Atoms and the Periodic Table Lecture Notes

Introduction to Elements and the Periodic Table

• Definition of an Element: An element is a pure substance that cannot be broken down into simpler substances through any chemical reaction.

• Element Identification: Each element is uniquely identified by a one- or two-letter symbol.

• Arrangement: Elements are organized strategically in the periodic table. The position of an element within this table provides significant information regarding its chemical properties.

• Common Elements and Their Symbols (Table 2.1):

  • Bromine: $Br$

  • Calcium: $Ca$

  • Carbon: $C$

  • Chlorine: $Cl$

  • Chromium: $Cr$

  • Cobalt: $Co$

  • Copper: $Cu$

  • Fluorine: $F$

  • Hydrogen: $H$

  • Iodine: $I$

  • Lead: $Pb$

  • Magnesium: $Mg$

  • Manganese: $Mn$

  • Molybdenum: $Mo$

  • Nitrogen: $N$

  • Oxygen: $O$

  • Phosphorus: $P$

  • Potassium: $K$

  • Sodium: $Na$

  • Sulfur: $S$

  • Zinc: $Zn$

Categories of Elements: Metals, Nonmetals, and Metalloids

• Metals:

  • Location: Found on the left side of the periodic table.

  • Appearance: Usually exist as shiny solids.

  • Conductivity: Good conductors of both heat and electricity.

  • State at Room Temperature: Solids, with the notable exception of mercury ($Hg$), which is a liquid at room temperature.

• Nonmetals:

  • Location: Found on the right side of the periodic table.

  • Appearance: Generally lack a shiny appearance.

  • Conductivity: Usually poor conductors of heat and electricity.

  • State at Room Temperature: Can exist in various states:

    • Solid: Examples include sulfur ($S$) and carbon ($C$).

    • Liquid: Example is bromine ($Br$).

    • Gas: Examples include nitrogen ($N_2$) and oxygen ($O_2$).

• Metalloids:

  • Location: Located along the solid staircase line starting at boron ($B$) and angling down towards astatine ($At$).

  • Properties: Possess properties intermediate between those of metals and nonmetals.

  • The Seven Metalloids: Boron ($B$), Silicon ($Si$), Germanium ($Ge$), Arsenic ($As$), Antimony ($Sb$), Tellurium ($Te$), and Astatine ($At$).

Focus on the Human Body: The Elements of Life

• Building-Block Elements: Four specific nonmetals—Oxygen ($O$), Carbon ($C$), Hydrogen ($H$), and Nitrogen ($N$)—make up $96\%$ of the mass of the human body.

Compounds and Chemical Formulas

• Definition of a Compound: A pure substance formed by the chemical combination of two or more elements.

• Chemical Formula Components:

  • Element Symbols: Indicate the identity of the elements present in the compound.

  • Subscripts: Indicate the specific ratio of atoms within the compound.

• Examples of Chemical Formulas:

  • Water ($H_2O$): Contains $2$ Hydrogen atoms and $1$ Oxygen atom.

  • Propane ($C_3H_8$): Contains $3$ Carbon atoms and $8$ Hydrogen atoms.

• Visual Representation: Compounds can be depicted via colored spheres where different colors represent different elements.

Structure of the Atom

• Atoms: All matter is composed of basic building blocks called atoms.

• Subatomic Particles: Atoms contain three primary particles:

  • Nucleus:

    • The dense core of the atom containing protons and neutrons.

    • Residency of most of the atom's mass.

  • Electron Cloud:

    • The area surrounding the nucleus where electrons are located.

    • Comprises most of the atom’s volume.

Quantifying the Atom: Atomic and Mass Numbers

• Atomic Number ($Z$):

  • Represents the number of protons in the nucleus.

  • Defines the identity of the element; every atom of a specific element has the same number of protons.

  • In a neutral atom, there is no net charge, meaning: $Z = \text{number of protons} = \text{number of electrons}$.

• Mass Number ($A$):

  • The total number of nucleons in the nucleus.

  • Formula: $A = \text{number of protons (Z)} + \text{number of neutrons}$.

• Calculations for Common Elements:

  • Hydrogen: Atomic Number = $1$, Mass Number = $1$, Protons = $1$, Electrons = $1$, Neutrons = $0$.

  • Carbon: Atomic Number = $6$, Mass Number = $12$, Protons = $6$, Electrons = $6$, Neutrons = $6$.

  • Nitrogen: Atomic Number = $7$, Mass Number = $14$, Protons = $7$, Electrons = $7$, Neutrons = $7$.

  • Oxygen: Atomic Number = $8$, Mass Number = $16$, Protons = $8$, Electrons = $8$, Neutrons = $8$.

Isotopes and Atomic Weight

• Isotopes: Atoms of the same element (same atomic number $Z$) that have a different number of neutrons, resulting in different mass numbers ($A$).

• Isotope Example (Chlorine):

  • ${}^{35}_{17}Cl$: Protons = $17$, Electrons = $17$, Neutrons = $35 - 17 = 18$.

  • ${}^{37}_{17}Cl$: Protons = $17$, Electrons = $17$, Neutrons = $37 - 17 = 20$.

• Atomic Weight: The weighted average of the masses of all naturally occurring isotopes of a specific element, reported in atomic mass units ($amu$).

Basic Features of the Periodic Table

• Rows and Columns:

  • Period: A horizontal row in the table.

  • Group: A vertical column in the table.

• Major Classifications:

  • Main Group Elements: Consist of the two columns on the far left and the six columns on the far right (Groups $1A$–$8A$).

  • Transition Metal Elements: Contained in the $10$ short columns located in the middle (Groups $1B$–$8B$).

  • Inner Transition Elements: Consist of the Lanthanides and Actinides (no group numbers assigned).

• Characteristics of Specific Groups:

  • Group 1A (Alkali Metals) & Group 2A (Alkaline Earth Elements): Soft and shiny metals, low melting points, good conductors of heat/electricity, react with water to form basic solutions.

  • Group 7A (Halogens): Exist as diatomic molecules (two atoms joined together), highly reactive, combine with many other elements to form compounds.

  • Group 8A (Noble Gases): Extremely stable, rarely combine with other elements.

The Unique Nature of Carbon

• Carbon exists in three primary elemental forms (allotropes):

  • Diamond: A three-dimensional network of Carbon atoms.

  • Graphite: Parallel sheets of Carbon atoms.

  • Buckminsterfullerene: A sphere (soccer-ball shape) containing $60$ Carbon atoms.

Electronic Structure: Shells, Subshells, and Orbitals

• Energy Levels (Shells):

  • Electrons occupy specific principal energy levels or shells identified by the quantum number $n$.

  • Shells are numbered $n = 1, 2, 3, 4, ...$.

  • Energy and Distance: Lower numbered shells are closer to the nucleus and lower in energy. Higher numbered shells are further from the nucleus and higher in energy.

• Electron Capacity of Shells:

  • Shell $n = 1$: Max $2$ electrons.

  • Shell $n = 2$: Max $8$ electrons.

  • Shell $n = 3$: Max $18$ electrons.

  • Shell $n = 4$: Max $32$ electrons.

• Subshells and Orbitals:

  • Shells are divided into subshells labelled $s, p, d, f$.

  • An orbital is a region of space where there is a high probability of finding an electron.

  • Capacity: Each single orbital can hold a maximum of $2$ electrons.

  • Orbital Counts per Subshell:

    • Subshell $s$: $1$ orbital ($2$ maximum electrons).

    • Subshell $p$: $3$ orbitals ($6$ maximum electrons).

    • Subshell $d$: $5$ orbitals ($10$ maximum electrons).

    • Subshell $f$: $7$ orbitals ($14$ maximum electrons).

• Orbital Shapes:

  • s orbital: Spherical shape.

  • p orbital: Dumbbell shape.

Electron Configuration

• Definition: Describes the arrangement of electrons within an atom's orbitals.

• Ground State: The lowest energy arrangement of electrons.

• Rules for Determination:

  • Rule 1: Electrons fill the lowest energy orbitals first (1s < 2s < 2p < 3s < 3p).

  • Rule 2: Each orbital holds a maximum of $2$ electrons.

• Configurations for Row 1 and Row 2 Elements:

  • $H (Z=1)$: $1s^1$

  • $He (Z=2)$: $1s^2$

  • $Li (Z=3)$: $1s^2 2s^1$

  • $C (Z=6)$: $1s^2 2s^2 2p^2$

  • $Ne (Z=10)$: $1s^2 2s^2 2p^6$

Valence Electrons and Periodic Trends

• Valence Shell: The outermost shell of an atom (the highest value of $n$).

• Valence Electrons: Electrons located in the valence shell; they determine the chemical properties of an element.

• Relationship to Group Number:

  • For main group elements (Groups $1A$–$8A$), the group number equals the number of valence electrons.

  • Exception: Helium ($He$) belongs to group $8A$ but only has $2$ valence electrons.

• Electron-Dot Symbols: Represent valence electrons as dots around the element symbol.

  • $1-4$ valence electrons: Represented as single dots.

  • > 4 valence electrons: Dots are paired.

• Periodic Trend: Atomic Size:

  • Increases down a column: Valence electrons are farther from the nucleus.

  • Decreases across a row: Increasing number of protons in the nucleus pulls electrons closer, making the atom smaller.

• Periodic Trend: Ionization Energy:

  • Definition: The energy required to remove an electron from a neutral atom ($Na + \text{energy} \rightarrow Na^+ + e^-$).

  • Decreases down a column: Valence electrons are farther from the positive nucleus and easier to remove.

  • Increases across a row: Increased nuclear charge (protons) holds electrons more tightly.