The periodic table is more than just a chart of elements. It can actually help you predict the properties of an element just by knowing where it is located on the table. This is because of something called periodicity, which means that certain trends in element properties repeat in regular intervals as you move along the table. In this article, we'll explore periodic trends in inorganic chemistry and explain what periodicity is. We'll also take a look at periodic trends as you move along a period in the periodic table and as you move down a group. By the end of this article, you'll be able to describe and explain trends in things like electron configuration, atomic radius, and first ionisation energy. We'll even cover density and electrical conductivity. So, if you want to know more about periodic trends, keep reading!
The periodic table is a fascinating creation that we still use today. It was first developed in 1869 by a Russian chemist named Dmitri Mendeleev, who built on the work of other scientists like John Newlands. Mendeleev initially ordered the elements by atomic mass, but he soon noticed that every eight elements or so shared certain properties. To reflect this, he arranged the elements in rows and columns, with similar elements placed above and below each other. However, there were some cases where the properties didn't match up perfectly, so Mendeleev left gaps in the table for undiscovered elements. His predictions were later proven correct when these elements were discovered. The periodic table remains an important tool for chemists and scientists alike to this day.
The periodic table is a systematic arrangement of elements based on their atomic number. It consists of rows, called periods, and columns, called groups, and is designed to show periodic trends. These trends are patterns that repeat themselves as you move across a period or down a group. Another word you might hear is periodicity, which refers to the repetition of properties after a certain interval. In the case of the elements, periodicity means the repetition of properties after a certain atomic number. Mendeleev was the first to notice periodicity in the table and arranged the elements accordingly. In this article, we'll take a closer look at some of the trends he observed and explain why they occur. We'll also delve into the history and structure of the periodic table. You can find more information on this topic in our article on the periodic table.
As we saw above, rows in the periodic table are known as periods. Columns in the periodic table are known as groups. All in all, there are seven periods and eighteen groups in the periodic table. This might seem like a lot to learn about, but as you'll see, they show regular periodic trends when it comes to certain properties. The properties we'll look at today include:
Electron configuration
Atomic radiusElectronegativity
First ionisation energy
Melting and boiling points
DensityElectrical conductivity
First up - electron configuration.
The period and group of an element on the periodic table tell you a lot about its electron configuration. The period of an element tells you how many electron shells it has - for example, elements in period 2 all have two electron shells. The group of an element tells you how many electrons it has in its outer shell. For example, all the elements in group 3 have three electrons in their outer shell. As you move across a period, the number of electrons increases by one each time. All elements in the same period have the same number of electron shells. The atomic number of an element is equal to its number of protons, and is also equal to its number of electrons in a neutral atom. This is shown in the periodic table by atomic number, which increases as you move along the period.
The number of electron shells increases as you go down a group. However, elements in the same group have exactly the same number of electrons in their outer shell - the only difference is the number of inner shells. For example, all the elements in group 1 have just one electron in their outer shell. But whilst hydrogen has only one electron shell in total, the next member lithium has two, and sodium has three.
Atomic radius refers to the distance between the center of an atom's nucleus and its outermost electron shell. This distance is measured in picometers (pm). The atomic radius of an element is influenced by factors such as its nuclear charge and the number of electron shells it has. When we look at atomic radius across a period, we observe that it decreases as we move along the period. To understand why this happens, we need to examine the atomic structure of the elements within a period.
The modern periodic table arranges elements by their atomic number, which is the number of protons present in the nucleus of an atom. As a result, each element has the same number of protons as it does electrons. These electrons are found orbiting the nucleus in electron shells. Although elements in the same period have different numbers of electrons, they have the same number of electron shells. For example, carbon has six electrons, while oxygen has eight electrons, yet both elements have only two electron shells. Thus, as we move across a period, the nuclear charge increases while the number of electron shells remains constant, causing the atomic radius to decrease.
Electron shells are attracted to the nucleus thanks to a strong electrostatic attraction between the positively charged nucleus and the negatively charged shells. This determines atomic radius. As you go across a period, the atomic number increases - each element has one more proton and one more electron than the element before it. It means that the nucleus' charge increases. The outermost electron shell experiences a stronger attraction to the positively charged nucleus, so the negative electrons are pulled in closer to the nucleus in the centre of the atom. This decreases atomic radius.
When we look at atomic radius down a group in the periodic table, we observe that it increases. This may seem counterintuitive since atomic number increases down a group, suggesting that nuclear charge also increases. However, the number of electron shells also increases down a group, causing the outermost electrons to be located further away from the nucleus. This increase in distance between the nucleus and the outermost electron shell results in an increase in atomic radius. For example, consider the elements lithium and sodium, which are both found in group 1 of the periodic table. Both of these elements have one electron in their outermost shell. However, sodium has more electron shells than lithium, and so its outermost electron is located further from the nucleus. This results in sodium having a larger atomic radius than lithium. Overall, the trend in atomic radius down a group is due to the increase in the number of electron shells, which causes the outermost electrons to be located further from the nucleus, resulting in a larger atomic radius.
Periodic trends in Electronegativity
Now let's look at trends in electronegativity.
Electronegativity is an atom's ability to attract a shared pair of electrons.
Shared pairs of electrons are always found in the outer shell of an atom. Electronegativity is all to do with the strength of the attraction between these bonded electrons and the atom's nucleus. Electronegativity depends on factors such as nuclear charge, number of electron shells, and shielding by inner electrons.
Electronegativity refers to an element's ability to attract electrons in a chemical bond towards itself. When we look at electronegativity across a period in the periodic table, we observe that it increases. This trend can be explained by the combined effects of nuclear charge, atomic radius, and electron shielding.
As we move along a period, nuclear charge increases, causing the bonded electron pair to be pulled closer to the nucleus. This results in a decrease in atomic radius. However, the effect of electron shielding must also be taken into account. Elements in the same period have the same number of electron shells, which means they have the same levels of shielding. If nuclear charge increases across a period but shielding remains constant, the overall nuclear charge felt by the bonded pair of electrons increases. This leads to an increase in electronegativity. For example, consider the elements carbon and oxygen. Carbon has six protons in its nucleus, with the charge of two of these protons being shielded by two inner shell electrons. Oxygen has eight protons in its nucleus, with the charge of two protons being shielded by two inner shell electrons as well. The outer shell electrons in carbon experience an overall effective nuclear charge of +4, while those in oxygen experience an overall effective nuclear charge of +6. Additionally, oxygen has a smaller atomic radius than carbon. As a result, oxygen has a higher electronegativity than carbon. In summary, electronegativity increases across a period due to the combined effects of increasing nuclear charge and decreasing atomic radius, which result in a higher overall effective nuclear charge felt by the bonded pair of electrons.
Electronegativity decreases down a group. Although nuclear charge increases, the number of inner electron shells also increases, and these inner electron shells shield the charge of the additional protons in the nucleus. As a result, the overall nuclear charge felt by the bonded pair of electrons remains the same. But we also know that atomic radius increases as you move down a group, meaning the bonded pair of electrons are further away from the nucleus. Thus, electronegativity decreases.
For example, earlier, we looked at lithium and sodium. Lithium has three protons in its nucleus, but the charge of two of these protons is shielded by two inner shell electrons. Sodium has eleven protons in its nucleus, but the charge of ten of these protons is shielded by ten inner shell electrons. Outer shell electrons in both elements experience an overall nuclear charge of +1. However, sodium has a larger atomic radius than lithium. Therefore, it has a lower electronegativity.
Check out Electronegativity for a further explanation.
Ionisation energy refers to the energy required to remove the outermost electron from one mole of gaseous atoms. This energy is affected by several factors, including nuclear charge, atomic radius, and electron shielding.
When we look at ionisation energy across a period in the periodic table, we observe that it increases. This is because nuclear charge increases across a period, causing the outermost electron to be more strongly attracted to the nucleus. Additionally, atomic radius decreases across a period, which means that the outermost electron is closer to the nucleus and therefore harder to remove. The number of inner electron shells remains the same across a period, which means that there is no increase in electron shielding to counteract the effect of increasing nuclear charge.
On the other hand, when we look at ionisation energy down a group in the periodic table, we observe that it decreases. Although nuclear charge increases down a group, the number of inner electron shells also increases. These inner electrons shield the outermost electrons from the increasing nuclear charge, making it easier to remove them. Furthermore, atomic radius increases down a group, meaning that the outermost electrons are further from the nucleus and therefore easier to remove. It is also worth noting that there are occasional anomalies in the trend of ionisation energy across a period. For example, boron and oxygen have lower ionisation energies than expected. These anomalies can be explained by the electronic configurations of these elements and their relationship to the effective nuclear charge experienced by their outermost electrons. Overall, the trend in ionisation energy can be explained by the combined effects of nuclear charge, atomic radius, and electron shielding.
Melting and boiling points are physical properties that depend on the strength of the bonds between atoms or molecules in a substance. Unlike other periodic trends, such as atomic radius or electronegativity, melting and boiling points do not exhibit a clear periodic trend as we move across a period.
The melting and boiling points of elements across a period can vary greatly due to the differences in their bonding and structure. For example, the melting and boiling points of the noble gases (group 18) are very low, as they exist as monatomic gases with only weak London dispersion forces between atoms. In contrast, the melting and boiling points of the elements in groups 1 and 2 are relatively high due to their metallic bonding, which involves strong electrostatic attractions between positively charged ions and negatively charged delocalized electrons. The elements in the middle of the periodic table (groups 13-17) have a variety of different bonding types, including covalent, ionic, and metallic. As a result, their melting and boiling points can vary significantly depending on the strength of the intermolecular forces between atoms or molecules.
Melting and boiling points down a group In general, the melting and boiling points of elements decrease down a group in the periodic table. This is because the atomic radius of the elements increases down a group, resulting in weaker intermolecular forces between atoms or molecules. Additionally, the increased number of electron shells down a group leads to greater electron shielding and a weaker attraction between the outermost electrons and the positively charged nucleus. However, there are some exceptions to this trend. For example, the melting and boiling points of the group 1 metals (except for francium) decrease down the group, but the melting and boiling points of the group 2 metals (except for beryllium) increase down the group. These trends can be explained by differences in the strength of metallic bonding and the size of the metal ions. Overall, while melting and boiling points do not show a clear periodic trend across a period, they do generally decrease down a group due to the influence of atomic size and electron shielding on intermolecular forces.
Melting and boiling points are physical properties that depend on the strength of the bonds between atoms or molecules in a substance. Unlike other periodic trends, melting and boiling points do not exhibit a clear periodic trend across a period. The melting and boiling points of elements across a period can vary greatly due to the differences in their bonding and structure. Sodium, magnesium, and aluminum have medium melting points because they bond using metallic bonding, forming giant metallic lattices held together by electrostatic attraction. To melt the metals, you need to overcome this metallic bonding. Aluminum has the highest melting point of the three as it has the strongest metallic bonding.
Silicon has a high melting point because it is a giant covalent macromolecule. All of its atoms are held together by strong covalent bonds, which require a lot of energy to be broken. Phosphorus, sulfur, and chlorine have low melting points because they are simple covalent molecules. Although there are strong covalent bonds within the molecules, the only forces between molecules are weak intermolecular forces which don't require much energy to break. Sulfur has a higher boiling point than phosphorous and chlorine because it forms larger molecules, which increases the strength of the intermolecular forces found between its molecules.
Argon has a very low melting point because it is a monatomic gas. It has extremely weak intermolecular forces between atoms which require hardly any energy to overcome. The trend for melting and boiling points down a group is not clear. For some groups, such as the halogens (group VII), melting and boiling points increase as you move down the group. But for other groups, such as the alkali metals (group I), melting and boiling points decrease as you move down the group. Density initially increases as you move along a period in the periodic table, but then drops dramatically when you hit the right-hand side of the table. This is because the elements become gases, and on average, the particles are found much further apart from each other. Density also increases as you move down a group.
Electrical conductivity shows a similar trend. Groups I-III, as well as the d-block elements, have a high electrical conductivity because they bond using metallic bonding and contain delocalized electrons which are free to move and carry a charge. In contrast, groups IV-VIII tend to have a low electrical conductivity because they bond using covalent bonding, and their outer shell electrons form part of a shared pair of electrons and are not free to carry a charge. However, some group IV elements are surprisingly good at conducting electricity. For example, carbon can form a solid called graphite, in which three of carbon's four outer shell electrons are covalently bonded to other carbon atoms, and the fourth outer shell electron is delocalized. This means that it is free to move through the substance and carry a charge. Other elements, such as silicon, are semiconductors, which means that they have properties somewhere between those of a conductor and an insulator. In summary, the periodic trends of melting and boiling points, density, and electrical conductivity depend on the bonding and structure of the elements and do not show clear periodic trends across a period. However, they do generally increase or decrease down a group due to the influence of atomic size and electron shielding on intermolecular forces.
Density initially increases as you move along a period in the periodic table, but then drops dramatically when you hit the right-hand side of the table due to the elements becoming gases. Density also increases as you move down a group. Electrical conductivity is high for groups I-III and the d-block elements due to metallic bonding, while it is low for groups IV-VIII due to covalent bonding. However, some group IV elements like carbon can conduct electricity due to the delocalized electrons in their structure. Overall, understanding periodic trends is important in predicting the properties and behavior of elements and compounds. The trends in atomic radius, electronegativity, first ionization energy, melting and boiling points, density, and electrical conductivity provide a basis for understanding the chemical reactivity and properties of different elements and compounds.
What are trends in the periodic table?
Periodic trends are repeating patterns found in the periodic table as you move across a period or down a group. Properties that show periodic trends include atomic radius, electronegativity, and first ionisation energy.
What is the periodic trend for ionisation energy?
Ionisation energy increases moving along a period and decreases going down a group.
What are the period and group trends in electronegativities?
Electronegativity increases across a period and decreases down a group.
