Without looking at the relevant tables, describe: a. the main horizontal trends in Pauling elect

step 1 Okay, so this question is about the elements in the periodic table. So before answering this que

Without looking at the relevant tables, describe: a. the...

Key Concepts

Vertical Trends in Common Oxidation Numbers

This concept deals with the variation in the common oxidation states of elements as one moves vertically in the periodic table. In a group, elements often display similarities in valence electron configurations and therefore tend to adopt similar oxidation numbers. However, subtle differences in atomic size, electron shielding, and the effective nuclear charge can result in variations in the stability and preferred oxidation states as one goes down a group.

Vertical Trends in Ionic Radii

This concept captures the tendency for ionic size to change when moving down a group. Ionic radii typically increase with increasing atomic number within a group because each successive element has an additional electron shell. Despite the concurrent rise in nuclear charge, the very effective shielding by inner electrons means that the outer electrons are less tightly bound, leading to larger ionic sizes.

Vertical Trends in Pauling Electronegativity

This concept focuses on how electronegativity values change as one moves down a group in the periodic table. Electronegativity generally decreases down a group because additional electron shells increase the distance between the nucleus and the valence electrons and introduce greater shielding effects. As a result, atoms become less effective at attracting bonding electrons as their size increases.

Horizontal Trends in Ionic Radii

This concept describes how the size of ions generally changes when moving across a period. Typically, the ionic radii decrease from left to right because the increasing positive nuclear charge more effectively pulls the electrons closer to the nucleus in isoelectronic series. Although the number of electrons may remain constant for certain ions, the enhanced effective nuclear charge results in a general contraction of ionic size across a period.

Horizontal Trends in Pauling Electronegativity

This concept refers to the systematic change of an element’s ability to attract electrons as one moves across a period in the periodic table. Generally, as the effective nuclear charge increases from left to right, atoms become better at attracting electrons, and therefore their Pauling electronegativities tend to increase. This trend reflects changes in atomic structure, including the number of protons in the nucleus and the shielding effect by inner electrons.

Horizontal Trends in Common Oxidation Numbers

This concept involves the periodic variation of the oxidation states most commonly adopted by elements within the same period. As one moves across a period, the nature of bonding and the tendency to lose or gain electrons change; metals tend to exhibit lower positive oxidation numbers while nonmetals show higher positive or even negative oxidation states. This shift reflects variations in electronic configuration and chemical reactivity across the period.

Transcript

00:01 Okay, so this question is about the elements in the periodic table.

00:10 So before answering this question, be sure you have the periodic table with you.

00:18 So let's start it.

00:20 The first one is ask you about the largest atomic radius in the group 1b, in the group 1.

00:31 So first, we're going to find what is group one.

00:36 The group one is hydrogen, lithium, sodium, sodium, potassium, and the largest atomic radios.

00:46 So we know in the same group, when you increase the atomic number, the radius will also increase.

01:06 So in the group 1, the largest radius should be the one has largest atomic number.

01:19 And if you look at a periodic table, it's definitely the front thing.

01:29 It's on the bottom.

01:32 Okay, so b, the smallest atomic radius in the period of 3, so now the element are in the same period.

01:48 So for period three, you check on the periodic table, it's sodium, magnesium, aluminum, silicon, phosphorus, sulfur, chlorine, and argon.

02:03 So smallest atomic radius.

02:07 So remember, exclude the noble gas element.

02:12 If you, in the same period, if you increase atomic number, the radius will also increase, will actually decrease.

02:39 So, it asks us the smallest atomic radius in the period three.

02:44 So definitely excluding the noble gas, she has the largest atomic number, which is actually chlorine.

02:55 Chlorine has a smallest radius in the period 3.

03:01 And say largest first ionization energy in group 2.

03:08 So this one group 2, our question lies in the same group.

03:14 In the same group, if you increase the atomic number.

03:29 Actually, the ionization energy, i would just use ei to replace, or actually decrease.

03:39 So the largest first ionization energy element in group 2 should have smallest atomic number, so which is beryllium, b .e.

04:00 Okay, so for the d, most electroneactive in group 16.

04:08 So this is now very similar.

04:16 So it's the same group.

04:21 Same group if you go increase atomic number.

04:26 From up to down, the electronactivity where actually goes down.

04:40 So it says most electronactive, which means it should have the smallest atomic number.

04:48 So group 16, which is oxygen has the largest oxygen is actually has actually smallest atomic number...

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