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Atomic Structure and Electron Configuration

The Structure of an Atom

An atom is the smallest particle of an element and cannot be broken down further without changing it.
All atoms have a similar structure, with a nucleus and orbitals.
Orbitals are electron clouds representing possible electron locations.
Subatomic particles (protons, neutrons, and electrons) are found in the nucleus or orbitals.

Subatomic Particles

Protons: Positively charged.
Electrons: Negatively charged.
Neutrons: Neutral (no charge).
An atom can have any combination of these particles.
In a neutral atom, the number of protons equals the number of electrons.
The number of neutrons can vary, even within the same element.

Location of Subatomic Particles

Nucleus: Contains protons and neutrons.
Orbitals: Contain electrons.
The number of each subatomic particle determines the type of atom.
Different elements have different particle counts and arrangements.
Example: Helium typically has two protons and two neutrons in the nucleus and two electrons orbiting the nucleus. When the protons and electrons are equal, the atom is neutral.
Larger atoms have more electrons and, therefore, more orbitals.

Mass of Subatomic Particles

The three types of particles do not have the same mass.
Each particle is about 6 (with 23 zeros) times smaller than a paperclip.
Scientists use the atomic mass unit (AMU) to measure the mass of subatomic particles on an atomic scale.
One atomic mass unit is defined as 1/12 the mass of a carbon-12 atom, which corresponds to 1.660538921 \times 10^{-24} grams.
Proton: 1.673 \times 10^{-24} g OR 1 amu.
Neutron: 1.675 \times 10^{-24} g OR 1 amu.
Electron: 9.109 \times 10^{-28} g OR 0.0006 amu.
The nucleus contains most of the mass of the atom because it contains all of the protons and neutrons.

Quarks

Smaller particles that make up protons and neutrons.

Atomic Number (Z)

All elements are different from each other because of the different number of each type of subatomic particle in each type of atom.
The atomic number (Z) is the number of protons in an atom, which differs for each element.
It is equal to the number of protons in the nucleus of every atom.

Periodic Table Organization

Elements are organized by atomic number on the periodic table.
The atomic number is written above the element symbol (e.g., Al -> 13), indicating the number of protons.
As you move right across the periodic table, the elements increase in atomic number.
To the right of aluminum is silicon, which has an atomic number of 14.

Ions and Charge

The number of electrons can vary and cannot be used to identify an atom.
Atoms can gain or lose electrons to form ions (atoms with a charge) and still be the same element.
Elements on the periodic table are in their elemental form (neutral atom).
In a neutral atom, the number of electrons is equal to the number of protons (atomic number).

Examples

Aluminum always has 13 protons. When it's neutral, it has 13 electrons. The net charge is 0 when there are 13 protons (+13) and 13 electrons (-13).
Elements can lose or gain electrons to form ions.
When aluminum loses three electrons, it has a +3 charge.
- It still has 13 protons in the nucleus because it's still aluminum.
- Now, it has only 10 electrons (-10), resulting in a +3 charge (13 - 10 = 3).
Atoms can also gain electrons.
- If aluminum gains five electrons, it would have a -5 charge due to the extra five electrons.
- 13 protons plus 18 electrons (13 + 5) results in a total charge of -5.
Some elements are much more likely to gain electrons than to lose them, and vice versa.

Mass Number (A)

The mass number provides the total mass of the nucleus or the total number of protons + neutrons.
Mass number is represented by "A".
The mass number can be different in different atoms, even in the same element.

Representation

Aluminum-27
Al-27
^{27}Al
The mass number is at the top left, and the atomic number is at the bottom left.

Calculating Neutrons (N)

N = A - Z
Example: For aluminum-27, N = 27 - 13 = 14 neutrons.
If you know the mass number and the atomic number, you can calculate the number of neutrons in the atom.

Isotopes and Atomic Mass

All atoms of a given element have the same atomic number.
However, atoms of the same element may have different mass numbers.

Isotopes

Atoms of the same element with different mass numbers are called isotopes.
Isotopes have the same number of protons and electrons but different numbers of neutrons.
The atomic number of an element is always the same, but the mass number can be different.

Example

Aluminum-30 has 3 more neutrons than aluminum-27, but they all have 13 protons and 13 electrons (if neutral).

Atomic Mass

The atomic mass of an atom is determined by the sum of all protons and neutrons in the nucleus.
This number is basically the entire mass of the atom.
The mass of an electron is much smaller than the mass of a proton or neutron and is usually not included in the atomic mass calculation.
The mass of a proton is approximately equal to the mass of a neutron.

Average Atomic Mass

In the periodic table, the number below the elemental symbol represents the average atomic mass of the element.

Formula

M(a) = (M1)(P1) + (M2)(P2) + … + (Mn)(Pn)
- M(a) = average atomic mass/weighted average mass of the element
  - Weighted mass means that more abundant isotopes affect the average more than rare ones.
- M1, M2 = atomic mass of isotope 1, isotope 2, etc.
- P1, P2 = fractional abundance of isotope 1, isotope 2, etc.
To find the average atomic mass, sum the products of the mass of each isotope by its fractional abundance.
To use the percent abundance, divide it by 100.
Multiply the mass of each isotope by its fractional abundance.

Electron Shells and Subshells

Electron Shell

A set of orbitals with the same principal quantum number (n).

Principal Quantum Number (n)

The number that describes the size of the orbital.
Represented by circles in diagrams but is not actually accurate.
Shells are filled consecutively in atoms, from the closest to the atom outward.
The most stable when they’re at the lowest energy level, fills in from the closest to the atom outward
Each shell can hold a different number of electrons.
- Example: N = 1 holds 2 electrons, but N = 2 holds 8 electrons.
A full electron shell is more stable than a partially filled shell.
- Example: Argon is unreactive because it has full electron shells.

Electron Subshell

A set of orbitals with the same principal quantum number (n) AND the same angular momentum quantum number (l).

Angular Momentum Quantum Number (l)

The number that describes the shape of an orbital.
- S ; l = 0
- P ; l = 1
- D ; l = 2
- F ; l = 3
Values of (m) indicate the number of subshells of each type.
- Example: There's only one S subshell for each electron shell because M can have only one value, 0.
- (m) can be anywhere from -l to +l.
- If (l) is 0, then (m) can only be 0.
- If there are 7 f subshells, (m) could be -3, -2, -1, 0, 1, 2 or 3 because f is always equal to three regardless of how many there are

Example: How many electrons can an N = 3 shell hold?

N = 3, so L = 0, 1, or 2.
Each value corresponds to a specific subshell:
- 3s (l = 0, m = 0)
- 3p (l = 1, m = -1, 0, 1)
- 3d (l = 2, m = -2, -1, 0, 1, 2)
From this subshell, you can determine the value (m).
- The ammount of orbitals can be now calculated by adding up the total combinations
- 3s + 3p + 3d = x
- 1 + 3 + 5 = 9
- Each orbital can hold two electrons, so multiply x by 2
- 9 x 2 = 18 electrons

Using the 2n^2 equation

2n^2
2 \times 3^2
2 \times 9 = 18

Subshell	No. of Orbitals	No. of Electrons
3s	1	2
3p	3	6
3d	5	10

Electron Configuration and Orbital Notation

Electron Configuration

Shows how electrons are positioned in an atom.

Orbital Notation

A diagram that uses lines and arrows to show shells, subshells, and orbitals for electrons in an atom.
The lines represent unoccupied orbitals.
- The numbers and letters on the bottom represent the orbitals name.
- EX. 1s --> quantum number one and second shell is s.
The arrows represent the electrons.
- The up arrow represents

Pauli Exclusion Principle

No two electrons can have identical quantum numbers and always must differ even by just one number.

Electron Spin Quantum Number (m_s)

The fourth quantum number and describes the angular momentum of an electron.
Can have a value of +½ or -½ no matter what the values of (n), (l) or (m) are.
Often indicated with up or down arrows because electrons repel each other.
Explains why only two electrons are allowed per orbital and, therefore, two possible values for (m_s).

Writing Orbital Notation

Identify the number of electrons from the periodic table.
- It must always have the same electrons and protons to be neutral.
Orbitals must be filled in order.
- Start with the lowest subshell first, which can only hold two electrons each.
- Keep going until you run out of electrons, each subshell only having two each

Hund's Rule

Electrons in the same sublevel (p), (d), or (f) are placed in individual orbitals before they are paired up in order to increase atomic stability

EX. a mother of three children would make sure all three have at least one thing before handing out seconds.

they will not occupy an orbital that already contains an electron if theres another orbital in the same subshell thats empty

as electrons are added to the orbitals, they will all have the same spin quantum number until they start filling half full orbitals.

Dot Structures

Circle: electron shell
Dot: electron
Does not show the size, shape, orientation of orbitals/electrons accurately.
A convenient way of showing the distribution of electrons in their shells.

Practice

2s → n = 2, l = 0
3d → n = 3, l = 2
1s → n = 1, l = 0
2p → n = 2, l = 1
4f → n = 4, l = 3

Noble Gas Configuration

Short hand notation for writing the electron configuration for an element, substituting the symbol of a noble gas.
Takes the elemental symbol of the last noble gas prior to the atom, followed by the configuration of the remaining electrons.
Helps separate the electrons of any atom into two groups:
- Those involved in a chemical reaction.
- Those not involved in a chemical reaction.
Example: For sodium, [Ne] 1s22s2p6 → [Ne]3s1

Core vs. Valence Electrons

Noble gases are the elements in the far right group of the periodic table.
Outer electrons typically involve reactions with other atoms.
Two sets of electrons are present in an atom: core electrons and valence electrons.
The more reactive electrons in an element are the valence electrons.
They are located outside of the noble gas core.
They are the electrons that are written out in detail whenever you write a noble gas configuration.
In periods 1, 2, and 3, the valence electrons are all located in the orbital with the highest value of (n) equaling 3.
Depending on the element, there might be other electrons in orbitals that are hidden underneath the outer valence electron shell.
The least reactive electrons in an atom are the core electrons.
They are the same ones that are abbreviated in the noble gas notation.
(Valence electrons are the ones that are written out.)
Noble gases are very nonreactive.
They are filled with electrons, which makes them very stable.
Core electrons are also very stable and nonreactive, unlike the valence ones.
All inner core electrons, nonvalence electrons, are core electrons.
Noble gases have a full valence shell, which is why they are nonreactive.
They usually have eight electrons in their valence shell, except for helium, which only has two.

Valence Electrons for Periods 4 and Higher

(s) and (p) orbitals with the highest n → valence electrons
Partially filled (d) and (f) orbitals → valence electrons
Full (d) and (f) orbitals → core electrons

Filling Orbitals

Electrons always fill the lowest energy orbitals first, like going up the stairs one at a time.
They fill the lower energy levels first and move up as each new electron is added.
Each time you move right within a row on the periodic table, it's a different element, and one more electron must be added to the next available orbital.
- A set of electrons is already present in the lower energy orbitals.
The order of filing orbitals is determined by the energy of the orbitals (lowest to highest).
This path of filling orbitals can be followed by going left to right on the periodic table.
- The first element is hydrogen → 1s
- Go to the right and add another electron → helium 1s2
- Go down one row and begin again with lithium → 2s

P-Block Elements

(p) block elements are in groups 13 - 18
Final electron goes into a (p) orbital, while valence electrons go into (s) and/or (p)
- Example: sulfur → noble gas before is neon → [Ne] 3s^2 3p^4
(p) block elements in periods 1 - 3 are straightforward and only fill (s) or (p) orbitals
(p) block elements in period 4 or higher are more complex
- Their valence electrons include (d) orbitals too
Orbitals fill order around this area in the order of 4s → 3d → 4p

D-Block Elements

(d) block includes group 3 - 12
Last added electron is always in a (d) orbital
Valence electrons are in (s) and (d) orbitals
- Example: iron → noble gas before is argon → [Ar] 4s^2 3d^6