Nuclear Symbols

1. Nuclear Symbols Nuclear Symbols are used to represent atoms and their isotopes • Example: uranium-238 or... U-238 The 238 is the atomic mass of the uranium atom • Why? Because uranium has other isotopes with different masses • Ex: uranium-235 and uranium-234 What does a nuclear symbol tell us? • The element(in this case uranium) • top number is the mass number # protons + # neutrons • bottom number is the atomic number # protons • You can determine the number of neutrons by subtraction mass number → atomic number → 238(mass #) - 92 (atomic #) = 146 (# of neutrons)

2. Ions • Ions – atoms that have lost or gained electrons • Positive ions have lost electrons. • Positive ions are called cations. • Examples: • Na+ lost one electron • Mg2+ lost two electrons • Al3+ lost three electrons • Negative ions have gained electrons • Negative ions are called anions. • Examples: • Cl-gained one electron • O2-gained two electrons • N3–gained three electrons

3. Orbits vs. Orbitals ► Bohr model places electrons in orbits(fixed paths) to represent principal energy levels The Electron cloud model puts electrons in orbitals - They are still in energy levels, but moving so fast their exact location cannot be predicted. There are four possible orbitals - “s” orbital – holds 1 pair of electrons – 2 total - “p” orbital – holds 3 pairs of electrons – 6 total - “d” orbital – holds 5 pairs of electrons – 10 total - “f” orbital – holds 7 pairs of electrons - 14 total 1st principal energy level has only one orbital: s 2nd principal energy level has two orbitals: s and p 3rd has three: s, p and d 4th has four: s, p, d and f 5th has four: s, p, d and f 6th has three: s, p, and d 7th has two: s and p

4. Electron Configuration • Where do these electrons go? • Not sure, they're moving fast! • but we can say where they probably are... • Electron configuration: The arrangement of electrons of an atom in its ground state into various orbitals around the nucleus. • There are three rules: • 1. Aufbau Principle: each electron fills the lowest energy orbital available 1st. • 2. Pauli's Exclusion Principle: a maximum of two electrons may occupy a single sub-orbital, but only if they have opposite spins. • 3. Hund’s Rule: single electrons with same spin must occupy each sub-orbital before additional electrons with opposite spins can occupy the same orbitals.

5. Electron Configuration • Determining Electron Configuration: • Step 1: Determine how many electrons the atom has • Step 2: Fill the lowest energy level first. • Step 3: Use superscripts to show how many electrons are in the orbital. • - Example: Helium has an atomic number of 2. This means it also has 2 electrons. • - The lowest energy level is 1s • - Helium's 2 electrons fit into the 1s orbital like so: • 1s2 • - Therefore, the electron configuration for Helium is: • 1s2 Carbon has ___ electrons: 6 Iron has ___ electrons: 26 1s2 2s2 2p63s2 3p64s2 3d6 1s2 2s2 2p2 Add up the superscripts, you get 26! Add up the superscripts, you get 6! Once you run out of electrons, stop.

6. Ground State Electron Configuration Examples: Helium – Boron – Aluminum – Copper – Ground State 1s2 1s22s22p1 1s22s22p63s23p1 1s22s22p63s23p64s23d9 Excited State 1s12s1 1s22s12p2 1s12s22p63s23p14s1 1s22s22p53s23p64s23d10 2 e- 5 e- 13 e- 29 e- - Electron configurations for the ground state always put electrons in the lowest possible energy level. - Excited state E.C.’s will have 1 or more electrons moved to a higher energy level.