Periodicity in elemental properties refers to the repeating patterns observed among elements when they are organised based on increasing atomic number on the periodic table. As we move across periods and down groups, there are specific trends that emerge in terms of atomic and ionic size, energy required to remove electrons, tendency to gain electrons, and ability to attract shared electrons.
Atomic Radius
- Definition: The atomic radius is half the distance between the nuclei of two like atoms bonded together.
- Trend across a period: As we move from left to right across a period, atomic radius generally decreases. This is due to:
- Increase in nuclear charge as the number of protons in the nucleus increases.
- Increased attraction between the outer electrons and the nucleus.
- Electrons being added to the same principal energy level, so shielding remains constant.
- Trend down a group: As we go down a group, atomic radius generally increases because:
- Addition of extra electron shells.
- Increased shielding effect from inner electrons.
- Weaker attraction between the nucleus and the outermost electrons.
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Ionic Radius
- Definition: The ionic radius is half the distance between the nuclei of two like ions in a crystal lattice.
- Trend across a period: Generally, cations (positive ions) are smaller than their parent atoms because they have lost outer electrons. Anions (negative ions) are larger than their parent atoms because of increased electron-electron repulsion in the expanded valence shell.
- Trend down a group: The ionic radius increases due to the same reasons as the atomic radius: addition of extra shells and increased shielding effect.
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Ionisation Energy
- Definition: Ionisation energy is the energy required to remove the outermost electron from a gaseous atom.
- Trend across a period: Increases from left to right because:
- Increased nuclear charge as we add protons.
- Decreased atomic radius, making the outer electron closer to the nucleus.
- Greater attraction between the outer electron and the nucleus.
- Trend down a group: Decreases as we move down because:
- Increased atomic radius.
- Increased shielding from inner electrons.
- Outer electron is further from the nucleus, hence easier to remove.
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Electron Affinity
- Definition: Electron affinity is the energy change associated with adding an electron to a gaseous atom.
- Trend across a period: Generally, electron affinity becomes more negative (more energy is released) as we move from left to right. This is because:
- Atoms are smaller with a greater nuclear charge, so they can attract and accommodate an additional electron more readily.
- Trend down a group: Electron affinity values become less negative (less energy is released) because:
- Larger atomic size.
- Reduced attraction between the incoming electron and the nucleus due to increased shielding.
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Electronegativity
- Definition: Electronegativity refers to the ability of an atom to attract shared electrons in a chemical bond.
- Trend across a period: Increases from left to right because:
- Atoms are smaller with an increased nuclear charge.
- Higher effective nuclear charge pulls shared electrons closer.
- Trend down a group: Electronegativity decreases as we go down a group due to:
- Increased atomic size.
- Decreased attraction for shared electrons.
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Understanding Periodicity
The periodic table's design reflects the periodicity in elemental properties, providing a tool to predict chemical behaviour. When moving across periods, there is a progressive change in properties due to increasing atomic number and changes in the electron configuration. Going down groups, elements share similar electron configurations in their outermost shell, leading to similar chemical behaviours. This organisation allows chemists to make informed predictions about an element's properties based on its position on the table.
FAQ
Electron affinity and electronegativity are related but distinct concepts. Electron affinity refers to the energy change when an electron is added to a neutral atom to form a negative ion. It's a measurable quantity, representing the energy released or absorbed during this process. Electronegativity, on the other hand, is a qualitative measure of an atom's ability to attract and bind with electrons in a chemical bond. It doesn't directly measure an energy change but rather provides a scale (like the Pauling scale) to compare atoms' relative abilities to attract electrons within a bond.
When an atom loses electrons to form a cation, the electron cloud shrinks. The loss of electrons results in a greater positive charge in the nucleus relative to the number of electrons, leading to a stronger pull on the remaining electrons and hence a smaller radius. Conversely, when an atom gains electrons to form an anion, the added electrons increase the electron-electron repulsion, causing the electron cloud to expand. Thus, the ionic radius increases. The magnitude of these changes also depends on the effective nuclear charge and the amount of shielding from inner electron shells.
Periodicity in the properties of elements is intrinsically linked to their electronic structure. As you move across a period, electrons are added to the same energy level, but the nuclear charge increases. This leads to trends like increasing ionisation energy and electronegativity. Similarly, as you move down a group, new electron shells are added, increasing atomic and ionic sizes. The repeating patterns or trends in properties observed when elements are arranged in order of increasing atomic number (which means the same order as their electronic configurations) is the essence of periodicity. This predictability in properties is primarily because of the systematic filling of electron orbitals.
The anomalous values in ionisation energy observed between certain pairs of elements can be attributed to electron configurations. For example, boron has an electron configuration ending in 2p1, whereas beryllium's ends in 2s2. It's easier to remove the p electron from boron than the s electron from beryllium because p orbitals are at a higher energy level. Similarly, the decrease in ionisation energy from oxygen to nitrogen is due to the repulsion between electrons paired in oxygen's 2p4 configuration. It's easier to remove an electron from the more repelled paired electrons in oxygen than from the singly occupied p orbitals in nitrogen.
As one moves from left to right across a period, the atomic radius decreases due to an increase in the nuclear charge without a substantial increase in shielding. With every added proton in the nucleus, the nuclear charge increases, pulling the electron cloud closer to the nucleus. Simultaneously, these additional electrons are added to the same main energy level, meaning that there isn't a substantial increase in shielding to counterbalance the increasing nuclear charge. Thus, even though more electrons are being added, they are pulled inwards more strongly, resulting in a smaller atomic radius.
Practice Questions
As one moves from left to right across a period in the periodic table, the ionisation energy generally increases. This increase is due to several reasons. Firstly, as we progress across a period, the atomic number increases, leading to an increase in the number of protons in the nucleus. This results in a stronger nuclear charge, which attracts the electrons more forcefully. Secondly, the atomic radius decreases across the period, making the outermost electron closer to the nucleus and harder to remove. Lastly, as electrons are added to the same energy level, the shielding effect remains constant, so the increase in nuclear charge isn't offset by increased shielding. Consequently, the energy required to remove the outermost electron (ionisation energy) increases.
Electronegativity decreases as one moves down a group in the periodic table due to a combination of factors. Primarily, as we move down a group, the atomic size or radius increases. This is because of the addition of extra electron shells. As the atomic size increases, the outer electrons are further from the nucleus. Consequently, the attraction between the nucleus and these outer electrons decreases. Additionally, the increase in the number of inner electron shells results in a greater shielding effect. This shielding reduces the effective nuclear charge experienced by the outermost electrons. As a result, atoms lower down in a group have a reduced ability to attract shared electrons in a chemical bond, leading to a decrease in electronegativity.
