chem exam flashcards

ORGANIC CHEMISTRY BASICS

HYDROCARBONS

• Types:

  • Alkanes – Single bonds (saturated)

  • Alkenes – One or more double bonds (unsaturated)

  • Alkynes – One or more triple bonds (unsaturated)

• Structures:

  • Aliphatic – Chains

  • Cyclic – Rings

  • Aromatic – Benzene ring

STRUCTURAL SHORTHAND

• C atoms = line intersections.

• H atoms attached to C often omitted.

ROOT NAMES (Carbon Chain Length)

ALKANE NAMING RULES

1. Find longest chain = root + -ane

2. Number the chain to give lowest numbers to branches

3. Name branches (alkyl groups): -yl (e.g., CH₃ = methyl)

4. Use position numbers, list in alphabetical order

5. Multiple branches: di-, tri-, tetra- (don’t affect alphabetical order)

6. Punctuation: commas between numbers, hyphens between numbers/letters

Examples:

• 2-methylbutane

• 2-chloro-3,6-dimethylnonane

CYCLIC ALKANES

• Prefix: cyclo- + root + -ane
  
  

• Ex: cyclopentane (C₅H₁₀), cyclohexane (C₆H₁₂)

HALOALKANES (R–X)

• Halogens: F, Cl, Br, I → named as fluoro, chloro, bromo, iodo

• Same rules as branched alkanes

• Ex: 1-bromopentane

REPRESENTATIONS

• Line-Angle: Each point = carbon

• H’s on C usually not shown
  
  

 QUICK NAMING CHECKLIST:

• Longest C chain
  Number for lowest substituent positions

• Name and number each substituent

• Alphabetical order (ignore di-, tri-, etc.)

• Proper punctuation: #, #-prefixname-rootane

Naming Alkenes, Alkynes, Haloalkanes, Cyclic and Aromatic Hydrocarbons

NAMING ALKENES & ALKYNES

Key Steps:

1. Find longest carbon chain with double or triple bond → parent chain.

2. Number the chain to give the lowest number to the double/triple bond.

3. Name branches as usual (alphabetical, position #).

4. Indicate position of double/triple bond.

5. Suffix:

   • –ene → double bond

   • –yne → triple bond

6. If 2+ double/triple bonds → use diene, triyne, etc.

Examples:

• pent-2-ene (not 2-pentene)

• 4-methylhex-2-yne

• pent-1,4-diene

NAMING HALOALKANES / HALOALKENES / HALOALKYNES

Halogens (Group 17) → use prefixes:

• F = fluoro

• Cl = chloro

• Br = bromo

• I = iodo

Steps:

1. Find the longest chain (with double/triple bond if present).

2. Number to give lowest locants to the bond or halogen.

3. Halogens treated as substituents.

4. Same punctuation & alphabetical order as before.
   
   

Example:

• 1-bromopentane
  
  

• 3-chlorobut-1-ene
  
  

NAMING BRANCHED ALKENES / ALKYNES

Same rules apply:

1. Longest chain with double/triple bond.

2. Number to give lowest bond position.

3. Name and number branches.

4. Root = position + bond + chain length.

5. Branches in alphabetical order.

Example:

• 3-ethyl-4-methylhex-1-yne

NAMING CYCLIC HYDROCARBONS

1. Add prefix “cyclo” to the root name.

2. Double/triple bond → still give lowest possible number.

3. If ring is smaller than side chain, ring is the branch.

Examples:

• 1-methyl-1-cyclopentene

• 4-cyclopentyloctane

NAMING AROMATIC COMPOUNDS (BENZENE)

• Benzene as parent chain → name substituents:

  • Chlorobenzene, ethylbenzene, etc.

• Benzene as side chain → called phenyl:

  • 2-phenylethane
    
    

QUICK NAMING SUMMARY:

Isomers

1. Structural Isomers

• Same molecular formula, different atom arrangement.

• Different physical/chemical properties (e.g. boiling point).

• Example: Pentane (C₅H₁₂)

  • n-pentane → bp 36°C

  • isopentane (methylbutane) → bp 28°C

  • neopentane (dimethylpropane) → bp 9.5°C

2. Geometric Isomers (Cis/Trans)

• Found in alkenes (double bonds).

• Double bonds can’t rotate → fixed shape.

• Each carbon in the double bond must have 2 different groups attached.

✔ Types:

• Cis: larger groups on the same side

• Trans: larger groups on opposite sides

✔ Naming:

• Prefix the name with cis- or trans-

• Examples:

  • cis-2-pentene

  • trans-2-pentene

  • trans-5-methyl-3-heptene
    
    

PROPERTIES OF HYDROCARBONS

Key Concept: Bigger + flatter = stronger IMF = higher bp

ALCOHOLS

• Functional group: –OH (hydroxyl group)

• Classified based on carbon attached to –OH:

  • Primary (1°): OH on a carbon with 1 other C

  • Secondary (2°): OH on a carbon with 2 other Cs

  • Tertiary (3°): OH on a carbon with 3 other Cs

NAMING ALCOHOLS

✔ Steps:

1. Find longest chain with the –OH group.

2. Number the chain to give lowest number to –OH.

3. Replace -e of alkane with -ol.

4. Indicate position of –OH (e.g., butan-2-ol).

5. Multiple –OH groups → keep the -e and use di-, tri-, etc. (e.g., butane-1,2-diol).

6. Name and number other branches, then put it all together.

✔ Examples:

• 1-propanol → OH on 1st carbon → primary

• 2-butanol → OH on 2nd carbon → secondary

• Cyclobutanol → OH on ring → secondary (only 2 Cs attached to the C with OH)
  
  

PROPERTIES OF ALCOHOLS

QUICK RECOGNITION:

• Ends in -ol

• Look for –OH group on the carbon chain

• Identify 1°, 2°, or 3° based on what the OH-bearing carbon is connected to
  
  

PRACTICE:

FUNCTIONAL GROUPS OVERVIEW

INTERMOLECULAR FORCES & PROPERTIES

• Alcohols, Carboxylic Acids, and Amides: Strongest H-bonding → highest melting/boiling points

• Ethers and Ketones: Polar, but can’t H-bond with themselves

• Esters: Low BP, pleasant smells

• Amines: Fishy smell, weak base, H-bonding (if N–H present)

NAMING RULES SUMMARY

 Alcohols:

• Replace -e with -ol

• Number from end closest to –OH

• Ex: Butan-2-ol (secondary alcohol)

 Haloalkanes:

• Halogen as prefix: fluoro-, chloro-, etc.

• Number to give halogen lowest position

• Ex: 2-chlorobutane

 Amines:

• Main chain = parent + -amine

• Extra groups on N = use N-

• Ex: N-methylethanamine

 Aldehydes:

• Always at end of chain → position 1

• Replace -e with -al

• Ex: Butanal

 Ketones:

• Carbonyl (C=O) inside chain

• Replace -e with -one

• Need number if 4+ carbons

• Ex: Pentan-2-one

 Carboxylic Acids:

• Replace -e with -oic acid

• Always start numbering from –COOH

• Ex: Propanoic acid

 Esters:

• 2 parts: alcohol (R′) + acid (R)

• Name: alkyl (R′) + acid-derived name with -oate

• Ex: Ethyl ethanoate

Amides:

• Based on carboxylic acid → replace -oic acid with -amide

• Substituents on N: N- or N,N-

• Ex: N-methylpropanamide

ISOMERS

Structural Isomers:

• Same formula, different structure

• Ex: C₅H₁₂ → pentane, isopentane, neopentane
  
  

Geometric (cis/trans) Isomers:

• In alkenes only (C=C)

• Cis = groups on same side

• Trans = groups on opposite sides
  
  

SOLUBILITY & POLARITY SUMMARY

SAMPLE QUESTIONS TO PRACTICE

1. Name the following alcohol and identify its class:
   CH₃CH(OH)CH₃ → 2-propanol, secondary alcohol

2. Draw and label cis-2-butene and trans-2-butene

3. Which has a higher boiling point: Butanol or Butane?
   → Butanol, due to hydrogen bonding

4. Identify the functional group:
   
   

   • CH₃COOH → Carboxylic acid
     
     

   • CH₃CH₂OCH₃ → Ether

Types of Reactions by Functional Group

Detailed Reaction Examples and Notes

Combustion of Alkanes

• Example:
  2 C₄H₁₀ + 13 O₂ → 8 CO₂ + 10 H₂O

Substitution

• Alkanes:
  CH₄ + Cl₂ → CH₃Cl + HCl

• Benzene:
  C₆H₆ + Cl₂ → C₆H₅Cl + HCl (with catalyst)
  
  

 Addition Reactions (Alkenes & Alkynes)

 Markovnikov’s Rule

• The H atom adds to the carbon already holding more H atoms.

• Applies to hydration and hydrohalogenation of alkenes.

 Dehydration of Alcohols

• Propanol → Propene + H₂O

• Catalyst: Sulfuric acid (H₂SO₄)

• Example for ethanol:
  CH₃CH₂OH → CH₂=CH₂ + H₂O

Oxidation of Alcohols

Oxidizing Agents:

• Potassium dichromate (K₂Cr₂O₇), hydrogen peroxide (H₂O₂), KMnO₄
  
  

Hydrogenation of Aldehydes/Ketones

• Reverse of alcohol oxidation

• Requires catalyst, heat, pressure
  | Reactant | Product |
  |-------------|-------------|
  | Aldehyde | Primary alcohol |
  | Ketone | Secondary alcohol |
  
  

Formation of Carboxylic Acids

• Aldehyde + [O] → Carboxylic Acid

• Used in Breathalyzer test with orange Cr₂O₇²⁻ turning green

Esterification

• Carboxylic acid + Alcohol → Ester + Water

• Condensation reaction, acid catalyst used

Hydrolysis & Saponification of Esters

Synthesis of Amines

• Alkyl halide + NH₃ → 1° Amine + HX

• Repeats with more alkyl halide to make 2° and 3° amines

Synthesis of Amides

• Carboxylic acid + Amine → Amide + H₂O

Hydrolysis of Amides

• Amide + H₂O → Carboxylic acid + Amine

• Requires acidic or basic conditions

Polymers

What is a Polymer?

• Long chains of repeating units called monomers.

Types of Polymers
Natural: DNA, starch, silk
Synthetic: Polystyrene, Polyester, Polyethylene, PVC, Nylon, PMMA, Polypropylene

Common Polymers & Uses

• Polystyrene: Styrofoam, cups

• Polyester: Clothes, bottles

• Polyethylene: Bags, bottles

• PVC: Pipes, raincoats

• Nylon: Ropes, toothbrushes

• PMMA (Plexiglas): Glass substitutes

• Polypropylene: Containers, money

Polymerization Types

1. Addition Polymerization

• Monomers with double bonds (alkenes)

• No by-product

• Example: Ethene → Polyethylene

2. Condensation Polymerization

• Two different monomers with functional groups

• By-product: Water

• Forms esters (polyester) or amides (nylon)

Key Points

• Addition = monomers add together (alkenes)

• Condensation = monomers link and release water

• Ester linkage = polyester

• Amide linkage = nylon

Atomic Theory and Structure

• Elements consist of atoms, which cannot be created, destroyed, or divided into smaller parts by chemical means.
  
  

• Atoms of the same element have identical size, mass, and properties.
  
  

Electron Discovery and Properties

• Electrons are negatively charged subatomic particles found in all atoms.
  
  

• Electrons have a charge-to-mass ratio that can be measured.
  
  

• Atoms are electrically neutral, so positive charge must balance the negative electrons.
  
  

Atomic Models

• Atom consists of a dense, positively charged nucleus containing protons and neutrons.
  
  

• Electrons move around the nucleus at relatively large distances.
  
  

Subatomic Particles

• Protons: positively charged particles in the nucleus.
  
  

• Neutrons: neutral particles in the nucleus, contributing to atomic mass.
  
  

Radioactivity

• Radioactivity is the spontaneous decay of an atomic nucleus, emitting energy or particles.
  
  

• Alpha, beta, and gamma radiation are types of emitted radiation.
  
  

Light and Electromagnetic Waves

• Light is an electromagnetic wave with electric and magnetic fields oscillating perpendicular to each other.
  
  

• Light exhibits a spectrum of wavelengths (continuous spectrum).
  
  

Quantum Theory

• Energy is quantized and can be absorbed or emitted only in discrete amounts called quanta.
  
  

• Energy of a quantum is given by:
  E=nhfE = nhfE=nhf
  where nnn is an integer, hhh is Planck’s constant, and fff is frequency.
  
  

Photoelectric Effect

• Light consists of particles called photons, each with energy proportional to its frequency.
  
  

• Electrons are emitted from a metal surface when struck by photons with energy above a threshold frequency.
  
  

• The kinetic energy of ejected electrons depends on the light’s frequency.
  
  

Atomic Spectra

• Atoms emit light only at specific wavelengths, producing a line spectrum.
  
  

• Electrons exist only in discrete energy levels (quantized energy states).
  
  

• When electrons move to lower energy levels, they emit photons with energy corresponding to the difference between levels.
  
  

Bohr Model

• Electrons orbit the nucleus in fixed energy levels or shells.
  
  

• Electrons absorb energy to move to higher levels and emit energy when returning to lower levels.
  
  

• The ground state is the lowest energy state of an atom.
  
  

Quantum Mechanics and Wave Nature of Electrons

• Electrons exhibit wave-particle duality and can be described as standing waves around the nucleus.
  
  

• Only certain orbits with whole-number multiples of wavelengths are allowed.
  
  

• The exact position and momentum of an electron cannot be known simultaneously (Heisenberg’s uncertainty principle).
  
  

• Electrons occupy orbitals—regions of high probability where electrons are likely to be found.
  
  

Orbitals and Electron Probability

• Orbitals represent probability distributions, not fixed paths.
  
  

• Electrons can transition between orbitals by absorbing or emitting quanta of energy.
  
  

• Orbitals can overlap, unlike discrete orbits in the Bohr model.

Quantum Numbers

Quantum numbers describe the orbital where an electron may be found:

1. Principal Quantum Number (n):

   • Describes the size and energy of an orbital

   • Allowed values: n=1,2,3,…n = 1, 2, 3, \ldots (up to infinity)

2. Angular Momentum Quantum Number (l):

   • Describes the shape of the orbital (subshell)

   • Allowed values: l=0l = 0 to (n−1)(n-1)

   • Each value corresponds to a letter:

     • 0 = s

     • 1 = p

     • 2 = d

     • 3 = f

3. Magnetic Quantum Number (m_l):

   • Describes the orientation of the orbital in space

   • Allowed values: ml=−lm_l = -l to +l+l

4. Spin Quantum Number (m_s):

   • Describes the spin of an electron

   • Allowed values: +12+\frac{1}{2} or −12-\frac{1}{2} (two opposite spin directions)

Additional Notes:

• Orbitals have areas of high electron probability separated by nodes (areas of zero probability).

• According to the Pauli Exclusion Principle, no two electrons in an atom can have the same set of all four quantum numbers (n,l,ml,ms)(n, l, m_l, m_s).

• Since electrons in the same orbital share n,l,mln, l, m_l, they must have opposite spins (ms)(m_s).

• Therefore, each orbital can hold a maximum of two electrons with opposite spins.

Energy Level Diagrams and Electron Configuration

Energy Levels – Rules

• Electrons fill lowest energy levels first (Aufbau Principle).

• Energy increases as nn increases.

• Within the same nn, energy order: s<p<d<fs < p < d < f.

• Sublevels can overlap in larger atoms (e.g., 4s fills before 3d).

Rules for Energy Level Diagrams

1. Aufbau Principle: Fill lowest energy orbitals first.

2. Pauli Exclusion Principle: No two electrons share the same set of quantum numbers (max 2 per orbital, opposite spins).

3. Hund’s Rule: Fill degenerate orbitals singly before pairing electrons.

Use boxes for orbitals, arrows for electrons.

Shorthand Electron Configuration

• Use noble gas symbol to shorten configuration, e.g., [He][He] replaces 1s21s^2.

Periodic Table & Electron Configuration

• Number of electrons in s, p, d, f orbitals matches the periodic table block columns.

• Transition metals fill d orbitals (10 electrons = 10 elements per period).

Explaining Ion Charges

• Electrons are removed from the highest energy level first.

• Examples:
  
  

  • Zn:[Ar]4s23d10\text{Zn}: [Ar] 4s^2 3d^{10} → Zn2+:[Ar]3d10\text{Zn}^{2+}: [Ar] 3d^{10} (loses 2 electrons from 4s)
    
    

  • Pb:[Xe]6s24f145d106p2\text{Pb}: [Xe] 6s^2 4f^{14} 5d^{10} 6p^2 →
    
    

    • Pb2+:[Xe]6s24f145d10\text{Pb}^{2+}: [Xe] 6s^2 4f^{14} 5d^{10} (loses 2 from 6p)

    • Pb4+:[Xe]4f145d10\text{Pb}^{4+}: [Xe] 4f^{14} 5d^{10} (loses 6s and 6p electrons)

Exceptions in Electron Configurations

• Cr: Expected [Ar]3d44s2[Ar] 3d^4 4s^2, Actual [Ar]3d54s1[Ar] 3d^5 4s^1 (half-filled 3d more stable).

• Cu: Expected [Ar]3d94s2[Ar] 3d^9 4s^2, Actual [Ar]3d104s1[Ar] 3d^{10} 4s^1 (fully filled 3d more stable).
  
  

Lewis Theory and Bonding

Octet Rule

• Atoms (except hydrogen) tend to have 8 valence electrons when bonded.

• Helps predict shapes of molecules.

• Exceptions:

  • Molecules with odd number of electrons.

  • Atoms with less than 8 electrons (incomplete octet).

  • Atoms with more than 8 electrons (expanded octet).

Ionic Bonds

• Form between a metal and a non-metal.

• Electrons are transferred (lost by metal, gained by non-metal).

• Examples: LiF, CaF₂.

Covalent Bonds

• Form between two non-metals.

• Electrons are shared.

• Examples: H₂O, CO₂.

How to Draw Lewis Structures

1. Count total valence electrons.

2. Draw the basic structure, pick the central atom.

3. Complete octets of outer atoms.

4. If central atom doesn’t have 8 electrons, make double or triple bonds.

Exceptions to Octet Rule

• Transition metals and heavier elements don’t always follow the rule because of d orbitals.

• Example: Tin (Sn) ions don’t have full octets.

• Iron (Fe) ions stabilize by half-filled d orbitals.

Expanded Octet

• Atoms in the 3rd period and beyond can have more than 8 electrons.

• Example: PCl₅.

Incomplete Octet

• Some atoms are stable with less than 8 electrons.

• Example: BF₃.

Bond Order

• Number of shared electron pairs between atoms.

• Higher bond order = shorter and stronger bond.

Resonance

• When more than one valid Lewis structure exists by moving electrons.

• Real structure is a hybrid of all forms.

• Example: CHO₂⁻, CO₃²⁻.

Formal Charge (FC)

• Helps find the best Lewis structure.

• Formula:
  FC=Valence electrons−Non-bonding electrons−Bonding electrons2FC = \text{Valence electrons} - \text{Non-bonding electrons} - \frac{\text{Bonding electrons}}{2}

• Structures with smallest formal charges are preferred.

• Example: SO₄²⁻ has resonance structures with formal charges to explain stability.

VSEPR Theory 

• Electron pairs (bonding or lone pairs) repel each other.

• Lone pairs repel more strongly than bonding pairs.

• Molecules arrange themselves in 3D shapes to keep electron pairs as far apart as possible.

How to predict shape:

1. Count total electron pairs around the central atom (single, double, triple bonds all count as 1).

2. Count lone pairs separately.

3. Use total pairs (called charge clouds) to find the shape.

Bond angles change mainly due to lone pairs pushing bonding pairs closer.

Covalent Bonding & Polarity

• Covalent bonds: atoms share electrons to fill valence shells.

• If both atoms have equal electronegativity (EN), the bond is nonpolar covalent.

• If EN difference > 1.7 → mostly ionic bond (electron transfer).

• If EN difference between 0.4 and 1.7 → polar covalent (unequal sharing).

• If EN difference < 0.4 → mostly nonpolar covalent.

Polarity of molecules

• Even if bonds are polar, a molecule can be nonpolar if it is symmetrical (e.g., CCl₄).

• Polar molecules have a net dipole moment due to uneven charge distribution.

Intermolecular Forces (IMFs)

• IMFs are forces between molecules (not bonds inside molecules).

• Stronger IMFs mean higher melting/boiling points, solubility in water, etc.

Types of IMFs:

1. Dipole-Dipole Forces
   Between polar molecules (+ and – ends attract).

2. Hydrogen Bonding
   Special strong dipole-dipole force when H bonds with N, O, or F.

3. London Dispersion Forces (LDFs)
   Weakest, occur between all molecules, especially nonpolar ones. Larger molecules/atoms have stronger LDFs.

Physical Properties of Liquids

• Surface tension: water molecules stick together (cohesion), allowing things like water drops or insects walking on water.

• Adhesion: water molecules stick to other surfaces, like glass.

• Capillarity: water climbs up narrow tubes due to adhesion and cohesion (important in plants).

Chemical Systems and Equilibrium

Equilibrium Systems

• Dynamic: reactions keep happening, but no net change.

• Reversible: can proceed forward or backward.

• At equilibrium, concentrations of reactants and products stay constant.

• Forward and reverse reactions happen at the same rate.

• Happens in a closed system (no matter in/out).

Types of Equilibrium

1. Solubility Equilibrium

• Happens when dissolved solute and undissolved solute exist together at a constant concentration.

• Rate of dissolving = rate of crystallizing.

• Example:
  CaSO4(s)⇌Ca2+(aq)+SO42−(aq)\text{CaSO}_4 (s) \rightleftharpoons \text{Ca}^{2+} (aq) + \text{SO}_4^{2-} (aq)

2. Phase Equilibrium

• Between phases like solid/liquid or liquid/gas in a closed system.

• Example: evaporation = condensation at equilibrium (constant vapor pressure).

• At melting/freezing point, rate of melting = rate of freezing.

• Example:
  H2O(s)⇌H2O(l)\text{H}_2\text{O}(s) \rightleftharpoons \text{H}_2\text{O}(l)

3. Chemical Reaction Equilibrium

• Happens in closed systems where reactions can reverse.

• Example:
  N2O4(g)+heat⇌2NO2(g)\text{N}_2\text{O}_4 (g) + \text{heat} \rightleftharpoons 2 \text{NO}_2 (g)

Reversible Reactions

• Equilibrium composition is the same whether starting with reactants or products.

• Example:
  N2O4(g)⇌2NO2(g)\text{N}_2\text{O}_4 (g) \rightleftharpoons 2 \text{NO}_2 (g)

Percent Reaction at Equilibrium

• Shows how much product forms compared to the maximum possible (theoretical yield).

• Example: In the H₂ + I₂ ⇌ 2HI system, you compare initial and equilibrium concentrations to find percent reaction.

Classes of Chemical Equilibrium

Equilibrium Constant, K

For:

• Only gases and aqueous species appear in the expression. Solids and liquids are not included because their concentration doesn’t change.

• KeqK_{eq} changes with temperature.

• Predicts how much reactants/products at equilibrium.

Heterogeneous vs. Homogeneous Equilibria

• Homogeneous: all species same phase (all gases or all aqueous).

• Heterogeneous: different phases (solid and aqueous).

• Example:

NaOH(s)⇌Na+(aq)+OH−(aq)\text{NaOH}(s) \rightleftharpoons \text{Na}^+(aq) + \text{OH}^-(aq) Keq=[Na+][OH−]K_{eq} = [\text{Na}^+][\text{OH}^-]

Magnitude of K

• K≫1K \gg 1: equilibrium favors products.

• K≪1K \ll 1: equilibrium favors reactants.

• K≈1K \approx 1: significant amounts of both reactants and products present.