Unit 4 Chem
Study Guide: Naming and Formula Writing in Chemistry
1. Outline of Naming Compounds
A. Ionic Compounds
Definition: Formed from the electrostatic attraction between ions. Typically consist of metals and nonmetals.
Naming Guidelines:
Name the metal (cation) first. If it's a transition metal, include the oxidation state in Roman numerals in parentheses.
Name the non-metal (anion) second, changing its suffix to -ide.
Example:
NaCl = sodium chloride
Fe2O3 = iron(III) oxide
B. Covalent Compounds
Definition: Made from the sharing of electrons between nonmetals.
Naming Guidelines:
Use prefixes to indicate the number of atoms. 1=mono, 2=di, 3=tri, 4=tetra, 5=penta, 6=hexa, 7=hepta, 8=octa, 9=nona, 10=deca.
The first non-metal retains its name, and the second non-metal takes the -ide suffix.
Example:
CO2 = carbon dioxide
N2O5 = dinitrogen pentoxide
C. Acids
Definition: Compounds that release H+ ions when dissolved in water.
Binary Acids (H + non-metal):
Use the prefix "hydro-" and the suffix "-ic".
Example:
HCl = hydrochloric acid
Oxyacids (H + polyatomic ion):
If the polyatomic ion ends in -ate, change the suffix to -ic. If it ends in -ite, change the suffix to -ous.
Example:
H2SO4 (sulfate) = sulfuric acid
H2SO3 (sulfite) = sulfurous acid
2. Writing Chemical Formulas
A. Ionic Compounds
Steps to Write Formulas:
Identify the cation and anion.
Determine the charges of the ions. (Cation usually has a positive charge, the anion a negative charge.)
Use the criss-cross method to balance the total charges. Write the cation first followed by the anion.
Example:
Calcium (Ca^2+) + Chloride (Cl^-)
Ca has a +2 charge, Cl has a -1 charge.
Need 2 Cl to balance Ca:
Formula = CaCl2.
B. Covalent Compounds
Steps to Write Formulas:
Use the prefixes from the name to determine the number of each element present.
Write the elements' symbols and the respective number of atoms as subscripts.
Example:
Dinitrogen tetroxide = N2O4 (2 nitrogen atoms and 4 oxygen atoms)
C. Acids
Steps to Write Formulas:
For binary acids: Write the hydrogen symbol (H) followed by the non-metal symbol and the appropriate subscript (if needed).
Example:
Hydrochloric acid = HCl (1 hydrogen, 1 chlorine).
For oxyacids, write H followed by the polyatomic ion's formula.
Example:
Nitric acid (from nitrate) = HNO3.
3. Summary of Key Points
The naming of ionic compounds involves the order of cation and anion, with proper suffixes.
Covalent compounds use prefixes to indicate atom count.
Acids have unique naming rules based on their composition.
The process of writing chemical formulas requires an understanding of the charges and names of the elements involved.
Overall, mastering these naming conventions and formula writing skills is essential for success in chemistry.
Study Guide: Bonding Principles in Chemistry
Types of Chemical BondsA. Ionic Bonds
Definition: Formed through the transfer of electrons from one atom to another, resulting in the attraction between positively and negatively charged ions.
Characteristics:
Typically occurs between metals and nonmetals.
High melting and boiling points.
Conduct electricity when dissolved in water or melted.
Examples: NaCl (sodium chloride), MgO (magnesium oxide).
B. Covalent Bonds
Definition: Formed when two or more nonmetals share electrons to achieve a full outer shell.
Characteristics:
Can be polar (unequal sharing of electrons, e.g., H2O) or nonpolar (equal sharing of electrons, e.g., O2).
Generally lower melting and boiling points compared to ionic compounds.
Do not conduct electricity in any state.
Examples: CO2 (carbon dioxide), CH4 (methane).
C. Metallic Bonds
Definition: Characterized by a sea of electrons that are shared among a lattice of metal cations.
Characteristics:
Conduct electricity and heat efficiently.
Malleable and ductile due to the mobile nature of the electrons.
High melting and boiling points.
Examples: Cu (copper), Fe (iron).
Bonding TheoriesA. Valence Shell Electron Pair Repulsion (VSEPR) Theory
Purpose: Predicts the geometry of molecules based on the repulsion between electron pairs in the valence shell of the central atom.
Key Shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.
B. Molecular Orbital Theory
Concept: Describes the molecular bonding in terms of orbitals that can be occupied by electrons.
Key Ideas:
Molecular orbitals form when atomic orbitals combine.
Bonding orbitals (lower energy) stabilize the molecule, antibonding orbitals (higher energy) destabilize it.
Electronegativity and Bond PolarityA. Electronegativity: The ability of an atom to attract electrons in a bond.
Fluorine is the most electronegative element, followed by oxygen and nitrogen.
Electronegativity differences determine bond type:
Nonpolar covalent: 0-0.4
Polar covalent: 0.5-1.7
Ionic: >1.7
HybridizationA. Definition: The mixing of atomic orbitals to form new hybrid orbitals that are degenerate (of equal energy).
Types of Hybridization:
sp (linear), sp2 (trigonal planar), sp3 (tetrahedral), dsp3 (trigonal bipyramidal), d2sp3 (octahedral).
Intermolecular ForcesA. Types of Intermolecular Forces:
London Dispersion Forces: Weak forces arising from temporary dipoles in molecules.
Dipole-Dipole Interactions: Occur between polar molecules due to positive and negative dipoles.
Hydrogen Bonds: Special case of dipole-dipole forces where hydrogen is bonded to highly electronegative atoms (N, O, F).
Key Points to Remember
The type of bond influences the properties of substances.
Understanding the concept of electronegativity is crucial for predicting bond polarity.
Hybridization is integral to understanding molecular geometry and bonding behavior in molecules.
Intermolecular forces affect boiling and melting points significantly.
Summary
Mastering bonding principles is essential for a comprehensive understanding of chemistry. It allows for predicting how substances will interact, their properties, and their behavior under different conditions.
Study Guide: Polyatomic Ions
Definition:
Polyatomic ions are ions that consist of two or more atoms bonded together, which carry a net charge due to the loss or gain of one or more electrons.
Common Polyatomic Ions:
Nitrate (NO3^-)
Charge: -1
Commonly found in fertilizers and explosives.
Sulfate (SO4^2-)
Charge: -2
Used in various industrial processes, including the production of sulfuric acid.
Phosphate (PO4^3-)
Charge: -3
Essential for energy transfer in biological systems (ATP).
Carbonate (CO3^2-)
Charge: -2
Found in minerals; a key component of rocks and shells.
Acetate (C2H3O2^-)
Charge: -1
Used in chemical synthesis; found in vinegar.
Hydroxide (OH^-)
Charge: -1
Important in acid-base chemistry; a component of bases like sodium hydroxide.
Bicarbonate (HCO3^-)
Charge: -1
Also known as hydrogen carbonate; acts as a buffer in biological systems.
Chlorate (ClO3^-)
Charge: -1
Used in bleaching and disinfecting.
Cyanide (CN^-)
Charge: -1
Highly toxic; used in industry and mining.
Naming Conventions:
-ate and -ite Suffixes:
The suffix -ate denotes a polyatomic ion with more atoms (e.g., nitrate NO3^-) compared to the -ite form which has one less oxygen (e.g., nitrite NO2^-).
Per- and Hypo- Prefixes:
The prefix per- indicates the ion has one more oxygen than the -ate (e.g., perchlorate ClO4^-).
The prefix hypo- indicates the ion has one less oxygen than -ite (e.g., hypochlorite ClO^-).
Key Points to Remember:
Polyatomic ions can form ionic compounds with metals.
They are often encountered in acid-base reactions and redox processes.
Understanding polyatomic ions is crucial for predicting chemical behavior and reactions in both inorganic and organic chemistry contexts.