Periodic Properties PDF

# Atomic & Ionic Sizes ## Definition: Atomic size or Atomic Radius is defined as the distance from the centre of the nucleus upto the centre of outermost electron. It is measured in Angstrom unit (A°). - **Nucleus** - **Valance Shell** - **Atomic Radies** ## The Ionic Size The ionic size is wh...

# Atomic & Ionic Sizes ## Definition: Atomic size or Atomic Radius is defined as the distance from the centre of the nucleus upto the centre of outermost electron. It is measured in Angstrom unit (A°). - **Nucleus** - **Valance Shell** - **Atomic Radies** ## The Ionic Size The ionic size is when the atom loses or gains electrons to become negatively charged (anion) or positively charged (cations) ions. When atoms lose or gain electrons, the size of the ion is not the same as the original atom. - When an atom donate an electron it becomes (true) ion. - Size of true ion < Neutral atom because of destroying outermost shell. - Size of cation is smaller than neutral atom because proton in nucleus > electrons in orbital, hence it feels more attraction than before. ## Ionization Energies ### Definition: The ionization energy or ionization potential may be defined as "the energy required to remove one electron from the outermost orbit of an isolated gaseous atom in its ground state." - In an atom, the energy required to remove first electron from a gaseous atom is called 1st ionization energy. - The energy required to remove one electron from unipositive ion to form bipositive ion is called 2nd ionization energy. Because after the removal of 1st electron, the electron is firmly bound to the nucleus & so more energy is required to remove 2nd electron. ## Factors on Which Ionization Energy Depends - **Size of Atom:** Larger the size of atom, lesser will be force of attraction & lower will be ionization energy. - **Nuclear Charge:** Greater the magnitude of nuclear charge, greater will be force of attraction, so higher will be ionization energy. - **Screening/Shielding effect:** The effect of reduction of force of attraction by the shells present between nucleus and valence electron is called screening effect. Greater the no of shells between nucleus and valence electron, lesser will be the electron nucleus attraction & less will be Ionization potential. s>p>d>f is order of screening or shielding effect. - **Penetration of sub-shells:** Ionization energy also depends upon the type of electron which are to be removed. s electrons are closer to the nucleus (more penetrated towards nucleus) and are more tightly held then p, d, or f electrons. Hence Ionization energy in the order s>p>d>f. - **Stability of Configuration:** Completely filled and exactly half filled subshells impart extra stability which results in higher ionization energy. # Electron Affinity ## Definition Electron affinity is the amount of energy released when an electron is added to the valence shell of a neutral, isolated gaseous atom. - Electron affinity is generally always negative because of tire electron opposes electron having negative charge ## Exceptions - Noble gases have zero electron affinity because they have completely filled orbitals and acquire inert electronic configuration. - Be and Mg of 2nd group have zero electron affinity because their s subshell is completely filled. ## Factors on Which Electron Affinity Depends - **Atomic Size:** Electron affinity is inversely proportional to atomic size. So when atomic size then electron are less attracted with increasing distance. - **Nuclear Charge:** Electron affinity is proportional to nuclear charge. As z increases, then Electron affinity increases because electron are more attracted. - **Stable Electronic Configuration:** Stable atoms do not accept electron. Example - half filled - p^3, d^5, f^7 Fully filled - p^6, d^10, f^14 Electron affinity is zero or very less. ## Trends in Electron Affinity - **Along a Period:** As nuclear charge increases then electron affinity increases because more energy is released. - **Down The Group:** As atomic size increases then electron affinity decreases because less energy is released. - **Group-17:** As we know, when atomic number increases then electron affinity decreases. So according to this rule, fluorine should have highest electron affinity but actually chlorine has highest electron affinity. Expected order: F>Cl>Br>I Real order: Cl>F>Br>I **Reason:** As we know electron affinity is tendency to accept electron. In case of fluorine, it attracts electron more but does not accept electron due to small size or low space in atom. While in case of chlorine, it attracts electron less than fluorine but accepts electron due to larger size or large space in atom. That's why electron affinity of chlorine is highest in periodic table. # Electronegativity ## Definition Electronegativity is the tendency of an atom to attract a shared pair of electrons in a covalent bond. ## Example - Fluorine is more electronegative than hydrogen or chlorine. - Hydrogen is less electronegative than chlorine. - Electronegativity is a relative value, it has no unit. ## Factors on Which Electronegativity Depends - **Size of atom:** Smaller the size higher will be electronegativity. - **Smaller size** - **Higher electronegativity** - **Nuclear charge:** Higher the nuclear charge, higher will be electronegativity. - **Charge on cation:** Higher the charge on a cation, lower will be electronegativity. - **Charge on anion:** Higher the charge on anion, higher will be electronegativity. - **Hybridization:** As the percentage of 's' character increases then electronegativity increases. Sp > Sp^2 > Sp^3 ## Trends of Electronegativity in Periodic Table - **Along the Period:** Along the period, electronegativity increases because nuclear charge on atoms increases. - **Down the Group:** Size increases, electron affinity decreases. # Polarizability ## Polarizing Power and Polarizability - If the ions are at distance then there will be no force of attraction. - But if they are closer enough, then they experience four types of forces.