GOB Chemistry Study Guide: Chapters 6, 7, and 8 (Ions, Compounds, Molecular Naming, Electronegativity, Gases)

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Chapter 6: Ionic and Covalent Compounds

Definitions and Formation of Ionic and Covalent Compounds

Understanding the difference between ionic and covalent (molecular) compounds is fundamental in chemistry. Ionic compounds are formed by the transfer of electrons from metals to nonmetals, resulting in the formation of cations and anions. Covalent compounds are formed by the sharing of electrons between nonmetals.

Ionic Compounds: Composed of positive and negative ions held together by electrostatic forces.
Covalent (Molecular) Compounds: Composed of atoms sharing electrons to achieve stability.
Example: NaCl (ionic), H2O (covalent)

Identification and Naming of Ionic and Covalent Compounds

It is important to distinguish between ionic and covalent compounds and to know their naming conventions.

Ionic Compounds: Name the cation first, then the anion. Use Roman numerals for transition metals to indicate charge.
Covalent Compounds: Use prefixes to indicate the number of atoms.
Example: CO2 is carbon dioxide; FeCl3 is iron(III) chloride.

Monatomic Cations and Anions

Monatomic ions are ions formed from single atoms. The following table summarizes common monatomic cations and anions:

Group	Type 1 Metal (Cation)	Anion
Group 1	Na+
Group 2	Mg2+
Group 13	Al3+
Group 17		Cl-
Group 16		O2-
Group 15		N3-

Additional info: Type 2 metals (transition metals) can have multiple charges, e.g., Fe2+ and Fe3+.

Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge. The following table lists common polyatomic ions:

Ion	Name
NH4+	ammonium ion
CN-	cyanide ion
CO32-	carbonate ion
NO3-	nitrate ion
SO42-	sulfate ion
PO43-	phosphate ion
OH-	hydroxide ion

Atomic Elements and Noble Gas Configuration

Atoms tend to gain, lose, or share electrons to achieve a noble gas electron configuration, which is associated with stability.

Rule of Seven: Refers to the tendency of atoms to seek eight electrons in their valence shell (octet rule), except for hydrogen and helium.
Electron Dot Structures: Used to represent valence electrons and predict bonding.

Chapter 7: Molecular Compounds and Electronegativity

Nomenclature of Molecular Compounds

Molecular compounds are named using prefixes to indicate the number of each type of atom present. The following table lists common prefixes:

Number	Prefix
1	mono
2	di
3	tri
4	tetra
5	penta
6	hexa
7	hepta
8	octa
9	nona
10	deca

Example: CO is carbon monoxide; N2O4 is dinitrogen tetroxide.

Electronegativity and Periodic Trends

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. It increases across a period and decreases down a group in the periodic table.

Polar Covalent Bonds: Formed when atoms have different electronegativities, resulting in unequal sharing of electrons.
Nonpolar Covalent Bonds: Formed when atoms have similar electronegativities, resulting in equal sharing of electrons.
Example: H2O is polar; O2 is nonpolar.

Molecular Shapes and Bond Angles

The shape of a molecule is determined by the arrangement of atoms and electron pairs around the central atom. The following table summarizes common molecular shapes:

Shape	Bonded Atoms	Lone Pairs	Example
Linear	2	0	CO2
Trigonal Planar	3	0	BF3
Tetrahedral	4	0	CH4
Trigonal Pyramidal	3	1	NH3
Bent	2	2	H2O

Intermolecular Forces

Intermolecular forces are forces of attraction between molecules. The three main types are:

Dipole-Dipole Interactions: Occur between polar molecules.
Hydrogen Bonding: A strong type of dipole-dipole interaction involving H bonded to N, O, or F.
Dispersion Forces (London Forces): Present in all molecules, especially nonpolar ones.
Example: H2O exhibits hydrogen bonding; CH4 exhibits dispersion forces.

Chapter 8: Properties of Gases

Gas Laws and Properties

Gases have unique properties and are described by several laws. Key properties include pressure, volume, temperature, and amount (moles).

Pressure: The force exerted by gas particles on the walls of a container.
Volume: The space occupied by a gas.
Temperature: A measure of the average kinetic energy of gas particles.
Amount: Measured in moles.

Units of Pressure

Pressure can be measured in several units. The following table summarizes common units and their equivalence:

Atmosphere	Unit Equivalent to 1 atm
Atmosphere (atm)	1 atm
Millimeters of Hg (mmHg)	760 mmHg
Torr	760 torr
Pascal (Pa)	101,325 Pa

Gas Laws

Several laws describe the behavior of gases under different conditions:

Boyle's Law: (at constant temperature)
Charles's Law: (at constant pressure)
Gay-Lussac's Law: (at constant volume)
Combined Gas Law:
Ideal Gas Law:

Additional info: R is the gas constant, .

Stoichiometry and Gas Calculations

Stoichiometry involves calculations based on balanced chemical equations. For gases, calculations often use the ideal gas law and molar volume at STP (Standard Temperature and Pressure).

Molar Volume at STP: 1 mole of gas occupies 22.4 L at STP.
Example: Calculate the volume of 2 moles of O2 at STP:

Endothermic and Exothermic Reactions

Reactions can absorb (endothermic) or release (exothermic) energy. The difference is determined by the energy of reactants and products.

Endothermic: Energy is absorbed; products have higher energy than reactants.
Exothermic: Energy is released; products have lower energy than reactants.

Summary of Key Concepts

Know how to calculate moles, molar mass, and convert between grams and moles.
Be able to balance chemical equations and identify reaction types.
Understand the differences between intermolecular forces and their effects on physical properties.
Be able to use gas laws for calculations involving pressure, volume, temperature, and moles.