Introduction to Atomic Structure
Welcome to your study notes for Atomic Structure! Everything in the universe, from the screen you are reading this on to the air you breathe, is made of tiny building blocks called atoms. Understanding how atoms are put together is the foundation for all of chemistry in Unit C1. Don't worry if science has felt overwhelming before; we will break down each concept step by step with clear rules, simple calculations, and easy memory tricks.
---1. Inside the Atom: Subatomic Particles
An atom is the smallest unit of an element that retains its chemical properties. It consists of a central nucleus surrounded by tiny particles orbiting in electron shells (energy levels).
The Three Subatomic Particles
Atoms are made up of three smaller particles known as subatomic particles:
• Proton: Found in the nucleus. Relative mass = \(1\). Relative charge = \(+1\).
• Neutron: Found in the nucleus. Relative mass = \(1\). Relative charge = \(0\) (neutral).
• Electron: Found orbiting in electron shells. Relative mass = \(\frac{1}{1840}\) (negligible / almost zero). Relative charge = \(-1\).
Why Are Atoms Electrically Neutral?
Every uncharged atom is electrically neutral. This happens because the number of positive protons in the nucleus is always equal to the number of negative electrons orbiting around it. The \(+1\) charges and \(-1\) charges cancel each other out completely, leaving an overall charge of zero.
Memory Trick:
• Proton = Positive (\(+1\))
• Neutron = Neutral (\(0\))
• Electron = Extremely small and negative (\(-1\))
Key Takeaway: The nucleus contains protons and neutrons (where nearly all the mass is concentrated), while electrons orbit in shells. Atoms are neutral because \(\text{number of protons} = \text{number of electrons}\).
---2. Atomic Number, Mass Number, and Standard Notation
To identify atoms and calculate how many particles they have, chemists use two numbers found on the Periodic Table:
• Atomic Number (\(Z\)): The number of protons in the nucleus of an atom. This number defines which element it is (e.g., any atom with 6 protons is always carbon).
• Mass Number (\(A\)): The total number of protons plus neutrons in the nucleus of an atom.
Nuclear Notation: \(^{A}_{Z}\text{X}\)
In standard nuclear notation, the Mass Number (\(A\)) is written at the top (it is the bigger number), and the Atomic Number (\(Z\)) is written at the bottom.
How to Calculate the Number of Each Particle:
1. \(\text{Number of protons} = \text{Atomic number } (Z)\)
2. \(\text{Number of electrons} = \text{Number of protons} = \text{Atomic number } (Z)\) (in a neutral atom)
3. \(\text{Number of neutrons} = \text{Mass number } (A) - \text{Atomic number } (Z)\)
Worked Example: Sodium (\(^{23}_{11}\text{Na}\))
• Protons: \(11\) (the bottom number, \(Z\))
• Electrons: \(11\) (equal to protons in a neutral atom)
• Neutrons: \(23 - 11 = 12\) (\(\text{Mass number} - \text{Atomic number}\))
Common Exam Pitfall: When calculating neutrons, always subtract the bottom number from the top number (\(\text{Mass} - \text{Atomic}\)). Never subtract the mass number from the atomic number!
Key Takeaway: Protons and electrons equal the atomic number (\(Z\)). Neutrons equal mass number minus atomic number (\(A - Z\)).
---3. Electronic Configuration (Elements 1 to 20)
Electrons do not just float around randomly; they arrange themselves in specific layers called electron shells or energy levels. For CCEA Double Award Science, you need to know how electrons fill shells for the first 20 elements (Hydrogen up to Calcium).
The Shell Filling Rules:
• 1st Shell (closest to nucleus): Holds a maximum of \(2\) electrons.
• 2nd Shell: Holds a maximum of \(8\) electrons.
• 3rd Shell: Holds a maximum of \(8\) electrons.
• 4th Shell: Takes the next electrons (for Potassium and Calcium).
Linking Electron Arrangement to the Periodic Table:
Your Periodic Table gives you huge clues about an atom's electron arrangement:
• Group Number (columns going down): Tells you the number of electrons in the outer shell (valence electrons). For example, Group 1 elements have \(1\) outer electron; Group 7 elements have \(7\) outer electrons.
• Period Number (rows going across): Tells you the number of occupied electron shells. For example, elements in Period 3 have \(3\) electron shells containing electrons.
Examples:
• Oxygen (Atomic number \(8\)): \(8\) electrons total \(\rightarrow\) Electronic configuration = \(2,6\) (Period 2, Group 6)
• Potassium (Atomic number \(19\)): \(19\) electrons total \(\rightarrow\) Electronic configuration = \(2,8,8,1\) (Period 4, Group 1)
• Calcium (Atomic number \(20\)): \(20\) electrons total \(\rightarrow\) Electronic configuration = \(2,8,8,2\) (Period 4, Group 2)
Key Takeaway: Shells fill in the order \(2, 8, 8, 2\). The outer shell electrons equal the Group number; the total number of occupied shells equals the Period number.
---4. Isotopes and Relative Atomic Mass (\(A_r\))
What is an Isotope?
Isotopes are defined as atoms of the same element with the same number of protons (same atomic number) but different numbers of neutrons (different mass numbers).
• Chemical Properties: Isotopes of an element have identical chemical properties because they have the exact same electron configuration.
• Physical Properties: Isotopes have slightly different physical properties (such as density and mass) because their masses are different.
Relative Atomic Mass (\(A_r\))
Relative Atomic Mass (\(A_r\)) is defined as the average mass of an atom of an element relative to \(\frac{1}{12}\text{th}\) the mass of a carbon-12 atom.
Because many elements exist naturally as a mixture of different isotopes, the \(A_r\) value takes into account both the mass of each isotope and how common (abundant) it is.
Formula for Calculating \(A_r\):
\(\text{Relative Atomic Mass } (A_r) = \frac{\sum (\text{Isotope Mass} \times \text{Percentage Abundance})}{100}\)
Worked Example: Calculating the \(A_r\) of Chlorine
Naturally occurring chlorine consists of \(75\%\) chlorine-35 (\(^{35}\text{Cl}\)) and \(25\%\) chlorine-37 (\(^{37}\text{Cl}\)).
Step 1: Multiply each isotope mass by its percentage abundance:
\((35 \times 75) = 2625\)
\((37 \times 25) = 925\)
Step 2: Add the results together:
\(2625 + 925 = 3550\)
Step 3: Divide by \(100\):
\(A_r = \frac{3550}{100} = 35.5\)
Common Exam Pitfall: In exam definitions, never just say "elements with different neutrons". You must state: "atoms of the same element with the same number of protons but different numbers of neutrons" to get full marks!
Key Takeaway: Isotopes share identical chemical behavior but have different masses. \(A_r\) is the weighted average mass compared to \(\frac{1}{12}\text{th}\) of a carbon-12 atom.
---5. Ion Formation: Cations and Anions
Atoms react to gain stability by achieving a full outer shell of electrons (like the noble gases). They do this by losing or gaining valence electrons, which turns them into charged particles called ions.
1. Cations (Positive Ions)
• Formed when metal atoms lose outer shell electrons.
• Because the atom loses negatively charged electrons, it is left with more protons than electrons, giving it an overall positive charge.
• Example: A sodium atom (\(\text{Na}: 2,8,1\)) loses \(1\) electron to form a sodium ion (\(\text{Na}^+: 2,8\)):
\(\text{Na} \rightarrow \text{Na}^+ + \text{e}^-\)
2. Anions (Negative Ions)
• Formed when non-metal atoms gain electrons into their outer shell.
• Because the atom gains extra negatively charged electrons, it now has more electrons than protons, giving it an overall negative charge.
• Example: A chlorine atom (\(\text{Cl}: 2,8,7\)) gains \(1\) electron to form a chloride ion (\(\text{Cl}^-: 2,8,8\)):
\(\text{Cl} + \text{e}^- \rightarrow \text{Cl}^-\)
Memory Trick:
• Ca+ion: The letter 't' looks like a plus sign (\(+\)) \(\rightarrow\) Positive ion.
• Anion: Sounds like "A Negative IOn" \(\rightarrow\) Negative ion.
Key Takeaway: Metals lose electrons to form positive cations; non-metals gain electrons to form negative anions.
---Quick Review Checklist
Before moving on to the next topic, make sure you can:
• State the relative mass, charge, and location of protons, neutrons, and electrons.
• Explain why neutral atoms have no overall electric charge.
• Calculate the number of protons, neutrons, and electrons from nuclear notation \(^{A}_{Z}\text{X}\).
• Write the electronic configuration for any of the first 20 elements.
• Define an isotope and calculate relative atomic mass (\(A_r\)) using percentage abundances.
• Explain how positive cations and negative anions are formed when atoms lose or gain electrons.