Inorganic Chem Exam

c=λν

c=speed of light

λ=wavelength

ν=frequency

λ=h/mv

λ=wavelength

h=Planck’s constant

m=mass

v=velocity

Used to calculate the wavelength of any object

E=hc/λ

E=energy of a photon

h=Planck's constant

c=speed of light

λ=wavelength.

This equation relates the energy of a photon to its wavelength.

E=hν

E=energy of a photon

h=Planck's constant

ν=frequency.

This equation relates the energy of a photon to its frequency.

E=-2.178E-18(1/nf²-1/ni²)

Equation for hydrogen atom energy levels, where nf and ni are the final and initial energy levels, respectively.

E=KQ₁Q₂/r

Equation for electrostatic potential energy between two charged particles, where K is Coulomb's constant, Q1 and Q2 are the magnitudes of the charges, and r is the distance between them.

negative-attract

positive-repel

P₁V₁=P₂V₂

pressure and volume of gas when temperature and number of moles is kept constant

V₁/T₁=V₂/T₂

Relationship between volume and temperature of a gas, assuming constant pressure and number of moles.

P₁/T₁=P₂/T₂

Relationship between pressure and temperature of a gas, assuming constant volume and number of moles.

P₁V₁/T₁=P₂V₂/T₂

Combined gas law

This equation combines the relationships of pressure, volume, and temperature of a gas, assuming the number of moles is constant.

PV=nRT

Ideal gas law, relating pressure, volume, temperature, and amount of gas.

R=0.0821Latm/kmol

grams to moles

divide

moles to grams

multiply

U_rms=√3RT/Mm

Root mean square

R=gas constant=8.314

Mm=Kg/mol

T=temperature

Sig fig calculation

Addition and Subtraction: rounded to the place of the last sig digit of the least accurate measurement

Multiplication and Division: rounded to the place of the measurement with the least number of total sig figs

Dalton’s Atomic Theory

1. All Matter is Made of Atoms

2. All atoms of a given element have the same mass, size, and chemical properties

3. Compounds are formed when atoms of two or more different elements combine in fixed, simple, whole-number ratios

4. Chemical reactions occur when atoms are rearranged, separated, or combined

• The Law of Conservation of Mass: Mass remains constant in chemical reactions because atoms are merely rearranged.

• The Law of Definite Proportions: A chemical compound always contains the same elements in the same proportion by mass.

• The Law of Multiple Proportions: When two elements form more than one compound, the ratios of the masses of one element that combine with a fixed mass of the other are simple whole numbers.

Atomic Radius

• Definition: Atomic radius is the distance from the nucleus of an atom to the outermost electron cloud. It represents the size of the atom.

• Trends in the Periodic Table:

  1. Across a Period (Left to Right): Atomic radius decreases.

     • Reason: Increased nuclear charge pulls electrons closer to the nucleus.

  2. Down a Group: Atomic radius increases.

     • Reason: Addition of electron shells increases the distance between the nucleus and the outermost electrons.

Ionization Energy

Types of Ionization Energy:

1. First Ionization Energy (IE1IE_1IE1​):

   • Energy needed to remove the first outermost (valence) electron.

2. Successive Ionization Energies (IE2,IE3,…IE_2, IE_3, \dotsIE2​,IE3​,…):

   • Energy required to remove additional electrons after the first.

   • Each successive ionization energy is higher because:

     • The remaining electrons experience a stronger attraction to the nucleus as the positive charge increases

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Trends in the Periodic Table:

1. Across a Period (Left to Right): Ionization energy increases.

   • Reason: Increased nuclear charge holds electrons more tightly.

2. Down a Group: Ionization energy decreases.

   • Reason: Outer electrons are farther from the nucleus and experience greater shielding.

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Key Notes:

• Large jumps in ionization energy occur when removing core (inner-shell) electrons after valence electrons are removed.

Energy and Frequency Relationship

direct

Energy and Wavelength

Indirect

Frequency and Wavelength

indirect

Complexity of configuration and energy of atom

direct

Pauli Exclusion Principle

No two electrons in the same atom can have the same set of four quantum numbers

Key Implications:

1. Electron Spin Pairing:

   • An orbital can hold a maximum of two electrons, and these electrons must have opposite spins (+12+\frac{1}{2}+21​ and −12-\frac{1}{2}−21​).

2. Electron Arrangement in Atoms:

   • The principle explains why electrons fill orbitals in a specific order (e.g., Aufbau principle) and why electron pairing occurs.

3. Stability of Matter:

   • The Pauli Exclusion Principle prevents electrons from collapsing into the same quantum state, ensuring the structure and stability of matter.

Aufbau Principle

Electrons occupy orbitals starting with the lowest energy level and progress to higher energy levels, following a specific sequence.

Hund’s Rule

When electrons fill orbitals of equal energy (degenerate orbitals), they occupy them singly with parallel spins before pairing up.

Naming compounds

1. Ionic Compounds:

   • Name the metal (cation) first. Use its full name (e.g., Sodium).

   • Name the non-metal (anion) second. Add "ide" to the root of its name (e.g., Chlorine → Chloride).

   • If the metal has multiple oxidation states (e.g., Iron), specify the charge using Roman numerals (e.g., Iron(III) Chloride).

2. Covalent Compounds:

   • Name the element furthest left on the periodic table first. If they're in the same group, the one with the lower atomic number comes first.

   • Use prefixes (mono-, di-, tri-, etc.) to indicate the number of atoms of each element.

   • Change the suffix of the second element to "ide" (e.g., CO₂ is carbon dioxide).

3. Acids:

   • For binary acids (hydrogen + non-metal): Start with "hydro-" followed by the non-metal with "ic" added to the root (e.g., HCl is Hydrochloric acid).

   • For oxyacids (hydrogen + polyatomic ion): If the ion ends in "-ate," change to "-ic" (e.g., H₂SO₄ is Sulfuric acid). If the ion ends in "-ite," change to "-ous" (e.g., HNO₂ is Nitrous acid).

Polarity

• Electronegativity Difference:

  • Nonpolar Covalent Bond: If the difference in electronegativity between two atoms is less than 0.4, the bond is nonpolar (electrons are shared equally).

  • Polar Covalent Bond: If the difference is between 0.4 and 1.7, the bond is polar (electrons are shared unevenly, creating a dipole).

  • Ionic Bond: If the difference is greater than 1.7, the bond is ionic (electrons are transferred from one atom to another).

• Molecular Geometry:

  • Symmetrical Molecules (e.g., CO₂): Even if the bonds are polar, the molecule itself may be nonpolar due to the symmetry (dipoles cancel each other out).

  • Asymmetrical Molecules (e.g., H₂O): If the molecule has an uneven distribution of charge, it is polar, with a positive side (δ+) and a negative side (δ-).

• Overall Polarity:

  • Nonpolar Molecule: No overall dipole moment (e.g., O₂, CH₄).

  • Polar Molecule: Has a dipole moment, where one end is slightly negative and the other slightly positive (e.g., H₂O, NH₃).

Energy and charge

direct

Energy and Radius

indirect

Potential Energy and Stability

indirect

Intermolecular Forces

1. London Dispersion Forces (LDF):

   • Present in all molecules, but especially important in nonpolar molecules.

   • Caused by temporary dipoles induced by the movement of electrons.

   • The weakest type of intermolecular force.

   • Stronger in larger molecules with more electrons.

2. Dipole-Dipole Interactions:

   • Occur between polar molecules.

   • The positive end of one molecule is attracted to the negative end of another.

   • Stronger than London Dispersion Forces, but weaker than hydrogen bonds.

3. Hydrogen Bonding:

   • A special type of dipole-dipole interaction.

   • Occurs when hydrogen is bonded to highly electronegative atoms like N, O, or F.

   • Stronger than regular dipole-dipole interactions, making substances like water have a higher boiling point.

4. Ion-Dipole Forces:

   • Occur between ionic compounds and polar molecules.

   • The ionic charge interacts with the dipole of the polar molecule.

   • Stronger than dipole-dipole interactions and hydrogen bonding.

Key Tip:

• Strength Order (Weakest to Strongest): London Dispersion Forces < Dipole-Dipole < Hydrogen Bonding < Ion-Dipole

cohesion

attraction between same molecule

ahesion

attraction between two different moelcules

States of matter and Energy

1. Solid:

   • Molecules: Tightly packed, vibrating in place.

   • Energy: Low energy, as molecules are not moving freely.

   • Shape & Volume: Fixed shape and volume.

2. Liquid:

   • Molecules: Closely packed but able to move past each other.

   • Energy: Moderate energy, allowing molecules to flow but not escape.

   • Shape & Volume: Fixed volume, but no fixed shape (takes the shape of the container).

3. Gas:

   • Molecules: Far apart, moving freely and rapidly.

   • Energy: High energy, as the molecules have enough energy to overcome intermolecular forces.

   • Shape & Volume: No fixed shape or volume (expands to fill the container).

Energy Changes:

• Melting (Solid → Liquid): Energy is absorbed (endothermic).

• Freezing (Liquid → Solid): Energy is released (exothermic).

• Vaporization (Liquid → Gas): Energy is absorbed (endothermic).

• Condensation (Gas → Liquid): Energy is released (exothermic).

• Sublimation (Solid → Gas): Energy is absorbed (endothermic).

• Deposition (Gas → Solid): Energy is released (exothermic).

Key Concept:

• Increasing energy (heating) leads to movement from solid to liquid to gas.

• Decreasing energy (cooling) leads to movement from gas to liquid to solid.

Properties of Liquids

• Definite Volume:

  • Liquids have a fixed volume but no fixed shape. They take the shape of their container but maintain their volume.

• Fluidity:

  • Liquids flow easily, meaning their molecules can move past each other. This allows them to take the shape of the container.

• Viscosity:

  • The resistance of a liquid to flow. High viscosity means the liquid flows slowly (e.g., honey), while low viscosity means the liquid flows easily (e.g., water).

• Surface Tension:

  • The force that causes the surface of a liquid to behave like a stretched elastic membrane. This is due to the cohesive forces between molecules at the surface (e.g., water forming droplets).

• Cohesion and Adhesion:

  • Cohesion: The attraction between molecules of the same substance (e.g., water molecules sticking together).

  • Adhesion: The attraction between molecules of a substance and different substances (e.g., water sticking to glass).

• Density:

  • Liquids typically have higher densities than gases but lower densities than solids. The density of a liquid is generally constant unless temperature or pressure changes significantly.

• Capillary Action:

  • The ability of a liquid to flow in narrow spaces without the assistance of external forces (e.g., water climbing up a thin tube).

• Evaporation and Boiling:

  • Evaporation: The process of liquid turning into gas at the surface when molecules gain enough energy.

  • Boiling: When a liquid turns into gas throughout the entire liquid, typically at a specific temperature (boiling point)

Kinetic Molecular Theory

1. Gas particles are so small compared to the distance between them that the volume/mass of each particle is negligible

   -volume of container=volume of gas

2. Gas particles are in constant motion until it hits something

   -cause of pressure

3. Gas particles exert no force on each other

   -IMFs don’t play a role, assume particles travel in straight line

4. The average kinetic energy of a collection of gas particles is directly proportional to the average kinetic energy of the Kelvin temperature of the gas

   -more particles=more heat

Lewis diagram and 3d drawing

1. Lewis Dot Diagram (Electron Dot Structure):

• Step 1: Determine the total number of valence electrons in the molecule or ion. Add electrons for negative charges and subtract for positive charges.

• Step 2: Write the symbols for the atoms. The least electronegative atom is usually placed in the center (except for hydrogen, which is always on the outside).

• Step 3: Place a pair of electrons between atoms to form bonds. Start with single bonds (one pair of electrons).

• Step 4: Distribute remaining electrons as lone pairs around atoms to fulfill the octet rule (8 electrons for most atoms, 2 for hydrogen).

• Step 5: Check if each atom has a full valence shell. If necessary, create double or triple bonds by sharing lone pairs between atoms.

2. 3D Structure (VSEPR Theory - Valence Shell Electron Pair Repulsion):

The shape of a molecule is determined by the repulsion between electron pairs around the central atom.

• Step 1: Count the electron pairs (bonding and lone pairs) around the central atom.

• Step 2: Determine the molecular shape using the VSEPR model:

  • Linear: 2 bonding pairs (180° angle) – Example: CO₂

  • Bent: 2 bonding pairs + 1 or 2 lone pairs – Example: H₂O

  • Trigonal Planar: 3 bonding pairs (120° angle) – Example: BF₃

  • Tetrahedral: 4 bonding pairs (109.5° angle) – Example: CH₄

  • Trigonal Bipyramidal: 5 bonding pairs – Example: PCl₅

  • Octahedral: 6 bonding pairs – Example: SF₆

• Step 3: Draw the 3D shape using lines and wedges:

  • Single line: Bond in the plane of the paper.

  • Wedge: Bond coming out of the paper (towards you).

Scientific notation

Converting a Number to Scientific Notation:

• Step 1: Identify the significant digits of the number (non-zero digits).

• Step 2: Move the decimal point to create a number between 1 and 10.

  • Example: 34,500 → 3.45

• Step 3: Count how many places you moved the decimal point.

  • If you moved it to the left, the exponent is positive.

  • If you moved it to the right, the exponent is negative.

• Step 4: Write the number as the product of the significant digits and 10 raised to the power of the count of decimal places moved.

SI units and scientific notations

Converting to SI Units (Example Process):

Step 1: Identify the prefix. Look for the prefix of the unit you're converting to.

• Example: Converting 3.2 kilometers (km) to meters (m):

  • 1 km = 10³m

Step 2: Convert the number. Multiply or divide by the appropriate power of 10 to convert to the SI base unit.

• 3.2 km = 3.2 x10³m = 3200 m

Step 3: Write in scientific notation. Express the result in scientific notation if necessary.

• Example: 3200 m = 3.2×10³ m

4. Converting Smaller Units (e.g., cm to m):

Step 1: Identify the conversion factor. 1 cm = 10−210^{-2}10−2 m. Step 2: Convert 150 cm to meters.

• 150 cm = 150×10^-2m=1.50cm

Electronegativity

Definition of Electronegativity:

• Electronegativity is the ability of an atom to attract shared electrons in a chemical bond. It is a measure of how strongly an atom can pull electrons toward itself in a covalent bond.

Trends in Electronegativity:

• Across a Period (Left to Right): Electronegativity increases as you move across a period on the periodic table. This is because atoms have more protons, which increases the effective nuclear charge, attracting electrons more strongly.

• Down a Group (Top to Bottom): Electronegativity decreases as you move down a group. This is due to the increased number of electron shells, which means the outer electrons are farther from the nucleus and are less attracted to it.

Electronegativity and Bonding:

• Nonpolar Covalent Bond: If two atoms have similar electronegativities, they share electrons equally (e.g., H₂, O₂).

• Polar Covalent Bond: If there is a moderate difference in electronegativity, the electrons are shared unevenly, creating a dipole (e.g., H₂O).

• Ionic Bond: If there is a large difference in electronegativity, the more electronegative atom will attract the electrons completely, resulting in electron transfer (e.g., NaCl)