Principles of Chemistry

What you'll learn

Principles of Chemistry forms the foundation of the entire CXC CSEC Chemistry syllabus. This section covers atomic structure, the periodic table, chemical bonding, formula writing, and equation balancing—skills that appear in every paper and underpin topics like acids, bases, salts, and organic chemistry. Mastery of these principles guarantees marks across multiple exam questions.

Key terms and definitions

Atom — the smallest particle of an element that can take part in a chemical reaction, consisting of a nucleus containing protons and neutrons surrounded by electrons in shells.

Element — a pure substance that cannot be broken down into simpler substances by chemical means, containing only one type of atom.

Compound — a substance formed when two or more elements are chemically combined in fixed proportions, which can only be separated by chemical means.

Ion — an atom or group of atoms that has gained or lost electrons to acquire a positive charge (cation) or negative charge (anion).

Relative atomic mass (Ar) — the average mass of an atom of an element compared to 1/12 the mass of a carbon-12 atom.

Valency — the combining power of an element, determined by the number of electrons an atom needs to lose, gain, or share to achieve a stable electron configuration.

Isotopes — atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different mass numbers.

Covalent bond — a chemical bond formed when two non-metal atoms share one or more pairs of electrons.

Core concepts

Atomic structure and notation

Every atom consists of three subatomic particles:

• Protons: positively charged particles in the nucleus, mass = 1 atomic mass unit (amu), charge = +1
• Neutrons: neutral particles in the nucleus, mass = 1 amu, charge = 0
• Electrons: negatively charged particles orbiting the nucleus in shells, mass = 1/1840 amu (negligible), charge = -1

The atomic number (Z) equals the number of protons in an atom and defines the element. The mass number (A) equals protons plus neutrons. Standard atomic notation shows mass number as a superscript and atomic number as a subscript before the element symbol: ²³₁₁Na represents a sodium atom with 11 protons and 12 neutrons.

Electron configuration describes how electrons are arranged in shells around the nucleus. The first shell holds a maximum of 2 electrons, the second holds 8, and the third holds 8 (for the first 20 elements). Examples:

• Carbon (6 electrons): 2,4
• Chlorine (17 electrons): 2,8,7
• Calcium (20 electrons): 2,8,8,2

The number of electrons in the outermost shell determines an element's chemical properties and reactivity. Elements with full outer shells (Group 0/18 noble gases) are unreactive.

The periodic table and periodic trends

The modern periodic table arranges elements in order of increasing atomic number. Elements with similar chemical properties appear in vertical columns called groups, while horizontal rows are called periods.

Key groups for CXC CSEC Chemistry:

• Group 1 (Alkali metals): lithium, sodium, potassium—highly reactive metals that form +1 ions
• Group 2 (Alkaline earth metals): magnesium, calcium—reactive metals forming +2 ions
• Group 7/17 (Halogens): fluorine, chlorine, bromine, iodine—reactive non-metals forming -1 ions
• Group 0/18 (Noble gases): helium, neon, argon—unreactive gases with full outer electron shells

Periodic trends appear as you move across periods or down groups:

1. Across a period (left to right): atomic radius decreases, metallic character decreases, non-metallic character increases
2. Down a group: atomic radius increases, reactivity of metals increases (Group 1 and 2), reactivity of non-metals decreases (Group 7)

The transition elements occupy the central block between Groups 2 and 3 and include industrially important metals like iron, copper, and zinc used in Caribbean industries including bauxite processing in Jamaica and Trinidad's Arcelor Mittal steel plant.

Chemical bonding

Atoms bond to achieve stable electron configurations, typically with 8 electrons in the outermost shell (octet rule) or 2 electrons for elements in the first period (duplet rule).

Ionic bonding occurs between metals and non-metals through electron transfer:

• Metal atoms lose electrons to form positive cations
• Non-metal atoms gain electrons to form negative anions
• Electrostatic attraction between oppositely charged ions creates the ionic bond
• Results in giant ionic lattice structures with high melting and boiling points
• Ionic compounds conduct electricity when molten or dissolved because ions are free to move

Example: Sodium chloride (NaCl)

• Sodium (2,8,1) loses 1 electron → Na⁺ (2,8)
• Chlorine (2,8,7) gains 1 electron → Cl⁻ (2,8,8)

Covalent bonding occurs between non-metal atoms through electron sharing:

• Atoms share pairs of electrons to fill outer shells
• Single bond = one shared pair, double bond = two shared pairs
• Results in simple molecular structures (low melting/boiling points) or giant covalent structures like diamond (high melting/boiling points)
• Most covalent compounds do not conduct electricity because they contain no charged particles

Example: Water (H₂O)

• Oxygen needs 2 electrons, each hydrogen needs 1
• Oxygen shares one pair with each hydrogen atom, forming two single covalent bonds

Metallic bonding occurs in metals:

• Metal atoms release outer electrons to form a "sea of delocalised electrons"
• Positive metal ions are held in place by electrostatic attraction to the electron sea
• Explains electrical conductivity (mobile electrons), malleability, and ductility of metals like the aluminum produced at Jamaica's Alpart refinery

Chemical formulae and equations

Empirical formulae represent the simplest whole-number ratio of atoms in a compound. To write formulae correctly:

1. Write symbols for both elements
2. Write valencies above each symbol
3. Exchange valencies as subscripts
4. Simplify to lowest terms if possible

Example: Aluminum oxide

• Al (valency 3), O (valency 2)
• Al₂O₃

For compounds containing polyatomic ions (groups of atoms with an overall charge), use brackets when more than one ion is needed:

• Calcium hydroxide: Ca(OH)₂
• Ammonium sulfate: (NH₄)₂SO₄

Common polyatomic ions:

• Hydroxide: OH⁻
• Nitrate: NO₃⁻
• Carbonate: CO₃²⁻
• Sulfate: SO₄²⁻
• Ammonium: NH₄⁺

Balancing chemical equations ensures equal numbers of each type of atom on both sides:

1. Write correct formulae for all reactants and products
2. Count atoms of each element on both sides
3. Add coefficients (large numbers in front) to balance—never change subscripts
4. Check all atoms balance and simplify coefficients if possible

Example: Combustion of propane

• Unbalanced: C₃H₈ + O₂ → CO₂ + H₂O
• Balanced: C₃H₈ + 5O₂ → 3CO₂ + 4H₂O

State symbols indicate physical states in equations:

• (s) = solid
• (l) = liquid
• (g) = gas
• (aq) = aqueous (dissolved in water)

Relative atomic mass and molar calculations

The relative atomic mass (Ar) allows chemists to compare atomic masses. The relative molecular mass (Mr) of a compound equals the sum of relative atomic masses of all atoms in the formula.

Example: Mr of calcium carbonate (CaCO₃)

• Ca = 40, C = 12, O = 16
• Mr = 40 + 12 + (3 × 16) = 100

The mole is the unit for amount of substance. One mole of any substance contains 6.02 × 10²³ particles (Avogadro's constant). The mass of one mole equals the relative atomic or molecular mass in grams.

For CXC CSEC Chemistry, mole calculations include:

• Mass = moles × molar mass
• Moles = mass ÷ molar mass
• Using molar ratios from balanced equations to calculate reacting masses

Physical and chemical changes

Physical changes alter the state or appearance without forming new substances:

• Melting, boiling, dissolving, freezing
• Easily reversible
• No change in mass or chemical composition

Example: Water freezing to ice or salt dissolving in water

Chemical changes produce new substances with different properties:

• Combustion, rusting, neutralization, thermal decomposition
• Difficult or impossible to reverse
• Energy changes occur (heat absorbed or released)
• Signs include color change, gas production, precipitate formation, temperature change

Example: Burning wood, limestone (CaCO₃) decomposing to quicklime (CaO) in Caribbean cement production

Worked examples

Example 1: Electron configuration and bonding

Question: Chlorine has atomic number 17. (a) Write the electron configuration of chlorine. [1 mark] (b) Explain how chlorine atoms bond with hydrogen atoms to form hydrogen chloride. Include a diagram showing electron arrangement in HCl. [4 marks]

Solution:

(a) 2,8,7 ✓

(b) Chlorine has 7 electrons in outer shell and needs 1 more electron to complete the octet ✓. Hydrogen has 1 electron and needs 1 more electron to complete the duplet ✓. They share a pair of electrons, forming a single covalent bond ✓.

Diagram showing: H with 1 outer electron and Cl with 7 outer electrons sharing one pair ✓

Example 2: Writing formulae using valencies

Question: Write the chemical formulae for: (a) iron(III) oxide [1 mark] (b) calcium nitrate [1 mark] (c) ammonium phosphate [1 mark]

Solution:

(a) Fe₂O₃ ✓ (iron valency 3, oxygen valency 2, exchange valencies)

(b) Ca(NO₃)₂ ✓ (calcium valency 2, nitrate ion NO₃⁻ valency 1, brackets needed)

(c) (NH₄)₃PO₄ ✓ (ammonium ion NH₄⁺ valency 1, phosphate ion PO₄³⁻ valency 3)

Example 3: Balancing equations and calculating relative molecular mass

Question: Methane burns in oxygen to produce carbon dioxide and water.

(a) Write a balanced equation with state symbols for this reaction. [3 marks]

(b) Calculate the relative molecular mass of methane (CH₄). [Ar: C = 12, H = 1] [2 marks]

Solution:

(a) CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ✓✓✓ (correct formulae, balanced, state symbols)

(b) Mr = 12 + (4 × 1) ✓ = 16 ✓

Common mistakes and how to avoid them

Confusing mass number with atomic number — The atomic number (bottom) tells you the number of protons; the mass number (top) is protons plus neutrons. To find neutrons, subtract atomic number from mass number.

Changing subscripts when balancing equations — Only add large coefficients in front of formulae. Never alter subscripts within a formula as this changes the substance itself (e.g., H₂O is water, but H₂O₂ is hydrogen peroxide).

Forgetting brackets for polyatomic ions — When more than one polyatomic ion is needed, place brackets around it before adding the subscript: Ca(OH)₂, not CaOH₂. The latter suggests only the hydrogen is doubled.

Mixing up ionic and covalent compound properties — Ionic compounds have high melting points and conduct when molten/aqueous; simple covalent compounds have low melting points and do not conduct. Remember: ionic = metal + non-metal, covalent = non-metal + non-metal.

Incorrect electron configurations — Always fill shells in order (2, 8, 8 for first 20 elements) and check your total electrons match the atomic number. For ions, add electrons for anions (negative) and subtract for cations (positive).

Confusing valency with charge — Valency is combining power (no sign), while charge shows the electrical state of an ion (with + or -). Sodium has valency 1 but forms Na⁺ ions with +1 charge.

Exam technique for Principles of Chemistry

Command word interpretation: "State" requires brief factual answers without explanation (1 mark each). "Explain" requires reasons using scientific terminology (2-3 marks). "Describe" needs a sequence or characteristics (2-3 marks). "Calculate" requires working shown, correct substitution into formulae, and units in the final answer.

Diagram requirements: When asked to "draw" or "show" electron arrangements, draw circles for shells, use dots or crosses for electrons, and ensure the correct number of electrons in each shell. Label shells if required. For dot-and-cross diagrams of covalent bonding, show only outer shell electrons.

Equation-writing marks: One mark typically awarded for correct formulae, one for balancing, one for state symbols. Write state symbols in brackets immediately after each formula, not at the end of the equation.

Definition accuracy: Learn definitions word-for-word from the syllabus, particularly for atom, element, compound, mixture, ion, isotope. Examiners mark against specific keywords, and missing one can lose the mark.

Quick revision summary

Atoms consist of protons, neutrons (in nucleus), and electrons (in shells). Atomic number = protons; mass number = protons + neutrons. Electron configuration determines chemical properties. The periodic table groups elements with similar properties vertically. Ionic bonding involves electron transfer between metals and non-metals; covalent bonding involves electron sharing between non-metals. Write formulae using valency exchange method. Balance equations by adding coefficients, never changing subscripts. Calculate Mr by summing Ar values of all atoms in a formula. Distinguish physical changes (reversible, no new substances) from chemical changes (new substances formed).