Empirical Formula, Moles, and Bonding

• These notes cover empirical formula, moles, combustion, Bunsen burners, ionic and covalent bonding, polarity, electronegativity, nomenclature, hydrogen bonding, and multiple bonds.

Empirical Formula

• Defined as a chemical formula showing the simplest whole-number ratio of elements or ions.

Mole (mol)

• A counting unit in chemistry where 1 mole contains 6.02 x 10^23 particles.
• 6.02 x 10^23 is Avogadro's Number, representing the number of molecules (or atoms) in one mole.
• Conversions between mass and moles are possible.
• Equal volumes of gases at the same temperature and pressure contain the same number of molecules.
• Moles are used to measure small quantities and to determine the ratios of atoms in a compound from mass data.
• The mole ratio of elements in a compound corresponds to their subscripts in the chemical formula.

Combustion

• Involves burning a substance with oxygen from the air, resulting in oxygen being bonded to at least one other atom in the products.
• For example, magnesium burns to form magnesium oxide (MgO).
• Removing the crucible lid during an experiment can increase burning due to available oxygen.

Bunsen Burners

• Safety procedures include tying back hair, avoiding loose clothing, and wearing goggles.
• In case of any issue, immediately turn off the gas supply at the source.
• Light the burner with a flint striker (or match).
• Adjust gas and air flow to obtain a flame with distinct inner and outer cones.

Determining Ionic/Empirical Formulas

• Ionic Bonding Model
• Cross Charges and Reduce or by Balancing the ion charges.
• Experimentally using mass data.

Experimental Determination of Empirical Formulas

• Determine the mass in grams of each element in the chemical.
• Divide each element's mass by its average atomic mass (g/mol) from the periodic table to convert to moles.
• Divide all mole values by the smallest mole value.
• Round the resulting mole values to the nearest whole number to determine subscripts in the chemical formula.

Ionic Bonding

• Involves electron transfer between atoms, leading to the formation of ions with stable valence electron configurations.
• Ions are held together by electrostatic attraction between opposite charges.
• Bonding occurs between a metal and a nonmetal.
• The smallest repeating unit is a "Formula Unit".
• Forms cations (positive ions) and anions (negative ions).
• Also known as "Salts," and those that dissolve in water are called "Electrolytes".

Covalent Bonding

• Involves sharing electrons between atoms to achieve stable valence electron configurations.
• Atoms' electron clouds overlap, and shared electrons belong to both atoms.
• The smallest repeating unit is a "Molecule."
• Bonding occurs between two or more nonmetals.
• Formed by sharing electrons until all atoms have a full valence level.

Ionic vs Covalent Compounds

Ionic

• Formed between a metal and a nonmetal.
• Involves electron transfer to form ions.
• Shows full charges on ions with electrons in final locations separated a little.

Covalent

• Formed between two or more nonmetals.
• Involves electron sharing.
• Shows electrons shared between atoms, with no charge or partial charges.

Water, Salts, and Other Compounds

• Water is not a salt and bonds covalently.
• Covalent sharing in water may be unequal leading to partial charges.
• Charges on salts (and water) affect particle attraction, influencing their state at room temperature.
• Charges determine how substances interact and dissolve.
• NaCl (ionic) has full charges (Na +1 and Cl -1) resulting in strong attractions and close proximity of ions.
• Salts are typically solids at room temperature due to strong attractive forces from full charges.
• Water has partial charges, enabling it to dissolve substances with either partial or full charges.
• Full charges exhibit stronger attractions than partial charges.
• Mineral oil appears to have a different bonding type without full or partial charges.

Covalent Bonding Model Steps

• Applies to compounds containing only nonmetals.
• Draw the Lewis Dot structure for each atom to determine electron sharing.
• Atoms with two electrons on one side will not bond.
• Atoms with a single electron on one side will share with another atom having a single electron.
• Circle the electrons that are going to share (use a different color).
• Shared electrons are shown directly between the element symbols in final model.

Polarity

• Uneven sharing of electrons across a molecule related to unequal attractive forces from protons in each atom's nucleus.
• Results in partial charges when polar bonds are asymmetrically distributed.
• Electrons are shared more by one atom but not fully transferred.
• Polarity explains cohesive and adhesive forces in substances like water.

Electronegativity (e.n.)

• The tendency of an atom in a covalent bond to attract shared electrons.
• Ranges on a 0-4 scale, with 4 indicating the strongest attraction.
• Higher e.n. value indicates the direction of electron shift in a polar bond.

Nomenclature

• A system of naming compounds based on their substance type.

Polar Bonds

• Unequal distribution of electrons, causing one atom to have the shared electrons more often.

Polar Molecules

• Unequal distribution of electrons across the molecule from one or more polar bonds that have not canceled due to symmetry.
• Results in partially negative (δ-) and partially positive (δ+) regions, leading to interactions with charged substances.

Nonpolar Bonds

• Atoms share electrons equally (or nearly equally).

Nonpolar Molecules

• No polar bonds, or symmetry cancels polar bonds.
• Most oils are nonpolar.

Hydrogen Bonding

• Forms between a hydrogen atom bonded to a highly electronegative atom (N, O, or F).
• Creates a strong polar bond and attractive force between partial charges of different molecules.
• It's an attractive force (cohesion), not an actual bond.
• Covalent bonds are within molecules, while hydrogen bonds are between molecules.

Multiple Bonds

• Covalent compounds can form double or triple bonds by sharing more than two electrons.
• Double bonds share 4 electrons.
• Triple bonds share 6 electrons.