Periodicity and the Periodic Table

Definition and History of Periodicity

• Periodicity Defined: Periodicity refers to the recurring trends observed in element properties.
• Mendeleev's Contribution: These trends first became apparent to Mendeleev when he arranged the known elements in order of increasing atomic mass. Based on the properties displayed by these elements, Mendeleev was able to predict the existence and properties of elements that were yet to be discovered at that time.
• Moseley's Discovery: Later experiments by Moseley using X-rays led to the accurate measurement of the atomic numbers of the elements.
• Atomic Number vs. Atomic Mass: Moseley found that the properties of elements are a periodic function of their atomic numbers, rather than their atomic mass.
• Construction of the Modern Periodic Table: The modern form of the periodic table is based on atomic number. It is constructed by:     * Arranging elements in increasing order of their atomic numbers.     * Organizing the elements so that those with similar physicochemical properties fall within the same vertical column.

Basis of the Modern Periodic Table

• Electronic Configuration: The modern periodic table is based on the electronic configuration of electrons in the outermost shells and the highest energy sub-shells.
• Chemical Properties: The reason for this organization is that only these outermost electrons determine the chemical properties of elements and their compounds, which in turn reflects the number of protons in an atom.

Fundamental Periodic Properties

• Ionization Energy: This is the energy required to remove an electron from an ion or a gaseous atom.
• Atomic Radius: This is defined as half the distance ($radius = \frac{d}{2}$) between the centers of two atoms that are touching each other.
• Electronegativity: This is a measure of the ability of an atom to form a chemical bond by attracting electrons to itself.
• Electron Affinity: This refers to the ability of an atom to accept an electron.

Basic Organization and Classification

• Three Basic Categories: The known elements are divided into three major categories:     * Metals     * Nonmetals     * Metalloids
• Organizational Criteria: The table is organized by atomic structure, atomic number, and both chemical and physical properties.
• Predictive Utility: The periodic table is a primary tool used to predict:     * Chemical behavior of the elements.     * Trends in properties.     * Physical and chemical properties of the elements.

Vertical Columns: Groups and Families

• Groups Definition: Vertical columns are called groups or families.
• Numbering: Groups are numbered from $1$ to $18$.
• Commonalities: Elements in the same group have the same number of valence electrons in their outer energy level.
• Chemical Behavior: Because they share valence electron counts, grouped elements behave chemically in similar ways.

Specific Group Characteristics

• Group 1: Alkali Metals:     * Composition: Metals.     * Valence Electrons: $1$.     * Reactivity: Very reactive.     * Properties: Solids, soft texture, shiny appearance, low density, and they react violently with water.
• Group 2: Alkaline-Earth Metals:     * Composition: Metals.     * Valence Electrons: $2$.     * Reactivity: Very reactive, though less so than Group $1$ metals.     * Properties: Solids, silver-colored, and more dense than alkali metals.
• Groups 3-12: Transition Metals:     * Composition: Metals.     * Valence Electrons: Typically $1$ or $2$.     * Reactivity: Less reactive than Group $1$ and Group $2$ metals.     * Properties: Higher density than alkali/alkaline-earth metals and good conductors of both heat and electricity.     * Placement: These are the short, center groups of the table.
• Groups 3-12: Rare Earth Elements (Below Main Table):     * Reason for Placement: These two rows are pulled out of sequence and placed below the main table to keep the overall table from becoming too wide.     * Lanthanide Series: Elements with atomic numbers $58$ through $71$. They follow Lanthanum ($Z=57$) in Period $6$.         * Valence Electrons: $3$.         * Reactivity: Very reactive.         * Properties: High luster but tarnish easily; high electrical conductivity; very small differences between individual elements.     * Actinide Series: Elements with atomic numbers $90$ through $103$. They follow Actinium ($Z=89$) in Period $7$.         * Valence Electrons: $3$ (but can be up to $6$).         * Reactivity: Unstable and all are radioactive.         * Origin: Most are made in laboratories.

Main Group Families (Groups 13-18)

• Metalloids (General): Located along a zig-zag line separating metals from non-metals (found in Groups $13-17$). They possess characteristics of both metals and nonmetals. They have low electrical conductivities that increase with temperature (semi-metals or semiconductors). Examples include $B$, $Si$, $Ge$, $As$, and $Te$.
• Group 13: Boron Family:     * Composition: $1$ metalloid and $4$ metals.     * Valence Electrons: $3$.     * Reactivity: Reactive.     * Physical State: Solid at room temperature.
• Group 14: Carbon Family:     * Composition: $1$ non-metal, $2$ metalloids, and $3$ metals.     * Valence Electrons: $4$.     * Reactivity: Varies.     * Physical State: Solid at room temperature.
• Group 15: Nitrogen Family:     * Composition: $2$ non-metals, $2$ metalloids, and $1$ metal.     * Valence Electrons: $5$.     * Reactivity: Varies.     * Physical State: All are solid at room temperature except Nitrogen ($N$).
• Group 16: Oxygen Family (Chalcogens):     * Composition: $3$ non-metals, $1$ metalloid, and $2$ metals.     * Valence Electrons: $6$.     * Reactivity: Reactive.     * Physical State: All are solid at room temperature except Oxygen ($O$).
• Group 17: Halogens:     * Composition: Nonmetals.     * Valence Electrons: $7$.     * Reactivity: Very reactive.     * Properties: Poor conductors of electric current; react violently with alkali metals to form salts; never found uncombined in nature.
• Group 18: Noble Gases:     * Composition: Nonmetals.     * Valence Electrons: $8$ (Note: Helium ($He$) has $2$).     * Reactivity: Unreactive (the least reactive group).     * Properties: Colorless, odorless gases at room temperature; outermost energy level is full; all are found in the atmosphere.

Hydrogen: The Unique Element

• Placement: Hydrogen ($H$) stands apart from the rest of the table because its properties do not match any single group perfectly.
• Valence Electrons: $1$.
• Reactivity: Very reactive, but it loses its single electron easily.
• Properties: Its physical properties are more similar to non-metals than to metals.

Horizontal Rows: Periods

• Periods Definition: Horizontal rows are called periods, numbered from $1$ to $7$.
• Energy Levels: All elements in a specific period have the same number of electron energy levels.     * Period 1: $1$ energy level (region of space for electrons).     * Period 2: $2$ energy levels.     * Period 5: $5$ energy levels.
• Pattern of Electrons: Moving from left to right across a period, each subsequent element has exactly one more electron in its outer shell than the element before it.
• Consequence: This arrangement lead to a regular pattern of change in chemical behavior across the period.

Detailed Periodic Trends: Atomic Radius

• Definition: The distance between the radii of two atoms ($radius = \frac{d}{2}$).
• Trend Across a Period (Left to Right): Radius DECREASES.     * Reasoning: As you move right, electrostatic attraction between the nucleus and electrons increases. Electrons are pulled closer to the nucleus, and valence electrons are more tightly held.
• Trend Down a Group: Radius INCREASES.     * Reasoning: Orbitals become larger, electrons are located farther from the nucleus, and valence electrons are less tightly bound.

Detailed Periodic Trends: Ionization Energy

• Successive Ionization: It requires more energy to remove each successive electron from an atom.     * First ionization energy: Energy to remove the first electron.     * Second ionization energy: Energy to remove the second electron, and so on.
• Quantum Leap: Once all valence electrons are removed, the energy required to remove the next (inner-shell) electron increases dramatically (a "quantum leap").
• Ionization Energy Values ($kJ\,mol^{-1}$) for Period 3 Elements:     * Sodium ($Na$): $I_1 = 495$, $I_2 = 4562$     * Magnesium ($Mg$): $I_1 = 738$, $I_2 = 1451$, $I_3 = 7733$     * Aluminum ($Al$): $I_1 = 578$, $I_2 = 1817$, $I_3 = 2745$, $I_4 = 11577$     * Silicon ($Si$): $I_1 = 786$, $I_2 = 1577$, $I_3 = 3232$, $I_4 = 4356$, $I_5 = 16091$     * Phosphorus ($P$): $I_1 = 1012$, $I_2 = 1907$, $I_3 = 2914$, $I_4 = 4964$, $I_5 = 6274$, $I_6 = 21267$     * Sulfur ($S$): $I_1 = 1000$, $I_2 = 2252$, $I_3 = 3357$, $I_4 = 4556$, $I_5 = 7004$, $I_6 = 8496$, $I_7 = 27107$     * Chlorine ($Cl$): $I_1 = 1251$, $I_2 = 2298$, $I_3 = 3822$, $I_4 = 5159$, $I_5 = 6542$, $I_6 = 9362$, $I_7 = 11018$     * Argon ($Ar$): $I_1 = 1521$, $I_2 = 2666$, $I_3 = 3931$, $I_4 = 5771$, $I_5 = 7238$, $I_6 = 8781$, $I_7 = 11995$
• Trends summary:     * Down a Column: Ionization energy decreases (less energy is required).     * Across a Row (Left to Right): Ionization energy generally increases (it gets harder to remove an electron).

Detailed Periodic Trends: Electronegativity

• Definition: The tendency of an atom to attract electrons to itself during chemical combination.
• Trend Across a Period (Left to Right): Electronegativity increases (low on the left, high on the right towards Fluorine).
• Trend Down a Group: Electronegativity decreases.

Questions & Discussion

• Exercise (Page 10): Identify at least $10$ elements that are metals, nonmetals, and metalloids on the Periodic Table.
• Tutorial (Page 15):     1. Write the ground state electronic configuration of the first five elements in group $13$. Indicate the valence electrons in the outer energy level of each of the elements.     2. What are the common chemical properties commonly exhibited by group $1$ elements?     3. Why are alkaline earth metals classified as group $2$ metals?
• Exercise (Page 46):     1. Explain why the chemical behavior of elements changes across a period.     2. Write the electronic configuration of elements in period $3-7$ and indicate the region of space available for electrons in any typical element in each period.