🧠 Chemistry Study Notes: Atomic Structure and the Periodic Table
Welcome to one of the most fundamental and important chapters in Chemistry! Don't worry if this seems tricky at first—we are essentially learning the 'alphabet' of chemistry. Understanding atoms and how they fit into the Periodic Table is like getting the instruction manual for the entire universe. Let's break it down!
Section 1: The Tiny World of the Atom
Everything around you—from the air you breathe to your phone—is made up of tiny particles called atoms.
The Structure of the Atom
Think of an atom like a mini solar system. It has a heavy centre, and tiny particles orbit around it.
The centre of the atom is called the nucleus. It contains two types of particles: protons and neutrons. The particles orbiting the nucleus are the electrons, which exist in specific energy levels called shells.
The Three Subatomic Particles
The identity and behaviour of every atom depend on its three subatomic particles:
1. Protons
- Charge: Positive (+1)
- Mass (Relative): 1
- Location: In the nucleus
- Crucial Role: The number of protons determines which element the atom is.
2. Neutrons
- Charge: Neutral (0)
- Mass (Relative): 1
- Location: In the nucleus
- Crucial Role: They add mass to the atom but have no electrical charge.
3. Electrons
- Charge: Negative (-1)
- Mass (Relative): Very, very small (we say effectively 0)
- Location: Orbiting the nucleus in shells
- Crucial Role: Electrons are involved in all chemical reactions.
Memory Aid: A common mistake is confusing the mass of protons and electrons. Remember, electrons are so light they are like tiny flies orbiting a heavy bowling ball (the nucleus)!
Key Takeaway for Section 1: Atoms have a positive nucleus (Protons & Neutrons) and negative electrons orbiting in shells. An atom is electrically neutral, meaning it has the same number of protons and electrons.
Section 2: Decoding the Atom – Numbers and Isotopes
Chemists use two key numbers to describe any atom: the Atomic Number and the Mass Number. You can find both of these numbers on the Periodic Table.
Atomic Number (Proton Number, Z)
The Atomic Number (Z) is the most important identifying feature of an element.
\[ \text{Atomic Number} = \text{Number of Protons} \]
If you change the number of protons, you change the element! For example, an atom with 6 protons is always Carbon (C). An atom with 7 protons is always Nitrogen (N).
Quick Tip: In a neutral atom (no charge): \[ \text{Number of Protons} = \text{Number of Electrons} \]
Mass Number (A)
The Mass Number (A) is the total number of particles found in the nucleus (protons and neutrons). Since electrons have negligible mass, we ignore them in this calculation.
\[ \text{Mass Number} = \text{Number of Protons} + \text{Number of Neutrons} \]
Calculating Subatomic Particles
To find the number of neutrons, simply subtract the atomic number from the mass number:
\[ \text{Number of Neutrons} = \text{Mass Number} - \text{Atomic Number} \]
Example: Lithium (Li) has an Atomic Number of 3 and a Mass Number of 7.
- Protons: 3 (from Atomic Number)
- Electrons: 3 (neutral atom, so Protons = Electrons)
- Neutrons: 7 - 3 = 4
Isotopes
Not all atoms of the same element are exactly identical!
Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons (and therefore a different mass number).
Example: Carbon-12 and Carbon-14.
- Both have 6 Protons (they are both Carbon).
- Carbon-12 has 6 Neutrons (12 – 6 = 6).
- Carbon-14 has 8 Neutrons (14 – 6 = 8).
Did you know? Carbon-14 is used in carbon dating to determine the age of ancient artifacts because it decays very slowly over time.
Key Takeaway for Section 2: Atomic Number tells you the element (Protons & Electrons). Mass Number tells you the weight (Protons + Neutrons). Isotopes are variations of an element with different numbers of neutrons.
Section 3: Electron Arrangement and Reactivity
Electrons orbit the nucleus in specific paths called shells or energy levels. The arrangement of these electrons determines how an element reacts chemically.
Rules for Electron Shells
For the first 20 elements (the ones you need to know for this level), electrons fill the shells from the inside out, following these simple rules:
- The first shell (closest to the nucleus) can hold a maximum of 2 electrons.
- The second shell can hold a maximum of 8 electrons.
- The third shell can hold a maximum of 8 electrons.
Step-by-Step Arrangement (Drawing the Dot-and-Cross):
Let's use Oxygen (Atomic Number = 8). We have 8 electrons to place.
1. Fill the first shell: 2 electrons used (8 - 2 = 6 left).
2. Place the remaining 6 electrons in the second shell.
The arrangement is 2, 6.
Valence Electrons
The electrons in the outermost shell are called valence electrons. They are the most important electrons because they are the ones involved in forming bonds and reacting with other atoms.
Atoms "want" to have a full outer shell, as this makes them stable (like the Noble Gases, see Section 5). They achieve this by gaining, losing, or sharing valence electrons.
Key Takeaway for Section 3: Electrons fill shells 2, 8, 8... The number of electrons in the outer shell (valence electrons) determines how reactive the atom is.
Section 4: The Periodic Table – The Master Key
The Periodic Table is an organised chart of all known elements. It’s organised based on the structure of the atoms.
Structure and Organisation
The table is arranged in order of increasing Atomic Number (number of protons).
1. Groups (Vertical Columns)
- Elements in the same vertical column are in the same Group.
- Group number tells you the number of valence electrons (outer shell electrons).
- Crucial implication: Elements in the same group have very similar chemical properties because they have the same number of outer electrons.
2. Periods (Horizontal Rows)
- Elements in the same horizontal row are in the same Period.
- Period number tells you the number of electron shells the atom has.
- Example: All elements in Period 2 have two electron shells.
Common Mistake to Avoid: Don't confuse groups (up and down, similar properties) with periods (left and right, increasing number of shells).
Key Takeaway for Section 4: Group number = outer electrons. Period number = total shells. This arrangement explains why elements have similar chemical behaviour.
Section 5: Properties of Key Groups
We focus on three specific groups whose chemical properties are very important.
Group 1: The Alkali Metals
(Lithium, Sodium, Potassium, etc.)
- Electron Structure: They all have 1 valence electron.
- Reactivity: Very reactive because they easily lose that single outer electron to become stable positive ions (+1 charge).
- Physical Properties: Soft (can be cut with a knife), low density, low melting points, stored under oil to prevent reaction with air/water.
Reaction with Water
Alkali metals react vigorously with water, producing a metal hydroxide (an alkali solution) and hydrogen gas (H₂).
Example: Sodium + Water → Sodium Hydroxide + Hydrogen
Reactivity Trend
Reactivity increases as you go down the group.
Why? As you go down, the outer electron is further away from the nucleus (more shells), meaning the positive nucleus has less pull on it, making it easier to lose.
Group 7: The Halogens
(Fluorine, Chlorine, Bromine, Iodine)
- Electron Structure: They all have 7 valence electrons.
- Reactivity: Very reactive non-metals because they easily gain 1 electron to complete their outer shell, forming negative ions (-1 charge).
- Physical Properties: Exist as diatomic molecules (e.g., Cl₂, Br₂). They show a clear transition in state and colour down the group (Fluorine/Chlorine are gases, Bromine is a liquid, Iodine is a solid).
Displacement Reactions
A more reactive halogen will displace a less reactive halide ion from its solution.
Reactivity Trend: Reactivity decreases as you go down the group.
Why? As you go down, the atom gets larger, making it harder for the nucleus to attract and capture an extra electron. Fluorine is the most reactive.
Group 0 (or Group 8): The Noble Gases
(Helium, Neon, Argon, etc.)
- Electron Structure: They all have a full outer shell (8 electrons, except Helium which has 2).
- Reactivity: They are unreactive (inert). Since their outer shell is full, they do not need to gain, lose, or share electrons.
- Uses: Argon is used in light bulbs (because it won't react with the hot filament). Helium is used in airships and balloons (because it is light and non-flammable).
Quick Review Box:
- Group 1: Loses 1 electron (Reactivity increases down the group).
- Group 7: Gains 1 electron (Reactivity decreases down the group).
- Group 0: Stable, inert (full outer shell).
You’ve covered the fundamentals of atomic theory and the structure of the Periodic Table! This knowledge is the foundation for everything else we study in chemistry. Great work!