Study Notes C2: Atoms, Elements and Compounds
Hello! Welcome to the foundation of all chemistry. This chapter is super important because it teaches you what everything is made of, from the air you breathe to the metal in your phone. Understanding atoms, elements, and compounds will make all future chemistry topics much easier. Don't worry if some of the diagrams look complicated—we’ll break them down step by step!
C2.2 The Structure of the Atom (Core)
Every single piece of matter is made up of tiny particles called atoms. While atoms are tiny, they have a specific structure made of even smaller, subatomic particles.
Atomic Structure: The Basics
The atom is structured like a miniature solar system (though this isn't exactly accurate scientifically, it's a useful model for IGCSE):
- A central, dense core called the nucleus.
- Particles orbiting the nucleus in specific paths called shells (or energy levels).
The three types of subatomic particles are:
| Particle | Location | Relative Charge | Relative Mass |
|---|---|---|---|
| Proton (p) | Nucleus | +1 (Positive) | 1 |
| Neutron (n) | Nucleus | 0 (Neutral) | 1 |
| Electron (e) | Shells/Orbits | -1 (Negative) | Negligible (about \( \frac{1}{1836} \) of a proton) |
In any neutral atom:
The number of Protons = The number of Electrons.
Defining Atoms by Numbers
We identify elements using two key numbers:
-
Proton Number (Atomic Number), Z:
This is the number of protons in the nucleus of an atom.
Why is this important? It determines which element the atom is. Every atom of Carbon has 6 protons, and every atom with 6 protons is Carbon. -
Mass Number (Nucleon Number), A:
This is the total number of protons plus neutrons in the nucleus.
To find the number of neutrons, you subtract the proton number from the mass number:
Number of Neutrons = Mass Number – Proton Number
Did you know? The term 'nucleon' simply means a particle found in the nucleus (protons and neutrons).
Electron Arrangement (Electronic Configuration)
Electrons fill the shells closest to the nucleus first. The shells have limits on how many electrons they can hold:
- 1st Shell (K-shell): Maximum 2 electrons
- 2nd Shell (L-shell): Maximum 8 electrons
- 3rd Shell (M-shell): Maximum 8 electrons (for the first 20 elements, which is the limit of our syllabus requirement)
We write the electronic configuration by listing the number of electrons in each occupied shell, separated by commas.
Example: Aluminium (Al) has 13 protons and 13 electrons.
1st shell: 2
2nd shell: 8
3rd shell: 3
Electronic configuration: 2, 8, 3
Connecting Structure to the Periodic Table
The Periodic Table is organized based on atomic structure:
1. Group Number (Columns I to VII):
The group number tells you the number of outer-shell electrons (or valence electrons).
Example: Potassium (K) is in Group I, so it has 1 electron in its outer shell (2, 8, 8, 1).
2. Period Number (Rows):
The period number tells you the number of occupied electron shells.
Example: Sodium (Na) is in Period 3, so it has 3 electron shells occupied (2, 8, 1).
3. Group VIII (The Noble Gases):
These elements (like Neon, 2, 8; or Argon, 2, 8, 8) are special because they have a full outer shell. This makes them very unreactive or stable. All other atoms try to achieve this stable configuration by reacting.
Quick Review C2.2 Takeaway: The atom is defined by its Proton Number. Electrons determine how the atom reacts, and their arrangement (configuration) is key.
C2.1 Elements, Compounds and Mixtures (Core)
Before we look at how atoms join together, we must clearly define the three main classifications of matter. This is a common source of confusion, so pay close attention!
1. Elements
An element is a substance that cannot be broken down into simpler substances by chemical methods.
- It is made up of only one type of atom.
- All atoms in an element have the same Proton Number.
- Examples: Gold (Au), Oxygen gas (O₂), Carbon (C).
2. Compounds
A compound is a substance formed when two or more different elements are chemically combined.
- The elements are held together by chemical bonds.
- The properties of a compound are very different from the elements that formed it.
- They can only be separated by chemical reactions (which are usually difficult).
- They combine in a fixed ratio.
- Examples: Water (H₂O), Salt (NaCl), Carbon Dioxide (CO₂).
3. Mixtures
A mixture is formed when two or more substances (elements or compounds) are physically combined but not chemically bonded.
- The substances retain their individual properties.
- They can be separated relatively easily using physical methods (like filtration, boiling, magnetism).
- They can be mixed in any ratio.
- Examples: Salt water, sand and iron filings, air (a mixture of gases).
Analogy: Think of LEGO bricks.
An Element is a pile of only red 2x4 bricks.
A Compound is a red 2x4 brick chemically locked onto a blue 2x4 brick.
A Mixture is a pile of red bricks lying next to a pile of blue bricks—you can easily separate them with your hand (a physical method).
Quick Review C2.1 Takeaway: Chemical bonds define compounds; the lack of bonds defines mixtures.
C2.3 Ions and Ionic Bonds (Core)
Atoms react to achieve a stable electronic configuration—usually a full outer shell (like a Noble Gas). They do this by either gaining or losing electrons. When an atom gains or loses electrons, it becomes a charged particle called an ion.
Formation of Ions
1. Positive Ions (Cations):
Formed when a metal atom loses one or more electrons. Since it loses negative charge, it becomes positively charged.
Memory Aid: Cations are Paws-itive (cat with paws).
2. Negative Ions (Anions):
Formed when a non-metal atom gains one or more electrons. Since it gains negative charge, it becomes negatively charged.
Example: Sodium (Group I, 2, 8, 1) loses 1 electron to become a Sodium ion, Na\(^+\) (2, 8). Chlorine (Group VII, 2, 8, 7) gains 1 electron to become a Chloride ion, Cl\(^-\) (2, 8, 8). Both now have a stable, full outer shell.
Ionic Bonds
An ionic bond is the strong electrostatic attraction between oppositely charged ions. These bonds typically form between metals (which form cations) and non-metals (which form anions).
The core syllabus requires you to describe the formation of ionic bonds between elements from Group I and Group VII using dot-and-cross diagrams.
Step-by-step for Sodium Chloride (NaCl):
- Start with a neutral Sodium atom (Na, 2, 8, 1) and a neutral Chlorine atom (Cl, 2, 8, 7).
- Sodium transfers its single outer electron (shown as a cross 'x') to the Chlorine atom.
- Sodium becomes Na\(^+\) (2, 8). Chlorine becomes Cl\(^-\) (2, 8, 8).
- The strong electrostatic force attracts the Na\(^+\) and Cl\(^-\) ions together, forming the ionic bond.
(Note: While diagrams cannot be rendered here, remember to show only the outer shells for the neutral atoms, and then show the full outer shells and charges for the resulting ions.)
Properties of Ionic Compounds (Core)
Ionic compounds (like salt) have specific properties due to the strength of their electrostatic bonds:
- High Melting Points and Boiling Points: A lot of thermal energy is needed to break the strong electrostatic forces holding the ions together.
- Electrical Conductivity:
- They are poor conductors when solid because the ions are fixed in position and cannot move.
- They are good conductors when molten or dissolved in water (aqueous). This is because the ions are free to move and carry the charge.
- Solubility: They are generally soluble in water, as the charged water molecules can pull the ions away from the structure.
Quick Review C2.3 Takeaway: Ionic bonds result from the transfer of electrons between metals (cations) and non-metals (anions), creating strong electrostatic attraction and high melting points.
C2.4 Simple Molecules and Covalent Bonds (Core)
When non-metals react together, they cannot transfer electrons to each other easily. Instead, they form a bond by sharing electrons.
Covalent Bonds
A covalent bond is formed when a pair of electrons is shared between two atoms, allowing both atoms to achieve the stable electronic configuration of a noble gas (a full outer shell).
The resulting unit is called a simple molecule.
Drawing Covalent Bonds (Dot-and-Cross Diagrams)
For covalent bonding diagrams, we focus on the shared electrons in the overlap region between the two atoms.
1. Hydrogen (\( \text{H}_2 \)):
- Hydrogen (1 electron) needs 1 more electron for a full shell (2 electrons total).
- Two H atoms share 1 pair of electrons.
The shared pair forms a single covalent bond.
2. Chlorine (\( \text{Cl}_2 \)):
- Chlorine (7 outer electrons) needs 1 more electron for a full shell (8 electrons).
- Two Cl atoms share 1 pair of electrons to complete both outer shells.
3. Water (\( \text{H}_2\text{O} \)):
- Oxygen (6 outer electrons) needs 2 electrons. Hydrogen (1 outer electron) needs 1 electron.
- Oxygen shares one electron pair with each of the two hydrogen atoms. This results in two single covalent bonds.
4. Methane (\( \text{CH}_4 \)), Ammonia (\( \text{NH}_3 \)), and Hydrogen Chloride (\( \text{HCl} \)):
- Methane: Carbon (4 outer electrons) shares one electron pair with each of the four hydrogen atoms, forming four single bonds.
- Ammonia: Nitrogen (5 outer electrons) shares one electron pair with each of the three hydrogen atoms, forming three single bonds. Nitrogen is left with one unshared pair of electrons (a lone pair).
- Hydrogen Chloride: Hydrogen (1 electron) shares one pair with Chlorine (7 outer electrons). This forms one single bond.
Don't worry if this seems tricky at first: The key rule is: count how many electrons the atom starts with, count how many it needs for a full shell (usually 8), and that tells you how many bonds (shared pairs) it must form.
Properties of Simple Molecular Compounds (Core)
Simple molecular substances (like water, methane, or chlorine gas) have very different properties from ionic compounds:
- Low Melting Points and Boiling Points:
- Inside the molecule, the covalent bonds are very strong.
- However, the forces between the individual molecules (called intermolecular forces) are very weak.
- When you melt or boil the substance, you only need a little energy to break these weak intermolecular forces, not the strong covalent bonds.
- Poor Electrical Conductivity:
They are poor electrical conductors (non-conductors) in all states (solid, liquid, or gas) because they do not contain free-moving charged particles (ions or delocalised electrons) to carry a current.
Quick Review C2.4 Takeaway: Covalent bonds involve sharing electrons between non-metals, forming simple molecules that are generally volatile (low MP/BP) and non-conductive due to weak forces between molecules.