Trends of the Periodic Table

Review! Periodic Table was first organized by… –Dmitri Mendeleev in the mid 1800’s –Mendeleev organized the elements by chemical reaction in rows, then by atomic mass in columns Henry Moseley then took Mendeleev’s table, kept the chemical reactivities together, but placed them in columns instead. He also ordered the elements by increasing atomic number in rows. When Moseley did this, all the periodic trends just fell into place. Remember: columns = groups/families, rows = periods

Periodic Trends

Electrons Electrons do not freely float in space Orbit around nucleus: Electron shells Each shell corresponds to an amount of energy.

Valence Electrons The valence electrons are the outermost electrons of an atom. The valence electrons determine the chemical properties Number of valence electrons equals the column number in the “A” columns Elements with the same number of valence electrons are very similar chemically – Alkali metals in Group 1A – 1 valence electron Li, Na, K, Rb, Cs – Halogens in Group 7A – 7 valence electrons F, Cl, Br, I

Atomic Radius What is Atomic Radii? Distance from the nucleus to the outermost level of e- (aka the valence shell) What trend do you see as you go across (left to right) the period? Atomic radius decreases Down the group? Atomic Radius increases WHY???

Explaining the Trend As you go L to R, the atomic radius decreases because as you go L to R, the amount of attraction between p + and e - increase. More attractions = smaller atomic radius As you go down a column, atomic radius increases because the e - are farther away from the nucleus. There are weaker attractions. Weaker attractions = larger atomic radius

Electronegativity What is Electro- negativity? An atom’s Luuuvvv for electrons! The tendency to attract another atom’s electrons What trend do you see as you go across the period? Electronegativity increases! Down the group? Electronegativity decreases! WHY???

Explaining the Trend As you go L to R, electronegativity increases because of the increase in protons. The more protons, the more able it will be to attract other atom’s electrons. More attractions (small radius) = large electronegativity As you move down a column, electronegativity decreases because of the increase in number electron an atoms already has. This means the atom will be less able to attract another atom’s electrons. Less attractions (large radius) = small electronegativity

Ionization Energy What is Ionization Energy? The energy needed to remove an electron What trend do you see as you go across the period? Ionization E increases Down the Group? Ionization E decreases WHY???

Explaining the Trend As you go L to R, the ionization energy increases because of the increase in the number of protons. The more protons, the more energy that is needed to remove an electron. More attractions (small radius) = large ionization energy As you go down a column, the ionization energy decreases because of the decrease in attractions. –Due to electron shielding –More electrons, leads to outer electrons less tightly held. The less attractions, the lower the energy that is needed to remove an electron. Less attractions (large radius) = small ionization energy

Ionization Energy Amount of energy required to remove an electron from the ground state of a gaseous atom or ion. –First ionization energy is that energy required to remove first electron. –Second ionization energy is that energy required to remove second electron, etc.

Ionization Energy It requires more energy to remove each successive electron. When all valence electrons have been removed, the ionization energy takes a quantum leap.

Electron Affinity What is Electron Affinity? The energy needed to add an electron As you go across the period electron affinity increases. Electron affinity decreases down the family WHY???

Explaining the trend As you go L to R, the electron affinity increases because of the increase in the number of protons. The more protons, the greater the attraction the protons have for electrons. More attractions (small radius) = large electron affinity As you go down a family, the electron affinity decreases because of the decrease in attractions. –Due to electron shielding –More electrons, leads to outer electrons less tightly held. The less attractions, the lower the electron affinity Less attractions (large radius) = small electron affinity

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