Advanced Chem
1. Matter
- Pure Substance
- element or compound
- uniform and defined chemical composition
- Mixture
- heterogeneous (can be separated)
- homogenous (can not be separated) Ways to separate mixtures:
- Chromatography - dissolving rate
- Distillation - boiling point
- Magnet - magnetism
- Water Sedimentation - solubility in water
- Filter - solubility/particle size
- Crystallization - solubility properties
- HCI - chemical reaction with HCI
2. Isotopes & Ions
![[Pasted image 20250921165051.png]] Never assume an atoms mass Isotopes: same element with different masses (different # of neutrons)
3. Radiation
Radiation: emission or transmission of energy in the form of waves or particles. Radioactive decay: is a type of radiation that occurs when an unstable atomic nucleus spontaneously loses energy by emitting particles or energy to become more stable Fission v.s. Fusion Fission: heavier nucleus splits into few smaller nucleuses (causes a chain reaction) Fusion: smaller nucleuses combine to become a heavier nucleus (needs extreme heat)
3.1. Types of Radioactive Decay
- Alpha Decay (α) In alpha decay, the nucleus emits an alpha particle, which is a helium nucleus ($^{4}_{2}$He). This particle consists of two protons and two neutrons. The process causes the parent nuclide's atomic number to decrease by two and its mass number to decrease by four. Alpha particles are heavy and have high ionizing power, but their penetrating power is low; they can be stopped by a sheet of paper.
- Beta Decay (β) Beta decay happens when a neutron in the nucleus turns into a proton, emitting an electron (a beta particle, $_{-1}^{0}$β). This decay increases the atomic number by one, but the mass number remains unchanged. Beta particles are much lighter than alpha particles, have less ionizing power, but greater penetrating power. They can be stopped by a thin sheet of aluminum.
- Positron Emission (β + ) In positron emission, a proton in the nucleus is converted into a neutron, and a positron ($_{+1}^{0}$β), which is an antimatter particle with the same mass as an electron but a positive charge, is emitted. This process decreases the atomic number by one, but the mass number stays the same.
- Electron Capture (EC) Electron capture is a decay process where a proton in the nucleus captures an inner-shell electron from its own atom. This captured electron combines with the proton to form a neutron, decreasing the atomic number by one while the mass number stays the same.
- Gamma Rays (γ) Gamma rays are not particles, but a form of high-energy electromagnetic radiation. Gamma decay occurs when a nucleus in an excited state returns to a lower, more stable energy state by releasing this energy. Since gamma rays have no mass or charge, the mass and atomic numbers of the nucleus do not change. Gamma rays have the greatest penetrating power and the least ionizing power, requiring several inches of dense material like lead for shielding. ![[Pasted image 20250922210641.png]]
3.2. Half Life
Radioactive half life: is the time required for exactly one-half of the nuclei in a sample of a specific radioactive isotope to decay. This is a probabilistic process, meaning we can't predict when a single nucleus will decay, but we can accurately predict the rate of decay for a large group of them.
The amount of a radioactive substance remaining after a certain time can be calculated using the decay equation: Nt=N0⋅(0.5)t1/2t - Nt = the amount of the substance remaining after time t - N0 = the initial amount of the substance - t = the elapsed time - t1/2 = the half-life of the substance
3.3 History of The Atom
Democritus - proposed everything was made from tiny particles surrounded by empty space Aristotle - proposed everything was made from earth, wind, water, and fire John Dalton - showed common substances broke down into same elements in the same proportions Ernest Rutherford - atoms consisted of large amounts of empty space, with concentration of mass in the center (nucleus) Niels Bohr - electrons orbited the nucleus James Chadwick - the nucleus had an uncharged particle, the neutron Werner Heisenberg - impossible to determine the position and speed of electrons
Unit 4:
4.1. Electromagnetic Spectrum
- Light waves are classified by their increasing energies (or wavelengths). The visible light spectrum is a small part of the entire electromagnetic spectrum.
- Long Wavelengths = Low Frequency → Low Energy (e.g., Radio, Microwaves, Infrared)
- Short Wavelengths = High Frequency → High Energy (e.g., Ultraviolet, X-ray, Gamma Ray)
- The separated bands of light viewed through a spectroscope give an atomic line spectra for that element, which is a "unique fingerprint" for that element.
- Example:Electrified Hydrogen gas produces a unique spectral banding pattern.
- Ground State: This is the lowest energy state possible for an electron. The electron is as close as possible to the positively charged nucleus.
- Excited State:
- When an atom absorbs energy (e.g., heat from a flame, electricity), its electrons can absorb this energy and become "excited," moving away from the nucleus to a new, higher energy level.
- The amount of energy required for an electron to jump a level is quantized (a specific, definite amount).
- The excited state is unstable and not where the electron normally stays.
- The electron will eventually lose the energy it gained and "fall back" to its lower energy level (closer to the nucleus), often back to the ground state.
- When the electron loses this excess energy, it is sometimes emitted as a packet of energy (photon) at a specific wavelength that falls within the "visible light" range.
- This visible light is what we see as the colored flame during a flame test.
- The specific drop in energy (ΔE) corresponds to a specific wavelength of light (λ or E=hν, where ν is frequency, λ is wavelength).### Electromagnetic Spectrum
- Light waves are classified by their increasing energies (or wavelengths). The visible light spectrum is a small part of the entire electromagnetic spectrum.
- Long Wavelengths = Low Frequency → Low Energy (e.g., Radio, Microwaves, Infrared)
- Short Wavelengths = High Frequency → High Energy (e.g., Ultraviolet, X-ray, Gamma Ray)
4.2. Electron Configurations
Electron Configuration: is a summary of where electrons are located around the nucleus of an atom. - The electron's distance from the nucleus Energy Level. - The type of subshell Sub- level the electrons reside in. - The number of electrons within each subshell.
| Sublevel | Shape | Max. Electrons | Number of Orbitals | Orbital Capacity |
|---|---|---|---|---|
| s | Spherical | 2 | 1 | 2 electrons |
| p | Dumbbell-shaped | 6 | 3 (on x,y,z axes) | 2 electrons each |
| d | Various shapes | 10 | 5 | 2 electrons each |
| f | Complex shapes | 14 | 7 | 2 electrons each |
| ![[Pasted image 20251003214906.png]] |
| Block Name | Sublevel | Location on P.T. |
|---|---|---|
| s-block | s | Groups 1 and 2 (plus Helium) |
| p-block | p | Groups 13 through 18 |
| d-block | d | Transition metals (Groups 3-12) |
| f-block | f | Inner transition metals (Lanthanides and Actinides) |
| ### Unit 4 | ||
| #### 4.1. Periodic Trends | ||
| Periodic trends are specific patterns in the properties of chemical elements that are revealed in the periodic table. These trends are explained by the electron configurations of the elements. | ||
| ##### Effective Nuclear Charge ($Z_{eff}$) | ||
| - Definition: The net positive charge experienced by an outer-shell (valence) electron. It is the pull the electron "feels" from the nucleus once the effect of shielding by inner (core) electrons is accounted for. | ||
| - Shielding: Core electrons block or "shield" the valence electrons from the full attractive force of the nucleus. | ||
| - Formula: $Z_{eff} = Z - S$ | ||
| - $Z$ = Atomic Number (# of protons) | ||
| - $S$ = Number of shielding (core) electrons | ||
| - Trend Across a Period (Left to Right): $Z_{eff}$ increases. | ||
| - Reason: The number of protons increases, but the number of shielding core electrons stays the same. This results in a stronger pull on the valence electrons. | ||
| - Trend Down a Group (Top to Bottom): $Z_{eff}$ remains relatively constant. | ||
| - Reason: While the number of protons increases, the number of shielding shells also increases, canceling out the effect of the added protons. | ||
| ##### Atomic Radius | ||
| - Definition: Half the distance between the nuclei of two identical atoms that are bonded together. Essentially, the size of the atom. | ||
| - Trend Across a Period: Atomic radius decreases. | ||
| - Reason: The effective nuclear charge ($Z_{eff}$) increases, pulling the electron cloud closer to the nucleus. | ||
| - Trend Down a Group: Atomic radius increases. | ||
| - Reason: A new principal energy level (shell) is added, placing the valence electrons further from the nucleus. | ||
| ##### Ionic Radius | ||
| - Definition: The radius of an atom's ion. | ||
| - Cations (Positive Ions): Are smaller than their parent neutral atom. | ||
| - Reason: The atom loses its valence electron(s), often resulting in the loss of an entire energy level. There is also less electron-electron repulsion. | ||
| - Anions (Negative Ions): Are larger than their parent neutral atom. | ||
| - Reason: The atom gains electron(s) in the valence shell. The added electron-electron repulsion causes the electron cloud to expand. | ||
| ##### Ionization Energy (IE) | ||
| - Definition: The minimum energy required to remove one valence electron from a neutral atom in its gaseous state. | ||
| - Trend Across a Period: Ionization energy increases. | ||
| - Reason: As $Z_{eff}$ increases, the nucleus holds onto its valence electrons more tightly, making them harder (and requiring more energy) to remove. | ||
| - Trend Down a Group: Ionization energy decreases. | ||
| - Reason: Valence electrons are in higher energy levels, further from the nucleus. Increased shielding makes them easier to remove. | ||
| - Successive IE: The energy to remove a second electron ($IE_2$) is always greater than the first ($IE_1$). A very large jump in IE occurs when trying to remove a core electron. | ||
| ##### Electronegativity | ||
| - Definition: A measure of the ability of an atom in a chemical bond to attract shared electrons to itself. | ||
| - Trend Across a Period: Electronegativity increases. | ||
| - Reason: A higher $Z_{eff}$ allows the nucleus to attract bonding electrons more strongly. | ||
| - Trend Down a Group: Electronegativity decreases. | ||
| - Reason: The bonding electrons are further from the nucleus and are shielded, resulting in a weaker attraction. | ||
| - Note: Noble gases are typically not assigned electronegativity values as they rarely form bonds. Fluorine (F) is the most electronegative element. | ||
| ##### Summary of Trends |
| Trend | Across a Period (→) | Down a Group (↓) | Reason |
|---|---|---|---|
| Effective Nuclear Charge ($Z_{eff}$) | Increases | Stays Constant | More protons, same shielding |
| Atomic Radius | Decreases | Increases | Increased $Z_{eff}$ pulls shells in; New shells are added |
| Ionization Energy | Increases | Decreases | Increased $Z_{eff}$ holds e- tighter; e- are further away |
| Electronegativity | Increases | Decreases | Increased $Z_{eff}$ attracts e- more; e- are further away |
| ### Unit 5 | |||
| #### 5.1. The Mole | |||
| - What is a Mole? | |||
| - A mole (mol) is the standard unit for measuring the amount of a substance in chemistry. It's a "counting unit," similar to how a "dozen" means 12 or a "ream" means 500. | |||
| - Since atoms and molecules are incredibly small, a mole represents a very large number of them, allowing us to work with amounts we can see and measure. | |||
| - Avogadro's Number: | |||
| - One mole of any substance contains $6.022 \times 10^{23}$ representative particles. This value is known as Avogadro's number. | |||
| - "Representative particles" can be: | |||
| - Atoms (for individual elements, e.g., Fe, C, He) | |||
| - Molecules (for covalent compounds, e.g., H₂O, CO₂) | |||
| - Formula Units (for ionic compounds, e.g., NaCl, MgO) | |||
| - Definition: The molar mass of a substance is the mass in grams of one mole of that substance. The units are grams per mole (g/mol). | |||
| - Finding Molar Mass: | |||
| - For Elements: The molar mass is numerically equal to the atomic mass found on the periodic table. | |||
| - Example: Carbon (C) has an atomic mass of 12.011 amu. Its molar mass is 12.011 g/mol. | |||
| - For Compounds: The molar mass is the sum of the molar masses of all the atoms in the chemical formula. | |||
| - Example: Water (H₂O) | |||
| - 2 Hydrogen atoms: $2 \times 1.008$ g/mol = 2.016 g/mol | |||
| - 1 Oxygen atom: $1 \times 15.999$ g/mol = 15.999 g/mol | |||
| - Molar Mass of H₂O = 2.016 + 15.999 = 18.015 g/mol | |||
| The mole is the central hub that connects the mass of a substance (something we can weigh) to the number of particles in it (something we can't count). We use conversion factors to move between these quantities. | |||
| #### 5.2. Acid Naming | |||
| - Binary Compounds: | |||
| - add prefix hydro | |||
| - replace ending (ide) to ic | |||
| - Ternary Compounds: | |||
| - change end (ate) to ic | |||
| - change end (ite) to ous | |||
| #### 5.3. Chemical Reactions | |||
| 6 indicators of a chemical reaction: | |||
| 1. production of gas | |||
| 2. production of light | |||
| 3. production of a precipitate (solid) | |||
| 4. temperature change | |||
| 5. color change | |||
| 6. odor change |
5 types of chemical reactions: 1. Synthesis - two or more substances combine to form a single new substance - A + B = AB 2. Decomposition - a compound breaks down into two or more simple substances - AB = A + B 3. Single Replacement - one element replaces a similar element in a compound - A +BC = AC + B - A must have a high oxidization energy than B in order for A to replace B in the compound 4. Double Replacement - positive and negative ions of two ionic compounds exchange places to form two new compounds - AB + CD = AD + CB - one product must be insoluble for a reaction to occur - NAG SAG PMS 5. Combustion - a reaction in which a substance reacts with oxygen gas - only uses carbon, hydrogen, and oxygen, resulting in carbon dioxide and dihydrogen monoxide
5.3. Oxidation-Reduction
OIL RIG - oxidation is loss (of electrons) - reduction is gain (of electrons) Oxidation and reduction always occur simultaneously, one cannot occur without the other Assigning Oxidation Numbers: 1. the oxidation number of a monatomic ion is equal to its charge 2. the oxidation number of hydrogen in a compound is always +1, UNLESS it is an ionic compound, where it is hydride, -1 3. the oxidation number of oxygen in a compound is always -2, UNLESS it is a peroxide, -1 4. the oxidation number of a free element is zero 5. for any neutral compound, the sum of the oxidation numbers is zero 6. for a polyatomic ion, the sum of the oxidation numbers must equal the charge of the ion
5.4. Balancing Redox Reactions
Half-Reaction Method: 1. write half reactions 2. balance all elements other than oxygen and hydrogen 3. balance oxygen by adding water 4. balance hydrogen by adding hydrogen ions 5. balance charge by adding electrons 6. multiple half reactions by factors to make the number of electrons lost equal to the number of electrons gained 7. add half reactions together 8. cancel out anything that is the same on both sides
5.5. Net Ionic Equations
- balance formula equations
- create the total ionic equation: determine which substances are strong electrolytes and write them in ionic form
- cancel ions that are common to both sides of the equations (these are the spectator ions)
- balanced net ionic equations: rewrite the equation focusing on the reacting species Okay, here are some notes on basic stoichiometry, formatted in a way that should align with your previous notes.
Unit 6
6.1. Stoichiometry
Stoichiometry is the study of the quantitative relationships between amounts of reactants and products in a chemical reaction. It is based on the law of conservation of mass and uses balanced chemical equations to predict the amounts of substances involved in a reaction. - Definition: The calculation of relative quantities of reactants and products in chemical reactions. - Foundation: Based on the Law of Conservation of Mass (matter cannot be created or destroyed in a chemical reaction). - Key Tool: Balanced chemical equations, which provide the mole ratios between substances. Mole ratios are derived from the coefficients in a balanced chemical equation and serve as conversion factors between different substances in a reaction. They are crucial for calculating the amount of one substance based on the known amount of another. - Source: Coefficients in a balanced chemical equation. - Purpose: To relate the amount (in moles) of one reactant or product to another. - Example: In the reaction $2H_2 + O_2 \rightarrow 2H_2O$, the mole ratio between $H_2$ and $O_2$ is $2:1$, and between $H_2$ and $H_2O$ is $2:2$ (or $1:1$).