Trends In Periodic Table 2xg1o

TRENDS IN PERIODIC TABLE Nicole Farquharson

General trend of effective nuclear charge (Zeff) • Going down the group, effective nuclear charge decreases. This occurs when there is maximum shielding (shielding of electrons from lower orbitals) increases. • Going across the period, effective nuclear charge increases. This occurs when there is medium – minimum shielding (electrons shielded by electrons in the same shell but different subshell or electrons shielded by electrons in the same subshell)

Atomic Radius • It is defined as half the distance between two bonded nuclei or the radius of an atom • Going down the group, atomic radius increases and the effective nuclear charge decreases because there is maximum shielding where each electrons has one more level of inner electrons that shield the outer electrons. When the effective nuclear decreases, the pull on the valence shell get weak. There is also additional energy levels required to accommodate the additional electrons and this tends to have a larger shells that are being occupied. • Going across the period, atomic radius decreases and the effective nuclear charge increases because there is medium (electrons are not so strongly shielded in the same shell but different subshells) less and minimum shielding (electrons are weakly shielded in the same subshell). There are electrons added to the same outer shell. The effective nuclear charge increases because it is felt by the valence electrons to draw them closer to the nucleus, this will pull the electron cloud in a little tighter. There is a greater attraction between the nucleus and the electron.

Other information • To understand this trend it is first important to realize that the more strongly attracted the outermost valence electron is to the nucleus then the smaller the atom will be. While the number of positively charged protons in the nucleus increases as we move from left to right the number of negatively charged electrons between the nucleus and the outer most electron also increases by the same amount. Thus you might expect there to be no change in the radius of the outermost electron orbital since the increasing charge of the nucleus would be cancelled by the electrons between the nucleus and the outermost electron. This, however, is not the case. The ability of an particular inner electron to cancel the charge of the nucleus for the outermost electron depends on the orbital of that inner electron. • Moving from left to right across a period, electrons are added one at a time to the outer energy shell. Electrons within a shell cannot shield each other from the attraction to protons. Since the number of protons is also increasing, the effective nuclear charge increases across a period. This causes the atomic radius to decrease.

Ionic radius •

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The radius of a positively or negatively charged ion. Going down the group ionic radius increases, cations are smaller and anions are bigger. Going across the period, there is a large increase from cation to anion. When a cation forms, electrons are removed from the outer shell. This results in effective nuclear charge to pull remaining electrons closer. Removing electrons decreases electron-electron repulsion so the electron clouds contract, so they can be close together. There is a great attraction between nucleus and electrons tend to pull the electron cloud in a little tighter. There would be medium – minimum shielding which increases net pull of electrons from nucleus. When an anion forms, electrons are added to the outer shell, the effective nuclear charge decrease because there is an increase in electron-electron repulsion which causes electrons to occupy more space and swells the electron cloud. There is less attraction between nucleus and electrons. There would be maximum shielding which decrease the net pull of electrons from nucleus. Summary: Ions may be larger or smaller than the neutral atom, depending on the ion's charge. When an atom loses an electron to form a cation, the lost electron no longer contributes to shielding the other electrons from the charge of the nucleus; consequently, the other electrons are more strongly attracted to the nucleus, and the radius of the atom gets smaller. Similarly, when an electron is added to an atom, forming an anion, the added

Ionization energy • The required energy in removing a mole of electrons from a mole of atoms in the gaseous state. The smaller the atom, the more tightly its electrons are held to the positively charged nucleus and the more difficult they are to remove. • Going down the group, ionization energy decreases, effective nuclear charge decreases. The electrons are further away from the nucleus and easier to remove the outermost shell. Effective nuclear charge increase because there is maximum shielding which shield the valence electrons more tightly. Electrons are removed from a higher energy level which is further from the nucleus. As the effective nuclear increases, the remaining electrons are more strongly attracted to nucleus making it harder to remove the next electron. • Going across the period, ionization energy increases and effective nuclear charge increase. As ionization energy increases the atoms gets smaller. The outer electrons are closer to the nucleus and more strongly attracted to the center. The attraction between nucleus and outer electrons increases, more energy is needed which makes an electron harder to remove. Removing successive electrons from the same energy level requires much more energy each time, but removing an electron from a lower energy level requires much more energy because the electron being removed is the so much closer to the positively charged nucleus.

Electron Affinity • This is the energy change when an electron is added to a gaseous atom to form an anion. • Going down the group, electron affinity decrease effective nuclear charge decrease because the nucleus is farther away from an electron being added. There is maximum shielding which means there is less pull on the nucleus based on the fact that the shell is stable. • Going across the period, electron affinity increase and the effective nuclear affinity increase. This is due to the great attraction between nucleus and electrons in which a valence shell that loses electrons easily will have little attraction for additional electrons. There is medium-minimum shielding making it easier to attract electrons so the value of electron affinity would be negative. The increasing effective nuclear charge attract the electron more strongly.

Electronegativity • The ability of an atom to attract an electron cloud towards itself in a chemical bond. It measures the ability of an atom to attract to itself the electron pair forming a covalent bond. The greater the electronegativity of an atom, the greater its attraction for electrons. • Going down the group, electronegativity decrease and effective nuclear charge increase. As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus. Decending a group electronegativity decreases because atomic radius increases due to electrons moving into new main energy levels (spdf). the electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius). • Going across a period, Electronegativity increase and because the atomic radius of elements decrease due to the nuclear charge increasing. The number of charges on the nucleus increases and this attracts the bonding pair of electrons more strongly.

• Electronegativity increases as you move across the periodic table from left to right. This occurs due to a greater charge on the nucleus, causing the  electron bonding pairs to be very attracted to atoms placed further right on the periodic table. Fluorine is the most electronegative element. • Electronegativity decreases as you move down the periodic table. This is caused by an increased amount of shielding, or screening, by the innermost electrons. As you move down the table more electrons are added between the nucleus and the bonding pair, causing the effective nuclear charge to be less. The increase in distance between the nucleus and the bonding pair decreases the attraction between the two.

• Electronegativity increases as you go across a period because, the next element along has an extra proton and an extra electron. The extra proton is in the nucleus, which means that the nucleus has a greater charge on it. The extra electron goes into the orbital with the lowest available energy, but because we are going across a period this means that valence electrons do not experience any extra shielding from the nucleus, so they feel a greater effective nuclear charge and are more strongly attracted by the nucleus. This stronger attraction means the electrons are pulled closer to the nucleus, meaning the atom gets smaller (atomic radius decreases across a period). Thus, a shared pair of electrons will be more strongly attracted to the nucleus and so the atom withdraws more electron density. It is more electronegative. As the number of electrons in an elements outer energy level nears a full octet (8) it gets increasingly difficult to remove an electron. • Electronegativity decreases as you go down a group. This is because as you go down a group, the principal number of the valence orbital increases, meaning that there is an extra 'shell' of electrons between the valence electrons and the nucleus. This means that the valence electrons experience greater shielding from the nucleus. This factor is more important than the increased number of protons in the nucleus and the increased charge on the nucleus. So despite the extra protons the valence electrons are less strongly attracted by the nucleus, and the electrons are not held as close to the nucleus (atom radius increases down a group). Thus, a shared pair of electrons will be less strongly attracted to the nucleus, so the atom withdraws less electron density. It is less electronegative. Inner energy level electrons block the attraction between the positively charged nucleus and the electrons in the outer energy level. This shielding effect makes it easier to remove a valence electron as the number of energy levels increases.

Lattice energy • The energy needed to break up the ions from the solid phase to gas phase. • Going down the group, The more anions, the lower the lattice energy to separate the ionic bonds. The atoms are not held tightly to the nucleus. The effective nuclear charge increases because there is maximum shielding and there is less pull on the nucleus. • Going across the period, lattice energy increases, the effective nuclear charge decreases. Lattice energy increases because the charge on the ion increases. Highly charged ions attract more strongly than ones with less charge. The more charge the greater the lattice energy. The more cations, the higher the lattice energy in order to separate the ionic bond. The cations are held tightly together because they are positively charged and they are losing electrons. The effective nuclear charge decreases because there is medium to minimum shielding. • As ionic radii increase, the electrostatic decrease so lattice energies of the compounds decrease. Lattice energy increases as charges on the ions increases as the value becomes more negative. When ions are closer together the lattice energy increases (becomes more negative).