General Chemistry: Atomic Mass, Chemical Bonding, and Organic Nomenclature Study Notes

Atomic Mass and Mass Spectrometry

Weighted Average Atomic Mass

The weighted average atomic mass of an element is calculated based on the relative abundance and mass of each isotope. This value is necessary for determining the average mass of atoms in a naturally occurring sample.

• Formula: $\text{Weighted Average} = (\text{fraction}_1 \times \text{mass}_1) + (\text{fraction}_2 \times \text{mass}_2) + \ldots$

• Example: $0.909 \times 20.0 + 0.003 \times 21.0 + 0.088 \times 22.0 = 20.18$ amu

Mass Spectrometry

Mass spectrometry is a technique used to measure the masses of atoms and the percentage abundance of isotopes. Atoms are ionized, accelerated, and deflected by a magnetic field, allowing for the separation and detection of isotopes based on their mass-to-charge ratio.

• Positively charged ions are separated and detected.

• Light particles are deflected more than heavier particles.

Mass Spectrum

A mass spectrum is a graphical representation of the results from a mass spectrometer, showing the relative abundance of detected ions as a function of their mass-to-charge ratio.

• Can be used to determine isotopic abundance.

Counting Atoms by Weighing Them

Mole and Avogadro's Number

The mole is a standard unit for counting atoms, defined as the amount of substance containing as many entities (atoms, molecules, ions) as there are atoms in exactly 12 grams of carbon-12.

• Avogadro's Number: $6.022 \times 10^{23}$ entities per mole

• Molar Mass: The mass (in grams) of one mole of a substance, numerically equal to the atomic or molecular mass in amu.

• Example: 1 mole of C-12 has a mass of 12.0 g and contains $6.022 \times 10^{23}$ atoms.

Entities and Particles

• Entities: Atoms, molecules, ions, or formula units depending on the substance.

• Example: 1 mole of NaCl contains $6.022 \times 10^{23}$ formula units of NaCl.

Molecules, Compounds, and Chemical Equations

Chemical Bonds

Chemical bonds are the attractive forces that hold atoms together in compounds. The main types are ionic, covalent, and metallic bonds.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a full outer shell of eight electrons.

• Exceptions: Hydrogen (2 electrons), Helium (2), Beryllium (4), Boron (6), and transition metals (variable).

Molecular Elements

• Some nonmetals exist naturally as molecules (e.g., $\text{H}_2$, $\text{O}_2$, $\text{F}_2$, $\text{Cl}_2$, $\text{Br}_2$, $\text{I}_2$, $\text{P}_4$, $\text{S}_8$).

Ionic Compounds

• Formed by the transfer of electrons from metals to nonmetals, resulting in positive and negative ions held together by electrostatic forces.

• Formula Unit: The smallest electrically neutral collection of ions in an ionic compound.

• Properties: High melting points, conduct electricity when molten or dissolved (electrolytes).

• Example: NaCl is an ionic compound; solid NaCl does not conduct electricity, but molten or dissolved NaCl does.

Lewis Dot Diagrams

• Visual representations of valence electrons as dots around the chemical symbol.

• Used to predict bonding in molecules and ions.

Drawing Lewis Structures

Steps for Drawing Lewis Structures (Example: Methanol, CH3OH)

1. Arrange the atoms with the least electronegative atom in the center (except hydrogen).

2. Count all valence electrons for all atoms in the molecule.

3. Place 2 electrons between atoms to form bonds.

4. Distribute remaining electrons to complete octets (or duets for hydrogen).

5. Check if all atoms have a full octet; if not, form double or triple bonds as needed.

• Polyatomic Ions: Add or subtract electrons based on the ion's charge.

Covalent Bonds

• Polar Covalent Bonds: Electrons are shared unequally between atoms (e.g., H2O).

• Nonpolar Covalent Bonds: Electrons are shared equally (e.g., O2).

Formulas and Naming for Ionic & Molecular Compounds

Naming Binary Ionic Compounds

• Name the metal first, then the nonmetal with its ending changed to "-ide" (e.g., NaCl = sodium chloride).

• For transition metals, indicate the charge with Roman numerals (e.g., FeCl2 = iron(II) chloride).

Formulas for Ionic Compounds

• Write the chemical symbols, cross the charges, reduce to simplest ratio (e.g., Na+ and O2− form Na2O).

Classical Naming for Metals

• Use "-ous" for lower charge, "-ic" for higher charge (e.g., Fe2+ = ferrous, Fe3+ = ferric).

Polyatomic Ions

• Common polyatomic ions include: $\text{NO}_3^-$ (nitrate), $\text{SO}_4^{2-}$ (sulfate), $\text{CO}_3^{2-}$ (carbonate), $\text{PO}_4^{3-}$ (phosphate).

• Names ending in "-ate" have more oxygen than those ending in "-ite" (e.g., sulfate vs. sulfite).

Bases and Acids

• Bases: Compounds that produce OH− in water (e.g., NaOH).

• Acids: Compounds that produce H+ in water. Binary acids (e.g., HCl) are named with the prefix "hydro-" and the suffix "-ic acid". Oxyacids (e.g., H2SO4) are named based on the polyatomic ion.

Hydrates

• Compounds with water molecules attached in a fixed ratio (e.g., CuSO4·5H2O).

• Named with a Greek prefix indicating the number of water molecules (e.g., pentahydrate).

Organic Chemistry: Hydrocarbons and Functional Groups

Structure of Hydrocarbons

Hydrocarbons are compounds composed only of carbon and hydrogen. They can be classified as alkanes, alkenes, alkynes, and aromatic hydrocarbons.

• Alkanes: Saturated hydrocarbons with only single bonds. General formula: $C_nH_{2n+2}$.

• Alkenes: Unsaturated hydrocarbons with at least one double bond. General formula: $C_nH_{2n}$.

• Alkynes: Unsaturated hydrocarbons with at least one triple bond. General formula: $C_nH_{2n-2}$.

• Aromatic Hydrocarbons: Contain benzene rings (C6H6 structure).

Naming Alkanes

• Use the prefix for the number of carbons (meth-, eth-, prop-, but-, etc.) and the suffix "-ane".

• Number the longest carbon chain to give the lowest possible numbers to substituents.

• Substituents are named as alkyl groups (e.g., methyl, ethyl) and their positions are indicated by numbers.

Naming Alkenes and Alkynes

• Indicate the position of the double or triple bond with the lowest possible number.

• Use the suffix "-ene" for alkenes and "-yne" for alkynes.

Cyclic Hydrocarbons

• Formed when carbon atoms join in a ring structure. Prefix "cyclo-" is used (e.g., cyclohexane).

• Number the ring to give substituents the lowest possible numbers.

Aromatic Hydrocarbons

• Benzene rings can be attached to parent chains and are referred to as phenyl groups.

• Substituents on benzene rings are named using ortho- (1,2-), meta- (1,3-), and para- (1,4-) positions.

Functional Groups

• Alcohols: Contain the -OH group. Named with the suffix "-ol" (e.g., ethanol).

• Halides: Contain halogen atoms (F, Cl, Br, I) attached to the carbon chain.

• Aldehydes: Contain the -CHO group. Named with the suffix "-al" (e.g., ethanal).

• Ketones: Contain the C=O group within the carbon chain. Named with the suffix "-one" (e.g., propanone).

• Carboxylic Acids: Contain the -COOH group. Named with the suffix "-oic acid" (e.g., ethanoic acid).

• Esters: Contain the -COOR group. Named with the suffix "-oate" (e.g., methyl ethanoate).

• Amines: Contain the -NH2 group. Named with the suffix "-amine" (e.g., ethylamine).

Table: Common Polyatomic Ions

Ion Name	Formula	Charge
Nitrate	NO3-	-1
Sulfate	SO42-	-2
Carbonate	CO32-	-2
Phosphate	PO43-	-3
Ammonium	NH4+	+1
Hydroxide	OH-	-1

Additional info:

• Some context and explanations have been expanded for clarity and completeness.

• Mnemonic devices and naming conventions are included for easier memorization.