General Chemistry Flashcards
(10 cards)
Periodic Law
Periodic law states that when elements are arranged in order of increasing atomic number, their chemical and physical properties recur periodically. This periodicity arises because valence‑electron configurations repeat in a fixed pattern. Consequently, groups of elements exhibit analogous reactivity and observable trends become predictable.
Names of Rows and Columns
In the periodic table, a horizontal line of elements is called a period, whereas a vertical column is called a group or family. Periods reflect the number of occupied electron shells, while groups collect elements that share identical valence‑electron configurations. Remember: period = row, group = column.
Atomic Radius and Electronegativity Trends
Atomic radius generally decreases from left to right across a period because nuclear charge grows faster than electron shielding, pulling electrons closer. It increases from top to bottom in a group because each step adds an electron shell, placing valence electrons farther away. Electronegativity shows the opposite trend: it rises across a period as atoms more strongly attract shared electrons and falls down a group as distance and shielding reduce this attraction.
Main Blocks of the Periodic Table
The periodic table is divided into s, p, d, and f blocks according to the subshell that receives the last electron. The s‑block comprises Groups 1–2, the p‑block Groups 13–18, the d‑block the transition metals (Groups 3–12), and the f‑block the lanthanides and actinides. Locating an element’s block links its position directly to its electron configuration.
Dipole and Determining Polarity
A dipole is a separation of electric charge within a molecule that makes one end slightly negative (δ−) and the other slightly positive (δ+). To find the negative pole, locate the more electronegative atom in each polar bond because electrons spend more time there. Draw an arrow pointing from the positive toward the negative centre or mark δ symbols to visualise the polarity.
Intermolecular Forces Defined
Intermolecular forces are the relatively weak attractions acting between separate molecules or formula units. They are distinct from the stronger covalent or ionic bonds within molecules and largely determine bulk properties such as boiling point, viscosity, and melting point. Their strength is typically one or two orders of magnitude smaller than that of chemical bonds.
Types of Intermolecular Forces
There are three principal intermolecular forces: London (dispersion) forces arising from temporary dipoles, dipole–dipole forces between permanent molecular dipoles, and hydrogen bonds, a strong subset of dipole–dipole interactions that occur when hydrogen is bonded to nitrogen, oxygen, or fluorine. Every molecule exhibits dispersion, but only polar molecules add the others. Hydrogen bonding can dramatically elevate boiling points.
Liquid Properties and Intermolecular Forces
The stronger the intermolecular attractions binding a liquid, the higher its boiling point, surface tension, and viscosity. Weak attractions let molecules escape easily, so the liquid evaporates readily and flows with low resistance. Macroscopic behaviour therefore directly mirrors microscopic forces.
Reversible Reactions
In a reversible reaction, products can recombine to regenerate reactants, so the process proceeds simultaneously in forward and reverse directions. The double arrow (⇌) symbol denotes this duality. Reaction conditions determine which direction predominates.
pH from [H₃O⁺] or [OH⁻]
pH is −log[H₃O⁺]; use a calculator’s log function and keep at least two decimal places. If only [OH⁻] is known, first convert to [H₃O⁺] using K₍w₎ and then take the log. A quick check is that pH+pOH = 14 at 25 °C.
