Atoms

5th century B.C.

The idea of an atom started as early as the ___ century B.C. when Democritus expressed his belief that matter is made up of smaller, indivisible particles he called ‘atomos,’ meaning indivisible.

John Dalton

formulated a precise definition of indivisible building blocks that we now call atoms in 1808. hIS work laid the foundation for the modern era of chemistry.

Hypothesis 1: All matter consists of indivisible particles called atoms

This hypothesis of Dalton simply states that atoms are the smallest particles, and it is impossible to divide atoms even further.

Of course, nowadays we know that this is not true, as many scientists have proven the existence of particles even smaller than the atom itself.

Hypothesis 2: All atoms of the same element are identical in terms of size, mass, and chemical properties. The atom of one element is different from the atom of another element

Simply put, this hypothesis of Dalton suggests that hydrogen atoms are all the same. If you obtain hydrogen atoms from different parts of the world and in outer space, all the atoms will be the same in all respects. In the same way, oxygen atoms are all the same; however, oxygen atoms differ from hydrogen atoms. This hypothesis explains why hydrogen atoms behave similarly regardless of the source but differently with respect to oxygen atoms.

Hypothesis 3: Atoms of different elements may combine in fixed proportions to form a compound

To visualize this theory (of Dalton), imagine a water molecule that is chemically written as H₂O. This chemical formula implies that ALL molecules of water are composed of two atoms of hydrogen and one atom of oxygen. Varying the proportion of at least one atom in the formula will give rise to a completely different compound.

Hypothesis 4: Chemical reactions involve reorganization of the atoms—changes in how they are bound together. The atoms themselves are not changed in a chemical reaction

This hypothesis of Dalton is another way of stating the law of conservation of mass. In other words, atoms can neither be created nor destroyed, regardless of the type of chemical reaction it undergoes. Because atoms remain unchanged in a chemical reaction, the total mass of the reactants is equal to the total mass of the products after the reaction.

Electron

It was the very first sub-atomic particle to be discovered. It was discovered in1897, Sir Joseph John Thomson through his cathode ray experiment, and was able to determine that its charge-to-mass ratio is – 1.76 x 10⁸ C/g.

Sir Joseph John Thomson

Who discovered the electron first through his cathode ray experiment andwas able to determine that its charge-to-mass ratio is – 1.76 x 10⁸ C/g.

He also proposed that the atom is a uniform, positive sphere with electrons embedded in it like a raisin, giving rise to his plum pudding model of the atom.

Robert Millikan

Who did the oil drop experiment that allowed him to calculate the charge of an electron, which he found to be – 1.6022 x 10^-19 C.

Ernest Rutherford

The plum pudding model was the widely accepted model of the atom for a long time until He made the most surprising observation with his gold foil experiment.

• He and his colleagues’ series of experiments, they bombarded a very thin layer of gold with a positive ɑ particle.

• If Thomson’s model is correct, the diffused positive charge of the atom should have caused the ɑ particle to pass through the foil with very little deflection.

• Instead, Rutherford observed that the majority of the particles passed through the foil either completely undeflected or with very minimal deflection, while others deflected at a large angle and in extreme cases, bounced back to the direction where they came from! This observation paved the way for the new atomic model. 

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He inferred that atoms are mostly empty space. This is why most of the ɑ particles passed through the foil either undeflected or slightly deflected.

• As for the particles that either largely deflected or bounced back, He proposed that contrary to Thomson’s model, an atom has its positive charges concentrated in its core, which he called the nucleus.

• As the positive ɑ particle approaches the positively-charged nucleus, it experiences repulsion due to similar charges, causing the particle to be deflected at such a large angle.

• Meanwhile, ɑ particles that directly hit the nucleus are deflected back towards the direction where they came from. His model gave rise to the atomic model that we know today!

James Chadwick

In 1932, He was able to account for this unexplained mass ratio:

(The atomic mass of hydrogen, an element with 1 proton and 1 electron, is 1.008 g/mole. Meanwhile, helium, an element with 2 protons and 2 electrons, has an atomic mass of 4.003. Therefore, the ratio of their atomic masses is about 1:4.

If protons and electrons are the only subatomic particles, their mass ratio should be 1:2 (the mass of electrons is usually omitted since it is 1840 times lighter than the proton!).

due to his discovery of neutrons. Succeeding experiments proved that the third subatomic particle is electrically-neutral, hence the name neutron.

Proton

1.67262 x 10^-24 g, 1.6022 x 10^-19 C, Charge unit: +1

Neutron

1.67262 x 10^-24 g, 0 C, Charge unit: 0

Electron

9.10938 x 10^-28 g, -1.6022 x 10^-19 C, Charge unit: -1

The answer is because of the difference in the properties of the atom, which can be ultimately attributed to the number of protons and neutrons, as well as the number and distribution of electrons in the space around the nucleus.

All the substances found on Earth are made up of atoms. If that’s the case, then how come a lot of substances vary significantly from one another?

atomic number

(designated as Z). In the modern periodic table, an atomic number is usually written on the upper left side corner of each element block.

• This number represents the number of protons in the atom’s nucleus. It also gives us an idea about the reactivity of the atoms.

• For instance, all atoms with an atomic number of at least 84 are radioactive.

mass number

If the number of neutrons is added to the number of protons, it becomes the _______ (no. of protons + no. of neutrons). The mass number is usually designated as A and is not written in the modern periodic table of elements.

atomic mass

This property is written below the element name in the periodic table. Due to huge differences between the mass of an electron, and proton and neutron, atomic masses are mostly attributed to the mass of protons and neutrons only (the mass of electrons is almost negligible compared to these two). - Also includes and calculates the average mass of its isotopes.

Neils Bohr

He intorduced the planetary model, wherein he proposed that electrons revolve around a positive nucleus in a predetermined orbit, just like how the planets in the solar system revolve around the sun.

Erwin Schrödinger

In 1926, He developed his famous equation, which gave birth to the quantum mechanical model. Contrary to Bohr’s planetary model, this model explains that electrons are most probably found in a three-dimensional space around the nucleus, which is known as the orbital.

The Principal Quantum Number (n)

It refers to the main energy levels (or shells) of an orbital. To better understand the principal QN, visualize an imaginary circle around the nucleus of an atom like what is shown below.

• The first circle corresponds to n = 1, and the electrons occupying the n = 1 _______ can be found anywhere within the space enclosed by the said energy level.

• This space enclosed by n = 1 is what we call the orbital.

• This can take values from 1 to ∞, and the higher the value of n, the higher is the energy of the orbital, and the farther the electron from the nucleus.

the larger the atom is

Another thing that the principal QN tells us is the size of the atom. Obviously, the higher the n value, __________.

Azimuthal quantum number

It also known as the angular momentum quantum number, pertains to the energy sublevels or subshells of the orbital.

• This quantum number has something to do with the shape of the orbital, and can only take integral values between 0 to n-1.

• Different shapes of orbitals are assigned different letters and ℓ values.

s (sharp) orbital

Its shape is sperical and its ℓ value is 0.

p (principal) orbital

Its shape is a dumb-bell and its ℓ value is 0

d (diffuse) orbital

Its shape is not easily visualized and its ℓ value is 2

f (fundamental)

Its shape is not easily visualized and its ℓ value is 3

Magnetic Quantum Number (mℓ)

Whereas ℓ has something to do with the shape of the orbital, this tells us about the orientation of orbitals in space.

• This quantum number takes the values between —ℓ to +ℓ, including zero.

• Therefore, for an s orbital, the only possible value of mℓ is 0, which implies that there is only 1 s orbital.

• Meanwhile, a p orbital can have an mℓ value of —1, 0, and +1, which implies that there are three types of p orbitals, each of which assumes a different orientation in space

Spin Quantum Number (ms )

This pertains to the electron spin, which can only be either clockwise or counterclockwise. As a result, there are only two possible values of ms , and these are +½ and —½.

Electron configuration

It is the distribution of electrons among the various orbitals in an atom, molecule, or ion

• the number of electrons present in a certain species

• the number of electrons each orbital can occupy

• the correct ordering of the orbitals

To be able to write electron configuration correctly, it is important that you know three things:

2 electrons

As for the number of electrons that each orbital can occupy, the s orbitals can occupy a maximum of

6 electrons

p orbitals can accommodate this many electrons

10 electrons

d orbitals can take this maximum amount of electrons

14 electrons

f orbitals can have this maximum number of electrons

The Aufbau Principle

It is a German term that means building up. This principle is named so because it dictates that as a proton is being added one by one to the nucleus to build up the element, electrons are likewise added to the orbital.

• Such principle gives rise to the question “As more electrons are added to the atom, what orbital will the electron fill up? Do we start with orbitals with the highest energy level, or do we start with the orbitals with the lowest?”

• The (n + ℓ) rule answers the dilemma established from this principle. According to the (n + ℓ) rule, the added electrons must occupy an orbital with a lower (n + ℓ) value first. In other words, the added electrons should occupy the orbital with lower energy first.

• This means that as electrons is being added, the 2s orbital (n + ℓ = 2 + 0 = 2) must be filled first before the electrons occupy the 2p (n + ℓ = 2 + 1 = 3) orbital.

• In cases wherein two different orbitals have the same (n + ℓ) value, say 3d (n + ℓ = 3 + 2 = 5) and 4p (n + ℓ = 4 + 1 = 5), the orbital with the lower principal QN must be filled first. These principles govern the mnemonics mentioned.

Hund’s rule of multiplicity

It states that every orbital in a subshell is singly occupied with one electron before any orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

Pauli's exclusion principle

It states that no two electrons can have the same set of four quantum numbers. thus two arrows with the opposite directions.

• If we consider the two electrons in the 2s orbital, both have the same principal QN (n = 2), azimuthal QN (ℓ = 0 for s orbital), and magnetic QN (mℓ = 0). What differentiates the two is the spin QN; one is +½, while the other one is -½.

• This difference is manifested in orbital diagrams as upward or downward pointing arrows. There is no convention as to which arrow pertains to +½ and -½ spin, although usually, +½ is given to the arrow pointing upward, and -½ to the arrow pointing downward.