Atomic Theories - Chapter 2

Chapter 2: Atomic Theories

• John Dalton's atomic theory is based on two laws: the law of conservation of mass and the law of definite proportions.
• In a chemical reaction, mass is neither created nor destroyed, according to the law of conservation of mass.
• The law of definite proportions stipulates that a chemical compound always contains the same proportion of elements by mass.
• Tiny particles known as atoms constitute all matter, according to Dalton's theory.
• Atoms are indestructible and can change their positions during reactions but remain intact.
• All atoms of a pure element have identical mass and properties.
• Atoms of different elements have different masses and properties.
• New and complex particles come into being when different elements combine to form compounds.
• Elements in a compound are always present in a fixed numerical ratio.
• The discovery of electrons led to the “Plum-Pudding model”, where negative charges (plums) are surrounded by a positive "cloud" (pudding).
• The Gold-Foil Experiment involved shooting alpha-rays through gold foil.
• Some alpha-particles were deflected, leading to the conclusion that an atom has two parts.
• An atom has a positively charged core (nucleus) that is 1/10,000th of the atom's size comprised of protons and neutrons.
• An atom contains a negatively charged shell that contains nearly mass-less particles (electrons) that make up the atom's size.
• Protons and neutrons have almost the same mass as 1/12th of a Carbon-12 atom (1 u) while electrons have a mass of 0.0005 u.
• The whole weight of an atom is essentially in the nucleus.

The Mole and Atomic Mass

• A mole is the unit for the amount of a substance, defined as 6*10^23 parts (Avogadro's Number).
• One mole of hydrogen has a mass of 1 gram.
• The atomic mass (M) is given in grams per mole (g/mol), defined by the formula M = m/n (atomic Mass = mass / amount).

Isotopes

• Isotopes are atoms with the same number of protons but a different number of neutrons, resulting in different masses.
• Out of 80 elements with isotopes, only 26 are stable.
• Deuterium is an isotope of hydrogen.
• Carbon-14 is important for age determination of fossils.
• Uranium-235 is used in reactors.

Calculating Isotopic Mass

• Chlorine has two isotopes: 75.77% is Cl-35 and 24.23% is Cl-37 with a calculation of (0.7577 * 35 + 0.2423 * 37 = 35.485 ) which demonstrates that the atomic mass of Cl is 35.485.
• Copper has two stable isotopes: Cu-63 and Cu-65 and can be calculated with: 63 * (x) + 65 * (1-x) = 63.546 resulting in x = 0.727 indicating that 72.7% of Copper is Cu-63.

Chapter 3: Electron Shells and the Quantum Mechanical Atom

• James Clerk Maxwell discovered the electromagnetic field.
• The energy of a photon is proportional to its radiation frequency: E = h * v (h…Planck’s constant).
• Niels Bohr's model describes electrons traveling in orbits around the nucleus.
• When an atom transitions from a higher to a lower energy level, energy is released as photons.
• Electrons in an atom are quantized, meaning they are restricted to certain energy levels.
• Bohr's theory is limited because it only applies to hydrogen.
• Heisenberg's uncertainty principle suggests that the position and momentum of physical properties cannot both be known at the same time.
• The Bohr Model’s two things that cannot be directly known at once are known orbits and a definite radius.
• Wave-particle duality suggests that electrons can be seen as both particles and waves.

Quantum Numbers

• Wave function 𝚿 can describe electrons.
• 𝚿^2 indicates the probability of finding an electron, defining regions called orbitals.
• The quantum number n indicates the size of the wave where n = 1 - ∞.
• Quantum number n is sometimes described with a capital letter: 1 = K, 2 = L, 3 = M, 4 = N, 5 = O.
• The secondary quantum number l defines subshells where larger shells contain more subshells.

Spin and Orbital Occupation

• Every shell holds 2n^2 electrons, and every orbital holds 2 electrons.
• Those 2 electrons differ by "spin": S= -1/2 or +1/2
• Orbital occupation rules:
  • Orbitals with lower energy are occupied first (E-minimum).
  • Each orbital first gets occupied with a single electron (Hund's rule).
  • Two electrons in one orbital must have differing spin (Pauli exclusion principle).

Chapter 4: The Periodic Table of Elements

• Mendelejew and Meyer published the first periodic table in 1869
• The first periodic table was missing the noble gasses.
• The number of valence electrons corresponds to the main group number.
• Metals are electric conductors, deformable, have metallic luster, and have high melting points.
• Non-Metals are often compounds, have poor conductivity, and are present as gas-to-solid states.
• Metalloids are semiconductors that occur between metals and non-metals.
• Alkali Metals (Group I) have a configuration of s1 and are very reactive.
• Alkaline-Earth Metals (Group II) have a configuration of s2 that are used in earth building compounds.
• Halogens (Group VII) have a configuration of s2p5 and are highly reactive.
• Chalcogens (Group VI) have a configuration of s2p4 and are quite reactive.
• Noble Gases (Group VIII) have a configuration of s2p6, a full shell making them non-reactive.

Atomic Properties

• Atomic radius, ionization energy, and electron affinity are especially important in the periodic table.
• Electronegativity is the ability to attract an electron; Fluorine (F) and Oxygen (O) have the highest electronegativity (EN).

Chapter 5: Formulas and Equations

• A chemical equation describes what happens during a chemical reaction where reactant states can be specified as solid (s), liquid (l), gas (g), and aqueous (aq).
• When sodium (Na) and chlorine (Cl) are in solution, they are always separated.
• Chemical equations have no atoms lost during the reaction, so there should be the same amount on each side.

Naming Conventions

• Gas molecules always consist of two atoms.
• Use “BrINClHOF” to remember the diatomic gasses: Bromine, Iodine, Nitrogen, Chlorine, Hydrogen, Oxygen, Fluorine.
• No prefixes are needed for naming ionic salt compounds because the compound explains the number of atoms.
• Cations (positive ions) are named easily, usually for metal ions.
• Sodium Ion = Na+
• Potassium Ion = K+
• Calcium Ion = Ca2+
• The charge can be provided in the PTE’s Main Group number.
• Copper(II)-Ion = Cu2+
• Iron(II)-Ion = Fe2+
• Iron(III)-Ion = Fe3+
• For simple anions, the naming is often the same as in molecules, using fluoride, chloride, oxide, sulphide where charge is derived by calculating 8 - Main Group number.
  • Polyatomic ions consist of more than one atom.
  • Just put the cation name in front of the anion name to name polyatomic ions.
  • Sodium chloride = NaCl
  • calcium oxide = Ca
  • potassium nitrate = KNO3
  • sodium hydroxide = NaOH
• Examples of naming a compound
  • Calcium chloride = CaCl2
  • Barium Nitrate = Ba(NO3)2
  • Potassium Hydroxide = KOH
  • Aluminum Carbonate = Al2(CO3)3
• A salt must be electrically neutral where positive charge equals negative.
• Sodium Chloride is NaCl where Na+ and Cl- result in NaCl
• Two ions of F- for one Mg2+ to keep it neutral so you need to write: MgF2

Concentration

• The concentration of a compound in an aqueous solution is given either in the units g/L or mol/L (most common) also written as M (molar mass).

Chapter 6: Energy and Chemical Change

• 1 cal = 1kcal = 4.184 J
• Heat Capacity (C) is the amount of heat an object must gain to raise its temperature for 1°C can be expressed as Q = C*(T2-T1) with the formula = C=Q/ΔT [ J / °C ].
• Specific Heat Capacity’s (s) C raises when the mass raises and is an extensive property with the formula s=C/m.
• Enthalpy ΔH = ΔH = Hproducts - Hreactants
• Any reaction in which heat is a product is said to be exothermic
• Any reaction that consumes energy is called endothermic.
• Standard Enthalpy ΔH° are values for ΔH for a reaction under standard conditions (25°C and 1 bar).

Conservation of Energy

• It is possible to switch sides of the equation. N2 + 3 H2  2 NH3 ΔH° = -92,38 kJ/mol and 2 NH3  N2 + 3 H2 ΔH° = +92,38 kJ/mol
• ΔH° is always related to the equation next to it.

Hess's Law

• The value of ΔH° for any reaction that can be written in steps equals the sum of the values of ΔH° of each of these individual steps.
• Download a table that explains predicting Enthalpy because all values of ΔH°f for the elements in their standard form are zero and is needed to explain standard enthalpy of a reaction (ΔH°R:)

Endothermic Reactions and Entropy

• Entropy is expressed as the disorder or randomness of the constituents of a thermodynamic system.
• A closed system always maximizes entropy.
• A high (positive) value for S means high system disorder, while a low negative value means low system disorder.
• ΔS = S(products) – S(reactants)
• A positive ΔS value means increasing disorder.
• A negative ΔS value means reducing disorder.
• The Gibbs free energy G is the combination of enthalpy and entropy and can be expressed as ΔG = ΔH – T * ΔS
• A negative value for ΔG indicates a spontaneous reaction.
• Strong exothermic reactions are always spontaneous.

Chapter 7: Chemical Bonds

• A low energy level is the key goal when two or more atoms react.
• The Octet Rule states that all atoms have 8 valence electrons that means having a noble gas state (configuration s2p6).

Iconic Bond

• Metal + Nonmetal
• Cations and Anions react with each other and build a lattice:
  • Ionization Energy of the metal is energy needed.
  • Electron affinity of the nonmetal is energy released.
  • Lattice energy (attraction) is energy released.
• High lattice energy comes from smaller ions that means more stable compounds.
• Predicts if an ionic compound will be good soluble in water or not where lower lattice energy results in better water solubility.
• Li, Na, K, and NH4 are almost always soluble.
• All metal nitrates and metal acetates are soluble.
• Iconic compounds typically take the formation of a
  • Chrystaline (high melting and boiling points)
  • Non-conductor
  • Melted or aqueous solutions are conductors
  • Brittle
• Covalent Bond results in
  • Nonmetal + Nonmetal
• The octet rule doesn’t work if Carbon can’t get 8 electrons and only has 4!

Hybridisation Process

• Spitting double used orbitals by rearranging the d-orbitals by combining an atom from the 3rd period and a strong bonding partner.
• Hybridisation happens with
  • C-Atom (1 free p-Orbital!)
  • Atoms from period 3 and higher in connection with strong electronegative partners (F, O, and Cl)
• Coordinate Bond results from
  • A bond established by both bond electrons coming from one partner.
  • Sometimes the usage of hybridisation.

Calculating Formal Charges

• Assign an electron from each bond pair to each of the partners (splitting them up) with a formula like O3
• Compare coordiate bond and hybridisation for the better system.

Rules of Resonance Structures

1. The only difference is the distribution of the electrons
2. Same charges on neighboring atoms are not allowed.
3. Those forms with the least total charge are the most important
4. The most electronegative atom shouldn’t have a positive formal charge

Chapter 8: Chemical Equilibrium

• Chemical reactions reach an equilibrium point with equal speed
• The concentration of reactants and products doesn’t change.
• To have equilibrium, no further concentration should change.
• The Law of Mass Action (k) has reactants transform into products or transforming products into reactants.
• Kc [Equilibrium Constant] is dependent on the temperature in which the reaction takes place.
• Chemical reactions take place on what’s not dissolved in water to remain.

Châtelier’s Principle

• When a system is at equilibrium it counteracts an opposing reaction with the factors of:
  • Change of temperature
  • Change of concentration
  • Change o pressure (volume)
• Reaction and product amounts at different temperatures:
  • Exothermic reaction (ΔH 0): opposite way!
• Applying pressure/lower volume favors reactions where a lower amount of gas is built.
• Releasing pressure/raise volume favors reactions where new additional gas molecules are built (higher need for space).
• Catalysts speed up the reactions without changing the system effect. Removing the product/supplying reactants can reduce side effects.
• When given temperature raise perform the exam question and reduce the pressure.
  • Endothermic reactions lead to more products.
  • Reversing a reaction equation causes the inversion of KC to 1/KC. Reaction * x increases K to the power of x.
• Make a balance and insert the numbers of a balanced reaction when you know its' "equilibrium" concentration to calculate Kc where 2NO + 2H2  N2 + 2H2O

Chapter 9: Solutions and Water

• Solutions are homogeneous liquids.
• Solubility comes from the ability of two substances to from a homogeneous solution.
• Adding salt to water results in a decreased melting point and a increased boiling point.

Dissolved Substances Properties

• Dissolving substances changes several liquid solution properties: melting/boiling point, viscosity, surface tension, ionic strength, and pH.
• Some compounds mix with other compounds instead of others regarding solubility where some compounds are soluble and are "insoluble" If:
• Have two compounds (A,B) leads to mol/L2.
• Have three compounds (A,B,C) leads to mol/L3.
• This can indicate a calculation for given m and indicate good solubility of the salt. - log (KA(B)) = pKA(B)

Solvation and Dilution Definitions

• Solvation is Process of generating a solvent molecules around creating a solvent solutions of either a ion or a normal solutions.
• For calculations
  • pH = ½ (pKA – log [concentration])
  • pOH = -log [OH-] | ph + pOH = 14
• VAPSE rule is Valance shell electron pair which can determine the dipole. dipole, and ionic or what will happen to a solution depending the properties of its' compound.

Chapter 11: Redox Reactions

• electrochemical reactions (Valence electrons move):

  • Redox equations

• electron’s reactions (Valence electrons do not move):

  • Acid-base reaction

• precipitation reaction

• formation of metal complexes Those reactions cause a flow of electrons.

• Reduction describes the process of releasing oxygen or taking up hydrogen.

• Oxidation describes the process of taking up oxygen or releasing hydrogen.

• oxidation: Compound releases electrons and increases its' number.

• reduction: Compound releases electrons, while decreasing its' number.