Covalent : Polar & Non-Polar

Figure I. Pauling Electronegativity Values 

American chemist Linus Pauling (1901 - 1994) devised a scale of relative electronegativity values by assigning fluorine, the most electronegative element, a value of 4. 0. The figure above displays the electronegativity values known to date.

Chlorine (3. 0) is more electronegative than hydrogen (2. 1). In a hydrogen chrloride molecule, the chlorine atom holds the shared electrons more tightly. This results in the chlorine end of the molecule being more megative than the hydrogen end.  When electrons are not equaly shared, the bond is assumed to be polar. Therefore, the bond in a hydrogen chloride molecule is described as a polar covalent bond, whereas the bond in a hydrogen molecule or a chlorine molecule is a non-polar covalent bond.

A polar covalent bond is not an ionic bond. In an ionic bond, an atom completely loses an electron. In a polar covalent bond, the atom at the positive end of the bond (hydrogen in HCl) still has some share in the bonding pair of electrons, as seen in figure II. To distinguish this arrangement from an ionic bond, the following notation is used:

Figure II. Representation of the Polar Hydrogen Chloride Molecule

The line between the atoms represents the covalen bond, a pair of shared electrons. The δ+ and δ- (read as "delta plus" and "delta minus") signify which end is partially positive and which is patially negative, wherein the word "partially" is used to distinguish this charge from the full charge on an ion. 

Using the electronegativity values for atoms in a compound, the type of bonding may be predicted. When the electronegativity difference is zero or is very small (<0. 5), the bond is non-polar covalent, with essentially equal sharing of electrons. When the electronegativity difference is large (>2. 0), complete electron transfer occurs and an ionic bond is formed, as in the case of sodium and chlorine. When the electron negativity difference is between 0. 5 and 2. 0, polar covalent bonds are formed. 

PRACTICE EXERCISES

Classify the type of bond seen between each pairs of atoms as non-polar, polar, or ionic.

1. Two hydrogen atoms
2. Oxygen and hydrogen 
3. Carbon and hydrogen
4. Nitrogen and three hydrogen atoms
5. Sodium and chlorine 

SOLUTIONS

1. Two H atoms have the same electronegativity; the electronegativity difference is 0; the bond is non-polar covalent. 
2. The electronegativity difference of 3. 5 - 2. 1 is 1. 4; the bond is polar covalent. 
3. The electronegativity difference of 2. 5 - 2. 1 is 0. 4; the bond is non-polar covalent. 
4. The electronegativity of N is 3. 0, that of H is 2. 1, so all three N--H bonds are polar covalent. 
5. The electronegativity difference of 3. 0 - 0.9 is 2. 1; the bond is ionic. 

Group I

Michaela Panaguiton

Nadine Leano

Ella Canuel

Jenny Gonzales

Published by Michaela Panaguiton

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. 

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Made by: Group 3

Marj Mendoza

Roxy Trillanes

Isabella Meily

Eula Manibog

Djoseth Macomb