Dalton’s Atomic Theory and the Periodic Table of Elements
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 (from the Greek word ‘atomos’, meaning uncuttable) and permanently unalterable.
- There are many kinds of atoms, but atoms of the same element have the same mass and properties.
- Atoms of one element are different from atoms of other elements.
- Atoms take part in chemical combinations, leading to 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 one 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.
- Two atoms of hydrogen combine with two atoms of oxygen in 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 (electrons) are embedded.
Rutherford’s Nuclear Atom
Based on his gold foil experiment, Ernest Rutherford proposed the following:
- The atom is mostly empty space.
- The positive charge is concentrated in a very small volume called the nucleus.
- The mass of the atom is concentrated in the nucleus.
Sub-atomic Particles
| Particle | Mass | Charge |
|---|---|---|
| Electron (e–) | 9.11 x 10-28 g | -1 |
| Proton (p+) | 1.67 x 10-24 g | +1 |
| Neutron (n) | 1.67 x 10-24 g | None |
Bohr’s Electronic Atom
Based on the quantum model, Bohr’s model proposed the following:
- Electrons in an atom move around the nucleus in circular orbits with definite energy levels (quantized orbits).
- Electrons can only have specific energy values.
Quantum Mechanical Atom
The Quantum Mechanical Model describes the following:
- The motion of electrons is related to standing waves.
- Heisenberg’s Uncertainty Principle states that 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
The Bohr model was a 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 model allowed electrons to occupy 3D space, requiring three coordinates, or three quantum numbers, to describe the orbitals in which electrons can be found.
Atomic Orbitals
Atomic orbitals are pictorial representations 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 describe the properties of atomic orbitals:
- Principal Quantum Number (n) – describes the size and energy level of the orbital.
- Angular Quantum Number (l) – describes the shape of the orbital.
- Magnetic Quantum Number (ml) – describes the orientation of the orbital in space.
These quantum numbers, along with the spin quantum number (ms), 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. Possible values are 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. Possible values are l = 0, 1, 2, … (n-1).
Magnetic Quantum Number, ml
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. Possible values are ml = +l … 0 … –l.
Orbital Shapes
The designations ‘s’, ‘p’, ‘d’, and ‘f’ for orbital shapes have nothing to do with their actual 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, ms
Describes the magnetic behavior of an electron, whether it is paired or unpaired (spin up or spin down). Possible values are ms = ± 1/2.
Periodic Table
The modern Periodic Table is adapted from Mendeleev’s original table. It 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. The properties of elements are periodic functions of their atomic numbers (Periodic Law).
Periodic Table – Representative Elements
The first two and last six groups represent a regular progression of chemical properties on which original periodic systems were based. These are called representative elements (formerly A groups) and are 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
Noble gases (Group 18 or VIIIA) have fully filled p orbitals (except He, with no p orbitals) and are relatively unreactive. They are sometimes designated as having zero valence.
Periodic Table – Transition Elements
Transition elements (Groups 3-12 or B groups) show a less marked progression of chemical properties. Their valence electrons (outermost) are in d orbitals (d block).
Periodic Table – Inner Transition Elements
Inner transition elements are located in the lower section of the Periodic Table. Their valence electrons are in f orbitals. They are made up of rare earth elements (lanthanides) and heavy rare earth elements (actinides). Lanthanides and actinides are extensions of Periods 6 and 7, respectively.
Other Names for Groups
- 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 is measured as atomic radius, which is half the distance from the center of like atoms when their surfaces are touching.
Periodic Trends in Atomic Size
- Decreases from left to right across a period: Due to an increase in attractive forces between protons and electrons.
- Increases from top to bottom down a group: Due to the addition of more shells at higher energy levels and increasing distance from the nucleus.
Anions are larger than the neutral atoms from which they are formed (e.g., F– > F). Cations are smaller than the neutral atoms from which they are formed (e.g., Mg2+ < Mg).
