Periodic Table – Organizing the Elements Chapter 5.4 & Chapter 14

Dmitri Mendeleev About 70 elements had been found by the mid 1800’s Mendeleev was the first to organize them in a systematic way

He listed the elements in order of increasing atomic mass Arranged the elements in columns so those with similar properties were side by side

He left blank spaces where nothing fit He predicted the physical properties of the missing elements He was mostly correct

Henry Moseley Moseley determined the atomic # of the elements and arranged the table by atomic number instead of atomic mass

The modern periodic table is arranged by atomic number The periodic table has atomic # increasing from left to right & top to bottom

Metals Vs Nonmetals The table can be divided into 3 main types of elements – metals, nonmetals & metalloids Metals reside on the left side of the table Nonmetals are on the upper right side of the table Metalloids are b/tw the metals & nonmetals

Metals are typically shiny (luster) & are good conductors of heat and electricity Metals are all solids, except Mercury (Hg) which is a liquid Nonmetals are usually dull (no luster) & are poor conductors of heat and electricity Nonmetals can be solid (C & S), liquid (Br) or gas (O & H)

Metalloids Metalloids are their own category of elements – there are only 7 metalloids! B, Si, Ge, As, Sb, Te, & At – they border the stair-step line that most periodic tables include – but not all. Metalloids have properties in b/tw the metals & nonmetals Si is used in computers b/c it is a semi-conductor

Periodic Law The horizontal rows on the periodic table are called periods Properties change as you move across a period

The properties repeat when you move from one period to the next Periodic Law: there is a periodic repetition of the chemical & physical properties of the elements

Groups Each vertical column is called a group or family Elements in the same group have similar properties

Groups have a number and a letter (pg 124) The group with Li, Na, K etc is called Group 1A Group 1A elements are called the alkali metals Alkali metals are the MOST active metals

Group 2A are the Alkaline Earth Metals Alkaline Earth Metals are not as active as the alkali metals, but react with many substances

Group 7A are called the Halogens (F, Cl, Br, I) Halogens are the most active nonmetals Group 0 (8A) are the noble gases (He, Ne, etc) The noble gases are mostly UNREACTIVE!

Groups 3A to 6A Groups 3A to 6A do not have specific names – they are usually named by the top element of the group. For instance, The Carbon family is Group 4A. These groups can contain metals, metalloids and nonmetals!

All group A elements are called the representative elements They exhibit a wide range of physical & chemical properties

Group B elements are the transition & inner-transition metals Gold & silver are transition metals Uranium is an inner-transition metal

The elements can also be classified by their electron configuration Electrons play the most important part in determining the properties of elements

Write the electron configurations for the Alkali Metals What similarities do you see? The Halogens? The Noble Gases?

The noble gases have their outermost s & p sublevels filled completely The Representative Elements have their outermost s & p sublevels partially filled

The Transition Metals – their outermost s & nearby d sublevels contain electrons The Inner Transition Metals – their outermost s & nearby f sublevels contain electrons

s block, p block, d block & f block Where are they? Blocks The Table can be broken up into blocks - tell you the outermost sublevels that are filled s block, p block, d block & f block Where are they?

“S” block Groups 1 & 2 Alkali Metals and Alkaline Earth Metals Electron Configuration ends in an S Sub-level.

“P” block Groups 13 thru 18 (or 3A - 8A) Includes halogens and noble gases (except He) Electron Configuration ends in a P Sub-level.

“D” block Groups 3 thru 12 Transition Metals Au, Ag, Fe, Pt, etc Electron Configuration ends in a D Sub-level.

“F” block “Inner Transition Metals” Includes Uranium Electron Configuration ends in an F Sub-level.

Periodic Trends An element’s placement in the periodic table determines characteristics like the size of the atom, its ability to attract electrons and the stability of its electron configuration.

Atomic Radius Size of atoms of each element: How will the size of atoms change as we proceed down a group? i.e. Compare the sizes of Li and Na. From Li to Na, we add an entire energy level, therefore the size increases.

How will the size of atoms change as we proceed across a period? Compare C, N and O. Which is largest? Oxygen has the most electrons. However, it also has the most protons. The outermost electrons of Oxygen are in the same sub-level as C and N.

Oxygen’s greater nuclear charge attracts the electrons, causing the atom to contract! Oxygen is the smallest of the three, Carbon is the largest. Atomic Radius decreases as we go across a period from left to right and increases going down a group.

Rank the following sets in order of decreasing Radius. Examples Rank the following sets in order of decreasing Radius. S, Cr, Se, Sr, Ne Sr, Cr, Se, S, Ne Fe, N, Ba, Ag, Be Ba, Ag, Fe, Be, N

Ionization Energy Amount of energy required to remove a valence electron from an atom. The more stable an element is, the harder it will be (more energy is required) to remove an electron. Some elements become more stable by losing an electron so they lose electrons easily (less energy needed).

How does ionization energy vary within a group (compare Li and Na)? The electron to be removed from Na is further from the nucleus than Lithium’s electron. Sodium’s electron is held more loosely and therefore easier (less energy) to remove.

How does ionization energy vary across a period How does ionization energy vary across a period? (Compare elements in 3rd period) Sodium attains a Noble Gas configuration by losing an electron, so little energy is required. Magnesium is somewhat stable due to a full 3s sub-level, so more energy is needed.

Ionization energy increases across a period & decreases down a group. Argon is a Noble Gas. Due to its stability, it is very difficult (much energy needed) to remove an electron. Chlorine has no stability in its configuration, so it is easier to remove an electron. Ionization energy increases across a period & decreases down a group.

Rank the following sets in order of decreasing Ionization Energy. Examples Rank the following sets in order of decreasing Ionization Energy. K, Zn, Cs, Ar, P Ar, P, Zn, K, Cs C, He, Ag, Pt, Sn He, C, Sn, Ag, Pt

2nd Ionization Energy After an element loses one electron, it may lose another. Sometimes it is easier to lose the second electron than the first. Sometimes the first electron is easier to remove This depends on the stability of the electron configuration after it loses the 1st electron

2nd Ionization Energy Magnesium will be more unstable after losing 1 electron, so the 2nd electron will be lost more easily Sodium becomes very stable after losing just 1 electron, so it will be more difficult to lose the second electron

Electronegativity Describes an element’s attraction for an electron in a covalent bond. Elements that need electrons to complete an energy-level will have a high electronegativity. Elements that want to lose electrons have low electronegativities.

How does Electronegativity vary within a group? (compare F and Cl) Both elements need an electron to complete a p sub-level. Fluorine’s p sub-level is closer to its nucleus, so it has a greater magnetic attraction for a free electron. F has a higher electronegativity!

How does electronegativity vary across a period? (period 2) Fluorine benefits the most by gaining an electron, so it has the highest electronegativity. Lithium, which wants to lose an electron has very little attraction for an additional electron.

Carbon can gain electrons but sometimes loses them as well, so its electronegativity is between F and Li. Noble Gases have no Electronegativity! Electronegativity increases across a period and decreases down a group (Noble Gases omitted).

Rank the following sets in order of decreasing Electronegativity: Examples Rank the following sets in order of decreasing Electronegativity: Cu, F, Mn, Sr, Si F, Si, Cu, Mn, Sr Al, Ca, S, Cl, Fe Cl, S, Al, Fe, Ca

Atomic Radius vs. Ionic Radius When atoms become ions, their size will change: Metals will lose electrons and become smaller than the neutral atom (Na+ is smaller than Na) Nonmetals will gain electrons and become larger than the neutral atom (F- is larger than F)