Unit 4: Element and Compound Properties – Periodic Trends, Bonding, and Nomenclature

Study Guide - Smart Notes

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Element and Compound Properties

Overview

This unit covers the fundamental properties of elements and compounds, focusing on periodic trends, types of chemical bonds, and the nomenclature of ionic, covalent, and acidic compounds. Mastery of these topics is essential for understanding chemical reactivity, molecular structure, and the classification of substances in general chemistry.

Periodic Trends

Atomic Radius

Definition: The atomic radius is the size of an atom, typically measured as half the distance between the nuclei of two identical atoms bonded together.
Trend Across a Period: Atomic radius decreases from left to right across a period due to increasing effective nuclear charge, which pulls electrons closer to the nucleus.
Trend Down a Group: Atomic radius increases down a group as additional electron shells are added, increasing the distance between the nucleus and valence electrons.
Example: Atomic radius: Li > Be > B > C > N > O > F (across period 2, radius decreases)

Shielding Effect

Definition: The reduction in effective nuclear charge on the electron cloud, due to a difference in the attraction forces among electrons in different shells.
As more energy levels are added (down a group), inner electrons shield outer electrons from the nucleus, increasing atomic size.

Ionic Radius

Cations: Positive ions are smaller than their parent atoms due to loss of electrons and reduced electron-electron repulsion.
Anions: Negative ions are larger than their parent atoms due to increased electron-electron repulsion.
Example: Na+ < Na, Cl- > Cl

Ionization Energy

Definition: The energy required to remove an electron from a gaseous atom.
Trend Across a Period: Increases from left to right due to increasing nuclear charge.
Trend Down a Group: Decreases due to increased atomic size and shielding effect.
Equation:

Electronegativity

Definition: The tendency of an atom to attract electrons in a chemical bond.
Trend: Increases across a period, decreases down a group. Fluorine is the most electronegative element.

Reactivity

Metals: Reactivity increases down a group (easier to lose electrons).
Nonmetals: Reactivity increases up a group (easier to gain electrons).

Melting/Boiling Points and Density

Melting and boiling points are influenced by the strength of intermolecular forces.
Density generally increases down a group.

Chemical Bonding

Types of Bonds

Type of Bond	Between	What Happens
Ionic	Metal and Nonmetal	Transfer of electrons
Covalent	Nonmetal and Nonmetal	Sharing of electrons
Metallic	Metal and Metal	Delocalized electrons ('sea of electrons')

Ionic Bonds

Formed by the electrostatic attraction between oppositely charged ions (cations and anions).
High melting and boiling points, conduct electricity when molten or dissolved in water.
Example:

Covalent Bonds

Formed by sharing electrons between nonmetals.
Can be single, double, or triple bonds depending on the number of shared electron pairs.
Lower melting and boiling points compared to ionic compounds.
Example: ,

Metallic Bonds

Involve a 'sea' of delocalized electrons moving freely among metal cations.
Account for properties such as electrical conductivity and malleability in metals.

Bond Polarity and Electronegativity

The difference in electronegativity between atoms determines bond type:

Bond Type	ΔEN (Difference in Electronegativity)
Ionic	High (>1.7)
Polar Covalent	Medium (0.4–1.7)
Nonpolar Covalent	Low (<0.4)

Nomenclature of Compounds

Naming Ionic Compounds

Name the cation first, then the anion (with -ide ending for simple anions).
For transition metals, indicate the charge with Roman numerals.
Polyatomic ions retain their names (e.g., sulfate, nitrate).
Example: is iron(III) chloride.

Naming Covalent Compounds

Use prefixes to indicate the number of each atom (mono-, di-, tri-, etc.).
The first element keeps its name; the second gets an -ide ending.
Example: is carbon dioxide.

Naming Acids

Binary acids: Use "hydro-" prefix and "-ic acid" suffix (e.g., HCl is hydrochloric acid).
Oxyacids: If the polyatomic ion ends in -ate, use "-ic acid"; if -ite, use "-ous acid" (e.g., H2SO4 is sulfuric acid, H2SO3 is sulfurous acid).

Lewis Dot Structures

Drawing Lewis Structures

Count total valence electrons for all atoms in the molecule.
Arrange atoms with the least electronegative atom in the center (except hydrogen).
Distribute electrons to satisfy the octet rule (or duet for hydrogen).
Use double or triple bonds if necessary to complete octets.
Example: has two double bonds between C and O.

Exceptions to the Octet Rule

Hydrogen (2 electrons), Boron (6 electrons), and expanded octets for elements in period 3 or higher.

Polyatomic Ions and Resonance

Polyatomic ions are groups of covalently bonded atoms with an overall charge.
Resonance structures occur when more than one valid Lewis structure can be drawn for a molecule or ion.

Polyatomic Ions to Know

Formula	Name	Formula	Name
CO32-	carbonate	NO3-	nitrate
SO42-	sulfate	NH4+	ammonium
PO43-	phosphate	OH-	hydroxide
ClO4-	perchlorate	CN-	cyanide
HCO3-	bicarbonate	NO2-	nitrite

Summary Table: Periodic Trends

Trend	Across a Period	Down a Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electronegativity	Increases	Decreases
Metallic Character	Decreases	Increases

Practice and Application

Practice naming and writing formulas for ionic, covalent, and acidic compounds.
Draw Lewis structures for molecules and polyatomic ions, including resonance forms and exceptions to the octet rule.
Apply periodic trends to predict and explain chemical properties and reactivity.

Additional info: This guide integrates textbook alignment, learning objectives, and practice problems to reinforce mastery of periodic trends, bonding, and nomenclature, as outlined in a typical General Chemistry curriculum.