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Topic 4 – The Periodic Table

1.3.1 – Classify the components of a periodic table (period, group, metal, metalloid, nonmetal, transition1.3.2 – Infer the physical properties (atomic radius, metallic and non-metallic characteristics) of an element based on its position on the Periodic Table. 1.3.3 – Infer the atomic size, reactivity, electronegativity, and ionization energy of an element from its position on the Periodic Table.

I. The Periodic Table In 1869, Dmitri Mendeleev, a Russian chemist, was the first

person to devise a table that organized the elements in a logical fashion, mostly by atomic mass, with a few exceptions made to accommodate certain elements based on their properties.

In the modern periodic table, elements are arranged according to increasing atomic #, so that elements with similar properties fall in the same column. periods: horizontal rows (1-7) groups (families): vertical columns (1-18, or 1A-8A and 1B-8B)

Representative (main group) elements are those in groups 1A–8A, or 1, 2 and 13-18.All main group elements in the same group have: the same number of valence electrons. the same oxidation number. similar chemical properties.

♦ Note the location of metals, non-metals and metalloids on the periodic table:

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II. Representative (Main Group) Elements Representative (main group) elements are those in groups 1, 2 and 13-18 (or 1A–8A).

Group 1 elements are known as alkali metals. one valence electron soft, silvery metals, easily cut with a knife very reactive, especially with air and water

Group 2 elements are known as alkaline-earth metals. two valence electrons They are harder, slightly less reactive, & have higher melting points than alkali metals.

(a) (b) (c)

Elements in groups 1 and 2 are very “active” (reactive) metals. Metals in group 1 are more reactive than metals in group 2 (because group 1 metals have lower

ionization energies). Alkali metals are too reactive to be found free in nature – they are always combined with some

other element (ex: NaCl). Alkaline earth metals are less reactive than alkali metals, but still too reactive to be found as free elements in nature.

♦ Video of sodium and potassium reactivity: http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/007_SODIUMANDPO.MOV)

♦ Group 17 (7A) elements are known as halogens. Halogens are the most reactive non-metals (they have high

electronegativities, and need only one electron to fill up their octets).

Shown here are flasks filled with chlorine, bromine and iodine (left to right). Note their states at room T.

Halogens often react with metals to form salts (such as the reaction shown above: Na(s) + Cl2(g) NaCl(aq)

♦ Group 18 (8A) elements are known as noble gases (or inert gases). These are the least reactive of all elements because they have a full outer shell. Helium is considered a noble gas even though it doesn’t have a full octet because it does have a

full outer shell of valence electrons (with 2 valence electrons).

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(a) The rxn. of potassium with water.(b) The rxn. of calcium with water. These photos illustrate that group 1 metals tend to be more reactive than group 2 metals. (c) The reaction of sodium metal and chlorine gas is very exothermic.

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III. Transition Elements Elements in the d-block are usually called transition elements (or transition metals). This includes

all elements in groups 3-12 (or the “B group elements”, 1B-8B). All of these elements are metals, with typical metallic properties. They are typically harder, denser, less reactive & have higher melting points (with exceptions

such as mercury, Hg) than groups 1 & 2 elements.

F-BLOCK ELEMENTS

F-block elements are neither representative (main group) nor transition elements. Lanthanides = top row of the f-block; also called rare earth elements Actinides = bottom row of the f-block; most are synthetic, all are unstable and radioactive

♦ Good short movie outlining periodic table trends:http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/044_PeriodicProper.MOV

Periodic Table Trends

I. Atomic Radius Atomic Radius – ½ the distance between the nuclei of two identical covalently-bonded atoms

within a group : Atomic radii increase as you go down a group. This is because the valence electrons occupy higher # (and energy) sublevels, farther

from the nucleus.

within a period : atomic radii decrease across a period You are adding a proton and an electron each time you go one element to the right, but the

added electrons are in the same energy level, and more protons means a stronger pull on the electrons in that energy level, thus pulling them in tighter.

valence electrons – electrons in the highest energy level There can only be a maximum of 8 valence electrons, since only the s and p sublevels (2e-

and 6e-, respectively) can be in the highest energy level. Valence e- are the only e- available to be lost, gained, or shared in the formation of chemical

compounds.

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♦ Good animation explaining the trend in atomic radius: http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/046_AtomicRadii.MOV

II. Ionic Radius Ionic Radius – ½ the diameter of an ion in an ionic compound

Ions are atoms (or groups of atoms) with a positive or negative charge. cation – positively charged ion (formed by losing e-)

When cations are formed, ionic radius decreases. Taking away electrons makes the radius smaller: the positive

charge of the nucleus is unchanged (you still have the same # of protons), so the remaining e- are drawn closer to the nucleus.

The more positive the charge on the cation, the smaller the cation.

anion – negatively charged ion (formed by gaining e-) When anions are formed, ionic radius increases. Adding an electron makes the radius larger: attraction from

nucleus is more “diluted” – same # of protons pulling on one extra electron; also greater repulsion from other e- (shielding).

The more negative the anion, the larger the cation.

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within a period : Ionic radius of cations and anions decreases from left to right. Note that this trend “recycles” when you move from anions (non-metals) to cations

(metals), because of the types of ions metals and non-metals tend to form. Metals form cations, and non-metals for anions (in reactions where ionic bonds occur).

within a group : Just like atomic radius, ionic radius increases as you go down a group.

Table Comparing Atomic and Ionic Radii of Selected Elements

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III. Ionization Energy Ionization Energy – minimum amount of energy required to strip away one electron from an

atom of that element in the gaseous state, producing a positive ion (cation)

The nucleus holds valence electrons less tightly. Outer electrons are shielded from the + charge of nucleus by all the inner shell e- that lie

between them and the nucleus.

within a period : Ionization energy increases to the right across a period, because the increasing # of protons more strongly attracts valence electrons in the same energy level.

Group 1 elements have the lowest ionization energies of any group. Within a given period, group 1 elements have the largest radius, putting their valence shell farthest from the nucleus, so it takes comparatively little energy to strip off one valence e–.

This trend is also a reflection of atomic radius. The larger the radius, the further the valence electrons are from the pull of the nucleus, and the less energy it would take to strip them away. Conversely, atoms in a group with smaller radii (like the halogens) have very high ionization energies, because their valence electrons are held so tightly by the nucleus. Much energy would be required to strip an electron away.

within a group : Ionization energy decreases as you go down a group, because the outermost electrons are farther from the nucleus (atomic radius increases down a group) and therefore the nucleus holds them less tightly. (Sort of like how the sun holds Neptune less tightly than all the other planets)

The valence electrons are shielded from the + charge of nucleus by all the e- that lie between them and the nucleus. Since large atoms have a greater numbers of e- shells, they experience greater shielding.

This trend helps to explain why the reactivity of metals increases as you go down a group for metals – their radii get larger, so their ionization energy decreases (and metals react by losing electrons, in order to achieve the stability of a full octet).

♦ Two good animations explaining ionization energies:http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/047_IonizationEner.MOV http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/048_PeriodicIoniza.MOV

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IV. Electronegativity Electronegativity – the ability of an atom to attract electrons in a chemical bond Remember, the goal for all atoms is to be STABLE.

Noble gases (Group 18) are the most stable elements because they have a full octet. Elements usually form ions such that the ion has a full octet. For example, calcium forms 2+

ions, because losing 2 electrons gives it a noble gas electron configuration.

The most electronegative element, fluorine, is arbitrarily assigned an electronegativity value of 4.0. All other values are calculated relative to F.

within a period : electronegativities increase left to right across a period Atomic radii decrease left to right across a period, you will recall, which means the valence

shell of electrons is closer to the nucleus. Thus, the nucleus is able to exert a greater attractive force on electrons.

within a group : electronegativities tend to decrease down a group Again, larger atoms have electron shells farther from the nucleus. The nuclei of larger atoms

cannot attract electrons as strongly as smaller atoms, whose nuclei are closer to the valence shell, and can therefore exert a greater pull on electrons.

There is also greater shielding in larger atoms which have more electron shells. This helps explain why reactivity for non-metals increases as you go up a group. The atomic

radii decrease, bringing the nucleus closer to the valence shell, which increases the ability of the nuclei to attract electrons (higher electronegativity). Remember that non-metals react (with metals) by gaining electrons.

Conversely, metals tend to have low electronegativities, because they have large radii (remember metals tend to lose electrons in chemical reactions – not gain them).

♦ Here is a good video segment explaining electronegativity:http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/052_Electronegativ.MOV

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V. Electron Affinity Electron Affinity – The energy change associated with adding an electron to an atom in its gaseous

state. All elements that will accept electrons have a negative (exothermic) electron affinity. Electron affinities generally become more negative (more exothermic) left to right across the PT,

although there are some exceptions.

For example, consider why N– has no electron affinity, but C– does. N– has an electron configuration of 1s22s22p4. Adding that extra electron to N means one p

orbital has 2 electrons, creating extra repulsion, making N– unstable. C– has an electron configuration of 1s22s22p3, which is relatively stable with 3 unpaired p

electrons; thus adding an electron to C to form C– doesn’t create any extra repulsion.

So what about O– then? O– is stable, unlike N–, because the greater nuclear charge of O– (the extra proton) is sufficient

to overcome the repulsion associated with putting an extra electron in an already occupied p orbital.

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Electron affinity generally becomes less negative (less exothermic) down a group. Larger atomic radii à weaker attractions between the nuclei and electrons. Note that there are some exceptions, though, such as in the halogens. Fluorine’s electron

affinity is abnormally low, probably because fluorine’s 2p orbitals are so small that electrons are very close together, creating unusually large electron-electron repulsions.

Electron affinity generally becomes more negative (more exothermic) L to R across a period. L to R across a period, nuclear charge (# of protons) increases, and radius decreases,

increasing the energy released when electrons are added. Metals tend to have low electron affinities, given that they tend to have large radii. This

should make sense if you remember that metals tend to lose electrons in chemical reactions – not gain them.

Note there are lots of exceptions to increasing trend for electron affinity across a period, though!

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