Intermolecular Force and Properties of Matter

Flashcards about intermolecular forces and properties of matter.

Intramolecular forces

Forces that occur within a molecule such as polar and nonpolar covalent bonds, ionic bonds, and metallic bonds.

Intermolecular forces

Forces that occur between molecules, including van der Waals interactions and hydrogen bonds.

van der Waals interactions

Keesom forces (between polar molecules), Debye forces (between a nonpolar molecule and a polar molecule), and London dispersion forces (between nonpolar molecules).

Hydrogen bond

Hydrogen on one molecule attached to (N, O, or F) atoms and either (N, O, or F) on another molecule.

Valence electrons

Represented as dots, each dot is one valence electron.

Polar molecules

Atoms do not share electrons equally, creating a dipole with slight positive and negative charges.

Examples of Polar molecules

HBr, HCl, HI, H2O, HF, NH3

Nonpolar molecules

Equal sharing of electrons between atoms.

Examples of Nonpolar molecules

O2, H2, Cl2 , I2, Br2 and Hydrocarbons: CH4 , C2H6 , C3H8.

Electronegativity

A measure of the tendency of an atom to attract a bonding pair of electrons towards itself.

Keesom Forces (dipole-dipole interactions)

Act between polar molecules (act between permanent dipoles). Oppositely charged ends attract and like ends repel resulting in a moderate strength interaction between permanent dipoles.

Examples of Keesom forces

HCl, HBr, HI

Keesom forces become stronger when

The difference of electronegativity between the atoms forming the molecule increases, the temperature decreases and the distance between molecules decreases.

Debye Forces

Occur between a nonpolar molecule and a polar molecule, where the permanent dipole in the polar molecule induces an electric dipole in the nonpolar molecule when it comes extremely close.

Examples of Debye Forces

Debye forces between O2 and H2O

London Dispersion Forces

An instantaneous dipole, that occurs accidentally in an atom or a molecule due to electrons movement, induces a similar dipole (induced dipole) in a neighboring atom or molecule. These forces occur in nonpolar molecules and are weak.

Examples of London Dispersion Forces

Homonuclear diatomic molecules: H2, Cl2, I2, O2, F2, Br2, Hydrocarbons: CH4, C3H8, C6H14 and Noble gases: He, Ne, Ar

London dispersion forces become stronger when

The size of atoms or molecules increases, the distance between molecules decreases, and the chain of carbon atoms in hydrocarbons becomes longer.

Which of the followings shows the highest London dispersion forces? a) C3H8 b) C4H10 c) C5H12 d)?

C6H14

Explain why F2 and Cl2 are gases, Br2 is a liquid and I2 is a solid at room temperature.

F2, Cl2, Br2 and I2 are nonpolar molecules, and they show London dispersion forces. As the size increased, London forces increased.

Hydrogen Bonding

Hydrogen on one molecule attached to a highly electronegative atom (N, O, or F) and either (N, O, or F) on another molecule. It is a strong force that occurs between Polar molecules.

Examples of Hydrogen Bonding

H2O, HF, NH3, CH3COOH

Explain how a covalent bond will form in N2?

Covalent bonding occurs when pairs of electrons are shared by atoms in order to gain more stability.

Which of the following bonds would require the most energy to break? A) Single bond B) Double bond C)

Triple bond

Bond energy

Energy required to break a particular covalent bond in a molecule in the gas phase.

Electronegativity Difference and Bond Type

The electronegativity difference (ΔEN) of two atoms determines their bond type: Ionic or Covalent. Covalent bonds can be either polar or nonpolar.

Nonpolar Covalent Bond

Non-metal atoms have no or little difference in electronegativity (EN) and electron pairs are shared equally.

Polar Covalent Bond

Non-metal atoms have a significant difference in electronegativity (EN) and electron pairs are shared unequally, drawn closer to the atom with higher EN.

Units of Temperature

Celsius (represented by °C). Based on freezing point of water as 0°C and boiling point of water as 100°C, Kelvin (represented by K), and Fahrenheit (°F). The SI unit of temperature is Kelvin (K)

Equations for Temperature Conversions

T (K) = T(°C) + 273 and T (°F) = 1.8 × T(°C) + 32

Balanced Chemical Equations

When the number of atoms involved in the reactants side is equal to the number of atoms in the products side (equal numbers of each type of atom on each side of the equation).