Characteristics of Solids and Liquids

Solids, which can be classified as:

Amorphous - No arranged structure of molecules

Crystalline - Molecules arranged into patterns

Liquids, which have special properties, such as:

• Surface Tension

• Viscosity

• Vapor Pressure

• Boiling Point

• Heat of Vaporization

Intermolecular Forces of Attraction

1. Dipole-Dipole — Created between polar molecules, always partially positive and partially negative.

   

2. Hydrogen Bond — A special type of dipole-dipole (hydrogen). It is 5x to 10x stronger than dipole-dipole. It is intermolecular in nature. It always bonds to Nitrogen, Oxygen, or Fluoride.

   

3. Ion-Dipole — Acts between an ion(+,-) and a polar molecule

   

4. London Dispersion Force — Between all atoms and molecules, Non polar (zero dipole moment)

   

Note: - - - - - - = Bond ————— = Connection

3 steps in the formation of liquids

1. “Expanding the solvent” - Overcoming the intermolecular forces in the solvent to give room for solute.

2. “Expanding the solute” - Separating the solute into it’s individual components/

3. Allowing the solute and solvent to interact.

Enthalpy of Solution

H_Solution = H₁+H₂+H₃

(+) - Endothermic (Heat Absorbed)

(-) - Exothermic (Heat Radiated)

Concentration

Percent by mass and volume and Mass-Volume

% by mass = (mass of solute(g) / mass of solution(mL)) × 100%

% by volume - (volume of solute(g)/volume of solution(mL))

% by mass-volume = (mass of solute (g) / volume of solution (mL)) × 100%

Mass of solute (m) = Molarity (M) × Volume (L) × Molar mass (g/mol)

Solution Stoichiometry

Molarity - # of moles of solute dissolved in a solution

Molarity (M) = moles of solute / liters of solution

Molality - # of moles of solute dissolved in 1 kg solution

Molality (m) = moles of solute / mass of solvent (kg)

1st Law of Thermodynamics

Law of Conservation of Energy - Energy cannot be created nor destroyed, only transformed.

Physics: U = Q - W; Work is done by the system

Chemistry: U = Q + W; Work is done on the system

Note: W = (-) done by the system

W = (+) done on the system

Q = + (absorbs Heat) Endothermic

Q = - (loses Heat) Exothermic

3 types of system

1. Open - Matter and Energy freely move from system to surroundings and vice versa. (Example: Open teakettle)

2. Closed - Energy freely moves from system to surroundings and vice versa. While Matter is restricted from moving. (Example: Closed teakettle)

3. Isolated - Energy nor matter cannot freely move from system to surroundings and vice versa. (Example: Vacuum Chamber)

3 types of heat transfer:

Conduction - Heat transferred through contact (solids)

Convection - Heat transferred through fluid movement (gas or liquid)

Radiation - Heat transferred through electromagnetic waves

Diatomic Elements

• Nitrogen (7) - N₂^-3 (14.01 g)

• Oxygen (8) - O₂^-2 (16.00 g)

• Fluorine (9) - F₂^- (19.00 g)

• Chlorine (17) - Cl₂^- (35.45 g)

• Bromine (35) - Br₂^- (79.90 g)

• Iodine (53) - I₂^- (126.90 g)

• Hydrogen (1) - H₂⁺ (1.01 g)

Halogen Group

• Fluorine (9) - F₂^- (19.00 g)

• Chlorine (17) - Cl₂^- (35.45 g)

• Bromine (35) - Br₂^- (79.90 g)

• Iodine (53) - I₂^- (126.90 g)

• Astatine (85) - At^- (210 g)

• Tennessine (117) - Te^- (294 g)

Important Elements / Compounds

• Sulfate - SO₄^-2 (96.06 g/mol)

• Ammonium - NH₄⁺ (18.04 g/mol)

• Aluminum - Al⁺³ (26.98 g/mol)

• Phosphorus - P^-3 (32.06 g/mol)

• Carbonate - CO₃² (60.01 g/mol)

• Nitrate - NO₃^- (62.01 g/mol)

• Chlorine - Cl₂^- (35.45 g/mol)

• Hydroxide - OH^- (17.01 g/mol)

• Hydrochloric Acid - HCl^-1 (36.46 g/mol)

• Sulfuric Acid - H₂SO₄⁺⁶ (98.08 g/mol)

• Carbon - C⁺⁴ (12.01 g/mol)

Properties of Periodic Table Groups

Alkali Metals (GROUP 1/1A)

• Less dense than other metals

• One loosely bound valence electron

• Highly reactive, with reactivity increasing moving down the group

• The largest atomic radius of elements in their period

• Low ionization energy

• Low electronegativity

Alkaline Earth Metals (GROUP 2/2A)

• Two electrons in the valence shell

• Readily form divalent cations

• Low electron affinity

• Low electronegativity

Transition Metals (GROUP 3-12/1-10B)

Lanthanides (rare earth) and actinides are also transition metals. The basic metals are similar to transition metals but tend to be softer and hint at nonmetallic properties. In their pure state, all these elements tend to have a shiny, metallic appearance. While there are radioisotopes of other elements, all actinides are radioactive.

• Very hard, usually shiny, ductile, and malleable

• High melting and boiling points

• High thermal and electrical conductivity

• Form cations (positive oxidation states)

• Tend to exhibit more than one oxidation state

• Low ionization energy

Metalloids or Semimetals

• Electronegativity and ionization energy intermediate between that of metals and nonmetals

• May possess a metallic luster

• Variable density, hardness, conductivity, and other properties

• Often make good semiconductors

• Reactivity depends on the nature of other elements in the reaction

Nonmetals

Halogens and noble gases are nonmetals, although they have their groups, too.

• High ionization energy

• High electronegativity

• Poor electrical and thermal conductors

• Form brittle solids

• Little if any metallic luster

• Readily gain electrons

Halogens

The halogens exhibit different physical properties from each other but do share chemical properties.

• Extremely high electronegativity

• Very reactive

• Seven valence electrons, so elements from this group typically exhibit a -1 oxidation state

Noble Gases

The noble gases have complete valence electron shells, so they act differently. Unlike other groups, noble gases are unreactive and have very low electronegativity or electron affinity.