GENERAL CHEMISTRY explained in 19 Minutes

Everything is composed of atoms, the smallest units that make up all matter. Atoms consist of a dense core containing protons and neutrons, surrounded by electrons. The number of protons in the core defines each element, which explains the unique composition of substances like water, made from hydrogen and oxygen.

Atoms are composed of multiple electron shells, with the outermost electrons known as valence electrons that govern chemical reactions. The explanation points out that while quantum mechanics provides a more detailed picture of electron distribution, the simplified shell model remains a useful analogy. This description underlines how the arrangement of valence electrons influences atomic behavior, exemplified by sodium.

Electrons determine the chemical nature of elements, as seen in the periodic table where elements in the same group share identical numbers of valence electrons. In main groups, valence electrons match the group number, with helium being an exception that, despite its two electrons, behaves like a noble gas. Similar electron configurations lead to comparable chemical reactions, while elements in successive rows display an increasing number of electron shells and mass. Even though transition metals follow a more complex pattern, the overall arrangement reveals a consistent order in chemical properties.

Elements continuously gain mass as protons, electrons, and neutrons are added, leading to larger isotopes. Variations in the number of neutrons create different isotopes within the same element, many of which are unstable. This instability results in the breakdown of these isotopes, releasing ionizing radiation that reflects the delicate balance of atomic structure.

Atoms remain neutral when electrons and protons are balanced. Adding extra electrons makes them negatively charged, while having fewer electrons results in a positive charge. Charged atoms are known as ions, with those bearing a negative charge specifically called anions.

Every cell in the periodic table functions like a dictionary entry, revealing an element's name, symbol, number of protons, and atomic mass—a sum of protons and neutrons. Understanding that positive ions are known as cations provides insight into how atoms interact and lose electrons. The table itself is organized with metals on one side and non-metals, mostly gases, on the other, demarcated by a distinct dividing line. This layout offers a structured approach to grasping chemical properties and periodic trends.

Atoms join together to form molecules, and when different elements bond, they create compounds with entirely distinct behaviors. The text introduces semimetals as substances exhibiting intermediate properties. Compounds often display characteristics that diverge sharply from those of the individual elements. An example highlights that combining an explosive metal with a toxic gas produces a compound that is even more explosive.

Using subscripts in a molecular formula counts atoms but can mask essential structural differences. Molecules sharing the same formula can be distinct isomers with unique arrangements that determine their properties. This structural variability is pivotal, as seen in how the arrangement of carbon atoms differentiates graphite from diamonds.

Lewis-Dot diagrams represent an atom by showing its valence electrons as dots and bonds as lines. This approach converts abstract atomic models into clear visuals that focus solely on electron placement and bonding. The method simplifies the understanding of chemical structure by eliminating unnecessary details.

Atoms form bonds to achieve a lower energy state by filling their outer electron shells. The universal principle of energy reduction, similar to a ball rolling downhill, drives this tendency. Atoms with a complete outer shell—typically eight electrons, except for hydrogen and helium, which require two—naturally reach a state of stability, mirroring the inertness observed in noble gases.

Atoms with full outer electron shells remain inert due to their inherent stability. Atoms lacking a complete shell naturally share electrons to achieve lower energy and greater stability. This electron sharing, much like a ball rolling downhill, establishes covalent bonds that minimize the system's energy.

A positively charged nucleus attracts electrons from another atom, defining the process known as electronegativity. This force determines the strength of chemical bonds as electrons are drawn from one atom to another. The periodic table reveals a clear trend with electronegativity increasing from the bottom left to the top right, making fluorine exhibit the strongest pull.

A significant difference in electronegativity, typically over 1.7, drives atoms to exchange electrons, establishing ionic bonds. In sodium chloride, chlorine eagerly accepts an electron from sodium, which seeks to shed its extra electron. The electron transfer creates a positively charged sodium ion and a negatively charged chlorine ion. This arrangement of oppositely charged ions forms the characteristic grid-like structure found in salts.

A grid of positive nuclei is interlaced with a sea of freely moving, delocalized valence electrons that forms metallic bonds. These electrons move unrestrainedly across the entire structure, which underpins the properties of metals. Their mobility imparts remarkable conductivity of heat and electricity, as well as significant malleability.

When atoms share electrons equally due to a low electronegativity difference, they form nonpolar covalent bonds. Increased differences between 0.5 and 1.7 cause one atom to draw electrons more strongly, creating a polar covalent bond without complete electron transfer. In water, oxygen’s high electronegativity skews electron sharing, leading to distinct partial charges and establishing two opposing poles.

Molecules with permanent electric dipoles attract charged entities by aligning so that opposite charges are adjacent. This natural arrangement occurs as the dipole’s positive and negative ends pull each other into a stable configuration. The resulting electrostatic interactions are the essence of intermolecular forces.

Hydrogen bonds exemplify a type of intermolecular force where hydrogen bonds with an atom that is highly electronegative. The bonding partner can be fluorine, oxygen, or nitrogen, emphasizing the role of strong electronegativity. This specific interaction reveals how hydrogen couples with elements having a high tendency to attract electrons.

Van der Waals forces arise from the attraction between strong dipoles, yet non-polar molecules can also experience electrostatic interactions. Random electron movements in atoms sometimes concentrate charge on one side, forming a fleeting dipole that induces a similar polarization in nearby particles. These transient dipoles generate brief molecular attractions that dissipate as electrons continue to move.

Water’s unique dipolar structure enables it to pull apart molecules by attracting opposite charges and maintaining separation through specific intermolecular interactions, often associated with Van der Waals forces. This capability allows water to dissolve charged and polar substances effectively while excluding nonpolar molecules like those found in oil. The phenomenon underscores the principle that like dissolves like, aptly summarized by the ancient maxim 'Similia Similibus Solventur.'

Soap molecules possess a polar head that attracts water and a nonpolar tail that bonds with oil. This dual structure enables them to encase nonpolar fat molecules in micelles when mixed with water. The micelle formation suspends grease, making it easier to wash away oily substances.

Forces sorted by their strength determine the removal of dirt particles. The sequence includes Ionic Bonds, Covalent Bonds, Metallic Bonds, Hydrogen Bonds, and Van der Waals forces. The ranking highlights how robust chemical interactions overcome adhesion and regulate particle transport.

Matter is classified into three states: solid, liquid, and gas. In solids, particles remain in a fixed, tightly packed structure where they only vibrate, requiring significant force to alter the arrangement. Liquids feature particles that can move freely while staying confined due to cohesive forces that maintain a constant volume. In gases, particles possess enough energy to move without restriction, filling any available space.

Temperature measures the average kinetic energy of particles, reflecting how much and how fast they move, while entropy gauges the degree of disorder within a system. Solids form at low temperatures or high pressures due to their orderly structure and minimal particle motion. In contrast, gases occur at high temperatures or low pressures as particles move more vigorously, creating a state of increased disorder.

Ionic bonds require a substantial amount of energy to break, resulting in high melting points. Salts, which often consist of these strong bonds, remain solid at room temperature. In contrast, substances with weaker bonding, like water, exhibit much lower melting points.

High-energy environments create plasma, an ionized gas existing at extreme temperatures or high electric potentials, distinguishing it from typical hydrogen-bonded liquids. Neon lights illustrate this concept when gas inside a tube is ionized by a high voltage, triggering collisions that elevate electrons to excited states. As electrons fall back to lower energy levels, they emit energy in the form of light. The color of the emitted light is determined by each element’s fixed energy levels, with different elements producing specific hues.

Elements emit a range of frequencies known as the emission spectrum, highlighting unique characteristics of matter. Matter is divided into pure substances, which consist of one element or compound, and mixtures that contain at least two pure substances. Mixtures may be uniform, like a salt solution where substances blend evenly, or non-uniform, such as a sand suspension with distinct regions. Milk exemplifies a colloid or emulsion, where larger particles remain evenly distributed despite not fully dissolving like in a true solution.

Chemical reactions can occur in a state between complete dissolution and suspension, reflecting an intermediate form. The reaction between sodium and water demonstrates how a process can abruptly release a substantial amount of energy, leading to rapid expansion. Four main types are outlined—synthesis, decomposition, single replacement, and double replacement—with distinct examples for each. The examples collectively emphasize that these reactions primarily occur for a singular underlying reason.

Chemical reactions rely on precise ratios to create stable compounds, with water synthesis requiring twice as much hydrogen as oxygen. Stoichiometry implements the conservation of mass to ensure that atoms are neither lost nor gained during a reaction. Balancing chemical equations involves matching each element by adjusting metals first, followed by nonmetals, and finishing with hydrogen and oxygen. This systematic process minimizes energy and leads reactions toward more stable states.

A balanced reaction requires twice as many hydrogen atoms as oxygen atoms, but using the same mass for different elements doesn’t guarantee the right number of particles. Instead, the atomic mass of each reactant is used to measure out 1 mole, ensuring an accurate count of atoms. This method guarantees that the reaction achieves the exact stoichiometric proportions necessary for lab experiments.

A measured amount of 1 mole illustrates that only chemical changes trigger reactions, not physical changes. A physical change alters a substance's appearance without affecting its composition, as seen in the example of hammering metal. In contrast, a chemical change transforms the substance itself, often indicated by the emergence of bubbles or a distinctive odor. Recognizing these differences is fundamental for understanding reaction mechanisms.

Chemical reactions require a threshold of energy to initiate, which prevents spontaneous, uncontrolled events like wood burning. The necessary energy barrier ensures that processes such as combustion occur only when sufficient energy is provided. Catalysts reduce this barrier, allowing reactions to proceed more easily and quickly. Additionally, they remain unchanged throughout the reaction, making them efficient facilitators in chemical processes.

Chemical reactions involve energy transformations that can be tracked using enthalpy, the measure of a system's internal heat content. A drop in enthalpy from the beginning to the end of a reaction signifies that heat has been released, defining it as exothermic. If the energy levels reverse, it implies that energy is being absorbed. This clear energy accounting highlights the central role of enthalpy in understanding reaction dynamics.

Chemical reactions are governed by both enthalpy and entropy, where energy release drives exothermic processes like a ball rolling downhill. Endothermic reactions, which absorb energy, can still be spontaneous if a significant increase in disorder offsets the required energy, much like melting ice above 0°C. Temperature is the critical factor that determines whether the energy changes create a spontaneous (exergonic) reaction or one that requires continuous energy input (endergonic).

In exothermic reactions, equilibrium is achieved when delta G is zero, halting any spontaneous net reaction. Reversible reactions occur at identical rates in both directions, ensuring constant concentrations of reactants and products. The system remains in a state of dynamic balance where internal activity contrasts with the external appearance of stasis.

Acid-base reactions are governed by the transfer of hydrogen ions, where acids donate and bases accept protons. The Brønsted-Lowry definition establishes that any molecule with a removable hydrogen functions as an acid while any molecule capable of accepting a hydrogen ion acts as a base. This proton exchange creates conjugate pairs, such that an acid transforms into its conjugate base and vice versa. The reversible nature of these reactions underpins the dynamic behavior observed in chemical systems.

Strong acids fully dissociate, releasing many protons that form hydronium ions, while weak acids dissociate less and produce fewer hydronium ions. The pH is the negative logarithm of the hydronium ion concentration, meaning each unit change represents a tenfold difference. Pure water is in equilibrium with a hydronium concentration of 1×10^-7, giving it a neutral pH of 7, with values below 7 indicating acidity and above 7 indicating basicity. A similar logarithmic approach applied to hydroxide ions produces the pOH scale that measures basic strength.

Neutralisation involves acids and bases that counterbalance each other, their combined properties summing to 14 so that knowing one value reveals the other. When a strong acid meets a strong base, they safely transform into water and a neutral salt rather than triggering an explosion. This balance is exemplified by the reaction of hydrochloric acid with sodium hydroxide to form water and table salt.

Table salt forms through ionic bonding when electrons are transferred between elements. In this redox process, sodium loses an electron and undergoes oxidation, while chlorine gains that electron and is reduced. This electron exchange establishes the fundamental roles of oxidants and reductants, resulting in a stable compound.

Oxidation numbers serve as hypothetical charges on atoms that reveal how redox reactions alter electron distribution. Fundamental rules designate hydrogen as +1, oxygen as -2, halogens as -1, and pure elements as 0, ensuring that atom sums equal the molecule's charge. A clear example is sulfuric acid, where the charge contributions from oxygen and hydrogen mandate a corresponding oxidation number for sulfur, illustrating how oxidation states dictate electron flow in chemical processes.

Ions and water balance charges and stoichiometry in redox reactions occurring in acidic or basic solutions. Electrons defy the idea of circular orbits, instead forming probabilistic clouds defined by four quantum numbers that determine their positions within shells, subshells, and orbitals. The quantum numbers—n, l, ml, and ms—set the energy level, shape, orientation, and spin, restricting each orbital to a maximum of two electrons through the Pauli exclusion principle. The Aufbau principle then orders the filling of these subshells, while noble gas shorthand simplifies the determination of an element’s valence electrons.