7.10: History of the Periodic Table

•Dmitri Mendeleev given credit for the first periodic table in 1869

•Grouped elements with similar chemical & physical properties in rows according to atomic mass

• He emphasized how useful the periodic table could be in predicting the existence and properties of still unknown elements.

•In 1913, modifications were made to rearrange the table according to atomic number

7.12: Periodic Trends

Trend #1: Atomic Radius

• An atoms size is determined by its atomic radius

• Atomic radius is

defined as half the distance between the nuclei of two like atoms

r

Atomic Radius of a Metal

Atomic Radius of a Nonmetal

Atomic Size Trend In general, going across a period, atomic radius decreases. •Moving left to right across the periodic table, no new energy levels are added.

•However, more protons are added increasing the nuclear charge.

•Electrons are drawn closer to the nucleus, decreasing the size of the atom.

In general, going down a group, atomic radius increases. •the number of energy levels (principle quantum levels) increases.

Which should be the larger atom? Why?

Na Cl

CONCEPT CHECK!

Na should be the larger atom

because the electrons are not

bound as tightly due to a smaller

effective nuclear charge.

Which should be the larger atom? Why?

Li Cs

CONCEPT CHECK!

Cs should be the larger atom

because of the increase in orbital

sizes in successive principal

quantum levels (to accommodate

more electrons).

Trend #2: Atomic Radius of IONS

Cations are smaller than their parent atom

• Cations lose electrons

and (usually) an energy level

• Since the nuclear charge remains the same, the electrons are pulled in closer.

Anions are larger than their parent atom.

• Anions gain electrons.

• The same number of protons attracting an increased number of electrons causes the electrons to not be bound as tightly to the nucleus resulting in an increase in size.

Which is larger: S or S-2. Why? S-2 is larger – Same number of protons, so gaining electrons causes less of an attraction to the nucleus (decreased nuclear attraction/pull)

Which is smaller: Fe or Fe+4. Why? Fe+4 is smaller – Same number of protons, but losing electrons causes more of an attraction to the nucleus – the remaining electrons are pulled in tighter.

(greater nuclear attraction/pull) Which is smaller: Na+ or Al3+? Why Al3+: They both have the same number of electrons, but aluminum has more protons. (greater nuclear pull)

Which is smaller: Be2+ or Na+. Why? Be2+ has fewer energy levels.

A neutron walks into a

‘restaurant’ and says, "Hey

bartender give me a drink."

The bartender gives him

one and says, “No charge

for you"

Trend 3: Electron Affinity Electron Affinity

•Energy change (in kJ/mol) associated with the addition of an electron to a gaseous atom. (In other words, the atom’s likelihood of gaining an electron)

•The more negative the value (since it is exothermic), the more of an affinity to electrons that atom has

X(g) + e– → X–(g)

In general as we go across a period from left to right, the electron affinity increases. •The values become more negative. •nuclear charge increases, increasing nuclear pull and the ability to attract electrons

In general, going down a group, electron affinity decreases. •The values become more positive •added energy levels ‘shield’ the power of the nucleus; nuclear pull decreases and it is harder to attract electrons

SHIELDING EFFECT:

• The process of the inner electrons shielding (repelling) the outer electrons causing the outer electrons to be less attracted to the nucleus.

• Shielding increases as you go down the periodic table because more electrons in more energy levels are added

• Shielding remains constant as you go across because all the electrons in a period are in the same energy level.

Trend 4: Ionization Energy •Energy required to remove an electron from a gaseous atom or ion. X(g) → X+(g) + e–

Mg → Mg+ + e– I1 = 735 kJ/mol (1st IE)

Mg+ → Mg2+ + e– I2 = 1445 kJ/mol (2nd IE)

Mg2+ → Mg3+ + e– I3 = 7730 kJ/mol *(3rd IE)

*Core electrons are bound much more tightly than valence electrons.

First ionization energy is always smaller because as each successive electron is removed, nuclear pull becomes stronger.

Trend #4: Ionization Energy Going across a period, ionization energy increases Same

energy level, increased nuclear pull causes the electrons to be pulled closer to the nucleus. More energy is needed to remove them the closer they are to the nucleus.

Going down a group, ionization energy decreases. Electrons are being removed further from the nucleus as more energy levels are added. ‘Shielding’ causes the electrons to be less attracted to the nucleus - little energy is needed to remove them.

Which atom would require more energy to remove an electron? Why?

Na Cl

CONCEPT CHECK!

Which atom would require more energy to remove an electron? Why?

Li Cs

CONCEPT CHECK!

Li would require more energy to

remove an electron because the

outer electron is on average

closer to the nucleus (so more

tightly bound).

Arrange the elements oxygen, fluorine, and sulfur according to increasing:

• Ionization energy S, O, F • Atomic size F, O, S

EXERCISE!

The Periodic Table – Final Thoughts 1. It is the number and type of

valence electrons that primarily determine an atom’s chemistry.

2. Electron configurations can be determined from the organization of the periodic table.

3. Certain groups in the periodic table have special names.

4. Basic division of the elements in the periodic table is into metals and nonmetals.

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Special Names for Groups in the Periodic Table

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