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Attractive and Repulsive Forces

• Attraction - opposite charge

• Repulsion - similar charge (+,+) (-,-)

Cohesive forces

Attraction between identical (similar) molecules

Adhesive forces

Attraction between different molecules

Intramolecular Forces of Attraction

Within a molecule; Na+, Cl- = Na-Cl

• Covalent Bond

• Coordinate covalent bond

• Ionic bond

Covalent Bond

Mutual sharing of electron

• Nonmetal + Nonmetal

Nonpolar Covalent Bond

• Bonding electrons shared equally (symmetrical distribution of electron) creating no charges on atoms

• NO or very slight EN difference (<0.4)

  • C-C (2.55 - 2.55 = 0)

• Diatomic molecules (gen-u-ine)

  • Oxygen, Hydrogen, Nitrogen

  • Chlorine, Iodine, Fluorine, Bromine (-ine except astatine)

• Inert gas (right most part of periodic table)

Polar Covalent Bond

• Bonding electrons shared unequally creating partial charges of atoms (asymmetrical)

• Slight EN difference (0.4–1.8)

• e.g., HCl, H2O

Coordinate Covalent Bond

Non-mutual sharing of electron

Ionic Bond

• Metal + Nonmetal

• aka: Electrovalent bond

• Full charge (cation + anion)

• Transfer of electron

• Large EN difference (>1.8)

True

1. HCl - polar covalent bond

2. CH4 - nonpolar covalent bond

3. NaCl - ionic bond

4. I2 - nonpolar covalent bond

5. KI - ionic bond

True

• Covalent - NM + NM

• Ionic - M + NM

• Metallic - M + M (alloy)

Intermolecular Forces of Attraction

• Between molecules (mixture)

• Responsible for most of the physical and chemical properties of matter

• The stronger the IMF = higher boiling point

Dipole

• Polar (partial charge)

• Permanent dipole moment (or separation of charge) due to uneven charge distribution

Induced Dipole

• Nonpolar (0 charge alone; can have induced polarity - temporary charge, temporary polar)

• Temporary dipole moment due to distortion of electron cloud induced by a polar molecule

Ion

• Salt

• Permanent charge

Van der Waals Forces

• NON-IONIC interactions between molecules

• WEAKEST of all intermolecular forces of attraction

Van der Waals Forces

• Dipole–Dipole

• Dipole–Induced Dipole

• Induced Dipole–Induced Dipole (London Dispersion Forces)

Dipole–Dipole

• aka: Keesom Forces

• MOA: Alignment or orientation

• FOA between: Polar + Polar

• Ex: HCl + HCl

  • Both have a permanent dipole. The slightly positive H of one is attracted to the slightly negative Cl of the other.

• Bond strength: 1–7 Kcal/mole

Dipole–Induced Dipole

• aka: Debye Forces

• MOA: Induction

• FOA between: Polar and Nonpolar

• Example: H2O (polar) + O₂ (induced dipole)

  • The polar water molecule, with its slightly negative oxygen and slightly positive hydrogen ends, causes a temporary shift in the electron cloud of the O₂ molecule, inducing a temporary dipole in the O₂.

• Bond strength: 1-3 kcal/mol

Induced Dipole–Induced Dipole

• aka: London

• MOA: Dispersion

• FOA between: Nonpolar + Nonpolar

• WEAKEST VAN DER WAALS

• Bond strength: 0.5-1 kcal/mol

• Example: Argon + Argon

  • The electrons in an argon atom can momentarily shift, creating a temporary dipole. This dipole then induces a dipole in a nearby argon atom, causing a weak, temporary attraction between them.

True

Ion–Dipole

• FOA between: Salt and Polar molecule

• H2O + NaCl

  • Na⁺ (positive ion) is attracted to the oxygen side of water (which is negative).

  • Cl⁻ (negative ion) is attracted to the hydrogen side of water (which is positive).

Ion–Induced Dipole

• FOA between: Salt + Nonpolar molecule

• I₂ + KI → KI₃

  • KI dissociates into ions in solution: K⁺ (cation), I⁻ (anion)

  • I₂ is a nonpolar molecule, meaning its electrons are equally shared.

  • When I⁻ (iodide ion) approaches I₂, it pushes the electrons in I₂ slightly, inducing a dipole.

  • This Ion (I⁻) + Induced Dipole (in I₂) creates a weak attraction → forms I₃⁻ (triiodide ion).

  • KI⁻ (ion) + I₂ (induced dipole) → KI₃⁻ (water soluble)

Hydrogen Bond

• aka: Hydrogen Bridge

• FOA between: Interaction between a molecule containing a hydrogen atom and a strongly electronegative atom (F, O, N)

• H₂O — H₂O (unique: intra and inter)

  • One water molecule has H–O–H

  • The oxygen is more electronegative than hydrogen, so:

    • O gets a partial negative charge (δ⁻)

    • H gets a partial positive charge (δ⁺)

  • The δ⁺ hydrogen of one water molecule is attracted to the δ⁻ oxygen of another water molecule

Intramolecular Forces

Intramolecular Forces

A. Polar Covalent
B. Nonpolar covalent
C. Ionic
D. Metallic bond

1. Transfer of electron → C. Ionic

2. Equal sharing of electron → B. Nonpolar covalent

3. Unequal sharing of electron → A. Polar covalent

4. Metal + Nonmetal → C. Ionic

5. Alloy formation → D. Metallic bond

Ionic

Transfer of electron

Nonpolar covalent

Equal sharing of electron

Polar covalent

Unequal sharing of electron

Ionic

Metal + Nonmetal

Metallic bond

Alloy formation

Intermolecular Forces of Attraction

Intermolecular Forces of Attraction

A. Van der Waals
B. Ion-dipole
C. Ion-induced dipole
D. Hydrogen bonds
E. Repulsive and attractive forces
F. Electrovalent forces

1. Keesom, Debye and London forces → A. Van der Waals

2. Solubility of ionic crystals in water → B. Ion-dipole

3. Formation of iodide complex, which accounts for the solubility of iodine in KI solution → C. Ion-induced dipole

4. Accounts for the unusual properties of water (high dielectric constant, abnormally low vapor pressure and high BP) → D. Hydrogen bonds

5. Forces necessary to cohere and forces necessary to prevent molecular interpenetration → E. Repulsive and attractive forces

6. Weak electrostatic force that brings about condensation of nonpolar gas molecules → A. Van der Waals (London dispersion forces)

Van der Waals

Keesom, Debye and London forces

Ion-dipole

Solubility of ionic crystals in water

Ion-induced dipole

Formation of iodide complex, which accounts for the solubility of iodine in KI solution

Hydrogen bonds

Accounts for the unusual properties of water (high dielectric constant, abnormally low vapor pressure and high BP)

Repulsive and attractive forces

Forces necessary to cohere and forces necessary to prevent molecular interpenetration

Van der Waals (London dispersion forces)

Weak electrostatic force that brings about condensation of nonpolar gas molecules

Physical Properties of a System

Additive, Constitutive, Colligative

Additive

• Depends on sum

• Mass or weight

• Molecular Weight

• Volume

Constitutive

• Depends on type and arrangement of molecules

• Solubility

• Refractive index (Refractometer)

• Opitcal activity (Polarimeter)

• Viscosity ⭐

Colligative

• Depends on the # of solute

• Concentration dependent

• ∆P - Vapor pressure lowering

• ∆Tb - Boiling point elevation

• ∆Tf - Freezing point depression

• π - Osmotic pressure

Physical Properties of a System

Physical Properties of a System

Properties of drug molecules

A. Electromagnetic radiation
B. Refractive index
C. Optical rotatory dispersion
D. Circular dichroism
E. Permanent dipole moment
F. Optical rotation

1. Characteristic frequency, wavelength or wavenumber → A. Electromagnetic radiation

2. Non-ionic phenomenon wherein the molecule has no net charge → E. Permanent dipole moment

3. Measurement of the angle of rotation → F. Optical rotation

4. n = sin i / sin r → B. Refractive index

5. Absorption spectroscopy based on differential absorption of left or right circularly polarized light → D. Circular dichroism

Electromagnetic radiation

Characteristic frequency, wavelength or wavenumber

• Increased frequency = shorter wavelength = Greater energy = GAMMA RAY

• Decreased frequency = longer wavelength = Lower energy = RADIOWAVES

Permanent dipole moment (Van der waals)

Non-ionic phenomenon wherein the molecule has no net charge

Optical rotation

Measurement of the angle of rotation

Refractive index

n = sin i / sin r

• (Incidence, Refraction); Snell's law

Circular dichroism

Absorption spectroscopy based on differential absorption of left or right circularly polarized light

First Law of Thermodynamics

• Defines the law of conservation of energy

• Energy cannot be created or destroyed

• Energy can be transformed; transferred; interconverted

• Energy can be interconverted, but the sum of energy must remain constant

• Closed system:

  • ∆E = EB – EA

  • ∆E is equal to ZERO for a cyclic process in a closed system

Second Law of Thermodynamics

• Defines entropy (S)

  • Measures the degree of randomness; disorderliness

  • Gas has greater entropy

  • Never decreasing; always increasing or constant

• The entropy of the system plus that of the surroundings must increase in an irreversible process and remains constant in a reversible process

• ∆Ssystem + ∆Ssurr ≥ 0

Third Law of Thermodynamics

The entropy of a pure substance is zero when that substance is in a perfectly crystalline state at absolute zero

Gibbs Free Energy

• A measure of chemical energy

• Represents the combined contribution of enthalpy and entropy

• Enthalpy - HEAT CONTENT

• Entropy - RANDOMNESS

• Formula: ∆G = ∆H – T∆S

Gibbs Free Energy

Interpretation:

• ∆G < 0 → Spontaneous (negative)

• ∆G = 0 → Equilibrium

• ∆G > 0 → Non-spontaneous (positive)

Gas Properties

• Indefinite shape

• Indefinite volume

• High compressibility = spacious

• Generally invisible

• Exhibit perfect elasticity

• Weak intermolecular force of attraction

• Molecules move randomly, in constant, rapid motion

Boyle’s Law

• Constant: Temperature (BoTe)

• P inversely proportional to V

• P1V1 = P2V2

Charles’ Law

• Constant: Pressure (ChaPre)

• Temperature is directly proportional to Volume

• Note: Temperature must be expressed in Kelvin (K = °C + 273)

• V1/V2 = T1/T2

• Ex. Hot air balloon

Gay-Lussac’s Law

• Temperature is directly proportional to Pressure

• Constant: Volume (Vice Ganda - Gay)

• Ex. Aerosol

Combined Gas Law

• Combination of Boyle’s law, Charles’ law, and Gay Lussac’s law

• Shows the relationship between pressure, volume, and temperature of a gas

• Formula: P1V1/T1 = P2V2/T2

Ideal Gas Law

• Formula: PV = nRT

• n = w ÷ MW

• R = ideal gas constant

  • 0.0821 L·atm/mol·K

  • 8.314 J/mol·K

  • 1.987 cal/mol·K (round off to 2)

  • Remember: J8 2cal

Standard PVT

• P = 1 atm; 760 torr/mmHg

• V = 22.4 L

• T = 273 K; 0°C

Real Gas Law

• Aka Van der Waals equation

• Modification of ideal gas law

Henry’s Law

• Constant: Temperature

• AMOUNT OF DISSOLVED GAS is DIRECTLY proportional to the PARTIAL PRESSURE of that gas in equilibrium with that liquid

• Formula: P = KH × C

• Soda is bottled under high pressure to keep CO₂ dissolved. When you open it, pressure drops → gas escapes → fizz!

Avogadro’s Law

• VOLUME is DIRECTLY proportional to AMOUNT OF GAS (mole)

• Constant: Temperature and Pressure (ATP)

Graham’s Law

• Effusion or diffusion rate is inversely proportional to the square root of molar mass (or density).

• High molar mass → slower effusion

• Low molar mass → faster effusion

• Helium - lighter, Air - heavier

Dalton’s Law of Partial Pressure

Ptotal = sum of partial pressures

• Formula: P1 + P2 + P3 + …..

• Where:

  • P1,P2,P3,… are the partial pressures of the gases.

  • The pressure a single gas would exert if it were alone in the container.

B. Van der Waals forces

Identify the weak intermolecular forces, examples of which are the Keesom-Debye, and London forces.

A. Ion-dipole, ion-induced dipole forces

B. Van der Waals forces

C. Hydrogen bonds

D. Repulsive and attractive forces

C. Ion-dipole, ion-induced dipole forces

Identify the forces that account in part for the solubility of ionic crystalline substances in water and are presumed to account for the solubility of iodine in a solution of KI.

A. Hydrogen bonds

B. Repulsive and attractive forces

C. Ion-dipole, ion-induced dipole forces

D. Electrovalent forces

B. Van der Waals forces (London)

Identify the weak electrostatic forces that bring about condensation of nonpolar gas molecules to form liquids and solids when molecules are brought quite close to one another.

A. Hydrogen bonds

B. Van der Waals forces

C. Ion-dipole, ion-induced dipole forces

D. Repulsive and attractive forces

A. Gibbs Free Energy

Identify the Law/Principle which states that when the reaction is reversible, ∆G = 0.

A. Gibbs Free Energy

B. Ideal Gas Law

C. Henry’s Law

D. Faraday’s Law

D. ∆E = EB + EA (minus)

Identify the statement/mathematical expression, which does NOT express the First Law of Thermodynamics.

A. States that energy is conserved

B. Forms of energy can be interconverted, but the sum of energies remains constant

C. ∆E is equal to zero for a cyclic process in a closed system

D. ∆E = EB + EA

A. 3rd Law of Thermodynamics

Select the concept that explains the thermodynamics state of a perfect crystal at absolute zero.

A. 3rd Law of Thermodynamics

B. 2nd Law of Thermodynamics

C. 1st Law of Thermodynamics

D. 1st, 2nd, 3rd Laws of Thermodynamics

A. 1.24

If 0.55 g of a gas dissolves in 1.0 L of water at 2 atm of pressure, how much will dissolve at 4.5 atm? (Henry - direct)

A. 1.24

B. 0.55

C. 0.01

D. NOTA

C. 819 mL

A 600 mL sample of nitrogen is warmed from 200 K to 273 K. Find its new volume if the pressure remains constant. (Charles - direct)

A. 150 mL

B. 500 mL

C. 819 mL

D. NOTA

C. Pressure

According to Boyle’s Law, the volume of a gas has an inverse relationship with ____.

A. Moles

B. Temperature

C. Pressure

D. Weight