJ/mole

Astatine (At) ----------> 917 kJ/mole 

*If we would check the periodic table, we would find these in a column from up to down. This shows that from the top to the bottom, ionization energy decreases.

Going across a period, ionization energy increases. The increase is not regular but a general trend can be seen. This can be observed in the elements in Period 3 :)

Sodium (Na) ----------> 492 kJ/mole
Magnesium (Mg) ----------> 733 kJ mole
Aluminum (Al) ----------> 577 kJ/mole
Silicon (Si) ----------> 782 kJ/mole
Phosphorus (P) ----------> 1013 kJ/mole
Sulfur (S) ----------> 1004 kJ/ mole
Chlorine (Cl) ----------> 1254 kJ/mole
Argon (Ar) ----------> 1525 kJ/mole

* In period 3, it does not completely follow the rule on ionization energy. If you'd observe you could see that the I.E. increases, then decreases then increases again and decreases from left to right.
* The ionization energy decreases from one atom to the next as one goes down a group. This energy decrease is due to the fact that the valence electrons at a high energy level or a farther distance from the nucleus, which means the attraction of the electrons to the nucleus becomes smaller. It is easier to remove the electrons. 

Practice:

In each set, arrange the elements from highest ionization energy to the lowest.

A. He, Ne, Ar, Kr

B.Mg, As, Al, F
C. K, Sc, Ge
D. Zn, Cr, K, Br
E. Ti, Al,In Ga

ELECTRON AFFINITY:

•   The amount of energy release when a gaseous atom gains an electron.
• The more positive an atom's electron affinity, the greater its tendency to accept an electron to form anions. 

* A negative electoron affinity means that energy is released when an electron is added to the atom; if energy is required, the electon affinity is positive. The more negative the E.A , the greater is the tendency of the atom to attract an electron.

* Halogens have the most negative electron affinities.

ELECTRON AFFINITY TRENDS:

*The trend of the E.A is not as regular as for ionic size and ionization energy.
*Electron Affinity becomes more negative from left to right and less negative from top to bottom. This translates to stronger affinity for electrons increasing from left to right and decreasing from top to bottom. 

Note: 
 * E.A of Nobles gases have not yet been measured. They no longer show tendencies to attract electrons because they are already what we call the IDEAL Gases.

* Ionization Energy and Electron Affinity  are very important properties that determine how atoms interact and bond with each other.
* Electron Affinity is related to the formation of anions while Ionization Energy is related to the formations of cations.

Remember: Electrons with high affinities for electrons are found on the right side of the periodic table. These are the non metals. Non metals have a strong tendency to form anions
* Halogens are an exception.

Submitted By Group 5:

Mia San Juan

Alex Vergara

Jhoanne Sanchez

Elma Tejada

Cierra Mortega