History of the Periodic Table. By 1860, more than 60 elements had been discovered and chemists had to learn them and all that they made. There was no method for accurately determining atomic or number of atoms in a particular compound. Enter Italian chemist Stanislao , who presented a method for accurately measuring the relative mass of an atom. This allowed chemists to search for relationships between atomic mass and other of the elements.

A. Mendeleev and Chemical Periodicity. Russian chemist Dmitri took new information on atomic mass, and for the textbook he was writing, included this along with his hopes to organize the elements according to their . When arranged in increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals. Such a pattern is referred to as . Mendeleev:
1. created a chart or table in which elements were grouped together - a table on the elements.
2. left empty spaces where elements were [?].
3. predicted

B. Moseley and the Periodic Law. English scientist Henry and Ernest (remember him?) worked with metal spectra. They discovered a previously unrecognized pattern. The pattern ranked elements in increasing order according to nuclear charge, or the number of in the nucleus. This led to the modern definition of atomic number and the recognition that atomic , not atomic mass, is the basis for the organization of the periodic table. Moseley's discovery was consistent with Mendeleev's ordering of the periodic table based on their properties. Mendeleev's principle is correctly stated in what is known as the [?]: The physical and chemical properties of the elements are periodic functions of their atomic .
The Periodic Law states that when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic . Horizontal rows () have predictable properties based on an increasing number of in the outer orbitals. Vertical columns () have similar properties because of their valence electron configuration.
C. The Modern Periodic Table. Elements are arranged in order of their atomic numbers so that elements with similar properties fall in the same , or group.
1. Noble Gases (Group ). These are the most elements. They were discovered mainly by Sir William Ramsay in the 1890's.
2. The (Period 6). These are elements so similar in chemical and physical properties, the process of separating and identifying was tedious and required efforts of many. [Atomic numbers 58 (Ce, cerium) to 71 (Lu, lutetium)
3. The (Period 7). These are the14 elements from 90 (Th, thorium) to 103 (Lr, lawrencium). The lanthanides and actinides fit between groups 3 and 4 and are set off below the main portion of the periodic table.
4. [?]. You can make very good guesses about the reactivity of an element by just knowing where the element is found on the periodic table. As you look across or up and down the periodic table, there are traits (properties) that follow predictable periodic trends. For example, in Groups 1 and 18, the differences between the atomic numbers of successive elements are 8, 8, 18, 18, and 32. Groups 2 and 13-17 follow a similar pattern.

Electron Configuration and the Periodic Table. The Group 18 elements ( gases) undergo few chemical reactions. This stability results from the gases' special configuration. The highest occupied levels are completely filled with electrons. In helium the level is full and subsequent noble gases contain stable . Generally, the electron configuration of an atom's highest occupied level governs the atom's chemical properties.

Periods and Blocks: Horizontal rows of the table are called . There are 7 periods in the modern table. The length of each period is determined by the number of that can occupy the sublevels being filled in that period. Based on the electron configurations of the elements, the periodic table can be divided into 4 (s, p, d, and f).

A. The s-Block Elements (Groups 1 and 2). These are noted as ns1 and ns2 where n = the number of the highest occupied energy level. These elements are very chemically reactive . Their outermost energy levels contain only [?] or [?] electrons. The ease with which these electrons are lost makes these elements extremely .
- Group 1 elements are known as metals. They have a group configuration of ns1. They are silvery in appearance and enough to cut with a knife. They are not found in nature as free elements. They combine vigorously with non-metals and melt at successively lower temperatures. Most are less dense than .
- Group 2 elements are known as the metals. They have a group configuration of ns2. They have a pair of in the outermost s level. They are harder, denser, and stronger than alkali metals. They have melting points and are reactive than alkali metals. However, they are still too reactive to be found in nature.

B. Hydrogen and Helium. Hydrogen exhibits the ns1 configuration, but it does not share the Group 1 . Hydrogen is unique, and has its own properties. Helium exhibits the ns2 configuration, but it does not share the Group 2 . It exhibits the noble gas properties of Group . This is due to its highest electron shell being full of two at electrons.

C. The d-Block Elements (Groups 3-12). These are known as the elements. The electrons are configured so that filling of ns is followed by additions to sublevel, or (n-1)dxns2, where x is an integer form 1 to 10 and corresponds to Group 3-12, respectively. There are some exceptions to this rule (nickel, palladium and platinum). Notice in each case the sum of the outer s and d electrons is equal to the number. d-Block elements have typical metallic properties (high luster, good of electricity, ). They are typically reactive than alkali and alkaline earth elements. Some are so unreactive they do not from compounds easily. Some exist in nature as free elements. Palladium, platinum, and gold are among the reactive of all.

D. The p-Block Elements (Groups 13-18). This group contains 2 electrons in the ns sublevel. Electrons enter the np sublevels. They have a group configuration of ns2npx, where x = integers 1- 6 corresponding to Groups 13-18. In Group 18, the [?] configuration of ns2np6 is reached. The total number of electrons in the highest occupied levels is equal to the number, minus . The groups contain , metalloids, and non-metals. Group 17 contains . Halogens are the most of the non-metals. They react vigorously with most metals to form . With electrons in the outermost energy level, they are almost noble. Along with the s-Block elements, the p-Block elements are known as elements.

E. The f-Block Elements. (Lanthanides and ). Lanthanides are between Groups 3 and 4 and involve filling the 4f sublevel. There are [?] elements in the block. They are shiny similar to Group 2 in reactivity. The Actinides are also between Groups 3 and 4 and involve filling the sublevel. These 14 elements are all . The first 4 are found naturally occurring; all the others are laboratory-made.

Electron Configuration and Periodic Properties. We have learned that the elements are arranged in the periodic table according to their atomic and that there is a rough correlation between the arrangement of the elements and their configurations. Let's look ate some characteristics of the elements and their period and group trends.

A. Atomic Radii. The size of an atom comes from the electron cloud around the . However, the boundary is fuzzy. If we look at two atoms bounded together and divide by two, the picture is clearer. So, atomic radii are defined as the distance between the of identical atoms that are bonded together.
- Period Trends: Believe it or not, the trend is to smaller atoms across a period because of the increasing charge of the nucleus. (This is due to "Coulombic Attraction" - attraction of the nucleus for the electrons). Electrons in the same shell don't have as much effect on one another as electrons from different shells. They don't "". Electron shielding becomes important when we talk about the effects of electrons from different shells, but moving across a period, electrons are added within the same shell and don't each other very well. Thus, shielding effect is within a period but increases within groups from top to bottom.
- Group Trends: As electrons occupy sublevels in successively higher main energy levels located further from the , the of the atoms increase. In general, the atomic radii of the main-group elements increase a group due to shell addition. is now important. Each shell shields the more distant shells from the nucleus and cancels some of the charge felt by the outer electrons. As you would expect, the valence electrons get farther and farther away form the nucleus, and so the atomic increases as we move down a group.

B. Ionization Energy. An electron can be removed from an atom is enough is supplied. "A+" represents an ion of element A with a single positive charge, referred to as a 1+ ion. An is an atom or group of bonded atoms that has a positive or negative charge. (Positive ions are and negative ions are .) For example, sodium forms an Na+ ion. Any process that results in the formation of an ion is referred to as . To compare the ease with which atoms of different elements give up electrons, chemists compare ionization . The energy required to remove the most loosely held electron from a neutral atom of an element is the ionization energy (or first ionization energy). Elements with ionization energy form positive ions (cations) easily. Elements with ionization energy form negative ions (anions) easily. Measurements of ionization energy are made on isolated atoms in the gas phase and are expressed in kilojoules per mole (kJ/mol).
- Period Trends: In general, ionization energies of the main-group elements across each period. There is greater coulombic attraction, making it harder to remove electrons.
- Group Trends: Among the main-group elements, ionization energies generally down the groups. The size increases as you go down a group, making it easier to remove the outer shell electrons, because they are further from the .
- Removing Electrons from Positive Ions. Energies required for removing additional from positive ions are referred to as secondary ionization energy, third ionization energy, etc. Each requires more and more energy due to the increasingly nuclear charge as successive electrons are removed. As electrons are removed, ionization energy gradually until a shell is empty-then it makes a big jump. That's because the next electron must come from a shell much closer to the nucleus.

C. Electron Affinity. Neutral atoms can also acquire electrons. The energy change that occurs when an electron is acquired by a neutral atom is called the atom's electron . Affinity means to like something. Electron affinity tells us how much an element likes to take on an extra electron by measuring the energy change when an electron is added. Most atoms energy when they acquire an electron.
The quantity of energy released is represented by a negative number. However, some atoms must be forced to gain an electron. The quantity of energy absorbed is represented by a number. An ion produced this way will be unstable and will lose the added electron spontaneously.
- Period Trends: (Group 17) gain electrons most readily.
- Group Trends: In general, electrons add with increasing difficulty as you go a group.

D. Ionic Radii. A positive ion is known as a cation and the formation of a cation the atomic radius because the removal of an electron results in smaller electron cloud (removes the entire outer shell!) and the effect of coulombic attraction. A negative ion is known as an anion and leads to an increase in the atomic radius due to the of an electron. The electrons are not drawn to the nucleus as strongly as before and the electron cloud spreads because of greater repulsion between the increased number of electrons. (Full shell, electrons repel!)
- Period Trends: Metals at the left tend to form cations and non-metals anions. Both cationic and anionic radii across a period.
- Group Trends: In both cations and anions there is a slight down a group. There is a big jump in ion size as you go a period because you go from cations (that are smaller that the neutral atom) to forming anions (that are larger than the neutral atoms).

E. Valence Electrons. Chemical compounds form because electrons are lost, , or shared between atoms. The electrons that interact in this manner are those in the highest energy levels. These are the electrons most subject to the influence of nearby atoms or ions. The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as electrons. Valence electrons are often located in incompletely filled main-energy levels. For example, the electron lost for the 3s sublevel of Na to form Na+ is a valence electron.

F. Electronegativity. is the attraction of an atom for electrons in a covalent bond. It's the measure of the ability of an atom in a chemical compound to attract . It is thought of in terms of [?] atoms in a bond and refers to how strongly the nucleus of an atom attracts the electrons of other atoms in a bond. For example, in the molecule HF, fluorine and hydrogen share and electron pair in a bond. But the fluorine atom attracts more than it fair share of the electron pair, and we say that fluorine is more that hydrogen. The most electronegative element is fluorine, and was arbitrarily assigned a number of . Values for the other elements are assigned based on their relation to this value.
- Trends: Electronegativity tends to across each period. There are exceptions and tends to or remain the same down a group. (See sample problem 5.7)

G. Periodic Properties of the d-Block and f-Block Elements. Properties of these elements vary less and with less regularity than main-group elements. Atomic radii generally across periods. Ionization energies generally across the periods and increase down groups. Electronegativity generally increases as radii decreases. The order in which electrons are removed is exactly reverse of the order given by the electron - notation. Removal is from the outermost s sublevel and results in formation of 1+, 2+, and 3+ ions.