Difficulty: Intermediate | Prerequisites: Atomic Structure and Electron Configurations (Ch. 4–5)
The periodic table is the map of chemistry. Its structure mirrors electron configurations: where an element sits tells you its valence electrons, its likely bonding behaviour and how its properties compare with its neighbours. This set of topics connects the table's layout to periodic trends (atomic radius, ionisation energy, electronegativity) and then to the two major types of chemical bonding, ionic and covalent. If you understand this material, naming compounds and predicting reaction behaviour in later chapters becomes far more straightforward.
The periodic table is arranged by increasing atomic number. Periods correspond to principal energy levels, and groups share the same number of valence electrons. Atomic radius increases down a group and decreases across a period; ionisation energy and electronegativity follow the opposite trend. Metals lose electrons to form cations, nonmetals gain electrons to form anions, and both types of ion seek the stable electron configuration of a noble gas (the octet rule). Ionic bonds form between metals and nonmetals; covalent bonds form between nonmetals sharing electrons, and bond polarity depends on the electronegativity difference.
Representative elements (Group A elements)
The elements in groups 1A through 8A of the periodic table. Their outermost electrons fill s and p orbitals.
Transition metals (Group B elements)
The elements in the d-block (groups 3B through 2B). Their distinguishing electrons fill d orbitals. Examples: copper, iron, zinc.
Period
A horizontal row in the periodic table. Each period corresponds to a principal energy level.
Group
A vertical column in the periodic table. Elements in the same group share similar chemical properties because they have the same number of valence electrons.
Atomic radius
Half the distance between the nuclei of two bonded atoms of the same element. In simple terms, it measures how big an atom is.
Ionisation energy
The energy required to remove an electron from a gaseous atom. Higher ionisation energy means the atom holds its electrons more tightly.
Electronegativity
A measure of an atom's ability to attract shared electrons in a chemical bond. Fluorine has the highest electronegativity; caesium has the lowest.
Cation
A positively charged ion, formed when an atom loses one or more electrons. Metals form cations.
Anion
A negatively charged ion, formed when an atom gains one or more electrons. Nonmetals form anions. An anion is larger than its neutral atom.
Valence electrons
The electrons in the outermost energy level of an atom. They determine bonding behaviour.
Octet rule
Atoms tend to gain, lose or share electrons to achieve a set of eight valence electrons (the electron configuration of a noble gas).
Ionic bond
The electrostatic attraction between a cation and an anion. Formed when electrons are transferred from a metal to a nonmetal.
Covalent bond
A bond formed when two atoms share one or more pairs of electrons. Formed between nonmetals.
Polar covalent bond
A covalent bond in which electrons are shared unequally because the atoms have different electronegativities.
Nonpolar covalent bond
A covalent bond in which electrons are shared equally (identical or very similar electronegativities).
Coordinate covalent bond
A covalent bond in which one atom contributes both electrons of the shared pair.
Unshared pair (lone pair)
A pair of valence electrons not involved in bonding.
Diatomic molecule
A molecule consisting of two atoms. Seven elements exist as diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂ (mnemonic: HONClBrIF).
The modern periodic table is arranged in order of increasing atomic number (not mass).
Each period (horizontal row) corresponds to a principal energy level. Phosphorus is in period 3; magnesium is in the same period.
Each group (vertical column) shares the same number of valence electrons and similar chemistry.
Representative elements = Group A (1A–8A). Their valence electrons fill s and p orbitals.
Transition metals = Group B. Their distinguishing electrons fill d orbitals. Copper (Cu) is a transition metal; caesium (Cs) is not.
Inner transition metals (lanthanides, actinides) fill f orbitals.
Elements filling p orbitals are classified in groups 3A through 8A.
Metals are good conductors, malleable, ductile, tend to lose electrons.
Nonmetals are poor conductors of electricity. Kr (krypton) is a nonmetal among Pt, V, Li and Kr.
Metalloids have properties intermediate between metals and nonmetals.
The electron that plays the greatest part in determining chemical properties is the electron (specifically, valence electrons).
Al, Mg, Li are all representative elements. Cu, Co, Cd and Ni, Fe, Zn are groups that include transition metals.
The element with configuration 1s² 2s² 2p⁶ 3s² 3p² is silicon (14 electrons).
The gallium ion (Ga³⁺, Z = 31, loses 3 electrons) has the configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰.
Atomic radius:
Increases going down a group (more energy levels, electrons are further from the nucleus).
Decreases going left to right across a period (more protons pull electrons closer).
In the second period, lithium has the largest atomic radius (leftmost). Neon has the smallest.
Among S, Cl, Se, Br: chlorine has the smallest atomic radius (highest and furthest right in the periodic table).
Ionisation energy:
The energy needed to remove an electron from a gaseous atom.
Increases across a period (left to right) and decreases down a group.
Electronegativity:
Increases across a period (left to right) and decreases down a group.
Fluorine has the highest electronegativity. Caesium has the lowest.
Elements on the right side of a period have higher electronegativities than those on the left.
Ion size:
A cation is smaller than its neutral atom (lost an electron, less electron-electron repulsion).
An anion is larger than its neutral atom (gained an electron, more repulsion).
Valence electrons:
Phosphorus (group 5A) has 5 valence electrons.
Metals in groups 1A, 2A and 3A lose electrons when they form ions.
The octet rule states that atoms tend to have the electron configuration of a noble gas (8 valence electrons, or 2 for elements near helium).
Ionic bonds form when a metal transfers electrons to a nonmetal. The resulting compound is called a salt.
Aluminium oxide has the formula Al₂O₃ (Al³⁺ and O²⁻ cross charges).
Properties of ionic compounds:
High melting points (potassium chloride has a relatively high melting temperature).
Solid at room temperature.
Conduct electricity when melted or dissolved in water.
Composed of metallic and nonmetallic elements.
"Low melting points" is NOT a characteristic of most ionic compounds.
Seven elements exist as diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂. Noble gases (Ne, He, Ar, etc.) do NOT form diatomic molecules.
Atoms share electrons in covalent bonds to attain a noble-gas electron configuration.
Unshared pairs (lone pairs) are pairs of valence electrons not involved in bonding.
O₂ is joined by a double covalent bond. N₂ has a triple bond. Cl₂ has a single bond.
A coordinate covalent bond is one in which a single atom contributes both bonding electrons.
A bond between a silicon atom and an oxygen atom is likely polar covalent (different electronegativities, but not different enough for a full ionic bond).
Among H–F, H–C, H–H and H–N: H–F is the most polar covalent bond (fluorine has the highest electronegativity).
H–H is nonpolar (identical atoms).
Water is a polar molecule. When placed between charged plates, the hydrogen region (partially positive) is attracted to the negative plate, and the oxygen region (partially negative) is attracted to the positive plate.
Electronegativity differences determine whether a material is a salt (ionic, like table salt), a polar molecule (like water) or a nonpolar substance (like cooking oil). This is why salt dissolves in water but oil does not.
Students often think atomic radius increases across a period. It decreases (more protons pull electrons inward).
"Low melting point" is a common wrong answer for ionic compounds. Ionic compounds have high melting points.
Neon does not form a diatomic molecule. It is a noble gas and exists as single atoms.
A cation is smaller than its parent atom, not larger. Losing electrons reduces electron-electron repulsion.
Potassium is in period 4, not period 3, despite having 3s and 3p electrons in its configuration. Its outermost electron is in 4s.
⚠️ Know the trend directions: atomic radius increases down a group, decreases across a period. Ionisation energy and electronegativity go the opposite way.
⚠️ Fluorine = highest electronegativity. Caesium = lowest.
⚠️ Be able to identify representative elements vs transition metals from a list of element symbols.
⚠️ Phosphorus has 5 valence electrons (group 5A).
⚠️ The octet rule refers to the noble-gas configuration, not literally 8 protons or 8 electrons total.
⚠️ H–F is the most polar covalent bond among the common pairings tested.
⚠️ Know the seven diatomic elements: H, N, O, F, Cl, Br, I.
True or False: Atomic radius increases from left to right across a period.
Fill in the blank: The element with the highest electronegativity is ______.
True or False: Neon exists as a diatomic molecule.
Fill in the blank: A positively charged ion is called a ______.
True or False: An anion is smaller than its neutral atom.
Q: How does atomic radius change from top to bottom in a group?
A: It increases. Each period adds a new energy level, placing electrons further from the nucleus.
Q: Which of the following is a transition metal: caesium, copper, tellurium, tin?
A: Copper (Cu). It is in the d-block.
Q: What is the charge of a cation?
A: Positive. A cation has lost one or more electrons.
Q: How many valence electrons does phosphorus have?
A: 5 (group 5A).
Q: Which of the following covalent bonds is most polar: H–F, H–C, H–H, H–N?
A: H–F. Fluorine's high electronegativity creates the largest difference.
Q: Why do atoms share electrons in covalent bonds?
A: To attain a noble-gas electron configuration.
Q: Which diatomic molecule is joined by a double covalent bond?
A: O₂.
Periodic trends explain why certain elements form ionic compounds and others form covalent ones, which connects to chemical nomenclature (Ch. 9). The octet rule and bonding concepts underpin molecular geometry (VSEPR) and intermolecular forces, if your course covers them. Electronegativity differences also predict whether a reaction is likely to occur in the activity series (Ch. 11).
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