The periodic table of elements has a total of entries. Elements are arranged in a series of rows periods in order of atomic number so that those with similar properties appear in vertical columns.
Elements in the same period have the same number of electron shells; moving across a period so progressing from group to group , elements gain electrons and protons and become less metallic. This arrangement reflects the periodic recurrence of similar properties as the atomic number increases.
For example, the alkali metals lie in one group Group 1 and share similar properties, such as high reactivity and the tendency to lose one electron to arrive at a noble-gas electron configuration. Modern quantum mechanics explains these periodic trends in properties in terms of electron shells. The filling of each shell corresponds to a row in the table.
In the s-block and p-block of the periodic table, elements within the same period generally do not exhibit trends and similarities in properties vertical trends down groups are more significant.
However, in the d-block, trends across periods become significant, and the f-block elements show a high degree of similarity across periods particularly the lanthanides. If we examine the physical state of each element, we notice that on the left side of the table, elements such as lithium and beryllium are metallic solids, whereas on the right, nitrogen, oxygen, fluorine, and neon are all gases. This is because lithium and beryllium form metallic solids, whereas the elements to the right form covalent compounds with little intermolecular force holding them together.
Therefore we can say that, in general, elements tend to go from solids to liquids to gases as we move across a given period. However, this is not a strict trend. As you move across a period in the periodic table, the types of commonly encountered bonding interactions change. For example, at the beginning of Period 2, elements such as lithium and beryllium form only ionic bonds, in general.
Moving across the period, elements such as boron, carbon, nitrogen and oxygen tend to form covalent bonds. Fluorine can form ionic bonds with some elements, such as carbon and boron, and neon does not tend to form any bonds at all. Another physical property that varies across a period is the melting point of the corresponding halide. A halide is a binary compound, of which one part is a halogen atom and the other part is an element or radical that is less electronegative or more electropositive than the halogen, to make a fluoride, chloride, bromide, iodide, or astatide compound.
Many salts are halides; the hal- syllable in halide and halite reflects this correlation. All Group 1 metals form halides that are white solids at room temperature. The melting point is correlated to the strength of intermolecular bonds within the element. First, we must analyze compounds formed from elements from Groups 1 and 2 e. To develop an understanding of bonding in these compounds, we focus on the halides of these elements. The physical properties of the chlorides of elements in Groups 1 and 2 are very different compared to the chlorides of the elements in Groups 4, 5, and 6.
All of the alkali halides and alkaline earth halides are solids at room temperature and have melting points in the hundreds of degrees centigrade.
The non-metal halide liquids are also electrical insulators and do not conduct electrical current. In contrast, when an alkali halide or alkaline earth halide melts, the resulting liquid is an excellent electrical conductor. This tells us that these molten compounds consist of ions, whereas the non-metal halides do not. This again demonstrates the type of bonding that these compounds exhibit: the left-most elements form more ionic bonds, and the further-right elements tend to form more covalent bonds.
The physical properties notably, melting and boiling points of the elements in a given group vary as you move down the table. In chemistry, a group is a vertical column in the periodic table of the chemical elements. There are 18 groups in the standard periodic table, including the d-block elements but excluding the f-block elements.
A physical property of a pure substance can be defined as anything that can be observed without the identity of the substance changing. The observations usually consist of some type of numerical measurement, although sometimes there is a more qualitative non-numerical description of the property.
Physical properties include such things as:. Within a group of the periodic table, each element has the same valence electron configuration. For example, lithium, sodium, potassium, rubidium, cesium, and francium all have a single electron in an s orbital, whereas every element in the group including fluorine has the valence electron configuration ns 2 np 5 , where n is the period.
This means the elements of a group often exhibit similar chemical reactivity, and there may be similarities in physical properties as well. Before a discussion of the melting points of various elements, it should be noted that some elements exist in different forms.
For example, pure carbon can exist as diamond, which has a very high melting point, or as graphite, whose melting point is still high but much lower than that of diamond. Different groups exhibit different trends in boiling and melting points.
For Groups 1 and 2, the boiling and melting points decrease as you move down the group. For the transition metals, boiling and melting points mostly increase as you move down the group, but they decrease for the zinc family. In the main group elements, the boron and carbon families Groups 13 and 14 decrease in their boiling and melting points as you move down the group, whereas the nitrogen, oxygen, and fluorine families Groups 15, 16, and 17 tend to increase in both.
The noble gases Group 18 decrease in their boiling and melting points down the group. These phenomena can be understood in relation to the types of forces holding the elements together. For metallic species, the metallic bonding interaction electron-sharing becomes more difficult as the elements get larger toward the bottom of the table , causing the forces holding them together to become weaker.
As you move right along the table, however, polarizability and van der Waals interactions predominate, and as larger atoms are more polarizable, they tend to exhibit stronger intermolecular forces and therefore higher melting and boiling points. Metallic elements are shiny, usually gray or silver in color, and conductive of heat and electricity. They are malleable can be hammered into thin sheets and ductile can be stretched into wires.
Some metals, such as sodium, are soft and can be cut with a knife. Others, such as iron, are very hard.
Non-metallic atoms are dull and are poor conductors. They are brittle when solid, and many are gases at STP standard temperature and pressure. Metals give away their valence electrons when bonding, whereas non-metals tend to take electrons. A metal and a non-Metal : On the left is sodium, a very metallic element ductile, malleable, conducts electricity.
On the right is sulfur, a very non-metallic element. Metallic character increases from right to left and from top to bottom on the table. Non-metallic character follows the opposite pattern. Advertisement Remove all ads. Short Note. Answer the following question. How does ionization energy vary down the group and across a period?
Solution Show Solution Variation of ionization energy down the group: On moving down the group, the ionization enthalpy decreases. This is because the electron is to be removed from the larger valence shell. The ionization energy of an element increases as one moves across a period in the periodic table because the electrons are held tighter by the higher effective nuclear charge.
This is because additional electrons in the same shell do not substantially contribute to shielding each other from the nucleus, however an increase in atomic number corresponds to an increase in the number of protons in the nucleus. The ionization energy of the elements increases as one moves up a given group because the electrons are held in lower-energy orbitals, closer to the nucleus and thus more tightly bound harder to remove.
Based on these two principles, the easiest element to ionize is francium and the hardest to ionize is helium. Boundless vets and curates high-quality, openly licensed content from around the Internet. This particular resource used the following sources:. PNG Wikipedia Public domain.
Skip to main content. Explanation: Two properties are important in determining ionization energies: i nuclear charge; and ii shielding by other electrons.
Related questions How does ionization energy relate to reactivity? What is ionization energy measured in? What are the first and second ionization energies? How does ionization energy increase? How does ionization energy change down a group?
How do trends in atomic radius relate to ionization energy?
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