Chapter 3: Periodic Table

Scientists needed a way to organize the growing list of elements. They grouped elements with similar properties. This made the elements easier to study and compare.

Dobereiner put some similar elements into groups of three, called triads. Add the atomic masses of the first and third elements, then divide by two. The result is close to the middle element’s atomic mass.

Limitation: only a limited number of elements could be grouped into triads.

Newlands arranged elements from lower to higher atomic mass. He noticed that every eighth element often had similar properties, like repeating musical notes.

Example: lithium and sodium show similar properties.

Limitation: the rule was successful mainly for lighter elements, up to calcium.

The physical and chemical properties of elements are periodic functions of their atomic masses.

Elements with similar chemical properties were placed in the same vertical groups.

He left gaps for undiscovered elements and successfully predicted some of their properties.

Some elements did not fit a strict atomic-mass order. The table also had no clear place for isotopes.

The physical and chemical properties of elements are periodic functions of their atomic numbers.

Atomic number is the number of protons. It also gives the electron count in a neutral atom. Arranging elements by atomic number explains the repeating patterns better than arranging them by mass.

For most elements, the period number tells us the highest occupied shell in the ground state. Ground state means the lowest-energy arrangement. Some transition elements are exceptions.

Moving one place across a period adds one proton. A neutral atom also has one more electron. The outer electron arrangement changes as we move across.

Period 1 is the shortest. Periods 6 and 7 are the longest, with 32 elements each. Period 7 is complete.

Elements in the same main group have similar outer electron arrangements. This is why they usually behave in similar ways.

For main-group elements, Group 1 usually means 1 valence electron and Group 2 means 2. Groups 13–18 usually have 3–8 valence electrons. Helium is an exception: it has 2.

Example: Li (2,1), Na (2,8,1) and K (2,8,8,1) each have one valence electron, so all belong to Group 1.

Li, Na, K, Rb, Cs, Fr. These reactive metals have one valence electron and commonly lose it to form +1 ions. They generally have low melting points. Hydrogen lies above this group but is not an alkali metal.

Be, Mg, Ca, Sr, Ba, Ra. They have two valence electrons and commonly form +2 ions. They are generally less reactive than Group 1 metals in the same period.

B, Al, Ga, In, Tl. They have three valence electrons. Boron is a metalloid; the other members are metals. +3 is a common oxidation state, but not universal.

C, Si, Ge, Sn, Pb. They have four valence electrons. Carbon is a nonmetal; silicon and germanium are metalloids; tin and lead are metals. Tin and lead commonly show +2 and +4 oxidation states.

N, P, As, Sb, Bi. They have five valence electrons. Metallic character increases down the group. Common oxidation states include −3, +3 and +5.

O, S, Se, Te, Po. They have six valence electrons and are also called chalcogens. Oxygen, sulfur and selenium are nonmetals; tellurium is a metalloid and polonium is usually classified as a metal.

F, Cl, Br, I, At. They have seven valence electrons, are reactive and commonly gain one electron to form −1 ions.

He, Ne, Ar, Kr, Xe, Rn, Og. Their valence shell is filled: helium has two electrons; the others typically have eight. These elements are generally very unreactive. Oganesson’s bulk physical properties are not yet experimentally established.

Most d-block elements are metals. Many form coloured ions or complexes (a central metal joined to surrounding atoms or groups). They often have more than one oxidation state, a signed value assigned using electron-counting rules.

A transition element has an incomplete d subshell in its atom or a common positive ion.

Periodicity is the recurrence of similar physical and chemical properties when elements are arranged in increasing atomic number. It occurs because similar valence-shell configurations recur.

Across a period, the nucleus gains protons. Inner electrons do not add much extra shielding. The nucleus therefore pulls more strongly on outer electrons. Down a group, extra shells put outer electrons farther away and increase shielding.

Across a period, protons increase while added electrons enter the same main shell. The stronger pull brings the electron cloud closer. Atoms usually become smaller.

Example (approximately, covalent radii): Li > Be > B > C > N > O > F.

Down a group, each new occupied shell puts the outer electrons farther from the nucleus. Inner electrons also shield them more. Atoms become larger.

Example: Li < Na < K < Rb < Cs.

Lanthanoids slowly become smaller across their series. Their 4f electrons do not shield the nucleus well. This decrease in size is called lanthanide contraction.

Across a period, atoms usually get smaller and hold outer electrons more strongly. Removing an electron usually takes more energy.

Down a group, outer electrons are farther away and more shielded. They are usually easier to remove.

Be needs more energy than B. Be loses a 2s electron, which is held more strongly than B’s 2p electron.

N needs more energy than O. N has a half-filled 2p³ arrangement. O has a paired 2p electron, which is easier to remove.

Helium has the highest first ionization energy of all elements.

Across a period, the stronger nuclear pull often makes electron gain more favourable. Noble gases and some stable subshell arrangements are exceptions.

Down a group, greater distance and shielding usually weaken the pull on an added electron.

Cl releases more energy than F when it gains an electron. F is very small, so its 2p electrons are crowded. They push the new electron away more strongly. Cl has a larger 3p subshell with more room.

Halogen order (energy released): Cl > F > Br > I.

Across a period, the number of inner occupied shells stays nearly the same. Electrons in the same shell shield one another less well. The pull felt by outer electrons usually increases.

Down a group, there are more inner shells to shield the outer electrons. K (2,8,8,1) has more shielding than Na (2,8,1).

More shielding and greater distance make an outer electron easier to remove. This helps explain the decrease in ionization energy down a group.

Across a period, atoms usually become smaller and the nuclear pull grows stronger. They attract shared electrons more strongly.

Down a group, greater size and shielding weaken the pull on shared electrons.

Keep these ideas separate: ionization energy concerns removing an electron; electron affinity concerns adding one; electronegativity concerns pulling shared electrons in a bond.

Across a period: stronger effective nuclear attraction → smaller atoms → generally higher ionization energy and electronegativity.

Down a group: more shells and shielding → larger atoms → generally lower ionization energy and electronegativity.

Special cases matter: Be/B and N/O for first ionization energy; Cl/F for electron affinity.