Periodic Table Basics: Atomic Number, Mass Number, Isotopes, Ions, and Groups

Groups and Periods on the Periodic Table

• Columns are called groups; rows are called periods.
• The periodic table is used to learn about elements, their symbols, and their properties.
• The table is referenced frequently in class, so keep it handy.

Atomic Number: What it is and Why It Matters

• Each element has a unique atomic number in the upper left corner of its box.
• The atomic number identifies the element and tells you the number of protons in the nucleus:
  • For hydrogen, the atomic number is 1.
  • For helium, it is 2.
  • For lithium, 3; for beryllium, 4; boron, 5; carbon, 6; nitrogen, 7; oxygen, 8; fluorine (9); neon (10).
• Elements are identified by symbol and name as well as by this unique atomic number.
• Examples from the lecture:
  • Sodium has atomic number 11 (symbol Na).
  • Manganese has atomic number 25 (symbol Mn).
  • Gadolinium is atomic number 64.
  • Bromine is atomic number 35.
• Important terminology:
  • The symbol for the atomic number is typically denoted by Z.
  • The atomic number equals the number of protons in the nucleus.
  • If you have a collection of atoms all with the same atomic number, they are the same element.

Mass Number and Nuclear Composition

• The mass number is not shown on the periodic table. It is denoted by A and represents the total number of protons plus neutrons in the nucleus:
  • $A = Z + N$ where Z is the atomic number (protons) and N is the neutron number.
• The mass number tells you how many nucleons (protons + neutrons) are in a specific isotope.
• Relation to the charge:
  • The mass number does not tell you the charge by itself; charge depends on the balance between protons and electrons.

Isotopes: Definition and Notation

• Isotopes are atoms of the same element (same Z) that differ in the number of neutrons (N) and thus in mass number A.
• Representation of isotopes:
  • Notation 1: Symbol − mass number, e.g., Na-23.
  • Notation 2: Mass number over atomic number with the symbol, e.g., $^{A}_{Z}X$, or the equivalent left-right form: symbol with superscript A or the left-side Z and top A.
• In sodium example from the lecture: Sodium-23 has A = 23 and Z = 11.
• For sodium-23: the notation on the left would be 23 over 11 next to Na, and on the right Na-23.
• Atomic mass vs mass number:
  • The atomic mass (on the periodic table) is a weighted average of all isotopes for that element.
  • The mass number A is specific to a given isotope and is not the same as the atomic mass.
• Most isotopes are radioactive; many are unstable, though some are stable. The speaker notes that there are stable isotopes for some elements, but most isotopes tend to be radioactive.

Worked Examples: Using A, Z, and N

• Example 1: Nickel-58
  • Atomic number Z = 28 (Nickel).
  • Mass number A = 58.
  • Neutron number N = A − Z = 58 − 28 = 30.
  • For a neutral nickel atom, number of electrons = Z = 28.
  • If the charge is not shown, the atom is neutral; if a charge were shown, electrons would differ from Z accordingly.
• Example 2: Nickel-59
  • Z = 28; A = 59.
  • N = 59 − 28 = 31.
  • Electron count for a neutral atom would be 28.
  • This is a different isotope of nickel than Ni-58.
• Example 3: Beryllium-9
  • Be is atomic number Z = 4.
  • A = 9.
  • N = A − Z = 9 − 4 = 5.
  • For a neutral Be atom, electrons = Z = 4.
• Example 4: Phosphorus with a -3 charge (P^{3−})
  • Element: Phosphorus, Z = 15.
  • Mass number A = 31 (given as top number in the isotope notation).
  • Neutron number N = A − Z = 31 − 15 = 16.
  • Charge: −3 means there are 3 more electrons than protons, so electrons = Z + 3 = 18.
  • Therefore, the ion is P^{3−} with 18 electrons, 15 protons, and 16 neutrons.

Ions, Atoms, and Charge

• An atom is neutral when it has the same number of protons and electrons.
• An ion occurs when protons and electrons are not equal.
• Cation: a positively charged ion (more protons than electrons) — formed by losing electrons.
• Anion: a negatively charged ion (more electrons than protons) — formed by gaining electrons.
• Etymology and connection to battery terminology:
  • Cathode attracts cations.
  • Anode attracts anions.
  • The terms cation and anion originate from this battery context.
• Summary:
  • If electrons are removed: you get a cation.
  • If electrons are added: you get an anion.

OILRIG mnemonic for Redox Behavior

• OILRIG stands for Oxidation Is Loss of electrons; Reduction Is Gain of electrons.
• Elements on the left side of the periodic table tend to lose electrons (oxidation) to form cations.
• Elements on the right side tend to gain electrons (reduction) to form anions.
• Group trends for electron loss/gain (as discussed in class):
  • Group 1A (alkali metals): lose 1 electron → form +1 cations.
  • Group 2A (alkaline earth metals): lose 2 electrons → form +2 cations.
  • Group 3A: lose 3 electrons → form +3 cations.
  • Group 4A: described as not performing this simple loss/gain pattern in the lecture (no fixed charge given).
  • Group 5A: gain 3 electrons (to form −3 anions).
  • Group 6A: gain 2 electrons (to form −2 anions).
  • Group 7A: gain 1 electron (to form −1 anions).
  • Group 8A (noble gases): generally do not react (no gain or loss).
• The distinction between left-side metals and right-side nonmetals, and the staircase/metalloid boundary, was mentioned as the commonly used separator for metals vs nonmetals (the “six z line” reference in the lecture).

Naming and Classifying Periodic Table Groups

• The groups are denoted as 1A, 2A, 3A, 4A, 5A, 6A, 7A, and 8A; the B groups (3B, 4B, etc.) are less central for this lecture.
• Names for some main groups:
  • Group 1A: alkaline metals
  • Group 2A: alkaline earth metals
  • Group 7A: halogens
  • Group 8A: noble gases (also called inert gases)
• Some groups (3A, 4A, 5A, 6A) may not have widely used special names in this lecture; the focus is on the named groups above.
• The term noble gases means inert; they do not readily react with other elements (e.g., helium, neon, argon, krypton, xenon, radon).
• Halogens (Group 7A): highly reactive nonmetals (e.g., chlorine is highly reactive).

Metals vs Nonmetals and the Metals–Nonmetals Boundary

• The lecture references separating metals from nonmetals using the zig-zag boundary on the periodic table (the metal/nonmetal division).
• The left side is largely comprised of metals; the right side contains nonmetals; the region around the boundary contains metalloids.

Atomic Mass vs Mass Number: Key Differences

• Mass number (A): the total number of protons and neutrons in a specific isotope.
• Atomic mass (atomic mass unit, amu): the weighted average mass of all isotopes of an element, taking into account their natural abundances.
• Important formulas:
  • $A = Z + N$
  • $N = A - Z$
  • The atomic mass is not equal to the mass number; it is a weighted average across isotopes.

Quick Recap: Notation and Calculation Rules from the Lecture

• When given a symbol with a mass number (A) and an atomic number (Z), compute:
  • Neutron number: $N = A - Z$
  • Number of electrons (for a neutral atom): $E = Z$
• When given a symbol with a charge, adjust electron count accordingly:
  • If charge is positive, electrons < protons; cation forms.
  • If charge is negative, electrons > protons; anion forms.
• If only A is given (without charge), you assume neutral unless specified otherwise; if a charge is specified, use it to determine the electron count.

Practice Connections: Why These Concepts Matter

• Isotopes and nuclear chemistry underpin nuclear reactions, radiography, medical tracers, dating methods, and more.
• Understanding ions and charges is essential for predicting chemical reactions, bonding, and electrolyte behavior in solutions.
• The group names and the OILRIG rule help predict common ion formation patterns in reactions.
• The atomic number uniquely identifies an element; the mass number differentiates isotopes of that element; the atomic mass relates to real-world measurements and isotopic abundances.

Practical Exam Tips and Common Questions

• Know the difference between atomic number (Z) and mass number (A).
• Be able to determine N from A and Z: N = A − Z.
• Be able to determine the electron count for a neutral atom: E = Z.
• Be able to recognize cations vs anions from a given charge and element.
• Memorize key groups: 1A (alkali metals), 2A (alkaline earth metals), 7A (halogens), 8A (noble gases).
• Remember OILRIG for oxidation vs reduction, and apply to left/right side behavior.
• Understand the difference between isotopic notation forms and how to read them in two common representations.
• Keep in mind that the atomic mass is a weighted average of isotopes, not the mass number of a single isotope.

Real-World Relevance and Safety Notes

• Nuclear chemistry concepts have wide applications in medicine, energy, and environmental science.
• Isotopes can be radioactive; handling and applications require appropriate safety protocols and regulatory guidelines.

Quick Reference Equations

• Mass number and neutron number relationship:
  • $A = Z + N$
• Neutron number from given A and Z:
  • $N = A - Z$
• Isotopic notation (two common forms):
  • $^{A}_{Z}X$ or X-$A$
• Atomic mass concept: weighted average of isotopes (not explicitly numeric here; see periodic table).

End of Key Points

• Review the periodic table for Z values and the group names.
• Practice calculating N and E from A, Z, and charge as shown in the examples.
• Familiarize yourself with isotope notation forms and the difference between A and atomic mass.