Download Nuclear Symbols

Nuclear Symbols
Nuclear Symbol - used to represent atoms and their isotopes
Isotopes – atoms with the same atomic number but different atomic masses due to
different number of neutrons
Nuclear symbols tell us:
mass number →
 The element (ex. uranium)
 top number is the mass number
# protons + # neutrons
atomic number →
 bottom number is the atomic number
# protons
 You can determine the number of neutrons by subtraction
238 (mass #) - 92 (atomic #) = 146 (# of neutrons)
Atomic Mass & Isotopes
**The Atomic masses in the Periodic Table are not mass numbers.**
 They are an average of all the isotopes of that element, weighted by
abundance.
·
Rounding the ave. atomic mass in the P.T. usually gives you the most common
isotope. For iron (mass 55.847 amu in P.T.) the most common isotope is iron-
56
_________
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
Orbits vs. Orbitals
► Bohr model: energy levels are “orbits”
► Modern Model: energy levels are broken into
orbitals
There are 4 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
The Electron Hotel
Electron configuration: The arrangement of electrons of an atom in its ground
state into various orbitals around the nucleus.
 “Electron Hotel”
 Electrons fill orbitals like people filling rooms in a hotel


Like a hotel, there are rules:
1. Electrons fill the lowest energy level available first. (Aufbau
Principle)

Electrons fill in the 1st hotel level first, before filling up the 2nd level
2. Single electrons with same spin must occupy each sub-orbital before
additional electrons with opposite spins can occupy the same
orbitals.(Hund’s Rule)
• Only one electron per room until all the rooms in that orbital have at least one
electron
3. No more than 2 electrons may occupy a single suborbital, and they
must have opposite spins. (Pauli’s Exclusion Principle)
• No more than 2 electrons in a room, and they must have opposite spins
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 ___
6 electrons:
1s2 2s2 2p2
Add up the superscripts, you get 6!
Once you run out of electrons, stop.
Iron has ___
26 electrons:
1s2 2s2 2p63s23p64s2 3d6
Add up the superscripts, you get 26!
Ground State Electron Configuration
Examples:
Helium –
2 e-
Ground State
1s2
Excited State
1s12s1
Boron –
1s22s22p1
1s22s12p2
5 e-
2
2
6
2
1
Aluminum – 13 e- 1s 2s 2p 3s 3p
Copper –
29
e-
1s22s22p63s23p64s23d9
1s12s22p63s23p14s1
1s22s22p53s23p64s23d10
- 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.