C2 Atoms, Elements and Compounds - IGCSE Combined Science

Atoms, Elements, and Compounds

Syllabus Overview

This note covers the Cambridge IGCSE Combined Science Syllabus topic C2: Atoms, Elements, and Compounds.

How to Use This Note

  1. Syllabus Checklist: Use the syllabus statements as a checklist to identify areas of confidence and weakness.

  2. Review: Consult the relevant sections in the coursebook (pages 274-300) to review the material for any statements you're not confident with.

  3. Video Tutorials: Utilize online resources such as Fuse School, FreeScienceLessons, and Cognito for additional explanations.

C2.3 Ions and Ionic Bonds

Core Concepts

  1. Ion Formation:

    • Cations: Positive ions formed by the loss of electrons.

    • Anions: Negative ions formed by the gain of electrons.

  2. Ionic Bond Definition: An ionic bond is a strong electrostatic attraction between oppositely charged ions.

  3. Formation of Ionic Bonds:

    • Ionic bonds typically form between elements from Group I and Group VII.

    • Dot-and-cross diagrams are used to illustrate electron transfer.

  4. Properties of Ionic Compounds:

    • (a) High Melting and Boiling Points: Due to strong electrostatic forces.

    • (b) Electrical Conductivity: Good when aqueous or molten (ions can move), poor when solid (ions cannot move).

    • (c) Solubility: Generally soluble in water.

Supplement Concepts

  1. Ionic Bonds Between Metals and Non-metals:

    • Formation of ionic bonds between ions of metallic and non-metallic elements.

    • Dot-and-cross diagrams are used to show electronic configurations.

  2. Giant Lattice Structure:

    • Ionic compounds form giant lattice structures.

    • Regular arrangement of alternating positive and negative ions. Example: Sodium chloride (NaCl).

C2.4 Simple Molecules and Covalent Bonds

Core Concepts

  1. Covalent Bond Definition: A covalent bond is formed when a pair of electrons is shared between two atoms.

    • This sharing leads to noble gas electronic configurations.

  2. Formation of Covalent Bonds:

    • Formation of covalent bonds in simple molecules like H2, Cl2, H2O, CH4, NH_3, and HCl.

    • Dot-and-cross diagrams are used to show electronic configurations.

  3. Properties of Simple Molecular Compounds:

    • (a) Low Melting and Boiling Points: Due to weak intermolecular forces.

    • (b) Poor Electrical Conductivity: No mobile ions or electrons.

Supplement Concepts

  1. Formation of Covalent Bonds (Continued):

    • Formation of covalent bonds in molecules like CH3OH, C2H4, O2, CO2, and N2.

    • Dot-and-cross diagrams are used to show electronic configurations.

Elements, Compounds, and Mixtures

  • Element: Consists of atoms of the same kind (e.g., oxygen (O), hydrogen (H), helium (He), carbon (C), potassium (K), sodium (Na)).

  • Compound: Consists of two or more kinds of atoms in a fixed ratio (e.g., table salt (NaCl), water (H_2O)).

  • Mixture: Examples include air, milk, and salad dressing.

Atomic Structure

  • Components: Protons, neutrons, and electrons.

  • Carbon Atom Example: 6 protons, 6 neutrons.

  • Location: Protons and neutrons in the nucleus; electrons outside the nucleus.

Subatomic Particles

Particle

Charge

Relative Mass

Location

Electron

-1

1/1840

Outside Nucleus

Proton

+1

1

Nucleus

Neutron

0

1

Nucleus

Periodic Table Information

Key Information Displayed

  • Atomic Number

  • Atomic Symbol

  • Name

  • Relative Atomic Mass

Examples

  • Hydrogen (H): Atomic number 1, relative atomic mass 1.

  • Helium (He): Atomic number 2, relative atomic mass 4.

  • Lithium (Li): Atomic number 3, relative atomic mass 7.

  • Beryllium (Be): Atomic number 4, relative atomic mass 9.

  • Sodium (Na): Atomic number 11, relative atomic mass 23.

  • Magnesium (Mg): Atomic number 12, relative atomic mass 24.

Atomic Number and Mass

  • Atomic Number: The number of protons in the nucleus of an atom.

  • Atomic Mass (Relative Atomic Mass): The average mass of the atoms in an element.

  • Element Symbol: Shorthand abbreviation for the element name, often derived from Greek or Latin.

Electronic Structure

  • Atoms are neutral (same number of protons and electrons).

  • Electrons arranged in energy levels (shells).

  • The nucleus contains protons and neutrons and is tiny compared to the atom.

Electron Shell Filling

  • Electrons fill shells closest to the nucleus.

  • Shell 1: Maximum of 2 electrons.

  • Shell 2: Maximum of 8 electrons.

  • Shell 3: Maximum of 8 electrons.

  • Shell 4: Maximum of 2 electrons (relevant for the first 20 elements).

Electronic Configurations of the First 20 Elements

  • H (1): 1

  • He (2): 2

  • Li (3): 2,1

  • Be (4): 2,2

  • B (5): 2,3

  • C (6): 2,4

  • N (7): 2,5

  • O (8): 2,6

  • F (9): 2,7

  • Ne (10): 2,8

  • Na (11): 2,8,1

  • Mg (12): 2,8,2

  • Al (13): 2,8,3

  • Si (14): 2,8,4

  • P (15): 2,8,5

  • S (16): 2,8,6

  • Cl (17): 2,8,7

  • Ar (18): 2,8,8

  • K (19): 2,8,8,1

  • Ca (20): 2,8,8,2

Ions and Electronic Structure

  • Ions have a different number of protons and electrons, resulting in an electrical charge.

  • Noble gas elements have stable electron arrangements.

  • Ions tend to achieve the electron structure of noble gases.

Electric Charge of Ions

To determine the electric charge of an ion, compare the number of protons (positive charge) to the number of electrons (negative charge).

Ionic Bonding

Formation of Ions

  • Positive ions (cations) and negative ions (anions).

  • Ionic bond: Strong electrostatic attraction between oppositely charged ions.

  • Ionic bonds often form between Group I and Group VII elements.

Properties of Ionic Compounds

  • High melting and boiling points.

  • Good electrical conductivity when aqueous or molten, poor when solid.

  • Generally soluble in water.

  • Giant lattice structure with regular arrangement of alternating positive and negative ions.

  • Example: Sodium chloride (NaCl).

Formation of Sodium Chloride (NaCl)

  1. Formation of Sodium Ions (Na⁺):

    • Sodium loses 1 electron to achieve a full outer shell.

    • Na arrow Na^+ + e^-

  2. Formation of Chloride Ions (Cl⁻):

    • Chlorine gains 1 electron to achieve a full outer shell.

    • Cl + e^- arrow Cl^-

  3. Overall Process:

    • Sodium loses an electron, and chlorine gains it.

    • Oppositely charged ions attract, forming a giant lattice.

Giant Ionic Lattice

Na^+ and Cl^- ions are held together by electrostatic forces in a 3D lattice structure.

Properties of Ionic Compounds

  • Formula indicates the ratio of ions (e.g., MgCl_2 implies 1 Mg atom for every 2 chlorine atoms).

  • High melting and boiling points due to strong electrostatic attractions.

  • Conduct electricity in molten or dissolved states because ions can move.

  • Usually soluble and solids at room temperature.

Bonding Examples

  1. Magnesium Oxide (MgO):

    • Magnesium (2.8.2) loses 2 electrons.

    • Oxygen (2.6) gains 2 electrons.

    • Mg^{2+} and O^{2-}.

  2. Sodium Fluoride (NaF):

    • Sodium (2.8.1) loses 1 electron.

    • Fluorine (2.7) gains 1 electron.

    • Na^+ [2.8] and F^- [2.8].

  3. Lithium Oxide (Li_2O):

    • Two lithium atoms each lose 1 electron.

    • One oxygen atom gains 2 electrons.

    • 2Li^+ and O^{2-}.

  4. Lithium Nitride (Li_3N):

    • Three lithium atoms each lose 1 electron.

    • One nitrogen atom gains 3 electrons.

    • 3Li^+ and N^{3-}.

  5. Aluminum Fluoride (AlF_3):

    • One aluminum atom loses 3 electrons.

    • Three fluorine atoms each gain 1 electron.

    • Al^{3+} and 3F^{-}.

Covalent Bonding

Definition

  • Covalent bond: Formed when a pair of electrons is shared between two atoms.

  • Typically occurs between non-metal + non-metal or hydrogen + non-metal.

  • Leads to noble gas electronic configurations.

Properties of Simple Covalent Compounds

  • Low melting and boiling points.

  • Often liquids or gases at room temperature.

  • Poor electrical conductivity.

  • Weak intermolecular forces (easy to break) compared to strong covalent bonds within the molecules.

Formation of Hydrogen (H_2)

  • Nuclei repel but are attracted to shared electrons.

Covalent Bond Formation

  • Nonmetals share valence electrons to achieve a noble gas configuration.

  • Example: Fluorine (F) has 7 valence electrons and shares one electron with another fluorine atom to form a covalent bond.

Single Covalent Bond

  • Sharing of two valence electrons.

  • Occurs between nonmetals and hydrogen.

  • Forms actual molecules, unlike ionic bonds.

  • Two specific atoms are joined.

How to Show Covalent Bond Formation

  1. Determine the final formula.

  2. Put the pieces together to achieve the right formula.

Example: Water (H_2O)

  • Each hydrogen has 1 valence electron and wants 1 more.

  • Oxygen has 6 valence electrons and wants 2 more.

  • They share to make each other "happy" (full valence shell).

Multiple Bonds

  • Atoms can share more than one pair of valence electrons.

  • Double bond: Atoms share two pairs (4) of electrons.

  • Triple bond: Atoms share three pairs (6) of electrons.

Example: Carbon Dioxide (CO_2)

  • Carbon (central atom) has 4 valence electrons and wants 4 more.

  • Oxygen has 6 valence electrons and wants 2 more.

  • Requires two double bonds.

Example: Ammonia (NH_3$$)

  • Nitrogen (N) has 5 valence electrons and wants 3 more.

  • Hydrogen (H) has 1 valence electron and wants 1 more.