Comprehensive Study Notes: Trends and Properties of the Periodic Table

The Genius of Mendeleev and the Periodic Table Structure

• Dmitri Mendeleev's Contribution

  • Mendeleev was the first scientist to recognize key trends in the properties of elements and subsequently produced the periodic table.

  • He was notably accurate in leaving empty spaces in his periodic table for elements that had not yet been discovered during his time.

• Foundational Organization

  • The Periodic Table organizes a total of $118$ elements: $92$ natural elements and $26$ synthetic (man-made) elements.

  • Elements are arranged in order of increasing atomic number.

  • Structure consists of $18$ vertical groups and $7$ horizontal periods.

  • Groups $1$, $2$, and $13$ through $18$ are specifically known as main group elements and are sometimes numbered using Roman numerals $I$ through $VIII$.

Classification of Elements: Metals, Metalloids, and Non-metals

• Metals

  • Physical Properties: Characterized by metallic luster (they are shiny), high density, and high thermal and electrical conductivity.

  • Mechanical Properties: Most metals are malleable and ductile, meaning they can be deformed or stretched without breaking.

  • Tendency: Predominantly found on the left side of the periodic table.

• Non-metals

  • Physical Properties: Possess little or no metallic luster; they are not shiny. They are poor thermal and electrical conductors.

  • Mechanical Properties: Brittle as solids; they are not malleable or ductile.

  • State of Matter: Many non-metals exist as gases at room temperature.

  • Tendency: Located to the right of the periodic table.

• Metalloids

  • Metalloids are separated by the "metalloid staircase."

  • They possess a mixture of both metallic and non-metallic properties.

  • The metalloid staircase runs from Boron ($B$) to Tellurium ($Te$).

  • Commonly recognized metalloids include: Boron ($B$), Silicon ($Si$), Germanium ($Ge$), Arsenic ($As$), Antimony ($Sb$), and Tellurium ($Te$).

Atomic Structure and the Valence Shell

• Electronic Organization

  • An element's properties and its position in the table are determined by the organization of its electrons.

  • Each Period: Represents an additional shell of electrons (e.g., period $3$ elements have $3$ occupied shells).

  • Each Group: Represents an additional electron in the outer shell ($1$ through $18$ numbering relates to specific valence counts).

• The Valence Shell

  • The outermost shell is called the valence shell.

  • The number of electrons in the valence shell determines the reactivity of the element.

  • Trends in Shells:

    • The number of electron shells stays the same across a period (from left to right).

    • The number of electron shells increases going down a group.

Trends in Atomic Radii

• Defining Atomic Radius

  • The radius is the distance from the nucleus to the outer shell electrons. It is influenced by two primary factors:

    1. The attractive force between the nucleus and the electrons (core charge).

    2. The total number of electron shells.

• Core Charge and Shielding

  • Shielding: Electrons in full inner shells shield the valence electrons from the positive charge of the nucleus.

  • Non-shielding: Electrons within the same valence shell do not shield each other.

  • Formula for Core Charge: $\text{Core Charge} = \text{\# of protons} - \text{\# of full shell electrons}$.

  • Moving to the right across a period: The number of protons increases while shielding electrons stay the same. This increases core charge, holding the valence shell more tightly and decreasing the atomic radius.

• Vertical Trend

  • Going down a group: The number of electron shells increases, which increases the distance between the nucleus and the valence electrons. This increases the atomic radius.

Ion Formation and the Octet Rule

• The Octet Rule

  • The first $20$ elements tend to gain or lose valence electrons to achieve a stable octet ($8$ electrons).

  • Stable octets are naturally found in Group $8$ (Group $18$), the noble gases (inert gases).

• Ions

  • An ion is an atom or group of atoms with an overall positive or negative electrostatic charge.

  • Charge is denoted with a superscript number and symbol (e.g., $Na^+$, $Fe^{2+}$, $Cl^-$, $O^{2-}$).

• Cations (Positive Ions)

  • Formed when an atom loses electrons.

  • Metals typically form cations to match the electron configuration of the nearest noble gas.

  • Example (Magnesium): $Mg$ [$1s^2 2s^2 2p^6 3s^2$] loses two electrons to match Neon ($Ne$) [$1s^2 2s^2 2p^6$].

• Anions (Negative Ions)

  • Formed when an atom gains electrons.

  • Non-metals typically form anions.

  • Example (Fluorine): $F$ [$1s^2 2s^2 2p^5$] has $7$ valence electrons and gains one to match the stable configuration of Neon ($Ne$) [$1s^2 2s^2 2p^6$].

Group Specific Properties

• Group 1: Alkali Metals

  • Have $1$ electron in the outer shell and low core charge.

  • Easily lose that electron to form ions with a $1+$ charge.

  • Highly reactive; only found naturally in compounds.

  • Reaction with Water: React violently to produce an alkali and $H_2(g)$. Reactivity increases down the group.

  • Hydrogen Exception: Hydrogen is a diatomic gas ($H_2$). It usually donates $1$ electron to other non-metals but is not an alkali metal.

• Group 2: Alkaline Earth Metals

  • Have $2$ electrons in the outer shell.

  • Lose $2$ electrons to form ions with a $2+$ charge (e.g., $Ca^{2+}$).

  • Less easy to lose electrons than Group $1$, so they are less rapidly reactive, but still produce an alkali and $H_2(g)$ when reacting with water.

• Group 14: Carbon Group

  • Includes a range of properties: Carbon (non-metal), Silicon and Germanium (metalloids), Tin and Lead (metals).

  • Carbon Allotropes:

    • Diamond: 3-D cubic lattice, transparent, hardest mineral, abrasive, electrical insulator, thermal conductor.

    • Graphite: 2-D layers, opaque, soft, dry lubricant, electrical conductor, thermal insulator.

    • Other Forms: Graphene, Nanotubes (zigzag, armchair, chiral).

  • Silicon Compounds: Make up $75\%$ of the Earth's crust; used as electrical semiconductors.

• Group 17: Halogens

  • Non-metals that form diatomic molecules (e.g., $F_2$, $Cl_2$, $Br_2$, $I_2$, $At_2$).

  • States at room temp: $F_2$ and $Cl_2$ are gases; $Br_2$ is a liquid; $I_2$ and $At_2$ are solids.

  • Have $7$ valence electrons; readily gain $1$ electron to form a $1-$ charge ion.

• Group 18: Noble Gases

  • Inert, colorless gases with a full outer shell.

  • Occur naturally in the atmosphere: $Ar ∼ 1\%$, $Ne ∼ 0.002\%$, $He ∼ 0.0005\%$.

• Transition Metals (Groups 3-12)

  • Generally hard with high melting points.

  • Malleable, ductile, and good conductors.

  • Form ions with a range of different positive charges.

Summary of Reactivity Trends

• Metal Reactivity

  • Increases down a group: Valence electrons are further from the nucleus and less strongly held.

  • Increases to the left across a period: Core charge decreases, so valence electrons are held less tightly.

• Non-Metal Reactivity

  • Increases up a group: Valence electrons closer to the nucleus allow incoming electrons to be bound more strongly.

  • Increases to the right across a period: Core charge increases, strengthening the bond to incoming electrons.

Vocabulary Check

• Group: A vertical column; also called a family. Elements have the same number of outer shell (valence) electrons.

• Period: A horizontal row. The period number represents the number of occupied electron shells.

• Metal: Atoms that tend to lose electrons to form cations. Lustrous, malleable, ductile, good conductors.

• Metalloid: Elements having properties of both metals and non-metals.

• Non-metal: Atoms that tend to gain or share electrons to form anions. Generally poor conductors and not lustrous.

• Atomic Radius: One-half the distance between the nuclei in a molecule of identical atoms.

• Metallic Character: The extent to which an element exhibits properties characteristic of metals.

Questions & Discussion

• Question: Which has the larger atomic radius out of carbon and silicon? Explain why.

  • Response: Silicon has the larger radius because it is in Period $3$ with three electron shells, whereas Carbon is in Period $2$ with only two electron shells.

• Question: Which has the larger atomic radius out of sodium and magnesium? Explain why.

  • Response: Sodium ($Na$) has the larger radius. Moving from sodium to magnesium ($Mg$) increases the core charge while the number of shells stays the same, pulling the electrons closer to the nucleus in magnesium.

• Question: Which has greater reactivity: magnesium or calcium? Explain why.

  • Response: Calcium ($Ca$) is more reactive because it is further down Group $2$. Its valence electrons are further from the nucleus and thus more easily lost.

• Question: Which has greater reactivity: sodium or magnesium? Explain why.

  • Response: Sodium ($Na$) is more reactive. It has a lower core charge than magnesium, making it easier to lose its single valence electron.

• Question: Which has greater reactivity: oxygen or fluorine? Explain why.

  • Response: Fluorine ($F$). As a non-metal, reactivity increases toward the right of a period because the higher core charge attracts incoming electrons more strongly.

• Question: Which has greater reactivity: chlorine or fluorine? Explain why.

  • Response: Fluorine ($F$). Non-metal reactivity increases up a group because the valence shell is closer to the nucleus, allowing for a stronger attraction of new electrons.