Chemical Bonding: Molecular & Ionic Compounds

• Unit 2 covers Chemical Bonding & Chemical Equations, specifically Chapters 3 & 4.

Molecular and Ionic Compounds

• Compounds are classified based on the type of bonds that hold the atoms together.

Molecular Compounds

• Atoms share electrons, forming covalent bonds.

Ionic Compounds

• Atoms are charged (ions), and the difference in charge holds them together.

• Some electrons are shared in molecular compounds, such as methane.

• Electrons are not shared in ionic compounds, like table salt.

• A molecule of methane (CH4) consists of a carbon atom bonded to four hydrogen atoms.

• A formula unit of sodium chloride (NaCl) is an ionic compound.

• Sodium atoms lose an electron to become sodium ions (Na+), while chlorine atoms gain an electron to become chloride ions (Cl-).

• The arrangement of Na+ and Cl- ions forms a crystal structure.

Chemical Formulas

• The number and type of atoms in a molecule are indicated using a chemical formula.
• A chemical formula serves as a form of notation.
• C3H6O indicates 3 carbon atoms, 6 hydrogen atoms, and 1 oxygen atom.
• Chemical formulas indicate the simplest structural unit.
• Sodium oxide, an ionic compound, is written as Na2O.
• Aluminum sulfate, consisting of two aluminum atoms and three sulfate groups, is written as Al2(SO4)3.
• This compound contains 3 sulfur atoms and 12 oxygen atoms (4x3).

Formulas and Names of Compounds

• A quick rule of thumb: nonmetal + nonmetal = molecular (covalent) compound and metal + nonmetal = ionic compound

Writing Ionic Formulas

• Charges must add to zero
• Cations (positive ions) are written first, and then anions (negative ions).
• Magnesium chloride: Mg+2 Cl-1
• Cobalt(III) oxide: Co+3 O-2

Naming Ionic Compounds

• The cation (positive ion) is named first
• If the element is a transition metal with several oxidation numbers, the charge is written as a Roman numeral.
• The anion (negative ion) is named second.
• If the compound is binary, end the name in -ide.
• Examples: NaCl, K2S, Al2O3, Cu2O

Naming Molecular Compounds

• If only one atom of the first element is present, it does not receive a prefix
• If there is more than one atom of the first element, it receives a prefix.
• When the compound has only two elements, it is binary and ends in -ide
• The second element always has a prefix.

Numerical Prefixes

• 1: mono-

• 2: di-

• 3: tri-

• 4: tetra-

• 5: penta-

• 6: hexa-

• 7: hepta-

• 8: octa-

• 9: nona-

• 10: deca-

• Examples: CO2, CO, P4S10, H2O

Writing Formulas of Molecular Compounds

• The less electronegative element (further left on the periodic table) is written first.
• If the first element has no prefix, there is only one atom, and no subscript is written.
• If the first element has a prefix, the proper number of atoms is written with a subscript.
• Determine the number of atoms of the second element and write the correct subscript.
• Examples: diarsenic pentabromide, carbon tetrachloride, sulfur hexafluoride, and iodine trioxide.

Polyatomic Ions

• Polyatomic ions are charged groups of covalently bonded atoms.
• Common polyatomic ions: NH4+ (Ammonium), OH- (Hydroxide), SO42- (Sulfate), NO3-1 (Nitrate), PO43- (Phosphate), CO32- (Carbonate)
• Writing Polyatomic Compound Formulas
• Charges MUST always add to zero
• Write the cation first, then the anion.
• If more than one polyatomic ion is needed to make the charges add to zero, use parentheses around the entire ion.
• Examples: Lithium hydroxide (Li+1 OH-1), Calcium nitrate (Ca+2 NO3-1), Aluminum phosphate (Al+3 PO4-3)

Writing Polyatomic Compound Names

• Name the cation first, then the anion.
• Examples: Na2SO4, Mg(OH)2, (NH4)3PO4, NiCO3

Formulas of Acids

• Acids begin with H
• Charges add to zero

Naming Acids

• For binary acids: hydro- -ic acid
• For acids that are not binary: -ic acid
• Examples: HCl, HI, H2S, H3PO4, H2SO4

Acids

• Acids are molecular compounds that form H+ when dissolved in water.
• (aq) is written after the formula to indicate the compound is dissolved in water.
• If not dissolved in water, it is not named as an acid
• Sour taste
• Acids dissolve many metals (like Zn, Fe, Mg; but not Au, Ag, Pt)
• Formula generally starts with H (e.g., HCl, H2SO4)
• Acids contain a H+1 cation and an anion in aqueous solution.
• Binary acids have a H+1 cation and a nonmetal anion.
• Oxyacids have a H+1 cation and a polyatomic anion.

Naming Oxyacids

• If the polyatomic ion name ends in -ate, then change the ending to -ic suffix.
• Example: Sulfate becomes sulfuric acid.
• If the polyatomic ion name ends in -ite, then change the ending to -ous suffix.
• Example: Sulfite becomes sulfurous acid.
• Write the word acid at the end of all names.

Writing Formulas for Acids

• Formulas start with H when the name ends in acid
• Write formulas as if ionic, even though it is molecular
• A hydro prefix means it is a binary acid; no prefix means it is an oxyacid.
• If the ending for the acid is -ic, the polyatomic ion ends in -ate; If ending is -ous, polyatomic ion ends in -ite.
• Example (binary acid): Hydrosulfuric acid (H2S)
• Example (oxyacid): Carbonic acid (H2CO3)

Acids You Should Know

• HF (hydrofluoric acid)
• HCl (hydrochloric acid)
• HBr (hydrobromic acid)
• HI (hydroiodic acid)
• HNO3 (nitric acid)
• H2SO4 (sulfuric acid)
• H3PO4 (phosphoric acid)
• H2CO3 (carbonic acid)

Lewis Structures and VSEPR Theory

• Octet Rule: ALL ATOMS (except for H) must have 8 electrons around them. Hydrogen has only 2 electrons around atoms.
• To make a Lewis dot structure:
• Determine the number of electrons.Arrange atoms in a molecule.
• Attach atoms with bonds (1 bond equals 2 electrons).
• Distribute non-bonding pairs.
• Use multiple bonds when necessary.
• Usually, the first atom of the formula is the central atom, but not always.

Writing Lewis Structures

• 1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.
• If it is an anion, add one electron for each negative charge.
• If it is a cation, subtract one electron for each positive charge.
• 2. The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.
• 3. Fill the octets of the outer atoms.
• 4. Fill the octet of the central atom.
• 5. If valence electrons runs out before the central atom has an octet, form multiple bonds until fulfilled.

Additional Tips & Further Considerations for Lewis Structures

• Elements from the 2nd period (C, N, O, and F) observe the octet rule
• Boron often has six electrons surrounding it
• Beryllium has 4 electrons
• 3rd period elements (and beyond) CAN exceed the octet rule using empty valence d orbitals.
• When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

Resonance

• When more than one valid Lewis structure can be written for a molecule, the actual structure is an average of all of them.

VSEPR (Valence Shell Electron Pair Repulsion) Model

• The structure of a molecule is determined by minimizing electron pair repulsions.

Predicting Molecular Geometry

• Draw the Lewis Structure
• Determine how many electron pairs are around the central atom.
• Count a multiple bond as one group

Polarity

• A bond is polar if there is a large difference in electronegativity between the atoms (it has a dipole moment).
• The greater the difference in electronegativity, the greater the polarity of the bond.
• Ionic bond: +A -B
• Polar covalent bond: δ+A δ-B
• Nonpolar covalent bond: A B
• The shape of the molecule must be considered to determine if the molecule as a whole is polar.
• Determining Molecular Polarity:
• O3: polar
• BF3: nonpolar
• CO2: nonpolar

Chapter 4: Chemical Equations & Calculations

• Molecular mass is the mass of one mole = atomic mass in grams.
• 1 mole = molar mass (g) = 6.02 x 10^23 particles

Examples

• What is the mass in grams of 2.25 mol of the element iron, Fe? (Answer: 126 g Fe)
• How many moles of calcium are in 5.00 g Ca? (Answer: 0.125 g Ca)
• What is the mass in grams of 7.5 x 10^15 atoms of nickel, Ni? (Answer is 7.3 x 10^-7 g Ni)
• Find the molar mass and the number of moles of 15.0 g of calcium hydroxide, Ca(OH)2. (Answers: 74.10 g/mol, 0.202 mol, 1.22 x 10^23 molecules)
• How many moles are there in 3.00 kg (NH4)2CO3 (Answers: 96.11 g/mol (NH4)2CO3, 31.2 mol (NH4)2CO3)
• How many carbon atoms are there in a diamond (pure carbon) with a mass of 52 mg? (Answer: 2.6 x 10^21 carbon atoms)

Chemical Reactions

• Visual evidence of a reaction: Release of a gas - CO2 is released when acid is placed in a solution containing CO32− ions Formation of a solid (precipitate) - A solution containing Ag+ ions mixed with a solution containing Cl− ions Heat is produced or absorbed - Acid and base are mixed together Color changes

Writing Chemical Equations

• Chemical equations utilize patterns that help understand and predict reactions.
• Combination: The joining of two or more elements or compounds, producing a product of different composition (A + B → AB). Examples: 2Na(s) + Cl2(g) → 2NaCl(s), MgO(s) + CO2(g) → MgCO3(s)
• Decomposition: Produces two or more products from a single reactant; reverse of a combination reaction. Examples: 2HgO(s) → 2Hg(l) + O2(g), CaCO3(s) → CaO(s) + CO2(g)
• Single-replacement: One atom replaces another in the compound, producing a new compound (A + BC → B + AC). Examples: Cu(s) + 2AgNO3(aq) → 2Ag(s) + Cu(NO3)2(aq), 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
• Double-replacement: Two compounds undergo a "change of partners". Two compounds react by exchanging atoms to produce two new compounds (AB + CD → AD + CB).

Precipitation Reactions

• Reactions between aqueous solutions of ionic compounds that produce an ionic compound that is insoluble in water are called precipitation reactions and the insoluble product is called a precipitate.

Solubility of Ionic Compounds

• Some dissolves very well in water at room temperature, other dissolves hardly at all in water at room temperature, which results in compounds to be soluble, while those that do not are said to be insoluble
• The degree of solubility depends on the temperature and even insoluble compounds dissolve, just not enough to be meaningful
• Molecular equations describe the chemicals put into the water and the product molecules.
• Complete ionic equations describe the actual dissolved species; aqueous strong electrolytes are written as ions.
• Solubility rules determine how to write equations for these ions
• Gas Evolving Reactions
• Some reactions form a gas directly from the ion exchange
• K2S(aq) + H2SO4(aq) → K2SO4(aq) + H2S(g)
• Other reactions form a gas by the decomposition of one of the ion exchange products into a gas and water
• K2SO3(aq) + H2SO4(aq) → K2SO4(aq) + H2SO3(aq) H2SO3 → H2O(l) + SO2(g)

Acid Base Reactions

• These reactions involve the transfer of hydrogen ion (H+) from one reactant (acid) to another base.
• H+ on HCl was transferred to the oxygen in OH-, giving H2O
• Writing chemical equations
• A chemical equation is the symbolic representation of a chemical reaction in terms of chemical formulas.
• reactants are starting substances in a chemical reaction and the arrow yields. The formulas on the right side of the arrow represent the products.
• the total number of atoms of each element on both sides of a chemical equation must match to be balanced

Stoichiometry Problems

• Steps: Balanced equation, convert to moles, find molar mass, write as ratio and convert to desired unit