7.5 Periodic Trends, Chemical Bonding, Reactions, Acids and Bases

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7.5 Periodic Trends, Chemical Bonding, Reactions, Acids and Bases

Key Takeaways Explain This Topic AI

Atomic radius and metallic character both increase down a group and decrease across a period, while ionisation energy and electronegativity do the exact opposite and peak at the top right, with fluorine highest at 3.98.

An electronegativity difference below about 0.4 gives a non-polar covalent bond, roughly 0.4 to 1.7 gives a polar covalent bond, and above about 1.7 gives an ionic bond.

Balancing an equation may change coefficients only — altering a subscript turns the substance into a different compound, as when H2O becomes hydrogen peroxide.

One mole is 6.022 x 10^23 particles, so 48 g of magnesium at 24 g/mol is 2 mol and yields 2 mol, or 80 g, of magnesium oxide in 2Mg + O2 gives 2MgO.

pH is a base-10 logarithm, so vinegar at pH 3 holds one thousand times the hydrogen-ion concentration of a sample at pH 6, and 'strong' describes full ionisation rather than high concentration.

Last updated: August 2026

7.5 Periodic Trends, Chemical Bonding, Reactions, Acids and Bases

Section 7.4 established what an atom contains and how many valence electrons it carries. This section is about what atoms do with those electrons: how the periodic table predicts behaviour, how bonds form, how reactions are written and quantified, and how acids and bases are recognised. The AdUCET (Adamson University College Entrance Test) science items at this level are conceptual and small-number, so every calculation below is designed to be done in your head.

How the periodic table is organised

Elements are laid out in order of increasing atomic number, and the layout itself is the information.

Periods are the seven horizontal rows. The period number tells you how many electron shells are occupied — a Period 3 element has three.

Groups (families) are the eighteen vertical columns. Members share a valence-electron count and therefore share chemistry. For main-group elements, Groups 1 and 2 have one and two valence electrons, and Groups 13 to 18 have three through eight.

Blocks describe which sub-level the outermost electrons occupy: the s block (Groups 1 and 2, plus helium), the p block (Groups 13 to 18), the d block (the transition metals, Groups 3 to 12) and the f block (lanthanides and actinides, printed separately beneath the main table).

The named families

Family Group Valence electrons Signature behaviour

Alkali metals 1 1 soft, low-density, the most reactive metals; form + 1 +1 + 1 ions; react vigorously with water; reactivity increases downward (Cs > K > Na > Li)

Alkaline earth metals 2 2 harder and less reactive than Group 1; form + 2 +2 + 2 ions; magnesium burns with a blinding white flame

Transition metals 3–12 variable dense, high-melting, multiple oxidation states, coloured compounds, widely used as catalysts (Fe, Cu, Zn)

Halogens 17 7 the most reactive non-metals; form − 1 -1 − 1 ions; reactivity decreases downward (F > Cl > Br > I)

Noble gases 18 8 (helium 2) outer shell already complete, so essentially unreactive

Hydrogen is printed above Group 1 for convenience but is a non-metal, not an alkali metal. Test writers like that detail.

Metals, non-metals and metalloids

A staircase line runs from boron down to astatine and splits the p block. To its left sit the metals — lustrous, malleable, ductile, good conductors, and inclined to lose electrons. To its right sit the non-metals — dull, brittle as solids, poor conductors, and inclined to gain or share electrons. Sitting on the line are the six common metalloids : boron, silicon, germanium, arsenic, antimony and tellurium. They behave in between, and silicon's semiconducting behaviour is the reason every phone has a chip in it.

The four periodic trends

Two forces explain all four. Effective nuclear charge is the net pull the protons exert on the outermost electrons; shielding is the screening done by filled inner shells. Move right along a period and protons are added without adding a shell, so the pull tightens. Move down a group and an entire shell is added, so the outer electrons sit farther out and are better shielded.

• Trend Across a period, left to right Down a group Underlying reason
• Atomic radius decreases increases added protons pull the same shell inward; a new shell adds distance

Ionisation energy (energy to remove the outermost electron) increases decreases a tighter, closer electron is harder to strip off

Electronegativity (pull on the electrons shared in a bond ) increases decreases the same nuclear pull, now acting on a shared pair

Metallic character (readiness to give electrons away) decreases increases the easier an electron leaves, the more metallic the element

The clean way to hold this: atomic radius and metallic character move together , both peaking at the bottom-left of the table, while ionisation energy and electronegativity move together in the opposite direction , peaking at the top-right. Fluorine is the most electronegative element at 3.98 on the Pauling scale; helium has the highest first ionisation energy. Noble gases are normally excluded from radius and electronegativity comparisons.

A quick application: rank sodium, magnesium and potassium by atomic radius. Sodium and magnesium share Period 3, and magnesium is further right, so magnesium is smaller. Potassium sits one period below sodium, so it is larger. The order is Mg < Na < K.

Bonding

The octet rule states that atoms gain, lose or share electrons until they hold eight valence electrons — a noble-gas arrangement. Hydrogen, helium, lithium and beryllium settle for two instead, the so-called duet. Which route an atom takes is predicted by the electronegativity difference , Δ E N \Delta EN Δ EN , between the two partners.

Δ E N \Delta EN Δ EN (Pauling) Bond type Worked example

0 to about 0.4 non-polar covalent Cl–Cl gives 3.16 − 3.16 = 0 3.16 - 3.16 = 0 3.16 − 3.16 = 0 ; C–H gives 2.55 − 2.20 = 0.35 2.55 - 2.20 = 0.35 2.55 − 2.20 = 0.35

about 0.4 to 1.7 polar covalent H–Cl gives 3.16 − 2.20 = 0.96 3.16 - 2.20 = 0.96 3.16 − 2.20 = 0.96 ; O–H gives 3.44 − 2.20 = 1.24 3.44 - 2.20 = 1.24 3.44 − 2.20 = 1.24

above about 1.7 ionic Na–Cl gives 3.16 − 0.93 = 2.23 3.16 - 0.93 = 2.23 3.16 − 0.93 = 2.23 ; Mg–O gives 3.44 − 1.31 = 2.13 3.44 - 1.31 = 2.13 3.44 − 1.31 = 2.13

Textbooks place these cut-offs slightly differently; treat them as a guide, not a law, and read the pattern rather than the decimal.

Ionic bonding

Metal plus non-metal, with electrons transferred outright. Sodium (2, 8, 1) hands its single valence electron to chlorine (2, 8, 7), producing Na + \text{Na}^+ Na + (2, 8) and Cl − \text{Cl}^- Cl − (2, 8, 8) — both now noble-gas-like. The product is not a molecule but a giant three-dimensional crystal lattice of alternating charges. Its properties follow directly:

High melting and boiling points , because millions of strong electrostatic attractions must be overcome.

Hard yet brittle. A sharp blow shifts one layer, bringing like charges face to face, and mutual repulsion splits the crystal.

Conducts only when molten or dissolved. In the solid, ions are locked; charge cannot flow until they are free to move.

Generally soluble in water , a polar solvent that surrounds and separates the ions.

Covalent bonding

Two non-metals share electron pairs instead of transferring them: one shared pair is a single bond ( H 2 \text{H}_2 H 2 ​ ), two is a double bond ( O 2 \text{O}_2 O 2 ​ , CO 2 \text{CO}_2 CO 2 ​ ), three is a triple bond ( N 2 \text{N}_2 N 2 ​ ). The result is a discrete molecule . Attractions between molecules are weak, so molecular substances melt and boil at low temperatures, conduct poorly, and are often insoluble in water when non-polar.

A crucial distinction: a bond is polar when Δ E N \Delta EN Δ EN is large enough, but a molecule is polar only when its polar bonds fail to cancel by symmetry. Carbon dioxide holds two polar C=O bonds pointing in exactly opposite directions along a straight line, so they cancel and the molecule is non-polar. Water holds two comparable polar bonds, but the molecule is bent, so they do not cancel and water is strongly polar. That contrast is one of the most frequently examined ideas in the whole topic.

Shapes follow from electron-pair repulsion: pairs around a central atom spread as far apart as they can, and lone pairs push harder than bonding pairs.

• Molecule Bonding pairs / lone pairs Shape Approximate bond angle
• CO 2 \text{CO}_2 CO 2 ​ 2 / 0 linear 180 ∘ 180^\circ 18 0 ∘
• BF 3 \text{BF}_3 BF 3 ​ 3 / 0 trigonal planar 120 ∘ 120^\circ 12 0 ∘
• CH 4 \text{CH}_4 CH 4 ​ 4 / 0 tetrahedral 109.5 ∘ 109.5^\circ 109. 5 ∘
• NH 3 \text{NH}_3 NH 3 ​ 3 / 1 trigonal pyramidal 107 ∘ 107^\circ 10 7 ∘
• H 2 O \text{H}_2\text{O} H 2 ​ O 2 / 2 bent 104.5 ∘ 104.5^\circ 104. 5 ∘
• Metallic bonding

A metal is a lattice of cations immersed in a delocalised sea of valence electrons that belong to no single atom. Mobile electrons carry charge and heat, giving high conductivity, and they reflect light, giving lustre. They also explain malleability and ductility : when a metal is hammered, layers of cations slide over one another and the electron sea simply flows around the new arrangement, so the metal deforms. Strike an ionic crystal the same way and it shatters, because sliding brings like charges together.

Hydrogen bonding and the oddity of water

A hydrogen bond is not a true bond but an unusually strong attraction between molecules, arising whenever hydrogen is bonded to nitrogen, oxygen or fluorine and is drawn towards a lone pair on a neighbouring N, O or F. Almost every strange thing about water traces back to it:

Water boils at 100 °C, while hydrogen sulfide — a heavier molecule with no hydrogen bonding — boils at about −60 °C.

Ice is less dense than liquid water (about 0.92 against 1.00 g/cm³) because the hydrogen-bond network freezes molecules into an open, roomy pattern. Ice therefore floats.

Water's high specific heat capacity of 4.18 J per gram per degree lets coastal air stay mild, and its high surface tension lets a water strider stand on a pond.

It dissolves ionic and polar substances so readily that it is called the universal solvent.

Formulas and names

To build an ionic formula, use the criss-cross rule : the size of each ion's charge becomes the other ion's subscript, then reduce to the simplest ratio.

Ca 2 + \text{Ca}^{2+} Ca 2 + with Cl − \text{Cl}^- Cl − gives CaCl 2 \text{CaCl}_2 CaCl 2 ​ .

Al 3 + \text{Al}^{3+} Al 3 + with O 2 − \text{O}^{2-} O 2 − gives Al 2 O 3 \text{Al}_2\text{O}_3 Al 2 ​ O 3 ​ .

Mg 2 + \text{Mg}^{2+} Mg 2 + with O 2 − \text{O}^{2-} O 2 − gives Mg 2 O 2 \text{Mg}_2\text{O}_2 Mg 2 ​ O 2 ​ , which reduces to MgO \text{MgO} MgO .

Na + \text{Na}^+ Na + with SO 4 2 − \text{SO}_4^{2-} SO 4 2 − ​ gives Na 2 SO 4 \text{Na}_2\text{SO}_4 Na 2 ​ SO 4 ​ .

Ca 2 + \text{Ca}^{2+} Ca 2 + with PO 4 3 − \text{PO}_4^{3-} PO 4 3 − ​ gives Ca 3 ( PO 4 ) 2 \text{Ca}_3(\text{PO}_4)_2 Ca 3 ​ ( PO 4 ​ ) 2 ​ — brackets are required whenever a polyatomic ion is multiplied.

Worth memorising: nitrate NO 3 − \text{NO}_3^- NO 3 − ​ , sulfate SO 4 2 − \text{SO}_4^{2-} SO 4 2 − ​ , carbonate CO 3 2 − \text{CO}_3^{2-} CO 3 2 − ​ , phosphate PO 4 3 − \text{PO}_4^{3-} PO 4 3 − ​ , hydroxide OH − \text{OH}^- OH − , bicarbonate HCO 3 − \text{HCO}_3^- HCO 3 − ​ and ammonium NH 4 + \text{NH}_4^+ NH 4 + ​ .

Naming follows the bond type:

Ionic, fixed-charge metal — metal name, then the non-metal stem with -ide : NaCl is sodium chloride, MgO is magnesium oxide.

Ionic, variable-charge metal — a Roman numeral states the charge: FeCl 2 \text{FeCl}_2 FeCl 2 ​ is iron(II) chloride, FeCl 3 \text{FeCl}_3 FeCl 3 ​ is iron(III) chloride.

Ionic with a polyatomic ion — keep the ion's own name: CaCO 3 \text{CaCO}_3 CaCO 3 ​ is calcium carbonate, ( NH 4 ) 2 SO 4 (\text{NH}_4)_2\text{SO}_4 ( NH 4 ​ ) 2 ​ SO 4 ​ is ammonium sulfate.

Covalent, two non-metals — Greek prefixes count the atoms, and mono- is dropped on the first element: CO is carbon monoxide, CO 2 \text{CO}_2 CO 2 ​ is carbon dioxide, N 2 O 4 \text{N}_2\text{O}_4 N 2 ​ O 4 ​ is dinitrogen tetroxide, SF 6 \text{SF}_6 SF 6 ​ is sulfur hexafluoride.

Prefixes never appear in ionic names. "Magnesium dichloride" is wrong; the formula MgCl 2 \text{MgCl}_2 MgCl 2 ​ is fixed by the charges, so no counting word is needed.

Balancing equations

Because mass is conserved, every element must appear in equal numbers on both sides. You may adjust coefficients in front of formulas; you may never change a subscript, since that would change the substance itself.

A reliable order: balance metals first, then non-metals other than oxygen and hydrogen, then hydrogen, and leave oxygen for last. If a fraction appears, multiply everything through.

Worked example 1 — burning ethane, C 2 H 6 + O 2 → CO 2 + H 2 O \text{C}_2\text{H}_6 + \text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O} C 2 ​ H 6 ​ + O 2 ​ → CO 2 ​ + H 2 ​ O .

Carbon: 2 on the left, so place 2 in front of CO 2 \text{CO}_2 CO 2 ​ .

Hydrogen: 6 on the left, so place 3 in front of H 2 O \text{H}_2\text{O} H 2 ​ O .

Oxygen on the right now totals ( 2 × 2 ) + 3 = 7 (2 \times 2) + 3 = 7 ( 2 × 2 ) + 3 = 7 atoms, which needs 7 2 O 2 \tfrac{7}{2}\text{O}_2 2 7 ​ O 2 ​ .

Clear the fraction by doubling every coefficient.

Final answer: 2 C 2 H 6 + 7 O 2 → 4 CO 2 + 6 H 2 O 2\text{C}_2\text{H}_6 + 7\text{O}_2 \rightarrow 4\text{CO}_2 + 6\text{H}_2\text{O} 2 C 2 ​ H 6 ​ + 7 O 2 ​ → 4 CO 2 ​ + 6 H 2 ​ O

Check: carbon 4 = 4 4 = 4 4 = 4 ; hydrogen 12 = 12 12 = 12 12 = 12 ; oxygen 14 = 8 + 6 = 14 14 = 8 + 6 = 14 14 = 8 + 6 = 14 .

The five reaction types

Type Pattern Balanced example

Synthesis (combination) A + B → AB \text{A} + \text{B} \rightarrow \text{AB} A + B → AB 2 Mg + O 2 → 2 MgO 2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO} 2 Mg + O 2 ​ → 2 MgO

Decomposition AB → A + B \text{AB} \rightarrow \text{A} + \text{B} AB → A + B CaCO 3 → CaO + CO 2 \text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2 CaCO 3 ​ → CaO + CO 2 ​

Single replacement A + BC → AC + B \text{A} + \text{BC} \rightarrow \text{AC} + \text{B} A + BC → AC + B Zn + 2 HCl → ZnCl 2 + H 2 \text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2 Zn + 2 HCl → ZnCl 2 ​ + H 2 ​

Double replacement AB + CD → AD + CB \text{AB} + \text{CD} \rightarrow \text{AD} + \text{CB} AB + CD → AD + CB AgNO 3 + NaCl → AgCl + NaNO 3 \text{AgNO}_3 + \text{NaCl} \rightarrow \text{AgCl} + \text{NaNO}_3 AgNO 3 ​ + NaCl → AgCl + NaNO 3 ​

Combustion fuel + O 2 → CO 2 + H 2 O + \text{ O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O} + O 2 ​ → CO 2 ​ + H 2 ​ O CH 4 + 2 O 2 → CO 2 + 2 H 2 O \text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O} CH 4 ​ + 2 O 2 ​ → CO 2 ​ + 2 H 2 ​ O

Recognise them by counting substances: one product means synthesis, one reactant means decomposition, a lone element swapping into a compound means single replacement, and two compounds trading partners means double replacement. Burning a fuel in oxygen is keyed as combustion even though it is also an oxidation.

The mole and simple stoichiometry

A mole is 6.022 × 10 23 6.022 \times 10^{23} 6.022 × 1 0 23 particles, a figure called Avogadro's number . It is defined so that one mole of an element weighs, in grams, the same number that appears as its atomic mass — the molar mass , in g/mol. Water is 2 ( 1 ) + 16 = 18 2(1) + 16 = 18 2 ( 1 ) + 16 = 18 g/mol; calcium carbonate is 40 + 12 + 48 = 100 40 + 12 + 48 = 100 40 + 12 + 48 = 100 g/mol; glucose is 72 + 12 + 96 = 180 72 + 12 + 96 = 180 72 + 12 + 96 = 180 g/mol. x molar mass (g/mol) x 6.022 x 10^23

• MASS (g) <----------------------- MOLES -----------------------> PARTICLES
• ----------------------> <-----------------------
• / molar mass / 6.022 x 10^23

So 36 g of water is 36 ÷ 18 = 2 36 \div 18 = 2 36 ÷ 18 = 2 mol, which is 1.2 × 10 24 1.2 \times 10^{24} 1.2 × 1 0 24 molecules.

Worked example 2 — from mass of reactant to mass of product. Magnesium ribbon burns in air by 2 Mg + O 2 → 2 MgO 2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO} 2 Mg + O 2 ​ → 2 MgO . How much magnesium oxide forms from 48 g of magnesium? Use Mg = 24 g/mol and O = 16 g/mol.

Convert mass to moles: 48 ÷ 24 = 2 48 \div 24 = 2 48 ÷ 24 = 2 mol of magnesium.

Read the mole ratio from the coefficients: 2 Mg produces 2 MgO, a ratio of 1 : 1, so 2 mol of magnesium oxide form.

Find the molar mass of the product: 24 + 16 = 40 24 + 16 = 40 24 + 16 = 40 g/mol.

Convert moles back to mass: 2 × 40 = 80 2 \times 40 = \mathbf{80} 2 × 40 = 80 g.

Verify with conservation of mass: 2 mol of Mg weighs 48 g and the 1 mol of O 2 \text{O}_2 O 2 ​ consumed weighs 32 g, and 48 + 32 = 80 48 + 32 = 80 48 + 32 = 80 g. The extra mass came from the air.

Acids, bases and pH

Two definitions matter at this level.

Arrhenius : an acid releases H + \text{H}^+ H + in water; a base releases OH − \text{OH}^- OH − . Simple, but confined to water.

Brønsted–Lowry : an acid is a proton donor ; a base is a proton acceptor . Broader, and it explains ammonia, which contains no OH group yet is plainly basic: NH 3 + H 2 O → NH 4 + + OH − \text{NH}_3 + \text{H}_2\text{O} \rightarrow \text{NH}_4^+ + \text{OH}^- NH 3 ​ + H 2 ​ O → NH 4 + ​ + OH − .

• Property Acid Base
• Taste (never tested in a laboratory) sour, as in calamansi and suka bitter
• Feel — slippery, soapy
• Litmus paper turns blue litmus red turns red litmus blue
• pH at 25 °C below 7 above 7
• With active metals such as Zn or Mg fizzes, releasing hydrogen gas no reaction
• With carbonates fizzes, releasing carbon dioxide no reaction
• In solution conducts electricity conducts electricity

The pH scale pH = − log ⁡ 10 [ H + ] and [ H + ] = 10 − pH \text{pH} = -\log_{10}[\text{H}^+] \qquad \text{and} \qquad [\text{H}^+] = 10^{-\text{pH}} pH = − lo g 10 ​ [ H + ] and [ H + ] = 1 0 − pH

Ordinary aqueous solutions at 25 °C run from about 0 to 14, with pure water at [ H + ] = 1 × 10 − 7 [\text{H}^+] = 1 \times 10^{-7} [ H + ] = 1 × 1 0 − 7 M, giving pH 7. No calculator is needed while the coefficient is 1: read the exponent and drop the minus sign, so 1 × 10 − 4 1 \times 10^{-4} 1 × 1 0 − 4 M gives pH 4.

Because the scale is a base-10 logarithm, each whole unit is a tenfold change in hydrogen-ion concentration. Moving from pH 4 to pH 2 means a hundredfold increase; suka at pH 3 carries a thousand times the H + \text{H}^+ H + of a sample at pH 6. Approximate values worth carrying: gastric juice 1.5–2, calamansi juice 2, suka 3, black coffee 5, pure water 7, baking-soda solution 8.5, milk of magnesia 10, and apog or limewater 12.

Strong versus weak

Strong means fully ionised in water, not concentrated. Hydrochloric, nitric and sulfuric acids are strong; the ethanoic acid in vinegar and the carbonic acid in softdrinks are weak, because most of their molecules stay intact. Sodium and potassium hydroxide are strong bases; ammonia and the magnesium hydroxide in milk of magnesia are weak. A concentrated weak acid can therefore read a higher pH than a dilute strong acid, and that reversal is a favourite trap.

Neutralisation and indicators

An acid and a base cancel each other, producing a salt and water: HCl + NaOH → NaCl + H 2 O \text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O} HCl + NaOH → NaCl + H 2 ​ O Mg(OH) 2 + 2 HCl → MgCl 2 + 2 H 2 O \text{Mg(OH)}_2 + 2\text{HCl} \rightarrow \text{MgCl}_2 + 2\text{H}_2\text{O} Mg(OH) 2 ​ + 2 HCl → MgCl 2 ​ + 2 H 2 ​ O

The second is exactly what an antacid tablet does to excess stomach acid.

• Indicator In acid In base
• Litmus red blue
• Phenolphthalein colourless pink, from about pH 8.3
• Methyl orange red yellow
• Bromothymol blue yellow blue, green at neutral
• Red cabbage extract red to pink green to yellow
• Turmeric, luyang dilaw yellow reddish brown
• What examiners test

A trend stated backwards. Atomic radius grows down a group but shrinks across a period; ionisation energy does the reverse.

Every hydrogen-containing compound called an acid. Methane and ammonia contain hydrogen and neither is acidic.

Subscripts altered while balancing. Writing H 2 O 2 \text{H}_2\text{O}_2 H 2 ​ O 2 ​ to balance oxygen produces hydrogen peroxide, a different substance.

"Strong" confused with "concentrated". Strength is about the degree of ionisation; concentration is about how much is dissolved.

pH treated as linear. A drop of three pH units is a thousandfold rise in acidity, not a threefold one.

Greek prefixes on ionic compounds. CaCl 2 \text{CaCl}_2 CaCl 2 ​ is calcium chloride, never calcium dichloride.

Polar bonds assumed to make a polar molecule. Carbon dioxide has two polar bonds and no overall polarity, because its linear shape cancels them.