Dalton’s Atomic Theory and the Periodic Table of Elements

Posted on Aug 27, 2024 in Chemistry

Dalton’s Atomic Theory

Accepted for much of the 19th century, Dalton’s Atomic Theory proposed the following:

• All matter is made up of tiny, invisible particles called atoms.
• Atoms are indivisible (atomos, Gr.), permanently unalterable.
• There are many kinds of atoms, but atoms of the same element have the same mass and properties.
• Atoms take part in chemical combinations.

This theory led to the development of the Law of Definite Composition (LDC) and the Law of Multiple Proportions (LMP).

LDC – Law of Definite Composition

When two or more elements combine chemically to form a compound, they will always combine in a fixed proportion by weight.

LMP – Law of Multiple Proportions

When two or more elements combine chemically to form more than one compound, the variable weights of the ones that combine with the fixed weights of the others are in a ratio of simple whole numbers. In other words, elements combine in a simple whole number ratio.

Law of Definite Composition

A compound is always composed of the same elements in the same proportion by mass, no matter how large or small the sample. Relative amounts are expressed as percent by mass, the ratio of the mass of each element to the total mass of the compound expressed as a percentage.

Law of Multiple Proportions

When two or more elements combine chemically to form more than one compound, the variable weights of the ones that combine with the fixed weights of the others are in a ratio of simple whole numbers.

For example, hydrogen and oxygen combine to form water (H₂O) and hydrogen peroxide (H₂O₂).

• Two atoms of hydrogen combine with one atom of oxygen in water, while two atoms of hydrogen combine with two atoms of oxygen in the case of hydrogen peroxide.
• The ratio of oxygen atoms combining with a fixed number of hydrogen atoms in the two compounds is 1:2.
• Every 2.016 g of hydrogen combines with 15.999 g of oxygen to form water and 31.998 g of oxygen to form hydrogen peroxide.
• The ratio is 15.999 : 31.998 or 1:2, which is a simple whole number ratio.

Plum Pudding Model

Developed by J.J. Thomson in 1904, the Plum Pudding Model described the atom as a sphere of positive charge where negative charges are embedded.

Rutherford’s Nuclear Atom

Rutherford’s Nuclear Atom proposed the following:

• The atom is mostly empty space.
• Positive charge is concentrated in a very small volume: the nucleus.
• The mass of the atom is concentrated in the nucleus.

Sub-atomic Particles

Bohr’s Electronic Atom

Based on the quantum model, Bohr’s Electronic Atom proposed the following:

• Electrons in an atom move around the nucleus in circular orbits with definite energy levels (quantized orbits).
• Electrons have energy with specific values only.

Quantum Mechanical Atom

The Quantum Mechanical Atom proposed the following:

• The motion of electrons is related to standing waves.
• Heisenberg Uncertainty Principle: It is impossible to know simultaneously the momentum and position of an electron precisely.
• An orbital defines the probability of finding an electron in a particular location.

Atomic Orbitals

Bohr Model

• 1D model that used one quantum number to describe the distribution of electrons in an atom.
• The only important information was the size of the orbit.

Schrödinger’s Atom

• Allowed electrons to occupy 3D space.
• Required three coordinates, or three quantum numbers, to describe orbitals in which electrons can be found.

Atomic Orbitals

Atomic orbitals are a pictorial representation of the solution to Schrödinger’s equation. They describe the probability of finding an electron in space.

Heisenberg’s Uncertainty Principle

The position and energy (momentum) of an electron cannot be measured accurately at any given time.

Quantum Numbers

Three coordinates from Schrödinger’s wave equations:

• Principal Quantum Number (n) – describes size
• Angular Quantum Number (l) – describes shape
• Magnetic Quantum Number (m_l) – describes orientation in space

Solutions to Schrödinger’s equation only allow specific values, meaning electrons may be treated as discrete particles. n, l, m_l, and m_s (spin quantum number) completely describe an electron.

Principal Quantum Number, n

Describes the size of the orbital and indirectly describes the energy level or energy shell the electron is in. It also indirectly describes atomic size. n = 1, 2, 3, 4…

Azimuthal Quantum Number, l

Also known as the angular or orbital quantum number, l identifies the type or shape of the orbital (subshell). Orbitals have shapes that are best described as spherical (l = 0), polar (l = 1), or cloverleaf (l = 2). They may even take on more complex shapes as the value of the angular quantum number becomes larger.

• n=1, l = 0 1s orbital
• n=2, l = 0 2s orbital, l = 1 2p orbital
• n=3, l = 0 3s orbital, l = 1 3p orbital, l = 2 3d orbital
• l = 0, 1, 2, . . . (n-1)

• n=4, l = 0 4s orbital, l = 1 4p orbital, l = 2 4d orbital, l = 3 4f orbital
• l = 0, 1, 2, . . . (n-1)

Magnetic Quantum Number, m_l

Specifies the orientation of orbitals in space. It was called the magnetic quantum number because the effect of different orientations of orbitals was first observed in the presence of a magnetic field.

• n = 1, l = 0, m_l = 0 1s orbital
• n = 2, l = 0, m_l = 0 2s orbital
• l = 1, m_l = +1, 0, -1

• n=3, l = 0, l =1, l = 2
• m_l = +2, +1, 0, -1, -2
• m_l = +l . . . 0 . . . –l

Orbital Shapes

Folklore: ‘s’ stands for ‘spherical’ and ‘p’ stands for ‘polar’ to designate the shapes of s and p orbitals. D then is for? Reality: designations have nothing to do with orbital shapes. They refer to groups of lines in the spectra of alkali metals. These line groups were called sharp (s), principal (p), diffuse (d), and fundamental (f).

Spin Quantum Number, m_s

Describes the magnetic behavior of an electron, whether paired or unpaired (spin up or spin down). m_s = ± 1/2

Periodic Table

Adapted from Mendeleev’s, the Periodic Table lists elements according to increasing atomic numbers rather than atomic weights.

• Vertical Columns = groups = family of elements
• Horizontal Rows = periods = series

At certain intervals (8, 18, or 32), there is a periodic recurrence of the detailed structure of atoms of elements. Properties of elements are periodic functions of their atomic numbers (Periodic Law).

Periodic Table

The first two and last six groups represent the regular progression of chemical properties on which original periodic systems were based.

• Representative elements – used to be A groups
• Related to the progressive addition of electrons in s and p orbitals
• Groups 1 & 2 (IA and IIA) – s block groups
• Groups 13-18 (IIIA – VIIIA) – p block groups
• Group 18 (0 or VIII A) – noble gases (p orbitals filled up with 8 electrons), except for He

Periodic Table – Noble Gases

• Group 18 (0 or VIIIA)
• Fully filled p orbitals (except He, with no p orbitals)
• Relatively unreactive
• Sometimes designated zero valence

Periodic Table – Transition Elements

• Groups 3-12 (B groups)
• Progression of chemical properties not so well marked
• Valence electrons (outermost) are in d orbitals – d block

Periodic Table – Inner Transition Elements

• Lower section
• Valence electrons in f orbitals
• Made up of rare earth elements (lanthanides) and heavy rare earth elements (actinides)
• Lanthanides and actinides = extensions of Periods 6 and 7 respectively

Other Names

• Group 1: alkali metals
• Group 2: alkaline earth metals
• Group 11: coinage metals (not an IUPAC approved name)
• Group 15: pnictogens (not an IUPAC approved name)
• Group 16: chalcogens
• Group 17: halogens
• Group 18: noble gases

Periodic Properties – Atomic Size

• Atomic size: measured as atomic radius
• Atomic radius: Half the distance from the center of like atoms when these atoms’ surfaces are touching

Periodic Properties – Atomic Size

• Decreases from left to right due to an increase in attractive forces between protons and electrons
• Increases from top to bottom due to the addition of more shells at higher energy levels and increasing distance from the nucleus

Periodic Properties – Atomic Size

• Anions are larger than neutral atoms from where they are formed (F^– > F)
• Cations are smaller than neutral atoms from where they are formed (Mg²⁺ < Mg)