Atomic Structure and The Periodic Table Study Guide

Early Models of the Atom and Dalton's Atomic Theory

Dalton's Atomic Theory:
- All elements are composed of tiny, indivisible particles called atoms.
- Atoms of the same element are identical in their properties and mass. The atoms of any one element are distinct from those of any other element.
- Atoms of different elements can physically combine or chemically unite with one another in simple whole-number ratios to form chemical compounds.
- Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another element as a result of a chemical reaction.
Definition of an Atom:
- An atom is the smallest particle of an element that retains the chemical and physical properties of that element.
Counting Atoms in Chemical Compounds:
- Carbon Dioxide ( $\text{CO}_2$ ): Contains $1$ carbon atom and $2$ oxygen atoms, totaling $3$ atoms.
- Water ( $\text{H}_2\text{O}$ ): Contains $2$ hydrogen atoms and $1$ oxygen atom, totaling $3$ atoms.
- Sulfuric Acid ( $\text{H}_2\text{SO}_4$ ): Contains $2$ hydrogen atoms, $1$ sulfur atom, and $4$ oxygen atoms, totaling $7$ atoms.

Subatomic Particles and Discovery of Atomic Structure

The Electron:
- A negatively charged subatomic particle located outside the atomic nucleus.
- Carries a relative electrical charge of $-1$ .
- Has a mass equal to approximately $\frac{1}{1840}$ the mass of a single hydrogen atom.
Cathode Ray Tube (CRT) Experiments:
- A CRT is a sealed glass vacuum tube device used to originate and study a flow of charged particles.
- Components include a negative cathode electrode, a positive anode electrode, and deflecting plates ( $+$ and $-$ ).
- Any metal can serve as the electrode material.

Cathode Ray Tube diagram showing deflecting plates and electron beam deflection

J.J. Thomson:
- Credited with the discovery of the electron.
- Used cathode ray tubes to measure the deflection of charged rays across magnetic and electric fields using various gases and different metal electrodes.
- Observed that the ray traveled from the cathode to the anode and deflected toward a positively charged plate, proving the ray consisted of negatively charged particles termed electrons.
- Calculated the charge-to-mass ratio ( $\frac{e}{m}$ ) of the electron.
Robert Millikan:
- Conducted experiments that led to the precise determination of the fundamental charge and mass of the electron in $1916$
E. Goldstein:
- Discovered positively charged rays (canal rays) traveling in the direction opposite to cathode rays in cathode ray tubes.
- Identified these positively charged subatomic particles, which were later named protons.
Proton:
- A positively charged subatomic particle carrying a relative electrical charge of $+1$ .
- Located inside the atomic nucleus.
Neutron:
- Discovered by James Chadwick in $1932$
- Carries no electrical charge ( $0$ neutral charge).
- Possesses a mass approximately equal to that of a proton.
- Located inside the atomic nucleus.
Thomson's "Plum Pudding" Model:
- The discovery of subatomic particles required modifying Dalton's solid indivisible sphere model.
- Thomson proposed that the atom consisted of negative electrons embedded uniformly within a sphere of continuous positive charge.

Plum pudding model showing negative electrons embedded in a positive sphere

Rutherford's Gold Foil Experiment and the Atomic Nucleus

Ernest Rutherford:
- Credited with discovering the atomic nucleus through the Gold Foil Experiment.
The Gold Foil Experiment:
- A beam of alpha particles (positively charged particles emitted from a radioactive source inside a lead shield) was directed at a very thin sheet of gold foil surrounded by a circular fluorescent detection screen.
- Observations: The vast majority of alpha particles passed straight through the gold foil without deflection. A small fraction of alpha particles were deflected at very large angles, and some bounced directly backward.

Rutherford's gold foil experiment showing alpha particle beam deflection

The Atomic Nucleus:
- The tiny, extremely dense, central core of an atom composed of protons and neutrons.
- Possesses a net positive charge.
- Contains more than $99\%$ of the atom's total mass.
- The atom consists mostly of empty space surrounding this central core.

Bohr's Quantum Model of the Atom

Niels Bohr:
- Applied quantum theory to Rutherford's nuclear structure.
- Assumed that electrons travel around the nucleus only in specific stationary circular orbits defined by their angular momentum.
- Calculated precise, quantized energy levels for these electronic orbits.
- Hypothesized that emission of light occurs when an electron transitions from a higher energy orbit to a lower energy orbit, releasing a photon corresponding to a specific line in an atomic emission spectrum.

Distinguishing Between Atoms: Numbers, Mass, and Isotopes

Subatomic Particle Summary:
- Protons ( $p^+$ ) and Neutrons ( $n^0$ ) reside in the nucleus, giving the nucleus a net positive charge.
- Electrons ( $e^-$ ) occupy the space surrounding the nucleus.
Atomic Number ( $Z$ ):
- The total number of protons in the nucleus of an atom.
- The number of protons defines the specific chemical identity of an element.
- In any neutral atom, the number of protons equals the number of electrons: $\text{Atomic Number} = \text{number of } p^+ = \text{number of } e^-$ .
- Examples:
- Aluminum ( $\text{Al}$ ): Atomic Number = $13$ , Protons = $13$ , Electrons = $13$
- Carbon ( $\text{C}$ ): Atomic Number = $6$ , Protons = $6$ , Electrons = $6$
- Bromine ( $\text{Br}$ ): Atomic Number = $35$ , Protons = $35$ , Electrons = $35$
- Iodine ( $\text{I}$ ): Atomic Number = $53$ , Protons = $53$ , Electrons = $53$
- Sodium ( $\text{Na}$ ): Atomic Number = $11$ , Protons = $11$ , Electrons = $11$
- Silver ( $\text{Ag}$ ): Atomic Number = $47$ , Protons = $47$ , Electrons = $47$
Mass Number ( $A$ ):
- The total combined count of protons and neutrons in the nucleus of an atom: $\text{Mass Number } (A) = \text{number of } p^+ + \text{number of } n^0$
- The number of neutrons is determined by subtracting the atomic number from the mass number: $\text{number of } n^0 = \text{Mass Number } (A) - \text{Atomic Number } (Z)$
- Neutron Calculation Examples:
- Aluminum ( $\text{Al}$ ): Mass Number = $27$ , Atomic Number = $13 \rightarrow 27 - 13 = 14\,n^0$
- Carbon ( $\text{C}$ ): Mass Number = $12$ , Atomic Number = $6 \rightarrow 12 - 6 = 6\,n^0$
- Bromine ( $\text{Br}$ ): Mass Number = $80$ , Atomic Number = $35 \rightarrow 80 - 35 = 45\,n^0$
- Iodine ( $\text{I}$ ): Mass Number = $127$ , Atomic Number = $53 \rightarrow 127 - 53 = 74\,n^0$
- Sodium ( $\text{Na}$ ): Mass Number = $23$ , Atomic Number = $11 \rightarrow 23 - 11 = 12\,n^0$
- Oxygen ( $\text{O}$ ): Mass Number = $16$ , Atomic Number = $8 \rightarrow 16 - 8 = 8\,n^0$
Isotopes:
- Atoms of the same element that possess identical numbers of protons but different numbers of neutrons (and thus different mass numbers).
- Isotope Notation Formats:
- $\text{element name} - \text{atomic mass}$
- Nuclear Symbol Notation: $\mathbf{^{A}_{Z}X}$ , where $A = \text{atomic mass}$ , $Z = \text{atomic number}$ , and $X = \text{element symbol}$ .
- Hydrogen Isotopes Example:
- Hydrogen-1: $\mathbf{^{1}_{1}H}$ ( $1\,p^+$ , $0\,n^0$ , $1\,e^-$ )
- Hydrogen-2: $\mathbf{^{2}_{1}H}$ ( $1\,p^+$ , $1\,n^0$ , $1\,e^-$ )
- Hydrogen-3: $\mathbf{^{3}_{1}H}$ ( $1\,p^+$ , $2\,n^0$ , $1\,e^-$ )
- Determining Neutrons in Specific Isotopes:
- $\mathbf{^{27}Al}$ : $27 - 13 = 14\,n^0$
- $\mathbf{^{12}C}$ : $12 - 6 = 6\,n^0$
- $\mathbf{^{80}Br}$ : $80 - 35 = 45\,n^0$
- $\mathbf{^{127}I}$ : $127 - 53 = 74\,n^0$
- $\mathbf{^{23}Na}$ : $23 - 11 = 12\,n^0$
- $\mathbf{^{108}Ag}$ : $108 - 47 = 61\,n^0$
Neutral Atom Data Table Solutions:
- Row 1: Element = Magnesium ( $\text{Mg}$ ), Atomic Number = $12$ , Mass Number = $25$ , Protons = $12$ , Electrons = $12$ , Neutrons = $13$
- Row 2: Element = Carbon ( $\text{C}$ ), Atomic Number = $6$ , Mass Number = $12$ , Protons = $6$ , Electrons = $6$ , Neutrons = $6$
- Row 3: Element = Potassium ( $\text{K}$ ), Atomic Number = $19$ , Mass Number = $40$ , Protons = $19$ , Electrons = $19$ , Neutrons = $21$
- Row 4: Element = Vanadium ( $\text{V}$ ), Atomic Number = $23$ , Mass Number = $51$ , Protons = $23$ , Electrons = $23$ , Neutrons = $28$

Atomic Mass Calculations

Atomic Mass:
- The weighted average mass of all naturally occurring isotopes in a sample of an element.
Atomic Mass Unit (amu):
- Defined precisely as one-twelfth ( $\frac{1}{12}$ ) the mass of a single carbon-12 atom.
Sample Calculation - Average Atomic Mass of Cesium:
- A given naturally occurring sample of cesium contains:
- $75\%$ Cesium-133 ( $\mathbf{^{133}Cs}$ )
- $20\%$ Cesium-132 ( $\mathbf{^{132}Cs}$ )
- $5\%$ Cesium-134 ( $\mathbf{^{134}Cs}$ )
- Step-by-step contribution calculation:
- Contribution from $\mathbf{^{133}Cs}$ : $0.75 \times 133 = 99.75\,\text{amu}$
- Contribution from $\mathbf{^{132}Cs}$ : $0.20 \times 132 = 26.40\,\text{amu}$
- Contribution from $\mathbf{^{134}Cs}$ : $0.05 \times 134 = 6.70\,\text{amu}$
- Total weighted average mass: $99.75 + 26.40 + 6.70 = 132.85\,\text{amu}$
- Accuracy Comparison: The calculated value ( $132.85\,\text{amu}$ ) closely matches the standard periodic table atomic mass for cesium ( $132.905\,\text{amu}$ ).

Development and Structure of the Periodic Table

Dmitri Mendeleev (~1872):
- First published a comprehensive periodic table.
- Arranged elements in order of increasing atomic mass.
- Placed elements with similar physical and chemical properties into the same vertical columns.
- Left blank spaces/lines within the table for elements he predicted existed but had not yet been discovered (e.g., spaces at atomic masses $44$ , $68$ , $72$ , and $100$ ).

Henry Moseley (1887–1915):
- British physicist who experimentally determined atomic numbers for the elements.
- Rearranged the periodic table in order of increasing atomic number rather than atomic mass, resolving inconsistencies in Mendeleev's table.
- Established the structural layout of the modern periodic table.
Periodic Law:
- When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
Periods and Groups:
- Periods: The horizontal rows on the periodic table (numbered $1$ through $7$ ).
- Groups (Families): The vertical columns of elements on the periodic table containing elements with similar chemical properties.

Block layout of the periodic table showing representative elements, d-transition, and f-transition elements

Group Designations:
- Representative Elements: Group A elements (Groups 1A through 7A).
- Alkali Metals: Group 1A (Group 1) elements.
- Alkaline Earth Metals: Group 2A (Group 2) elements.
- Halogens: Group 7A (Group 17) elements.
- Noble Gases: Group 0 / Group 8A (Group 18) elements.
- Transition Metals: Group B elements located in the center block ( $d$ -transition elements).
- Inner Transition Metals (Rare Earth Metals): Placed below the main body of the periodic table ( $f$ -transition elements).
Broad Classification of Elements:
- Metals: Found on the left side and center of the periodic table. Good conductors of heat and electricity, lustrous, malleable, and ductile.
- Nonmetals: Found on the upper right-hand side of the periodic table. Generally poor conductors of heat and electricity, non-malleable, and brittle in solid form.
- Metalloids (Semi-metals): Located along the diagonal stair-step line separating metals from nonmetals. Possess properties intermediate between those of metals and nonmetals.

Periodic table showing classification into metals, semimetals, and nonmetals

Key Vocabulary Terms

Atom: The smallest particle of an element that retains the chemical and physical properties of that element.
Electron: A negatively charged subatomic particle located outside the atomic nucleus, carrying a relative charge of $-1$ and a mass equal to approximately $rac{1}{1840}$ the mass of a single hydrogen atom.
Cathode Ray Tube (CRT): A sealed glass vacuum tube device used to originate and study a flow of charged particles.
Proton: A positively charged subatomic particle located in the nucleus of an atom carrying a relative charge of $+1$ .
Neutron: A subatomic particle located in the nucleus of an atom that possesses no electrical charge and has approximately the same mass as a proton.
Nucleus: The dense, central core of an atom composed of protons and neutrons, accounting for more than $99\%$ of an atom's mass.
Atomic Number ( $Z$ ): The total number of protons in the nucleus of an atom, which defines its specific chemical identity.
Mass Number ( $A$ ): The total combined count of protons and neutrons in the nucleus of an atom: $\text{Mass Number } (A) = \text{number of } p^+ + \text{number of } n^0$ .
Isotopes: Atoms of the same element that possess identical numbers of protons but different numbers of neutrons (and thus different mass numbers).
Atomic Mass: The weighted average mass of all naturally occurring isotopes in a sample of an element.
Atomic Mass Unit (amu): A unit of mass defined precisely as one-twelfth ( $\frac{1}{12}$ ) the mass of a single carbon-12 atom.
Periodic Law: The principle stating that when elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
Periods: The horizontal rows of elements on the periodic table.
Groups: The vertical columns of elements on the periodic table sharing similar chemical properties.
Representative Elements: Group A elements on the periodic table.
Transition Metals: Group B elements located in the middle section of the periodic table.
Metals: Elements found on the left side of the periodic table that are good conductors of heat and electricity.
Nonmetals: Elements found on the right side of the periodic table with physical and chemical properties distinct from metals.
Metalloids (Semimetals): Elements located along the stair-step line of the periodic table with properties intermediate between metals and nonmetals.