Properties of Water and States of Matter

Properties of Water

• Water has a high capacity to absorb and release thermal energy with minimal temperature changes.
• Ice floats on liquid water due to its lower density, forming a six-sided crystal structure.
• Water exhibits a high surface tension due to strong attractive forces between surface molecules.
• Water molecules can form concave meniscus and exhibit capillary action because of their cohesive and adhesive properties.
• Water acts as a universal solvent because its positive and negative ends attract other substances.

States of Matter and Solubility

• Solid (s) - represents a substance in a solid state.
• Liquid (l) - represents a substance in a liquid state.
• Gas (g) - represents a substance in a gaseous state.
• Aqueous (aq) - signifies a substance dissolved in water.

Determining States of Matter

• Elements - Determined by their state as noted on the periodic table.
• Molecular Compounds - Smaller molecules tend to be gases, while larger molecules are typically liquids or solids.
• Ionic Compounds - Usually solid at room temperature and pressure.
• Ionic Compounds in Aqueous Environments - Refer to the solubility table for their state:
  • If very soluble, the state is aqueous (aq).
  • If only slightly soluble, the state is solid (s).

Acids and Bases

• pH - Measures the concentration of hydrogen ions (H+) in a solution.
• Acidity - Every decrease of 1 on the pH scale represents a ten-fold increase in hydrogen ion concentration.
• Acid-base Indicators - Chemicals that change color based on pH, like litmus paper and phenolphthalein.

Properties of Acids and Bases

Acids

• Molecular compounds that ionize in water to release hydrogen ions.
• Recognized by the presence of hydrogen (H) on the left side or the –COOH group on the right side of the chemical formula.
• Become acidic only in aqueous environments.
• Solid, liquid, or gas states follow ionic compound naming rules.

Naming Acids in Aqueous Environments

Non-metals

• Solid, liquid, or gas states.
• Poor conductors of heat and electricity.
• Found to the right of the staircase on the periodic table.

Metalloids

• Properties lie between metals and non-metals.

Ions and Isotopes

• Isotopes - Atoms of the same element with different numbers of neutrons, leading to different mass numbers.
  • Example: Hydrogen-2.
• Atomic Molar Mass - The average of all an element's isotopes' masses found on the periodic table.
• Most Common Isotope - Found by rounding the mass number on the periodic table.
• Ionization - The process where an atom gains or loses electrons to form an ion.
• Cation - A positively charged ion formed when an atom (usually a metal) loses one or more electrons.
  • Example: Sodium ion (Na+) loses one electron, resulting in a 1+ charge.
• Anion - A negatively charged ion formed when an atom (usually a non-metal) gains one or more electrons.
  • Example: Chlorine atom (Cl-) gains one electron, resulting in a 1- charge.

Ionic Compounds

• Formed when electrons transfer from one atom to another, resulting in a complete valence energy level/stable octet for each atom.
• Attraction between cation and anion through electrostatic forces, forming a neutral ionic compound joined by an ionic bond.
• Formula Unit - Represents the simplest ratio of ions in an ionic compound.

Naming Binary Ionic Compounds:

1. Name the cation first. Use the element's name (typically a metal).
2. Name the anion second. Use the first part of the element's name (typically a non-metal) and change the ending to -ide.

Writing Formulas for Binary Ionic Compounds:

1. Write the cation symbol followed by the anion symbol.
2. Determine the number of each ion needed to balance the charges.
3. Use subscripts to indicate the ratio of cations to anions.
   • No subscript is necessary for a single atom.

Multivalent Elements

• More than one stable ion charge. The first charge listed on the periodic table is the most common.
• Roman numerals in brackets after the element name indicate the ion's charge.
  • Example: Iron(II) bromide.
• Determine the cation charge from the anion charge if provided only with the chemical formula.
  • Example: FeBr2 is iron(II) bromide.

Polyatomic Ions

• Several non-metallic atoms joined by covalent bonds.
• Naming rules are similar to binary ionic compounds, but the polyatomic ion's ending is NOT changed.
• Brackets are used around polyatomic ions when a subscript is required.

Properties of Ionic Compounds

• High melting point/solid at room temperature - Strong and continuous attraction between cations and anions holds them tightly in a crystal lattice.
• Crystalline structure - Retains its crystal shape, even when ground into powder.
• Solubility in water - Cations attract the negative end of water molecules, and anions attract the positive end, leading to solubility.
• Form electrolytic solutions - Dissociate into ions. The greater the ion concentration, the higher the conductivity.

Molecular Elements and Compounds

• Covalent bond - Forms when atoms share a pair of valence electrons to achieve a stable octet in each atom.
• Molecule - An independent unit composed of fixed numbers of non-metallic atoms held together by covalent bonds.
  • Examples: H2, CO2, CH4, PO43-.