Atom is the smallest part of an element. It cannot be divided further into smaller parts.

Modern atomic theory:

• All matter is composed of atoms.

• Atoms of any given element share the same chemical properties while atoms of different elements have different properties.

• Chemical compounds consist of atoms combined in specific ratios. That is, only whole atoms can combine—one A atom with one B atom, or one A atom with two B atoms, and so on.

• Chemical reactions change only the way that atoms are combined in compounds. The atoms themselves are unchanged and do not disappear

Atoms are composed of tiny subatomic particles called protons, neutrons, and electrons.

The protons and neutrons are packed closely together in a dense core called the nucleus. Surrounding the nucleus, the electrons move about rapidly through a large, mostly empty volume of space.

Measurements show that the diameter of a nucleus is only about 10-15 m, whereas that of the atom itself is about 10-10 m.

Each atom has a specific number of protons, neutrons, and electrons, and the identity of the element is determined by the number of protons within the nucleus, also called the element’s atomic number (Z).

In a periodic table, elements are listed in order of increasing atomic number, beginning at the upper left and ending at the lower right.

The sum of the protons and neutrons in an atom is called the atom’s mass number (A).

Atomic number and mass number can be written using chemical symbols by showing the element’s mass number (A) as a superscript and its atomic number (Z) as a subscript in front of the atomic symbol.

Atoms with identical atomic numbers but different atomic mass are called isotopes.

Hydrogen has three isotopes:

• The most abundant hydrogen isotope, called protium, has one proton but no neutrons and thus has a mass number of 1.

• The second hydrogen isotope, called deuterium, also has one proton, but has one neutron and a mass number of 2.

• The third isotope, called tritium, has two neutrons and a mass number of 3.

Atomic mass = sum of all[ (isotopic abundance) * (isotopic mass)].

In 1869, the Russian chemist Dmitri Mendeleev organized the elements in order of increasing mass and then organized elements into groups based on similarities in chemical behaviour. His table is a forerunner of the modern periodic table. The table has boxes for each element that give the symbol, atomic number, and atomic mass of the element.

Of the 118 currently known elements

• 94 are classified as metals— aluminum, gold, copper, and zinc, for example. Metals are solid at room temperature (except for mercury), usually have a lustrous appearance when freshly cut, are good conductors of heat and electricity, and are malleable rather than brittle.

• 18 elements are non-metals. All are poor conductors of heat and electricity. Eleven are gases at room temperature, six are brittle solids, and one is a liquid. Bromine is the only liquid non-metal.

• The metalloids are located in a zigzag band between the metals on the left and non-metals on the right side of the periodic table.

Another way of classifying the elements in the periodic table is based on similarities in chemical behaviour. Beginning at the upper left corner of the periodic table, elements are arranged by increasing atomic number into seven horizontal rows, called periods, and 18 vertical columns, called groups.

• The two large groups on the far left and the six on the far right are called the main group elements and are numbered 1A through 8A.

• The 10 smaller groups in the middle of the table are called the transition metal elements and are numbered 1B through 8B.

• Alternatively, all 18 groups are numbered sequentially from 1 to 18.

• The 14 groups shown separately at the bottom of the table are called the inner transition metal elements and are not numbered.

Many elements in the periodic table show remarkable similarities and thus are clubbed together as one subgroup.

• Group 1A—Alkali metals: Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr) are shiny, soft metals with low melting points.

  • All react rapidly (often violently) with water to form products that are highly alkaline, or basic—hence the name alkali metals.

  • Because of their high reactivity, the alkali metals are never found in nature in the pure state but only in combination with other elements.

• Group 2A—Alkaline earth metals: Beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra) are also lustrous, silvery metals but are less reactive than their neighbours in group 1A. Like the alkali metals, the alkaline earths are never found in nature in the pure state.

• Group 7A—Halogens: Fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At) are colourful and corrosive non-metals. All are found in nature only in combination with other elements, such as with sodium in table salt (sodium chloride, NaCl). In fact, the group name halogen is taken from the Greek word hals, meaning salt.

• Group 8A—Noble gases: Helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn) are colourless gases.

  • The elements in this group were labeled the “noble” gases because of their lack of chemical reactivity—helium, neon, and argon do not combine with any other elements, whereas krypton and xenon combine with a very few.

Different electrons have different amounts of energy and thus occupy different regions within the atom. Furthermore, the energies of electrons are quantized or restricted to having only certain values.

The electrons in an atom are grouped around the nucleus into shells, according to the energy of the electrons. The shell is designated using the letter n; n = 1 for the first shell (period 1), n = 2 for the second shell (period 2), and so on.

• The farther a shell is from the nucleus, the larger it is, the more electrons it can hold, the higher the energies of those electrons, and thus the easier they are to remove because they are the farthest away from the positively charged nucleus.

• The first shell (the one nearest the nucleus) can hold only 2 electrons, the second shell can hold 8, the third shell can hold 18, and the fourth shell can hold 32 electrons.

Within shells, electrons are further grouped into subshells of four different types, identified in order of increasing energy by the letters s, p, d, and f. The first shell has only one subshell, s. The second shell has two subshells: an s subshell and a p subshell. The third shell has an s, p, and d subshell. The fourth shell has an s, p, d, and f subshell.

Within each subshell, electrons are grouped into orbitals, regions of space within an atom where the specific electrons are most likely to be found. There are different numbers of orbitals within the different kinds of subshells.

• A given s subshell has only one orbital, a p subshell has three orbitals, a d subshell has five orbitals, and an f subshell has seven orbitals.

• Each orbital can hold only two electrons, which differ in a property known as spin.

summarizing the electron distribution theory, we get-

• The first shell has a maximum capacity of only two electrons. The two electrons have different spins and are in a single 1s orbital.

• The second shell has a maximum capacity of eight electrons. Two are in a 2s orbital, and 6 are in the three different 2p orbitals (two per 2p orbital).

• The third shell has a maximum capacity of 18 electrons. Two are in a 3s orbital, 6 are in three 3p orbitals, and 10 are in five 3d orbitals.

• The fourth shell has a maximum capacity of 32 electrons. Two are in a 4s orbital, 6 are in three 4p orbitals, 10 are in five 4d orbitals, and 14 are in seven 4f orbitals.

The exact arrangement of electrons in an atom’s shells and subshells is called the atom’s electron configuration and can be predicted by applying three rules:

• RULE 1: Electrons occupy the lowest-energy orbitals available, beginning with 1s.

• RULE 2: Each orbital can hold only two electrons, which must be of opposite spin.

• RULE 3: Two or more orbitals with the same energy are each half-filled by one electron before any one orbital is completely filled by the addition of the second electron.

The periodic table can be divided into four regions, or blocks, of elements according to the electron shells and subshells occupied by the subshell filled last.

• The main group 1A and 2A elements on the left side of the table (plus He) are called the s-block elements because an s subshell is filled last in these elements.

• The main group 3A–8A elements on the right side of the table (except He) are the p-block elements because a p subshell is filled last in these elements.

• The transition metals in the middle of the table are the d-block elements because a d subshell is filled last in these elements.

• The inner transition metals detached at the bottom of the table are the f-block elements because an f subshell is filled last in these elements.

Because the valence electrons are the most loosely held, they are the most important in determining an element’s properties.

In an electron-dot symbol (also called Lewis symbols), dots are placed around the atomic symbol to indicate the number of valence electrons present.

• the dots are distributed around the four sides of the element symbol, singly at first until each of the four sides has one dot.

• As more electron dots are added they will form pairs, with no more than two dots on a side.