Measuring acidity

The acid concentration is $0.1{\text{ mol dm}}^{-3}$. Since it is a strong acid, we treat it as fully dissociated, so $[{\text{H}}^{+}]$ is also $0.1{\text{ mol dm}}^{-3}$.

Using the pH definition:

$\text{pH}=-{\log }_{10}[{\text{H}}^{+}]$

Substitute $0.1$:

$\text{pH}=-{\log }_{10}(0.1)=1$

Because the pH scale is logarithmic, for a strong acid each tenfold decrease in $[{\text{H}}^{+}]$ makes the pH go up by 1.

• Dilution by $10^1$ (10×) $\to$ pH rises by 1
• Dilution by $10^2$ (100×) $\to$ pH rises by 2
• Dilution by $10^3$ (1000×) $\to$ pH rises by 3

The dilution is stated as a factor of 1000.

Write this as a power of ten:

$1000=10^3$

A $10^3$ dilution corresponds to an increase of 3 pH units.

Start from pH 1 and add the 3-unit increase:

$\text{Final pH}=1+3=4$

Compare with the options:

• Option A (7.0) would need a much bigger dilution, like $1,000,000$ ($10^6$).
• Option C (2.0) matches only a 10× dilution.
• Option D (3.0) matches only a 100× dilution.

So, diluting a pH 1 strong acid solution by 1000 times increases the pH by 3 units, giving a final pH of 4.0.

Because the acid is strong, we assume it dissociates completely. That means the hydrogen ion concentration equals the acid concentration at the start:

$[{\text{H}}^{+}{]}_{\text{initial}}=0.1{\text{ mol dm}}^{-3}$

A dilution factor of $1000$ means the volume becomes $1000$ times larger, so the concentration becomes $1000$ times smaller:

$\begin{array}{rl}[{\text{H}}^{+}{]}_{\text{final}}& =\frac{[{\text{H}}^{+}{]}_{\text{initial}}}{1000}\\ & =\frac{0.1}{1000}\\ & =0.0001{\text{ mol dm}}^{-3}\end{array}$

This is the same as $10^{-4}{\text{ mol dm}}^{-3}$.

Apply the pH relationship:

$\text{pH}=-{\log }_{10}[{\text{H}}^{+}(aq)]$

Insert $10^{-4}$:

$\begin{array}{rl}\text{pH}& =-{\log }_{10}(10^{-4})\\ \text{pH}& =4.0\end{array}$

Now compare with the answer choices:

• Option A is wrong because pH 7.0 is neutral, and even after a $1000\times$ dilution this solution is still acidic.
• Option C would correspond to only a $10\times$ dilution ($0.01{\text{ mol dm}}^{-3}$).
• Option D would correspond to only a $100\times$ dilution ($0.001{\text{ mol dm}}^{-3}$).

Our calculated pH is 4.0, which matches Option B.

A titration is a method for getting numerical data about a solution. It is used to find an unknown concentration by reacting it with another solution whose concentration is already known.

For environmental monitoring, this lets scientists measure how acidic or basic the wastewater is by determining how much reacting substance is needed. This links to Topic 3: Types of chemical reaction, especially pH, titrations and indicators.

Here’s why two choices don’t fit the purpose of titration:

• Option A: An indicator may change color during a titration, but that color shift is only a signal for measurement. The aim is not to permanently recolor the factory waste.
• Option C: A titration tells you about chemical concentration in a small sample. It cannot be used to calculate the physical volume of an entire river or stream.

Option B is incorrect because adding more $H^{+}$ ions makes the solution more acidic. More acidity generally increases harm to aquatic organisms, so the goal is usually to reduce acidity (move pH toward 7), not raise it.

The best reason is Option D. If the wastewater is acidic, releasing it could damage ecosystems. A titration allows the scientist to determine the acid concentration accurately, and from that they can calculate the exact amount of alkali (base) needed to neutralize the waste so it can be discharged more safely.

To answer this, focus on the purpose of a titration. It is a quantitative lab procedure used to determine an unknown concentration in a solution. This is a core idea in MYP Chemistry within Types of chemical reaction: pH, titrations and indicators.

Indicators can change color in titrations, but that change is only used to show the endpoint has been reached. Environmental testing is about chemical safety and correct pH, not improving appearance.

A higher $text{H}^{+}$ concentration means a lower pH, so the solution becomes more acidic. Strongly acidic water can damage fish and plants, so deliberately increasing acidity would not usually increase safety.

Titration is done on a measured lab sample (for example, a $25{\text{ cm}}^3$ portion). It provides information about chemical concentration, not the size or volume of a natural waterway.

A titration reveals how acidic the wastewater is. Using that result, the scientist can calculate precisely how much alkali (a base) is needed to neutralize the acid so the water reaches a safer pH before release.

We are asked about the colour bromothymol blue shows in a solution with a $pH$ of $7.0$. A $pH$ of $7.0$ means the solution is neutral (neither acidic nor basic).

Bromothymol blue is a standard indicator we learn about in Topic: pH, titrations and indicators. Its colours depend on pH:

• Acidic form ($pH<6.0$): Yellow
• Basic form ($pH>7.6$): Blue
• Transition range ($pH$ $6.0$ to $7.6$): An in-between colour made from both

Because $pH=7.0$ lies inside the transition range, the indicator is partway between its acidic and basic forms.

Using what we know about the indicator:

• Option A cannot be right because blue is only seen clearly in more basic solutions ($pH>7.6$).
• Option C cannot be right because yellow is the colour in acidic conditions ($pH<6.0$).
• Option D cannot be right because bromothymol blue doesn’t behave as a colourless indicator (that idea fits indicators like phenolphthalein in acids, not this one).

In the transition region, the observed colour looks like a mix of the two end colours. Combining yellow (acid colour) with blue (base colour) gives green.

So at $pH=7.0$, bromothymol blue appears green.

The question tells us how bromothymol blue looks at the extremes:

• Acidic solutions: Yellow
• Basic solutions: Blue

We need the colour at pH 7.0. On the pH scale, $7.0$ is neutral—right between values below $7.0$ (acidic) and above $7.0$ (basic).

At the neutral point ($pH=7.0$), the indicator is not fully in its acidic form or fully in its basic form. That means the colour we see is a blend of the two main colours rather than one extreme.

From the blending idea:

• Option A is wrong because blue corresponds to basic conditions ($pH>7.0$).
• Option C is wrong because yellow corresponds to acidic conditions ($pH<7.0$).
• Option D is wrong because bromothymol blue is not known for being colourless like phenolphthalein.

Since $pH7.0$ is the midpoint where the indicator is transitioning, the expected colour is green—the combined appearance of yellow and blue.

The prompt asks about bromothymol blue in a neutral solution, which is defined as having a $pH=7.0$. This is covered in Topic 3: Types of chemical reaction when discussing pH and neutral solutions.

So we’re looking for the colour at the point between the acidic and basic behaviours described.

We’re told bromothymol blue is yellow in acidic solutions. Since $pH$ 7.0 is neutral (not acidic, meaning not $pH<7$), we can remove the yellow option.

We’re also told the indicator is blue in basic solutions. A $pH$ of 7.0 is neutral (not basic, meaning not $pH>7$), so the blue option should be eliminated too.

Even though some indicators can be colourless in certain conditions, bromothymol blue is known for showing visible colours around neutral pH. So “colourless” is not a sensible match here.

With the extreme colours and the unlikely property removed, the remaining answer is green.

That also matches the idea that the neutral point is between yellow (acid) and blue (base), giving $yellow+blue=green$.

A buffer is identified by what it does: it helps keep the $pH$ from changing much when small quantities of acid or base (the titrant) are added. Knowing this idea is key for understanding Types of chemical reaction: Acids and bases.

On a titration curve, the fastest pH change happens in the steep, almost vertical part close to the equivalence point. Because a buffer is meant to reduce pH changes, a “most rapid change” description does not match a buffer region.

A lasting indicator color change is used to show the end point of a titration, which is usually chosen to be near the equivalence point. That observation is useful for deciding when to stop, but it is not what defines the buffering part of the curve.

The smallest $pH$ is usually found right at the beginning if the analyte starts as a strong acid. The buffer region is not the start; it appears later (especially in titrations involving a weak acid or weak base), when the mixture contains a weak species together with its conjugate partner.

Using the definition of a buffer, the buffer region matches the ‘plateau-like’ section of the curve. In that part, the $pH$ changes only slightly because the weak acid/base and its conjugate form react with and “mop up” small added amounts of titrant.

In the topic of pH, titrations and indicators, graphs help us see how $pH$ changes as titrant is added. To spot the 'buffer region', we focus on how steep or flat the curve is (its gradient).

A buffer region is commonly seen when a weak acid is titrated with a strong base (or the reverse). On a plot with pH on the $y$-axis and volume of titrant on the $x$-axis, the buffer region is the section that looks fairly level rather than steep.

With Graphical Analysis, we search for the part where the pH changes the least for a given added volume, meaning $

$pH$ is small while $

$V$ is the same.

Mathematically, the gradient is defined as: $\text{gradient}=\frac{dpH}{dV}$

In a buffer region, both a weak acid and its conjugate base are present together. This combination reduces the effect of added acid/base, so on the graph the curve shows a small gradient (a flatter section).

Option A describes the pH changing extremely quickly. On the graph, that would be a very steep (almost vertical) part of the curve. That steep jump happens at the equivalence point, not in the buffer region, so Option A should be removed.

Consider the other incorrect choices:

• Option C: The permanent indicator color shift signals the 'end point', which is chosen to match the equivalence point area (the steep part), not the flatter buffer section.
• Option D: The lowest pH (if starting with an acid) occurs right at the beginning. That is a single starting condition, not a whole region where the curve stays relatively flat.

Neither option matches the low-slope part of the curve associated with buffering.

Using the curve shape and elimination, the buffer region is the stable part of the graph. It corresponds to the nearly horizontal section where adding more titrant causes only a tiny change in pH (very little vertical movement).

For process of elimination, start by remembering what happens when a weak acid is titrated with a strong base. The buffer region is found after some titrant has been added but before the equivalence point. At this stage, there are significant amounts of both the weak acid ($HA$) and its conjugate base ($A^{-}$), which helps the solution resist changes in $pH$.

Option A refers to the part of the curve where the pH rises (or falls) very quickly, giving an almost vertical slope. That behavior belongs to the equivalence point region, not the buffer region, which is the flattest part of the curve.

Option C is about the end point—the visible indicator change used to decide when to stop the titration. Although the end point is chosen to be close to the equivalence point, it does not describe the plateau-like buffering section.

Option D is describing the starting pH of the analyte. If the titration begins with an acid, the pH is lowest at 0 mL titrant. Once base is added, the pH begins increasing, so the buffer region cannot be the minimum pH point.

After ruling out the others, Option B remains. This matches what a buffer does: it keeps $pH$ from changing much when small amounts of acid or base are added. On the titration curve, this shows up as a flatter section where the pH stays almost constant.