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General Chemistry
A matter of study that deals with Mass, Volume, States (Solid, Liquid, Gas), Volume, Molecular motion, Plasma, Ionized Gas, Enthalpy, Phase Change, Matter Classification, Fundamental Chemistry Laws, and Atomic Structure.
Matter
Anything that has mass and occupies space.
Mass
The amount of matter in an object.
Volume
The amount of space occupied by an object.
States (Solid, Liquid, Gas)
The three physical forms of matter.
Solid
Has a definite shape and volume, and is non-compressible.
Liquid
Assumes the shape of its container, has a definite volume, and is non-compressible.
Gas
Assumes the shape and volume of its container, and is compressible.
Molecular Motion
Refers to the vibration and gliding of molecules in different states of matter.
Plasma
The ionized gas which is the fourth state of matter, and is the most abundant state of matter.
Ionized Gas
Consists of particles with positive and negative charges, thus greatly affected by magnetic fields.
Enthalpy
The heat reaction energy of a system.
Phase Change
The transition of a substance from one state to another, such as melting, freezing, evaporation, and condensation.
Matter Classification
Includes Pure substance (Element, Compound) and Mixture (Homogeneous, Heterogeneous).
Pure Substance
Either an element (the simplest form of a substance) or a compound (two or more chemicals united and separable through chemical means).
Element
The simplest form of a substance that cannot be separated into simpler substances by chemical means.
Compound
Two or more chemical elements united in fixed ratios and separable through chemical means.
Mixture
A combination of two or more substances where individual substance identities are retained and separable through physical means.
Homogeneous
A mixture with a single phase, resulting in a clear, colored solution.
Heterogeneous
A mixture with two phases, resulting in a suspension or colloid, ex. milk.
Classification Based on Dependent to the Amount of Matter Present
Includes Extrinsic Property (Dependent) and Intrinsic Property (Independent).
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Neutrons
Neutrons have no charge and are found in the nucleus. They contribute to the atomic mass of an element.
Atomic mass
The atomic mass of an element is the total mass of protons and neutrons in the nucleus of an atom.
Nucleon
A nucleon is a proton or a neutron present in the nucleus of an atom.
James Chadwick
James Chadwick discovered the neutron in 1932.
Atomic number
The atomic number of an element is the number of protons in the nucleus of an atom. It determines the element's identity.
Charge
The charge of an element is determined by the difference between the number of protons and electrons.
Eugene Goldstein
Eugene Goldstein discovered anode rays and made significant contributions to the understanding of atomic structure.
Electrochemistry
Electrochemistry is the study of chemical processes that cause electrons to move.
Capillary electrophoresis
Capillary electrophoresis is a separation technique based on the electrophoretic mobility of compounds.
Anode
The anode is the positively charged electrode where oxidation occurs in an electrochemical cell.
Cathode
The cathode is the negatively charged electrode where reduction occurs in an electrochemical cell.
Redox
Redox refers to oxidation-reduction reactions where one substance loses electrons (oxidation) and another gains electrons (reduction).
Valence
Valence refers to the combining capacity of an element, expressed as the number of electrons it can lose, gain, or share in a chemical bond.
Oxidizing agent
An oxidizing agent is a substance that oxidizes other substances by accepting their electrons.
Isotopes
Isotopes are atoms of the same element with different numbers of neutrons, leading to variations in atomic mass.
Isobars
Isobars are atoms of different elements that have the same mass number but different atomic numbers.
Isomers
Isomers are compounds with the same molecular formula but different structural arrangements.
Chemical bonds
Chemical bonds are the attractive forces that hold atoms together in compounds.
Ions
Ions are charged particles formed when atoms gain or lose electrons.
Van der Waals forces
Van der Waals forces are weak electrostatic forces of attraction between molecules, including dipole-dipole interactions and London dispersion forces.
H bonding
Hydrogen bonding is a type of dipole-dipole interaction between a hydrogen atom and an electronegative atom (N, O, F).
Dipole
A dipole is a molecule with a positive and a negative end due to an unequal distribution of electron density.
Covalent bonding
Covalent bonding involves the sharing of electron pairs between atoms, leading to the formation of molecules.
Ionic bonding
Ionic bonding involves the transfer of electrons from a metal to a nonmetal, resulting in the formation of ions held together by electrostatic forces.
Valence shell electron pair repulsion (VSEPR) theory
VSEPR theory predicts the three-dimensional geometry of molecules based on the repulsion between electron pairs around the central atom.
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Valence bond theory
States that bonds are formed by sharing of electrons from overlapping atomic orbitals, leading to the formation of covalent bonds. It involves the creation of s-spherical sigma bonds (stronger) through head-on overlap, and p-dumbbell pi bonds (weaker) through sideways overlap.
Molecular orbital theory
States that bonds are formed from the interaction of atomic orbitals to create molecular orbitals. This theory describes bonding as a lower energy, stable state, and antibonding as a higher energy, unstable state.
Synthesis Reaction
A type of reaction where two or more substances combine to form a new, more complex substance. It is also known as a combination or direct union reaction (A + B o AB).
Decomposition Reaction
A type of reaction in which a complex substance breaks down into two or more simpler substances. It is also known as analysis or disintegration reaction (AB o A + B), such as complete combustion (CH4 + O2 o CO2 + H2O) or incomplete combustion (CH4 + O2 o CO + H2O).
Single Displacement Reaction
A reaction in which an element or ion moves out of one compound and into another. It is also known as a substitution or replacement reaction (AB + X o AX + B), such as in metal reactivity series reactions.
Double Displacement Reaction
A type of reaction in which two compounds exchange ions or bonds to form new compounds. It is also known as a metathesis or exchange reaction (AB + CD o AC + BD), such as in neutralization reactions (NaOH + HCl o NaCl + H2O) or precipitation reactions (AgNO3 + NaCl o AgCl + NaNO3).
Reactivity Series
A list of metals arranged in order of their decreasing reactivity. It helps predict the products of single displacement reactions and the likelihood of a metal displacing another in a compound. For example, the reactivity series for metals: Li > K > Ba > Ca > Na > Mg > Al > Mn > Zn > Cr > Fe > Cd > Co > Ni > Sn > Pb > H > Cu > Ag > Hg > Pt > Au.
Nomenclature of Inorganic Compounds
The systematic naming and classification of chemical compounds, such as covalent compounds (e.g., CO - Carbon monoxide, SiO2 - Silicon dioxide) and ionic compounds (e.g., Pb(NO3)2 - Lead(II) nitrate, Pb4+ NO3-). It includes naming conventions for monoatomic ions, polyatomic ions, and their corresponding salts and acids.
Avogadro's Number
A constant used to express the number of atoms or molecules in a mole of a substance. It is approximately 6.022 x 10^23, and is used to calculate the number of atoms/molecules in a given mass or quantity of a substance.
Molarity (M)
A measure of the concentration of a solute in a solution, expressed in moles of solute per liter of solution. It is denoted by 'M' and is used in chemical calculations to determine the amount of a substance present in a solution.
Formality (F)
A measure of the concentration of a substance in a solution, expressed as the number of formula weights of a solute per liter of solution. It is used in certain chemical stoichiometry calculations, particularly in the context of acid-base reactions and neutralization.
Critical Thinking Question
What are the key concepts of valence bond theory and molecular orbital theory, and how do they differ in their approach to understanding chemical bonding? Compare and contrast their implications for the formation and stability of chemical compounds.
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Electronic Configuration Principle
Atoms may be built by progressive filling of energy of main energy sub level, i.e., levels of lower energy levels are occupied first.
Electronic Configuration Example
Calcium: Atomic # 20, Atomic mass 40. Electronic Configuration: 1s^2 2s^2 2p^6 3s^2 4s^2
Shortcut for Noble Gases
He: 2, Ne: 10, Ar: 18, Kr: 36, Xe: 54, Rn: 86
Quantum Numbers - Principal Quantum Number
Symbolized by 'n', ranges from 1 to 7, representing the main energy level and size of orbital electron cloud.
Azimuthal Angular Momentum
Symbolized by 'l', ranges from 0 to 3, representing the shape of the orbital subshell: 0 (s), 1 (p), 2 (d), 3 (f).
Magnetic Quantum Number
Symbolized by 'm', represents the orientation of the orbital in space (e.g., m = -1, 0, 1 for n = 2, l = 1).
Degenerate Orbitals
Orbitals with the same energy level (e.g., n = 3: 0, 1, 2; n = 4: 0, 1, 2, 3).
Magnetic Spin
Represents the magnetic moment and rotation spin of electrons (e.g., ms = +1/2, -1/2 for a single electron).
Quantum Theories - Pauli's Exclusion Principle
No two electrons in an atom can have the same set of quantum numbers; they must be exclusive.
Quantum Theories - Heisenberg's Uncertainty Principle
It's impossible to accurately predict or determine the particle's velocity, position, and momentum simultaneously.
Quantum Theories - Hund's Rule
Orbitals are filled up singly before pairing up. The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins.
Gas Laws - Boyles/Mariotte
P1V1 = P2V2 (temperature in Kelvin)
Gas Laws - Charles
V1/T1 = V2/T2 (pressure in atm)
Gas Laws - Gay-Lussac
P1/T1 = P2/T2 (volume in L)
Gas Laws - Combined Ideal Gas Law
PV = nRT (R = 0.08205 L.atm/mol.K at STP)
Avogadro's Law
Equal volumes of different gases at the same temperature and pressure contain the same number of moles. (N_A = 6.022 x 10^23)
Dalton's Law of Partial Pressures
The total pressure in a mixture of non-interacting gases is equal to the sum of the partial pressures of each gas.
Graham's Law
The rate of effusion/diffusion and speed of gas are inversely proportional to the square root of their density, given the same temperature and pressure for 2 gases.
Fick's 1st Law
The diffusion rate/flux of liquid or gas is directly proportional to the concentration gradient from high concentration to low concentration.
Henry's Law of Gas Solubility
The solubility of gas decreases with a decrease in temperature and increases with an increase in pressure, i.e., in a sealed container, more CO2 is dissolved in water at higher pressure.
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Van der Waals
An equation of state for a fluid composed of particles that have a non-zero volume and do experience an attractive force.
Internal Pressure per mole
Represents the deviation of pressure from ideal gas behavior due to the non-zero volume and attractive forces between molecules. It is denoted as nb.
Incompressibility
The property of a substance that does not change in volume when subjected to pressure. It is associated with the ideal gas law.
Raoult's Law
A law that states the partial vapor pressure of each component in an ideal mixture of liquids is proportional to the mole fraction of that component in the mixture.
Mole Fraction
The ratio of the number of moles of one component in a mixture to the total number of moles in the mixture.
Temperature Conversion
The conversion between different temperature scales, including Celsius (C), Fahrenheit (F), and Kelvin (K). The formulas are: F = C x (9/5) + 32, C = (F - 32) x (5/9), and K = C + 273.15.
Absolute Temperature
Temperature measured in Kelvin (K), where 0K represents absolute zero, the lowest possible temperature.
Solute
The substance being dissolved in a solution. It is typically present in smaller quantity compared to the solvent.
Solvent
The substance that dissolves the solute to form a solution. It is typically present in larger quantity compared to the solute.
Colligative Properties
Properties of a solution that depend only on the number of solute particles present and not on the nature of the particle.
Vapor Pressure Lowering
The reduction in vapor pressure of a solvent due to the presence of a solute, based on Raoult's law.
Boiling Point Elevation
The increase in the boiling point of a solvent in the presence of a non-volatile solute.
Ebullition
The process of boiling, where a liquid is converted into vapor at a temperature below its boiling point.
Freezing Point Depression
The lowering of the freezing point of a solvent in the presence of a solute, based on colligative properties.
Osmotic Pressure
The pressure needed to stop osmosis, which is the movement of solvent molecules through a semipermeable membrane from a dilute solution to a more concentrated solution.
Thermodynamics
The study of energy conversion and transformation in the universe, including the study of systems, surroundings, state functions, and laws.
Scholarly Assistant's Insights
A flashcard deck covering General Chemistry, Matter, Atomic Structure, Chemical Reactions, Bonding, Gas Laws, Quantum Numbers, and Thermodynamics.
General Chemistry
Atomic Structure
Chemical Reactions
Quantum Mechanics
Gas Laws
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