How is the concentration of hydroxide ions affected when excess base is dissolved in a solution of

so at the end (last problem), the PH value for neutral (equal concentration of H+ and OH-) water @ 50°C is 6.64. I've always assumed that only PH=7 is neutral. Does the neutral PH value of everything depend on temperature? In other words, does the [H+]=[OH-] concentration equilibrium change depending on the temperature? So no matter the temperature, the same rule applies for neutral water in that the number of hydronium ions equals the number of hydroxide ions. A pH (little p) of 7 only means neutral water at 25°C, but for other temperatures this means a pH above or below 7. This is due to the changing value of water's self-ionization constant, Kw, with temperature. At 25°C we know it to be 1.0 x 10^(-14), but at something like 50°C it's about 5.5 x 10^(-14). When we do the math using the Kw at 50°C to find the pH and pOH we find that they are both lower than 7, simply meaning that there more hydronium and hydroxide ions in neutral water at higher temperatures. Hope that helps. If you mean how does he solve the equation around that time, he's using antilog. An antilog is how you would undo a logarithm by making both sides of the equation exponents to a number equal the value of the logarithm's base, in this case 10. After that, 10^(-4.75) is just a math operation so we can have a calculator do for us. Hope that helps. Jay is exponentiating the equation to solve it. Exponentiation undoes a logarithm since exponential and logarithmic functions are inverses of each other. Exp...

The concentration of hydroxide ions will increase when excess base is dissolved in a solution of sodium hydroxide because the amount of hydroxide ions per unit volume increases. This happens only when base added dissolves in water. If the base is not soluble in water, the concentration of hydroxide ions remains constant.

\( \newcommand\) • • • • • • • • • • • • • • • • • • • • • • • • • • • • • • • Learning Objectives • Assess the relative strengths of acids and bases according to their ionization constants • Rationalize trends in acid–base strength in relation to molecular structure • Carry out equilibrium calculations for weak acid–base systems We can rank the strengths of acids by the extent to which they ionize in aqueous solution. The reaction of an acid with water is given by the general expression: \[\ce \nonumber \] where the concentrations are those at equilibrium. As noted in the section on equilibrium constants, although water is a reactant in the reaction, it is the solvent as well, soits activityhas a value of 1, which does not change the value of \(K_a\). Note It is a common error to claim that the molar concentration of the solvent is in some way involved in the equilibrium law. This error is a result of a misunderstanding of solution thermodynamics. For example, it is often claimed that K a= K eq[H 2O] for aqueous solutions. This equation is incorrect because it is an erroneous interpretation of the correct equation K a= K eq(\(\textit\) = 1 for a dilute solution, K a= K eq(1), orK a= K eq. The larger the \(K_a\) of an acid, the larger the concentration of \(\ce \] Because the ratio includes the initial concentration, the percent ionization for a solution of a given weak acid varies depending on the original concentration of the acid, and actually decreases with increasing ...

\( \newcommand\) • • • • • • • • • • • • • • • • • • • • • • • • • • Learning Objectives • To describe the concentrations of solutions quantitatively Many people have a qualitative idea of what is meant by concentration. Anyone who has made instant coffee or lemonade knows that too much powder gives a strongly flavored, highly concentrated drink, whereas too little results in a dilute solution that may be hard to distinguish from water. In chemistry, the concentration of a solution is the quantity of a solute that is contained in a particular quantity of solvent or solution. Knowing the concentration of solutes is important in controlling the stoichiometry of reactants for solution reactions. Chemists use many different methods to define concentrations, some of which are described in this section. Molarity The most common unit of concentration is molarity, which is also the most useful for calculations involving the stoichiometry of reactions in solution. The molarity (M) is defined as the number of moles of solute present in exactly 1 L of solution. It is, equivalently, the number of millimoles of solute present in exactly 1 mL of solution: \[ molarity = \dfrac\): Preparation of a Solution of Known Concentration Using a Solid Solute Calculations Involving Molarity (M): [youtu.be] Concentrations are also often reported on a mass-to-mass (m/m) basis or on a mass-to-volume (m/v) basis, particularly in clinical laboratories and engineering applications. A concentration expres...