Introduction to Atomic Structure
- Considered the building blocks of matter
- Made up of three main subatomic particles:
- All protons are identical, regardless of the element in which they are found. This is also true for neutrons and electrons.
- Atoms have a small, dense nucleus in the middle but most of the atom is nothing more than empty space!
- Extremely small in size relative to the atom but it contains 99.99% of the atom’s mass.
- If the atom were enlarged to the size of a football stadium, the nucleus would be the size of a bee!
- Contains 2 of the 3 subatomic particles
- Protons (+ charged particles)
- Neutrons (neutrally charged particles)
Atomic Mass Units
- Mass of proton = 1 amu (atomic mass unit)
- Neutrons have slightly more mass than protons but they are considered equal.
- Mass of neutron still = 1 a.m.u.
- Mass of electron is tiny even on an atomic scale (0.0006 amu).
- # of protons in nucleus determines the atomic number of an element.
- This identifies the element!
- Hydrogen--Atomic #1 = 1 proton
- Helium--Atomic #2 = 2 protons
- Oxygen--Atomic # 8 = 8 protons
- All atoms with 8 protons MUST be oxygen!
- The # of protons in an atom CANNOT be changed (otherwise, you will have a different element!)
- Even though the # of protons will never change, the # of neutrons can vary from atom to atom.
- Atoms of the same element that have the same # of protons, but different number of neutrons are called isotopes .
- Every single atom is going to be an isotope of that element!
- Hydrogren has 3 different isotopes (all of them have 1 proton)
- Every atom of “H” is going to be one of these isotopes!
- Mass # of an atom is the sum of the protons and neutrons in its nucleus.
- The electrons are ignored because they are so small!
- The mass number varies for different isotopes of an element.
- Carbon has two known isotopes:
- Carbon-12 (6 protons and 6 neutrons) (6 electrons)
- Carbon-14 (6 protons and 8 neutrons) (6 electrons)
- Any sample of an element as it occurs in nature contains a mixture of isotopes.
- The atomic mass of an element is the average mass of all the isotopes of that element. Therefore, the atomic mass is usually not a whole number.
- For example, the atomic mass for carbon is 12.011
- Remember, there are two naturally-occurring isotopes: C-12 and C-14
- Since the atomic mass is much closer to “12”, this tells you that there are MANY more atoms of Carbon-12 than there are Carbon-14.
- Electrons have a negative charge and are found outside of the nucleus.
- In an uncharged atom, the # of electrons = # of protons
- Always assume that the atom has a zero charge and is neutral.
- Sometimes, atoms will either lose or gain electrons.
- When this happens, they do not equal the # of protons, and the atom becomes positively or negatively “charged”.
- These charged atoms are called ions .
- Space in which electrons are found.
- Arranged in energy levels:
- Lower energy level = closer to the nucleus
- Higher energy level = further from nucleus
- 1 st energy level - 2 electrons.
- 2 nd energy level -8 electrons
- 3 rd energy level -18 electrons
- An atom’s bonding ability is determined by the arrangement of electrons in the outermost energy level.
- These are called valence electrons.
- Atoms that have only one electron or those that only need one electron are much more likely to bond.
- Some elements have a complete set of electrons and will not bond (Noble Gases)
- It is possible to get even smaller than these three subatomic particles.
- Current theory states that protons and neutrons are themselves made up of even smaller particles known as quarks.
- There better our technology gets, the more subatomic particles there are being discovered.