GENERAL-CHEMISTRY

Chemical Compounds and Ions in Chemistry

Page 3: Chemistry

• Science that deals with the properties of matter

• Changes matter undergoes

• Natural laws that describe these changes

Page 5: Units of Measurement

• Mass: Kg (fundamental unit)

• Weight: N (derived unit) = Kg × m s2

• Volume: mL (derived unit) = cm3

• Energy: Joules (derived unit) = Nm = (Kg × m s2)(m)

Page 7: Classification of Matter

• Phase/State: Solid, Liquid, Gas

• Solid: Holds shape, fixed volume

• Liquid: Shape of container, fixed volume

• Gas: Shape of container, volume of container

Page 8: Phase Change

• Melting: Fusion, liquefaction, thawing

• Condensation: Rain

• Sublimation: Naphthalene balls, Iodine crystals

• Deposition: Dry ice / Cardice

Page 10: Mesophases

• Types: Smectic, Nematic, Cholesteric

• Critical Points: Critical Pressure, Critical Temperature

Page 11: Phase Equilibrium

• Melting, Freezing, Vaporization, Condensation, Sublimation, Deposition

• Variable: Temperature and Pressure

Page 12: Classification of Substances

• Pure substances: Elements, Compounds

• Mixtures: Heterogeneous, Homogeneous

Page 15: Separatory Techniques

• Filtration, Evaporation, Distillation, Sublimation, Crystallization, Separatory funnel, Chromatography, Magnetic separation

Page 16: Properties of Matter

• Thermodynamic properties: Intensive/Intrinsic, Extensive/Extrinsic

• Physical Properties: Additive, Constitutive, Colligative

Page 18: Example Problems

• Determining molecular weight, concentration, and identity of solutions

Page 23: Fundamental Chemistry Laws

• Conservation of Mass

• Law of Definite/Constant Proportions

• Law of Multiple Proportions

• Law of Reciprocal Proportion

• Gas Laws: Boyle’s, Gay-Lussac’s, Charles’, Avogadro’s

Page 24: Law on Mass Conservation

• Total mass of products = Total mass of reactants

• Mass cannot be created or destroyed

• Basis of Stoichiometry

Page 27: Balancing Chemical Reactions

• Adjust coefficients, not subscripts

• Balance atoms that occur once on each side

• Balance polyatomic ions as a whole

• Balance pure elements last

• Rewrite H2O as H(OH) if condensation is observed

Chemical Compounds and Ions in Chemistry

Page 28:

• Example Problems:

  • Al + S8 → Al2S3

  • Al2(SO4)3 + Ca(OH)2 → Al(OH)2 + CaSO4

  • C3H8 + O2 → CO2 + H2O

  • H3PO4 + NaOH → Na3PO4 + H2O

Page 29:

• Proust’s Law:

  • A chemical compound always contains exactly the same proportion of elements by mass.

Page 30:

• Law of Multiple Proportions:

  • When chemical elements combine, they do so in a ratio of small whole numbers.

Page 31:

• Law of Reciprocal Proportions:

  • Law of combining weights:

    • Elements combine in the ratio of their combining weights or chemical equivalents.

    • Or in some simple multiple or sub-multiple of that ratio.

  • Also called the Law of Equivalents

Page 32:

• Example Problems:

  • Determine the ratio of hydrogen to carbon in methane and oxygen to carbon in carbon dioxide.

  • Prove that the law of reciprocal proportion holds in water.

Page 33:

• Example Problems:

  • Prove the law of reciprocal proportion:

    • P PH3 PCl3 H CI HCI Sr.

    • Compounds Combining Combining No. elements weights

    • I PH3 P H 31 3 2

    • PCl3 P Cl 31 106.5

Page 34:

• BGCA = Ideal Gas Law:

  • Pressure:

    • 1atm = 760mmHg = 760torr = 101.3kPa

  • Temperature:

    • 9℃ = 5℉ − 160

    • K = ℃ + 273.15

  • Volume

  • Moles

  • R (gas constant)

• Standard Temperature and Pressure (STP):

  • T = 273.15K

  • P = 1atm

  • V = 22.2L

Page 35:

• Example Problems:

  • A gas occupies a volume of 2.5 liters at a temperature of 26.85°C. If the temperature is increased to 260.33°F while keeping the pressure constant, what will be the new volume of the gas, in daL?

  • Suppose you have a gas confined in a syringe at an initial pressure of 1520 mmHg and an initial volume of 50 mL. If you decrease the volume to 0.025 L while keeping the temperature constant, what will be the final pressure, in atm, of the gas?

Page 36:

• Example Problems:

  • A gas occupies a volume of 2.0 liters with 3 moles of molecules. If additional molecules are added, and the number of moles increases to 10 moles while keeping the temperature and pressure constant, what will be the new volume of the gas, in cL?

  • A gas is initially at a pressure of 1.7 atmospheres and a temperature of 80.33°F. If the temperature is increased to 500 Kelvin while keeping the volume constant, what will be the new pressure, in torr, of the gas?

Page 37:

• Example Problems:

  • A gas sample has an initial pressure of 1.0 atmospheres, an initial volume of 1.6 liters, and an initial temperature of 50 Celsius. If the volume is increased to 7111 mL, and the temperature is raised to 212 Fahrenheit, what will be the final pressure of the gas?

Page 38:

• Ideal Gas vs Real Gas:

  • Real gases do not behave well.

  • Ideal State Requirements:

    • Low Pressure → repulsion

    • High Temperature → attraction

    • Large Volume → negligible volume

  • Real Gas:

    • Van der Waals Equation

Page 39:

• Example Problems:

  • A 0.7mol sample of ammonia (NH3) gas occupies a volume of 2.5L at a temperature of 300K. The Van der Waals constant for ammonia are a = 4.0L2atm/mol2 and b = 0.04L/mol. Calculate the pressure exerted by the gas in the container in torr.

Page 40:

• Fundamental Laws:

  • Gas interactions (in mixtures):

    • Dalton’s law of Partial Pressures (gas in gas)

    • Raoult’s law (vapor pressure of solvent)

    • Henry’s law on solubility (gas in liquid)

  • Movement:

    • Graham’s law (molecular weight)

    • Fick’s 1st law

Page 41:

• DALTON’S LAW PARTIAL PRESSURE:

  • Total pressure in a mixture is equal to the sum of the partial pressures of each gas.

  • PT = PN2 + PO2 + PCO2 … + PX

  • PX PT = nx nT (this is called χ)

  • PX = PT χ

Page 42:

• Example Problem:

  • Determine the mole fraction of sucrose aqueous solution with the following reading:

  • A mixture of gases contains empyreal air and mephitic air. The partial pressure of the former is 0.4atm and the latter is 0.5atm. Calculate the total pressure of the gas mixture.

Page 43:

• Example Problems:

  • A gas mixture contains CO2, CH4, and N2. The partial pressure of CO2 and CH4 are 1.5 atm and 2.0 atm, respectively. If the total pressure of the gas mixture is 4.5 atm, find the partial pressure of N2.

Page 44:

• Example Problems:

  • If the total pressure of a canister of gas is 800 torr, determine the pressure, in atm, imparted by 0.2 mole of carbonic acid gas if the total amount of gas inside is 0.6 mol.

  • The following is the label content of a canister of gas. Determine the partial pressure imparted by helium if the barometer reads 1.7atm. Determine the total volume of the canister if the thermometer reads 95°F.

    • Content Amount (mol)

      • N2 0.1

      • He 0.2

      • CO2 0.3

Page 45:

• RAOULT’S LAW ON VAPOR PRESSURE:

  • Vapor pressure of a solvent above a solution is equal to the vapor pressure of the pure solvent at the same temperature scaled by the mole fraction of the solvent present:

    • Psolution = (χsolvent)(Psolvent 0)

    • ∆𝑃 = (χsolute)(Psolvent 0)

Page 46:

• HENRY'S LAW ON GAS SOLUBILITY:

  • Increasing the vessel pressure will increase gas solubility.

  • P1 P2 T3 T2 T1

  • Solubility of O2 in water

  • Partial Pressure A B of O2

Page 47:

• Solid VS Gas solubility (TEMP):

  • methane KNO 2.0

  • oxygen carbon 100 monoxide 1.0 nitrogen NaCI

  • helium

  • Temperature (°C)

Page 48:

• Solid/Liquid VS Gas solubility (P):

  • Solubility

  • Gas

  • Solid or Liquids

  • Pressure

Page 49:

• Example Problems:

  • The Henry’s constant for oxygen in water at a certain temperature is 1.2 x 10-3 mol/Latm. If the partial pressure of O2 in air is 0.25atm, calculate the concentration of O2 in the water. Determine the amount of oxygen, in g, in 1L of that solution.

Chemical Compounds and Ions in Chemistry

Page 50:

• Example Problems:

  • Determine the Henry's constant for He.

  • Calculate how much N2 will escape out of a 2.5L solution if the pressure is reduced from 0.8atm to 0.5atm.

  • Determine whether the solution is unsaturated, saturated, or supersaturated: 0.8mM O2 under 0.5atm partial pressure.

Page 51:

• GRAHAM'S LAW:

  • Rate of diffusion and speed gas are inversely proportional to the square root of their density.

• Example Problem:

  • HCl (36.46g/mol) + NH3 (17.03g/mol) → NH4Cl (white ppt.)

Page 52:

• Related Terms:

  • Diffusion = the gradual mixing of molecules of one gas with the molecules of another gas by virtue of their kinetic properties.

  • Effusion = passage of a gas under pressure though a small opening.

Page 54:

• Example Problems:

  • Calculate the ratio of the molar masses of helium to methane if the rate of diffusion of helium is 3 times faster than that of methane.

  • Determine the molar mass of an unknown gas if the ratio of the rate of diffusion of the unknown gas is 4.5 times faster than that of carbon dioxide.

Page 55:

• FICK'S FIRST LAW (FLUX, J):

  • Movement of particles (diffusion flux) is proportional to the concentration gradient (from high concentration to low concentration).

• Equation: J = -D(d𝜑/dx)

• J = Flux, D = Diffusivity, 𝜑 = Concentration gradient, x = Path length

Page 56:

• Example Problems:

  • Calculate the diffusion flux of ions across a glass membrane with different concentrations on each side.

  • Determine the diffusivity of a solid material based on the rate of gas diffusion and concentration gradient.

Page 58:

• Atomic Structure:

  • Democritus proposed the concept of "atomos" and indivisibility.

  • John Dalton introduced the billiard ball model and the concept of multiple proportions.

  • JJ Thomson proposed the plum pudding model and conducted the cathode ray experiment to discover electrons and protons.

  • Ernest Rutherford developed the nuclear model through the gold film experiment and discovered the nucleus (including neutrons).

  • Niels Bohr proposed the planetary model and introduced the concept of electron configuration and orbits.

  • Erwin Schrodinger developed the quantum model and introduced the concept of orbitals (s, p, d, f).

Page 61:

• Subatomic Particles:

  • Proton (p+): +1 charge, 1 mass, discovered by E. Rutherford.

  • Electron (e-): -1 charge, 0 mass, discovered by JJ Thomson and RA Millikan.

  • Neutron (n0): 0 charge, 1 mass, discovered by J. Chadwick through the Millikan Oil Drop Experiment.

Page 62:

• Atomic Mass Units:

  • Weighted average mass of naturally occurring isotopes of an atom.

  • Mass number = #p+ + #n0

Page 63:

• Example Problems:

  • Calculate the atomic mass units of carbon atoms given the abundance of C-12 and C-13 isotopes.

  • Determine the atomic mass units of chlorine atoms given the abundance of Cl-35 and Cl-37 isotopes.

  • Find the relative abundance of Li-6 based on the abundance of Li-7 and the average atomic mass of naturally occurring Lithium.

Page 64:

• Nuclide Writing:

  • Isotope Symbols: Mass number, Charge, Element Symbol, Atomic number

  • Mass Number (A) = #p+ + #n0

  • Charge = #p+ - #Electrons

Page 68:

• Electron Configuration:

  • s-block, p-block, d-block, f-block

  • Example: Li (3), S (16), Ar (18)

Page 69:

• Quantum Numbers:

  • Principal (n), Azimuthal/Angular (l), Magnetic (ml), Spin (ms)

Page 70:

• Electron Configuration:

  • Example: Li (3), S (16), Ar (18)

Chemical Compounds and Ions in Chemistry

Page 71: Rules in Electron Configurations

• Aufbau Principle

  • Lower energy levels are filled up first

• Hund’s Rule

  • Orbitals are filled up singly before pairing up

• Pauli’s Exclusion Principle

  • No two electrons can have the same set of quantum numbers

  • Orbitals can only occupy 2 electrons because ms should have only 2 values (+1/2 or -1/2)

Page 72: Diamagnetic and Paramagnetic

• Diamagnetic

  • No unpaired electrons

  • Very weakly repelled by magnets

  • Field bends slightly away from the material

• Paramagnetic

  • At least one unpaired electron

  • Attracted to magnets

  • Field bends slightly toward the material

Page 73: Periodic Table

Page 74: History of the Periodic Table

• Jons Jakob Berzelius

  • Element symbols

• Johann Dobereiner

  • Law of triads

  • Middle element is average of the 1st and 3rd

• John Alexander Newlands

  • Law of octaves (periods)

  • Pattern reoccurs every 8th element

• Dmitri Mendeleev

  • Father of Modern Periodic Table

  • Atomic Mass/Weight

  • Periodicity (the Periodic Law)

  • Vertical arrangement (eka)

• Lothar Meyer

  • Atomic Mass/Weight

  • Arranged by valency

• Moseley

  • Atomic number

  • First modern periodic table

Page 75: History of Chemical Compounds and Ions

• Glenn Seaborg

  • Discovered transuranic elements

• Bismuth

  • Heaviest stable atom

• Latest addition: Og (oganesson)

Page 76: Practice Problems

• Predict the atomic mass of Na using the known mass of Li and K.

Page 77: Groups and Periods in the Periodic Table

Page 78: Valence Electrons and Valency (Ox. State)

• Valency Table

Page 79: Group Classification in the Periodic Table

Page 80: Trends in the Periodic Table

• Atomic size

• Ionization energy

• Electron affinity

• Electronegativity

• Metallic property

• Non-metallic property

• Metalloids

Page 81: Example Problems

• Classify the following by the concept each are representing: 𝛿−, 𝛿+, 0, E, Na, Cl

Chemical Compounds and Ions in Chemistry

Atomic Size (Page 82)

• Atomic radius decreases from H to Rn

• Supporting details:

  • H, He, Li, Be, B, C, N, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Ga, Ge, As, Se, Br, Kr, Rb, Sr, In, Sn, Sb, Te, I, Xe, Cs, Ba, TI, Pb, Bi, Po, At, Rn

Ionic Size (Page 83)

• Group 1A ions are larger than Group 2A ions

• Group 6A ions are larger than Group 7A ions

• Supporting details:

  • Group 1A: Li+, Na+, K+, Rb+, Cs+

  • Group 2A: Be2+, Mg2+, Ca2+, Sr2+, Ba2+

  • Group 3A: B3+, Al3+, Ga3+, In3+

  • Group 6A: O2-, S2-, Se2-, Te2-

  • Group 7A: F-, Cl-, Br-, I-

Ionization Energy (Page 84)

• Ionization energy increases from H to Lr

• Supporting details:

  • H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, As, Se, Br, Kr, Rb, Sr, Y, Zr, Nb, Mo, Tc, Ru, Rh, Ag, Cd, In, Sn, Sb, Te, I, Xe, Cs, Ba, La, Hf, Ta, W, Re, Os, Ir, Pt, Au, Hg, Tl, Pb, Bi, Po, At, Rn, Fr, Ra, Rf, Db, Sg, Bh, Hs, Mt, Ds, Rg, Cn, Uut, Uuq, Uup, Uuh, Uus, Uuo

Electron Affinity (Page 85)

• Electron affinity increases from H to Xe

• Supporting details:

  • H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Ga, Ge, As, Se, Br, Kr, Rb, Sr, In, Sn, Sb, Te, Xe

Electronegativity (Page 86)

• Electronegativity increases from H to Rn

• Supporting details:

  • H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, As, Se, Br, Kr, Rb, Sr, Y, Zr, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Cd, In, Sn, Sb, Te, I, Xe, Cs, Ba, La, Hf, Ta, W, Re, Os, Ir, Pt, Au, Hg, Tl, Pb, Bi, Po, At, Ra

Metals, Metalloids, and Non-metals (Page 87)

• Metals, nonmetals, and metalloids are listed

• Supporting details:

  • Metals: H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, As, Se, Br, Kr, Rb, Sr, Y, Zr, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Cd, In, Sn, Sb, Te, I, Xe, Cs, Ba, La, Hf, Ta, W, Re, Os, Ir, Pt, Au, Hg, Tl, Pb, Bi, Po, At, Rn, Fr, Ra, Ac, Rf, Db, Sg, Bh, Hs, Mt, Ds, Rg, Uub - Uuq

  • Metalloids: B, Si, Ge, As, Sb, Te

  • Nonmetals: Ce, Pr, Nd, Pm, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, Lu, Th, Pa, U, Np, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No, Lr

Chemical Reactions: Electronic Stability (Page 88)

• Octet Rule, Duet Rule, and Expanded Octet are mentioned

• Supporting details:

  • Octet Rule: atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with 8 valence electrons (except for H and He)

  • Duet Rule: atoms tend to gain, lose, or share electrons to achieve a stable electron configuration with 2 valence electrons (only for H and He)

  • Expanded Octet: atoms in period 3 and above can have more than 8 valence electrons by utilizing empty p-orbitals

Radioactivity: Nuclear Instability (Page 89)

• Nuclear instability is mentioned

• Supporting details:

  • Radioactive elements: B, Y, Paper, Aluminium, Lead

Radioactivity: Nuclear Stability (Page 90)

• Nuclear stability and radiation are mentioned

• Supporting details:

  • Radiation particles: Alpha (helium nucleus) and Beta (high-energy electron)

  • Penetration: Alpha particles have low penetration (stopped by paper or aluminium), while Beta particles have medium penetration (stopped by lead)

Chemical Compounds and Ions in Chemistry

Page 91:

• Type of decay represented by the following:

  • Co-60 → Co-60 + γ

  • Po-218 → Pb-214 + α

  • Rn-222 → Po-218 + α

  • U-234 → Th-230 + α

  • K-40 → Ar-40 + β

Page 92:

• Molecular Bonding

Page 93:

• Intramolecular Forces of Attraction

  • Chemical bonds

    • Covalent bond

      • NM-NM

      • Polar covalent

      • Non-Polar covalent

    • Ionic bond

      • M-NM

      • Cation-Anion

Page 94:

• Electronegativity Values

  • Determine the type of bond present in the following compounds:

    • NaCl

    • CS2

    • LiBr

    • PH3

    • H3C-Na

    • O-C

Page 95:

• Bonding Theories

  • Valence or Lewis bond Theory

    • Unpaired electrons of atoms will pair up to complete their octet

    • Atomic orbitals of reactants will overlap forming molecular orbitals

    • Sigma bond = single bond

    • Pi bond = double bond

  • Molecular Bonds

Page 96:

• Valence Bond Theory

  • Steps to determine the geometry of a molecule

  • Example: SO2

Page 97:

• Practice Problems

  • Draw the Lewis structures of:

    • NO-

    • N2

  • Draw Lewis structures of:

    • Hypochlorite ion, OCI-

    • Ethane, C2H6

Page 98:

• Charge

  • Partial charges

    • Neutral covalent

    • Polar covalent

    • Ionic (Formal charges)

  • Formal charges

Page 99:

• Practice Problems

  • H-C=N=N-H

  • H3C-Ö-N=O

Page 100:

• Answers to Practice Problems

  • H-C=N=N-H

  • H3C-Ö-N=O

Page 101:

• Number of Electron Pairs and Molecular Geometry

Page 102:

• Practice Problem

  • Determine the geometry of the following bimolecular compounds/molecules:

    • CO2

    • BH3

    • SnCl2

    • CH4

    • NH3

    • H2O

    • PCl5

    • SF6

Page 103:

• Bonding Theories

  • Molecular Orbital Theory

    • Bonding electrons are shared across the entire molecule

    • Possibility of Antibonding molecular orbitals

      • Sigma star (δ*)

      • Pi star (π*)

  • Molecular Bonds

Page 104:

• Bonding Theories

  • Molecular Orbital Theory

    • Conservation of orbitals

    • Predict diamagnetism vs paramagnetism

    • Predict presence of double/triple bonds

    • Bond Order = Bonding electrons - Antibonding electrons

  • Molecular Bonds

Page 105:

• Molecular Orbital Diagrams

  • MO diagrams for Nitrogen (N) and Oxygen (O)

Page 106:

• Practice Problems

  • Construct the molecular orbital diagrams of the following and determine the bond type between atoms and its magnetic property:

    • H2

    • N2

    • F2

    • NO

Page 107:

• Intermolecular Forces of Attraction

  • Forces between molecules or compounds

  • Influenced by charge interaction and polarizability

  • Van der Waals forces

    • Dispersion

    • H-bonding

    • Keesom

    • Debye

    • London

Page 108:

• Comparison of Intermolecular and Intramolecular Forces

• Weak, moderate, strong, and very strong forces

• Types of forces: dispersion, H-bonding, ion-ion, dipole-dipole, covalent bonds, ion-dipole

Page 109:

• Practice Problem

  • Classify the following animations in terms of the intermolecular force they represent:

    • He

    • He

Chemical Compounds and Ions in Chemistry

Page 110:

• Practice Problem: Classify the following animations in terms of the intermolecular force they represent.

Page 111:

• Practice Problem: Classify the following animations in terms of the intermolecular force they represent.

Page 112:

• Chemical Formulas: MOLECULAR BONDING

  • Formula Type: Kekule/Lewis

    • Description: All atoms, Bonds, Lone electrons

  • Formula Type: Structural

    • Description: All atoms, Bonds

  • Formula Type: Skeletal

    • Description: Heteroatoms, Bonds

  • Formula Type: Condensed

    • Description: All atoms Bonds(double, triple)

• ORGANIC COMPOUNDS

Page 113:

• Practice Problems: Convert the following structure to other formula type mentioned in the previous slide.

  • Structure: CH2=CHCH2OH

Page 114:

• Chemical Formulas: MOLECULAR BONDING

  • Formula Type: Molecular

    • Description: Summary of atoms present

    • Compound Type: Covalent

    • Method: Valence bond theory

  • Formula Type: Empirical

    • Description: Subscripts are reduced

    • Compound Type: Ionic

    • Method: Criss-cross

• INORGANIC COMPOUNDS

  • Example: MgO

    • Mg + O → MgO

Page 115:

• Practice Problems:

  • A compound is composed of 52.14% C, 13.13% H, and 34.73% O by mass.

    • What is the empirical formula?

    • What is the molecular formula if the molar mass of the compound was determined to be 138.204g/mol?

  • A compound consists of 20.32g of C, 5.12g of H, and 7.9g of N.

    • What is the empirical formula?

    • What is the molecular formula if the molar mass of the compound is 236.448g/mol?

Page 116:

• Practice Problems: Combustion Analysis

  • A 3.480g sample contains only C and H. In a combustion reaction, it produced 10.63g of CO2 and 5.22g of H2O. Determine the empirical formula of the compound.

Page 117:

• Chemical Nomenclature

Page 118:

• LOOK ON BACK FOR Naming

• Naming Compounds Flowchart

• Examples of each!

• Is there a METAL in the compound?

  • Algorithm:

    • YES (type 1, 2)

      • Is METAL a TRANSITION METAL?

        • NO (type 1)

        • YES (type 2)

    • NO (type 3)

      • YES (Acid)

• Count different elements

• Metal / Nonmetal

• Metal / Poly lon

• Tran Metal / Nonmetal

• Tran Metal / Poly lon

• H + Polyatomicion

• lonic Binary

• lonic Ternary

• Poly ion ending in -ate

• Rootanion + -ic acid

• Poly ion ending in -ite

• Name metal first

• Rootanion + -ous acid

• Name transition metal first

• Name nonmetal

• Determine charge of transition metal and write as superscript(2)

• Name first element, using Greek prefixes (except "mono-")

• Name same as periodic table

• Name nonmetal second and change ending to 'ide"

• Only 1 element

• 2 elements

• Name metal first

• Diatomic molecule

• Covalent Binary (type 3)

• Name polyatomic ion second

• Name transition metal first

• DO NOT change ending

• Determine charge of transition metal and write as Roman Numeral after name

• "-ide"

• Name nonmetal second and change ending to "-ide"

Page 119:

• Metal-containing

  • Type 1 (Non-transition metal containing)

    • Ionic binary [metal nonmetal-ide]

    • Ionic ternary [metal polyatomic ion]

  • Type 2 (Transition metal containing)

    • Ionic binary [T.metal (OxS) nonmetal-ide]

    • Ionic ternary [T.metal (OxS) polyatomic ion]

• Stock/Systematic

• Common/Classical

• Transition Metal (OxS)

• Transition metal-ous

• Transition metal-ic

Page 120:

• Atomic lons

• Most common form on top

• Non-metal Containing

  • Type 3 (no leading H)

    • 1 element containing (diatomic molecule)

      • [element name]

    • 2 element containing (covalent binary)

      • [prefix-NM prefix-NM-ide]

Page 121:

• Non-metal Containing

  • Acids

    • H + element

      • [Hydroanion-ic acid]

    • H + polyatomic ions

      • Oxyanion-ate ─[Anion-ic acid]

      • Oxyanion-ite ─[Anion-ous acid]

Page 122:

• Example Problems: Determine the type of compound according to IUPAC nomenclature and provide their corresponding names.

  • Na2SO4

  • MnCl2

  • Sn(CN)4

  • Mg(HCO3)2

  • CaF2

  • P2O5

  • HCl

  • CH3COOH

Page 123:

• TABLE OF POLYATOMIC IONS

Page 124:

• Example Problems: Determine the type of compound according to IUPAC nomenclature and provide their corresponding names.

  • Na2SO4

  • MnCl2

  • Sn(CN)4

  • Mg(HCO3)2

  • CaF2

  • P2O5

  • HCl

  • CH3COOH

Chemical Compounds and Ions in Chemistry

Page 125:

• Example Problems:

  • Convert the following chemical names to their correct chemical formula:

    • Hydrobromic acid

    • Nitrous acid

    • Sodium arsenate

    • Plumbous acetate

    • Chromous nitride

    • Dibromide monoxide

    • Xenon tetrafluoride

    • Tetraphosphorus hexoxide

  • Notice the vowel elisions

Page 126:

• Chemical Reactions

Page 127:

• Reaction Types:

  • Double Replacement Reactions

  • Single Replacement Reactions

  • Combustion Reactions

  • Synthesis Reactions

  • Decomposition Reactions

  • Precipitation Reactions

  • Acid-Base Reactions

  • Oxidation-Reduction Reactions

Page 128:

• Reaction Types:

  • Reduction-Oxidation (RedOx) Reactions

  • Oxidation: A compound loses electrons

  • Reduction: A compound gains electrons

Page 129:

• RedOx Reactions:

  • Oxidation Number / Oxidation State / Valency / Charge

  • Rules for determining oxidation numbers:

    • Uncombined elements: 0

    • Neutral compound: sum is 0

    • Ionic compound: sum is equal to charge

    • Fluorine: always -1 (other halogens can have -3, -5 in some cases)

    • Oxygen: always -2 (+1/+2 with F, -1 in peroxides O2 -2)

    • Hydrogen: always +1 (-1 in metal hydride LiH)

    • Group 1A = +1; Group 2A = +2; Al = +3

Page 130:

• Practice Problems:

  • Cl2

  • Na+

  • NO2

  • MnO2

  • NaH

  • H2O2

  • Cr2O7-2

  • BaF2

  • CH4

  • SO3-2

  • Na2S

  • CN-1

  • CO

  • O2

  • HClO4

  • Fe2O3

  • FeO

  • NH4Cl

  • K3PO4

Page 131:

• Reaction Types:

  1. Synthesis / Combination / Direct Union: A + B → AB

  2. Decomposition / Analysis: AB → A + B

  3. Single Replacement: AB + X → AX + B (Redox Type)

Page 132:

• Single Displacement Reaction Activity Series:

  • Li - Leo

  • K - Kissed

  • Ba - Betty's

  • Ca - Cheek

  • Na - Never

  • Mg - Minding

  • AI IS Alex

  • Zn - Zinc's

  • Cr - Cross

  • Fe - Face

  • Cd - Closely

  • Co - Coming

  • Ni - Near

  • Sn - Steaming

  • Pb - Perhaps

  • H - He

  • Cu - Could

  • Hg - Hang

  • Ag - Around

  • Pt - Patricia

  • Au - Again

Page 133:

• Reaction Types: 4. Double Displacement / Metathesis / Exchange: AB + CD → AC + BD

  • Neutralization [acid + base → salt + water]

  • Precipitation [refer to solubility rules] (Non-Redox Type)

Page 134:

• Acids and Bases

Page 135:

• Common Properties of Acids and Bases:

  • Acids:

    • Taste: Sour

    • pH solution: <7

    • Oxides: NM oxides

    • Litmus test: Blue to Red

    • Metal Reaction: Formation of H2

  • Bases:

    • Taste: Bitter

    • pH solution: >7

    • Oxides: M oxides

    • Litmus test: Red to Blue

    • Metal Reaction: Formation of CO2

Page 136:

• Definition Systems:

  • Arrhenius:

    • Acid: Increases [H+]

    • Base: Increases [OH-]

  • Bronsted-Lowry:

    • Acid: Proton donor

    • Base: Proton acceptor

  • Lewis:

    • Acid: Electron pair acceptor

    • Base: Electron pair donor

Page 137:

• Definition Systems:

  • Lewis Theory:

    • Lewis Acid/Electrophile: Metal cation (electron poor)

    • Lewis Base/Nucleophile: Nonmetal anion (electron rich)

  • Pearson’s HSAB (Hard-Soft Acid-Base) Concept:

    • Hard-Hard / Soft-Soft: Stronger interaction due to similarity

    • Hard-Soft / Soft-Hard: Weaker compared to HH/SS

Page 138:

• Definition Systems:

  • Pearson’s HSAB Concept:

    • Acids & Bases:

      • Hard:

        • Ionic radius: Small

        • Oxidation state: High

        • Polarizability: Low

        • Electronegativity: High

      • Soft:

        • Ionic radius: Large

        • Oxidation state: Low

        • Polarizability: High

        • Electronegativity: Low

    • Examples:

      • Hard: G1A and G2A cations, NH4+, Ti4+, Cr3+

      • Soft: OH-, F-, Cl-, CO3-2, CH3COO-

      • Heavy metals: Ag+, Au+, Hg2 2+/Hg2+, Cd2+

      • H- (hydrides), I-, SCN-

Page 139:

• pH:

  • Strong acids and bases:

    • pH = log [H+]

    • pOH = log [OH-]

  • Weak acid/base buffers:

    • Henderson–Hasselbalch equation:

      • pH = pKa + log[A-]/[HA] (if acid)

      • pH = pKa + log[B]/[BH+] (if base)

  • Weak acid/base:

    • RICE table

Page 140:

• pH calculations:

  • Determine the pH of a 0.07M HCl solution.

  • Determine the pH of a base buffer containing 0.1M NH3 (pKb = 4.76) and 0.20M NH4Cl solution.

  • Determine the pH of 0.2 M acetic acid (pKa = 4.8).

  • ICE Table

Page 141:

• Common Ion Effect:

  • Addition of compounds with ions that are common to the already dissolved substances in solution.

  • This leads to:

    • Shift in equilibrium

    • Suppressed ionization of dissolved substances (esp. WA/WB)

    • pH change

Page 142:

• Common Ion Effect Practice problems after Chemical Equilibrium - Add CH3CO2 CH3CO2H H+ + CH3CO2 [H+] pH Reaction shifts left

Page 143:

• Stoichiometry

Page 144:

• Stoichiometry:

  • The determination of the proportions in which elements or compounds react with one another.

Page 145:

• Expression of Concentrations

Page 146:

• Example Problems:

  • Acetylene gas C2H2 undergoes combustion to form carbon dioxide and water when it is used in oxyacetylene torch for welding.

  • Balance the reaction and answer the following questions:

    • How many grams of water can form if 113g of acetylene is burned? [g → g]

    • How many moles of carbon dioxide can form if 150g of acetylene is burned? [g → mol]

Chemical Compounds and Ions in Chemistry

Page 147:

• Example Problem: Sodium metal burns in air according to a reaction

• Balance the equation and answer the following:

  • If 2 mol of sodium is consumed, how many grams of sodium oxide will be formed?

  • If only 0.7 mol of O2 is available, how many mol of sodium oxide can be formed?

Page 148:

• Example Problem: Reaction of calcium chloride with phosphoric acid

• Calculate the grams of calcium phosphate produced

• Assuming the reaction volume remains the same, determine the %w/v, molarity, normality, and osmolarity of calcium phosphate.

Page 149:

• Practice Problem: Reaction of HNO3 with barium hydroxide solution

• Determine the %w/v, M, and N of the nitric acid solution.

Page 150:

• Thermodynamics

Page 151:

• Functions in thermodynamics

• State and non-state functions

• Path dependence

• Examples of functions: Enthalpy, Internal energy, Gibbs Free Energy, Entropy, Work, Heat

Page 152:

• Hess's Law in thermodynamics

• Total enthalpy change for a reaction is the sum of all changes

• Enthalpy (H) and internal energy (U) in a thermodynamic system

Page 153:

• Thermodynamic Laws

• Zeroth Law: Thermal Equilibrium

• First Law: Mass and Energy Conservation

• Second Law: Entropy

• Third Law: Definition of the Kelvin Scale

Page 154:

• Spontaneity of Reactions

• Spontaneity based on entropy and heat

• Gibbs free energy (AG)

• Spontaneous and non-spontaneous reactions at different temperatures

Page 155:

• Spontaneity of Reactions (continued)

• Spontaneous and non-spontaneous reactions based on entropy (AS)

• Spontaneous reactions at low temperatures and non-spontaneous reactions at high temperatures

Page 156:

• Reaction System Types: Open, Closed, Isolated

• Heat Flow: Endothermic and Exothermic

• Change in Enthalpy: Endergonic and Exergonic

Page 157:

• Practice Problems: Determining the spontaneity of reactions based on entropy change (AS), enthalpy change (AH), and temperature (T)

Page 158:

• Chemical Kinetics

Page 159:

• Chemical Kinetics: Study of reaction rates and reaction mechanism

• Reaction Rate Unit: Rate = moles/s

• Average Reaction Rate and Instantaneous Reaction Rate

Page 160:

• Chemical Kinetics (continued)

• Instantaneous Reaction Rate and Rate Laws

• Relationship between reaction rate constant (k) and initial reactant concentration

Page 161:

• Reaction Rate Theories

• Collision Theory: Reaction rate is proportional to the number of collisions per time

• Activation Energy (AE) and proper orientation for reactions

• Transition State Theory: Rate depends on the AE required to form intermediate/transition states

Page 162:

• Reaction Coordination Graph

• Activation Energy (Eact) and Heat of Reaction (AH)

• Reactants, products, and transition state in a reaction

Page 163:

• Factors affecting Reaction Rate

• Reactant's nature, concentration, catalyst, surface area, and temperature

Page 164:

• Practice Problems: Determining the average rate of disappearance of oxygen and the instantaneous reaction rate based on rate law equation

Page 165:

• Chemical Equilibrium

Page 166:

• Chemical Equilibrium (continued)

• Law of Mass Action: Reaction rate is proportional to the product of the concentration of reactants raised to their balanced-equation coefficients

Page 167:

• Law of Mass Action (continued)

• Equilibrium direction based on Keq value

Chemical Equilibrium and Le Chatelier's Principle

Page 168

• Chemical Equilibrium

  • Definition: A state in which the forward and reverse reactions occur at equal rates.

  • Equilibrium is achieved when the concentrations of reactants and products remain constant over time.

• Le Chatelier's Principle

  • Definition: If an external stress is applied to a system at equilibrium, the system adjusts to partially offset the stress and reach a new equilibrium.

  • The equilibrium shift depends on the type of stress applied.

• External Stress

  • Concentration:

    • If the concentration of a reactant or product is changed, the system will shift to restore equilibrium.

    • Increasing the concentration of a reactant will shift the equilibrium towards the product side.

    • Decreasing the concentration of a reactant will shift the equilibrium towards the reactant side.

  • Pressure and Volume:

    • Changing the pressure or volume of a system will only affect the equilibrium if there is a difference in the number of moles of gas on each side of the reaction.

    • Increasing the pressure will shift the equilibrium towards the side with fewer moles of gas.

    • Decreasing the pressure will shift the equilibrium towards the side with more moles of gas.

  • Temperature:

    • Changing the temperature of a system will affect the equilibrium.

    • Increasing the temperature will shift the equilibrium in the endothermic direction (absorbing heat).

    • Decreasing the temperature will shift the equilibrium in the exothermic direction (releasing heat).

  • Catalyst:

    • Adding a catalyst does not affect the equilibrium position.

    • A catalyst only speeds up the rate of the forward and reverse reactions, allowing