In a solution, the slightly charged part of a H2O molecule?
a)attracts the ions with ionic compounds. b)forms ionic bonds with ions in ionic compounds. c)attracts nonpolar molecules. d)forms ionic bonds with other water molecules.
In an aqueous solution of ionic salt, the oxygen atom of the water molecule is attracted to the ...?
1 negative ion if the salt , due to oxygen's partial positive charge 2 negative ion of the salt, due to oxygen's partial negative charge 3 positive ion of the salt, due to oxygen's partial positive charge 4 positive ion of the salt,due to oxygen's partial negative charge How do you get the answer? Thanks!
How does the ion exchange chromatography used in the separation of amino acids providing examples?
Ion exchange chromatography is based on the mutual attraction of oppositely charged particles.If you fill a chromatography column with a resin whose outside is covered with negative charges, e.g. R-SO3- groups and you pass a liquid containing positive and negative ions, the the positive ions will be attracted by the R-SO3- groups and the negative ions will be repelled. Thus the positive ions are retained in the column and the negative ions will flow with the liquid and leave the column.Alternatively, if you would fill the column with a resin whose outside is covered with positive charges, e.g. R-NH3+ groups, these would retain the negatively charged ions form the solution.Aminoacids have both amino-groups, that can be protonated to -NH3+ groups, and carboxylic acid groups, that can be de-protonated to -COO- groups.Wether or not amino acid particles have a net positive or negative charge depends on the pH of the solution they’re in.The pH at which the amino acid particles have a netcharge of 0 is called the isoelectric point, or pI.e.g. ALANINE has a pI = 6,11, that means that at a pH < 6,11 the particles will be more and more protonated, leading to positively charged ALA-ions (ALA+), and at a pH > 6,11 the ALA particles will be de-protonated, leading to negatively charged ALA-ions (ALA-).e.g. ARGININE has a pI = 10,76, so at pH< 10,76 there will be positively charged ARG-ions (ARG+) and at pH> 10,76, there will be negatively charged ARG-ions (ARG-)Now suppose we have a mixture of ALA and ARG, solved in a buffer solution of pH 5, then there will be ALA+ and ARG+ ions in this solution.When we introduce this mixture in a R-SO3- resin containing column, both the ALA+ and ARG+ will bind to the column.If we now flush the column with a buffer solution of pH 8, the ALA+ will change to ALA- (for pH 8 is above the pI 6,11 of ALA) and these negatively charged ions no longer will bind to the negatively charged resin particles, so they will leave the column along with the buffer solution.The ARG+ ions will keep a positive charge (pH 8 < pI 10,76 for ARG) and will still bind tot the negatively charged resin particles in the column.Thus the ALA is separated from ARG.When you now flush the column with a buffer pH 12, then the ARG- ions will also be washed out of the column and you can collect them.It’s a little simplistic, but this is the principle.
What are the factors that effect the degree of hydration of a metal ion?
The heat energy released when new bonds are made between the ions and water molecules is known as the hydration enthalpy of the ion. The hydration enthalpy is the enthalpy change when 1 mole of gaseous ions dissolve in sufficient water to give an infinitely dilute solution. Hydration enthalpies are always negative. Factors affecting the size of hydration enthalpyHydration enthalpy is a measure of the energy released when attractions are set up between positive or negative ions and water molecules.With positive ions, there may only be loose attractions between the slightly negative oxygen atoms in the water molecules and the positive ions, or there may be formal dative covalent (co-ordinate covalent) bonds.With negative ions, hydrogen bonds are formed between lone pairs of electrons on the negative ions and the slightly positive hydrogens in water molecules.The size of the hydration enthalpy is governed by the amount of attraction between the ions and the water molecules.The attractions are stronger the smaller the ion. For example, hydration enthalpies fall as you go down a group in the Periodic Table. The small lithium ion has by far the highest hydration enthalpy in Group1, and the small fluoride ion has by far the highest hydration enthalpy in Group 7. In both groups, hydration enthalpy falls as the ions get bigger. The attractions are stronger the more highly charged the ion. For example, the hydration enthalpies of Group 2 ions (like Mg2+) are much higher than those of Group 1 ions (like Na+). The alkali metal ions are highly hydrated. The smaller the size of the ion, the greater is the degree of hydration. Thus, Li+ ion gets much more hydrated than Na+ ion which in turn is more hydrated than K+ion and so on. The extent of hydration decreases down the group.As a result of larger hydration of Li+ ion than Na+ ion, the effective size of Li+ion is more than that of Na+ion. Further the ionic radii in water (called hydrated ionic radii) decreases in the order:Li+ > Na+ > K+ > Rb+ > Cs+As a result, the hydrated Li+ ion being largest ionic size, has the lowest mobility in water. On the other hand, the hydrated Cs+ ion being smallest in size has the highest mobility in water.
What properties of water enable its molecules to interact with ions in solution?
Water molecules have a permanent dipole (the O atom is slightly negative and the H atoms have slight positive charge). This allows the ions in a solution to become attracted to water molecules due to electrostatic attraction. This also allows Hydrogen bonding between other molecules.
Why does hydration enthalpy decrease on going down a group?
The size of the hydration enthalpy is governed by the amount of attraction between the ions and the water molecules.The attractions are stronger the smaller the ion. For example, hydration enthalpies fall as you go down a group in the Periodic Table. The small lithium ion has by far the highest hydration enthalpy in Group1, and the small fluoride ion has by far the highest hydration enthalpy in Group 7. In both groups, hydration enthalpy falls as the ions get bigger.The attractions are stronger the more highly charged the ion. For example, the hydration enthalpies of Group 2 ions (like Mg2+) are much higher than those of Group 1 ions (like Na+).For more details referenthalpies of solution and hydration
What is the ionic mobility order of the following cations?
I don’t know much about Francium as it is a very rare and radioactive metal and its properties have not been studied.For other ions in polar solvent like water, [math]Cs^+>Rb^+>K^+>Na^+>Li^+[/math]In water, these ions are surrounded by water molecules because water has a permanent dipole which is attracted towards positive charge of these ions.[math]Li^+[/math] has very small size compared to others and hence, its charge to radius ratio is very high. Due to this it attracts more number of water molecules and hence, due to more numbers of water molecules surrounding it, it cannot move easily in water. Hence, it has least ionic mobility.In non-polar solvents, the order is reversed: [math]Cs^+
Why can’t the OH group in NaOH release the H+ ion only instead of releasing OH- as a whole?
In water, NaOH exists mostly in dissociated form, with Na+ and OH- ions each separately surrounded by water molecules. So your question amounts to the following: why can’t the hydroxide ion dissociate, releasing a proton and leaving behind an oxide ion?The answer is that there are all sorts of reactions we don’t talk about because they are so unfavourable that they never occur to an appreciable extent. This is one of them. For a stable species will only ever relinquish a proton if there is another species that “wants” it more. The Na+ ions are already positively charged, so they will repel H+ ions, and while H+ is attracted to the oxygen atom in a water molecule, it would be much more strongly attracted to the doubly negative oxide ion that it would have to leave behind. In other words, the equilibrium constant for this reaction:OH- + H2O - - - - -> O{2-} + H3O+is extremely small; the equilibrium lies far to the left.In fact, if there were any oxide ions around, they would just tear the protons away from surrounding neutral water molecules:O{2-} + H2O - - - - -> 2OH-An ion as small and highly charged as oxide basically can’t exist in any solvent. If the solvent had dissociable protons, like water, it would immediately protonate oxide to form a more stable species. If it’s an aprotic solvent, it will never be able to exert a strong enough pull to remove protons from OH- to leave behind oxide, or to remove oxide ions from an ionic lattice such as that of sodium oxide or calcium oxide.