Lecture Notes Chemistry PDF

# Lecture 1 - Recognize and identify the following: atom, element, compound and molecule - **Element:** Fundamental substance that can't be broken down any smaller. - **Atom:** Smallest unit of an element that retains properties and chemical behavior. - **Diatomic Elements:** H<sub>2</...

# Lecture 1 - Recognize and identify the following: atom, element, compound and molecule - **Element:** Fundamental substance that can't be broken down any smaller. - **Atom:** Smallest unit of an element that retains properties and chemical behavior. - **Diatomic Elements:** H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂ - **Top 3 Elements:** Oxygen, Silicon, & Aluminum. - **Compounds:** Substance with two or more elements - **Covalent:** Two Non-metals - **Ionic:** Metal + Non-metal - **Cation** loses e- - **Anion** gains e- - Recognize and identify isotopes of the same element - $A/Z$ = Element - A = mass - Z = atomic # - Identify by name and symbol different elements of the periodic table. - 1-36: Sr, Pb, Ag, Cd, Sn, I, Xe, Ba, Pt, Au, Hg, and Pd. - Identify common oxidation charge/ - 1-20 - 31-36 - Identify metals, non-metals, and metalloids. - Metals - Metalloids - Non-metals # Lecture 2 - Distinguish between the bonds of molecular and ionic compounds. - Identify molecular and ionic compounds based on formula/name - Cations: Na+ metals - Anions: Cl- Non-metal - **Compound with metal:** Ionic - Polyatomic ions (we should know!! ) - Nitrate: NO₃^- - Nitrite: NO₂^- - Carbonate: CO₃^2- - Bicarbonate: HCO₃^- - Autate: CH₃COO^- - Ammonium: NH₄⁺ - Perchlorate: ClO₄^- - Sulfate: SO₄^2- - Sulfite: SO₃^2- - Phosphate: PO₄^3- - Hydroxide: HO^- or OH^- - Write chemical formula: Ionic, binary molecular, and hydrated (from the name) - **Monoatomic Ions:** Sodium chloride - **Polyatomic Ions:** Potassium Autate - recognize the name. - **Transition Metals:** Iron (II) chloride - Cation - Anion - Charge - **Molecular:** Sulfur Dioxide. - S - O₂ - In order: Mono, Hepta, Octa, Nona... Tri, Tetra, Penta... - **Hydrates:** Copper (II) sulfate pentahydrate. - CuSO₄ * 5H₂O # Lecture 3 - Define the terms: mass, mole, molarity, atomic mass, molar mass, density, ppm, ppb. - **Covalent bonds:** Molecular mass. - **Ionic:** Formula mass. - **Molarity:** *moles of solute* / *vol (L) of solution* = *mol/L* - **Mass of solute:** *mol solute* = *M x L solution* - *mass solute* = *molar mass* x *mol solute* - **Density:** Ratio of *mass of sample* / *volume* = *g/mL (cm³) or g/L* - **ppm:** *(mass solute)/(mass solution)* x 10⁶ # Lecture 4 - Apply the rules of stoichiometry to balance chemical equations. - Determine the initial number of moles of reactant and the final # or r and p. - Identify a limiting reactant and calculate the yield. - Calculate the percent yield or use percent yield to predict the product form a reaction. - *grams of known* x *mol/g* = *mol* - *mol* of *known* using mole ratio = *moles of desired* - *moles of desired* x 1/*(molar mass)* = *mol. - **Limiting Reactant**: Least amount of product (excess = other reactant) - **Theoretical Yield:** Maximum possible yield from a reaction. - **Actual Yield:** Yield from experiment. - **Percent Yield:** *(actual/theoretical)* x 100 - **Example:** 2H₂ + O₂ → 2H₂O - 3 mol 2 mol - H₂O/ H₂ = 2/2 = x = 3 mol H₂O (limiting!) - H₂O/O₂ = 2/1 = x = 4 mol H₂O (theoretical yield max!!) # Lecture 5 - n = 3 *mL* = 0 *p* = ____ *a* = _____ - l = 1 *ms* = -1/2 *l* = 1 *d* = 5 - l = 2 *ms* = +1/2 *l* = 2 *f* = 7 - Circle an electron with these four quantum numbers. - *n* = size 1-shape *ml* = orientation *ms*= spin - Explain nodes, how many nodes in an atom? - **Nodes:** A point where the electron possibility is 0. - Shapes of the orbitals: - *s* - *p_x* - *p_y* - *p_z* - **Penetration:** The ability of an electron in a given sub shell to penetrate within shells or sub shells to get close to the nucleus. - Degenerate Orbitals ~same energy level - **Orbitals:** Regions of space with a high probability of finding an electron. # Lecture 6 - Anomalies: - **Cr:** 4s¹3d⁵ - **Cu:** 4s¹ 3d¹⁰ - Electron configurations - Ions: - Remove valence electrons (two have to be removed) - Transition: Pull out *p* first. - Shielding and effective nuclear charge. - **Effective:** total amount of attraction an e- feels to nucleus: z-s = (2 for O) - Ground State = Normal elctron configuration - Exited State = Not the normal electron configuration - more electron displaced. - **Shielding:** The repulsion from electrons closer to the nucleus reduce the outer electrons attraction to it. - Valence electrons - outermost shell. - **Example:** O: 1s² 2s² 2p⁴ vs 1s² 2s² 2p³ 3s¹(excited) - Ga : [Ar] 4s² 3d¹⁰4p¹ - [Ar] is shielding Se by 18. # Lecture 7 - **Electron Affinity:** Opposite of ionization. The energy of the process of adding an electron to a gaseous atom to form an anion. - **Trends in size:** - **Cations:** - Na - Bigger Z_eff - Smaller radius - From compression - Na⁺ - Smaller Z_eff - Smaller radius - From compression - **Anions:** - F^- - Larger Z_eff - Smaller radius - From compression - O^2- - Larger Z_eff - Larger radius - From compression - **Trends in Ionization Energy**: Larger Z_eff = smaller ionization energy. - **Given data and match to elements** --> **Noble gas configuration** - Easier to remove from a p orbital - Easier to remove paired e- - Harder to remove from a noble gas electron configuration. # Lecture 8 - **Electronegativity:** - More electronegative atom will always have the delta negative. - H-F. Dipole Present - H-H - F-F. Covalent - **Lattice Energy:** Predict which one has larger or smaller lattice energy. - **Charge:** The larger the charge, the greater the lattice energy. - **Size:** The closer the ions are, the greater the lattice energy. - **Lattice Energy:** Energy required to break apart an ionic bond. - Larger / Lattice = Stronger ionic bond. # 3D Structures - Non-polar - Non-polar - Non-polar - **Polar:** A molecular structure that isn’t balanced with degrees is polar. - Any other arrangement would introduce more than 90 degrees of interaction with lone pairs of electrons. - Double bonds have the most repulsion - Keep away from each other - Keep away from lone pairs # Hybridization - **Sigma bonds:** Head to head overlap. - Strong - Single bond - **Pi bonds:** Parallel overlap - Weak - Double and triple bonds - **Reality:** All C-H bonding is equal. - **Theory:** All C-H orbital-orbital overlap must be identical. The valence orbitals, 2s and 2p must mix (hybridize) to become equivalent. - **Hybridization:** Mixing atomic orbitals to form new hybrid orbitals suitable for pairing of electrons to form chemical bonds in valence bond theory. # Inter Molecular Forces - **Attraction between molecules:** - Inter Molecular Forces (IMF) of liquids is higher than gases. - This means liquids will have a higher boiling point. - **Dispersion Force:** - All molecules have it. - Big molecule = More dispersion force. - **Temporary Dipoles:** Occur because of the movement of electrons within molecules. This creates a temporary dipole which can induce a dipole in neighboring molecules. - **Larger atoms:** More polarizable. - **Polarizability:** How easily a molecule's electron cloud can be distorted by an external electric field. - **Larger the molecule:** The larger the magnitude of the dipoles. - **More electronegativity:** Not relevant. - **Symmetrical:** More points of contact = bigger IMF. - Strongest IMF. - **Ion-Dipole attraction:** - Ion to dipole. - **Dipole-Dipole Forces:** - Attraction between permanent dipoles (electronegative). - Stronger than dispersion forces. - **Higher boiling point:** The stronger the force, the higher the temperature required to overcome it, leading to a higher boiling point. - **Hydrogen bonds:** A subcategory of dipole-dipole where a hydrogen atom is bonded to a highly electronegative atom (usually oxygen, nitrogen, or fluorine). - **Solubility:** "Like dissolves like" - **Miscible Liquids:** Dissolve in each other. - **Hydrophilic:** Polar solvents - **Hydrophobic:** Non-polar solvents - **Immiscible:** Doesn’t dissolve. - Long and non-polar - water (small and highly polar) # Enthalpy of Solution - **Enthalpy of solution** is negative when heat is released and positive when heat is absorbed. - **ΔH_solution= ΔH_solute + ΔH_solvent + ΔH_mix** - ΔH_mix is usually very small - ΔH_hydration can be negative or positive in the case of water - **ΔH_hydration:** Always negative which means heat is **released** when ions are surrounded by water molecules. - **ΔH_lattice:** The energy required to separate one mole of an ionic compound into its gaseous ions. - **Reactions:** - **1) NaCl(s) → ΔH_lattice Na⁺(g) + Cl^-(g)** (endothermic) - **2) Na⁺(g) + Cl^-(g) --> H₂O Na⁺(aq) + Cl^-(aq)** (exothermic) - **Lewis Structures**