what is the meaning of weak base and weak acid and what is the relation between it and the absorbtion ?
610740
Your blood keeps its pH range stable.
1034417
the term 'besylate form' means that a substance has basic properties ('basic': from the word 'base', opposite to 'acid', 'acidic') but it is formulated in a way to be a neutral salt with a reaction with another substance owning particular structure which called 'besylate'. So after the reaction, the former substance will be a neutral salt in 'besylate form'.\n\ne.g.: \namlodipine (a hypotensive agent) is a base, and so it is difficult to dissolve in water. this property is undesirable in phamaceutics because it makes the absorption of the drug to be minimum. \nto counter this negative property, scientist made a dissolvable salt of this, using another substance with besylate structure. this combination makes the former drug more stable and dissolve well in the stomach so the absorption also reach the optimum level.
378003
pH for a salt of strong base and weak acid is given by:\n         pH = 7 - (1/2)pKb - (1/2)log C\nwhere pKb = dissociation const of strong base\n       C  = concentration of Acid.\n\nKb = Kw / Ka = (10^-14) / (6.5 x 10^-5) = 1.53x10^-10\npKb = -log (1.53x10^-10) = 9.8154\nlog C = log 10^-1 = -1\n\npH = 7 - (1/2)x 9.8154 - (1/2)x-1\n   = 7 - 4.9077 + 0.5\n   = 2.5923
2508
Weak acids and organic acids in water.\n\nBut remember the solution depends on what the solvent is.  In water, the seven strong acids essentially completely ionize...i.e. can be considered completely ionized.
394317
acids are sour and bases are bitter.\nacids turn blue litmus red and its the other way round in bases
451857
Chemicals used in the reaction is important. What is the pKa of the reactant? Research on which buffer and what its pH buffering region is important, because it should cover the pKa region of the chemicals used. Once knowing which buffer to make, we can take a buffer by making a mixture that have a weak acid with its salt of conjugate base, or weak base with its conjugate acid. It should be noted that here buffer solutions are presented in terms of the Brønsted-Lowry notion of acids and bases.
829401
Basically, water in pure water the water molecules will undergo autoionisation which will produce to types of ions namely H+ and OH- ions. Anyway it is a reversible reaction where H2O exist in equilibrium with H+ and OH- ions. \n\nWhen a strong acid is dissolved in water, it will disscociate completely to produce H+ ions and the anions. This is a forward and non-reversible reaction. Whereas when a weak acid is dissolve in water only part of it will dissociate and the degree of dissociation depends on is solubility constant. This is a reversible reaction.\n\nWhen an acid is mix with water both the acid and the water will contribute its proton. ( acid dissociate completely and water partially.) But the proton from water is so little and can be neglected so normally the concentration of the H+ ions of the acid is used for calculation.\n\nHope this helps.... Good luck..
719257
Acid rain is defined as any type of precipitation with a pH that is unusually low (Brimblecombe, 1996). Dissolved carbon dioxide dissociates to form weak carbonic acid giving a pH of approximately 5.6 at typical atmospheric concentrations of CO2 (Seinfeld and Pandis, 1998). Therefore a pH of <5.6 has sometimes been used as a definition of acid rain [3]. However, natural sources of acidity mean that in remote areas rain has a pH which is between 4.5 and 5.6 with an average value of 5.0 and so rain with a pH <5 is a more appropriate definition
1029104
acid has a ph value below 7\nand alkali has a value above 7\n\nAn acid (often represented by the generic formula HA) is traditionally considered any chemical compound that when dissolved in water, gives a solution with a pH of less than 7. That approximates the modern definition of Brønsted and Lowry, who defined an acid as a compound which donates a hydrogen ion (H+) to another compound (called a base). Common examples include acetic acid (in vinegar) and sulfuric acid (used in car batteries). Acids generally taste sour;\n\nIn chemistry, an alkali (from Arabic: al-qaly القالي) is a specific type of base, formed as a carbonate, hydroxide or other basic (pH greater than 7) ionic salt of an alkali metal or alkali earth metal element. The word alkali or the adjective alkaline are frequently used to refer to all bases, since most common bases are alkalis, although strictly speaking this is inaccurate.
698593
You are not giving us all of the info.\nYes the total or initial concentration of the acid is C=(8.345*10^-3)/0.05 = 0.1669 M\n\nFor a monoprotic acid Ka=[H+][A-]/{HA]\n[A-]=[H+]=10^-pH=10^-4\n[HA]=C-[H+]= 0.1669-0.0001= 0.1668\nso you have calcualted Ka correctly.\n\nKa depends on the nature of the acid and temperature (for small molecules whereas for bigger molecules like proteins it is also affected by the nature of the surrounding chemical groups).\n \nThus addition of base doesn't affect Ka, only the pH.\nIs the base monoprotic or not? If not, what's the valency? Is it strong or weak? If it's weak what's the Kb? Is it the conjugate base of the weak acid? Without these info we can't calculate the pH.
969526
Bases to neutralize acids,,;-)=
117085
Buffering, in chemical terms, is a weak acid and its cooresponding salt or a weak base and its cooresponding salt.  The idea is that the weak acid will go to an equilibrium like the following example.  Suppose that I have a weak acid HA and the salt with the anion A-.  The acid will go into equilibrium:\nHA <---> H+ + A-\nWhat will happen is that if you add H+ ions to the buffer solution (acids will add H+ ions), they will be consumed into the reaction to become more HA.  If you add a base (adding OH- ions), it will react away H+ to form water and force more A- to form.  Either way, this buffer system will resist a change in pH.  That's what buffers do. \nBuffers are essential in our body because even though we consume all of these orange juices and sodas, which are both acidic, our body needs to maintain a very precise slightly basic pH, or else we could die.  This is done with CO2 in solution in our bodies.
279663
Well, Sulfuric acid is a diprotic acid. The first proton is a strong acid, and the second proton is a fairly strong weak acid.\n\nBecause it's a strong acid, there will be 18 M of H3O+ ions.\n\n-log(18) = -1.255\n\nTo the correct number of significant figures, the pH is -1.26.\n\nEDIT: dr_futuredentist is very wrong. The second hydrogen\ndoes not dissociate fully at all.
618488
Look at their conjugate bases and whichever is more stable will have the weakest acid. PKA values are also easy to prove this becuase it is a realtion of the H+ content.
1057594
I think this is how you do it:\nacid and base is measured in terms of pH and pOH. as you know, the H makes it acidic and the OH makes it basic, and if you mix an H and OH, you get water, which is to say it was neutralized.\nSo, let's say you have HBr with a pH of 4. You that pH is a measure of the concentration of H present in the mixture. So, to get the actual number you do -10ph, and you'll get a number like 10^-5 or something. So, to neutralize that many Hs, that's how many OH's you need. so, you'll need a pOH 4, which is the same thing is a pH of 10 (pOH = 14 - pH). \nhope that helps, good luck. and read the book, and do some of hte exercised in the book.
609380
The Bronsted Lowry Theory of acids and bases is the receiving (base) or the donating (acid) of a "proton" - in effect a Hydrogen ion H+\n\nThe Bronsted-Lowry definition is named for Johannes Bronsted and Thomas Lowry, who independently proposed it in 1923. A Bronsted-Lowry (BL) acid is defined as any substance that can donate a hydrogen ion (proton) and a Bronsted-Lowry base is any substance that can accept a hydrogen ion (proton). Thus, according to the BL definition, acids and bases must come in what is called conjugate pairs. \n\nFor example, consider acetic acid dissolved in water: \n\nCH3COOH (l) + H2O (l) <=> CH3COO-(aq) + H3O+(aq)\n\nBL acid       BL Base     Conjugate base  Conjugate Acid\n\nThe presence of H3O+(aq) in concentrations greater than OH-(aq) is characteristic of a solution (in water) which has a pH less than 7 when you are dealing with normal acid-base theory\n\nBronsted Lowry does not consider the concentrations of either; simply the transfer of the H+.\n\nCheck the link for more info - it is easier to read and a good explanantion
18816
An acid, according to the Bronsted definition of acids.\n\n"Proton" is semi-correct. When a Bronsted acid gives up a hydrogen ion, in the balance there is a proton missing, but at no time the hydrogen ion will have been without electrons in the orbital. Usually, the hydrogen ion is transferred to a solvent molecule which then spreads the positive charge to the adjacent atom which donated two electrons to create a covalent bond with the hydrogen.