Acid And Base Litmus Paper Essay

CONCEPT

To an extent, acids and bases can be defined in terms of factors that are apparent to the senses: edible acids taste sour, for instance, while bases are bitter-tasting and slippery to the touch. The best way to understand these two types of substances, however, is in terms of their behavior in chemical reactions. Not only do the reactions of acids and bases result in the creation of salts and water, but acids and bases can be defined by the ways in which they participate in a reaction—for instance, by donating or accepting electron pairs. The reaction of acids and bases to form water and salts is called neutralization, and it has a wide range of applications, including the promotion of plant growth in soil and the treatment of heartburn in the human stomach. Neutralization also makes it possible to test substances for their pH level, a measure of the degree to which the substance is acidic or alkaline.

HOW IT WORKS

Phenomenological Definitions of Acids and Bases

Before studying the reactions of acids and bases, it is necessary to define exactly what each is. This is not as easy as it sounds, and the Acids and Bases essay discusses in detail a subject covered more briefly here: the arduous task chemists faced in developing a workable distinction. Let us start with the phenomenological differences between the two—that is, aspects relating to things that can readily be observed without referring to the molecule properties and behaviors of acids and bases.
Acids are fairly easy to understand on the phenomenological level: the name comes from the Latin term acidus, or “sour,” and many sour substances from daily life—lemons, for instance, or vinegar—are indeed highly acidic. In fact, lemons and most citrus fruits contain citric acid (C6H8O7), while the acidic quality of vinegar comes from acetic acid (CH3COOH). In addition, acids produce characteristic colors in certain vegetable dyes, such as those used in making litmus paper.
The word “base,” as it is used in this context, may be a bit more difficult to appreciate on a sensory level. It helps, perhaps, if the older term “alkali” is used, though even so, people tend to think of alkaline substances primarily in contrast to acids. “Alkali,” which serves to indicate the basic quality of both the alkali metal and alkaline earth metal families of elements, comes from the Arabic aZ-qiZi. The latter refers to the ashes of the seawort plant, which usually grows in marshy areas and, in the past, was often burned to produce soda ash for making soap.
The reason chemists of today use the word “base” instead of “alkali” is that the latter term has a narrower meaning: all alkalies are bases, but not all bases are alkalies. Originally referring only to the ashes of burned plants containing either sodium or potassium, alkali was eventually used to designate the soluble hydroxides of the alkali and alkaline earth metals. Among these are sodium hydroxide or lye; magnesium hydroxide (found in milk of magnesia); potassium hydroxide, which appears in soaps; and other compounds. Because these represent only a few of the substances that react with acids in the ways discussed in this essay, the term “base” is preferred.


THE FORMATION OF SALTS

chemistry evolved, and physical scientists became aware of the atomic and molecular substructures that make up the material world, they developed more fundamental distinctions between acids and bases. By the early twentieth century, chemists had applied structural distinctions between acids and bases—that is, definitions based on the molecular structures and behaviors of those substances.
An important intermediary step occurred as chemists came to the conclusion that reactions of acids and bases form salts and water. For instance, in an aqueous solution, hydrochloric acid or HCl(aq) reacts with the base sodium hydroxide, designated as NaOH(aq), to form sodium chloride, or common table salt (NaCl[aq]) and H2O. What happens is that the sodium (Na) ion (an atom with an electric charge) in sodium hydroxide switches places with the hydrogen ion in hydrochloric acid, resulting in the creation of NaCl and water.
Ions themselves had yet to be defined in 1803, when the great Swedish chemist Jons Berzelius (1779-1848) added another piece to the foundation for a structural definition. Acids and bases, he suggested, have opposite electric charges. In this, he was about eight decades ahead of his time: only in 1884 did his countryman Svante Arrhenius (1859-1927) introduce the concept of the ion. This, in turn, enabled Arrhe-nius to formulate the first structural distinction between acids and bases.

The Arrhenius Acid-Base Theory

Arrhenius acid-base theory defines the two substances with regard to their behavior in an aqueous solution: an acid is any compound that produces hydrogen ions (H+), and a base is one that produces hydroxide ions (OH-) when dissolved in water. This occurred, for instance, in the reaction discussed above: the hydrochloric acid produced a hydrogen ion, while the sodium hydroxide produced a hydroxide ion, and these two ions bonded to form water.
Though it was a good start, Arrhenius’s theory was limited to reactions in aqueous solutions. In addition, it confined its definition of acids and bases only to those ionic compounds, such as hydrochloric acid or sodium hydroxide, that produced either hydrogen or hydroxide ions. But ammonia, or NH3, acts like a base in aqueous solutions, even though it does not produce the hydroxide ion. These shortcomings pointed to the need for a more comprehensive theory, which came with the formulation of the Br0nsted-Lowry definition.

The Br0nsted-Lowry Acid-Base Theory

Developed by English chemist Thomas Lowry(1874-1936) and Danish chemist J. N. Br0nsted (1879-1947), the Bronsted-Lowry acid-base theory defines an acid as a proton (H+) donor, and a base as a proton acceptor, in a chemical reaction. Protons are represented by the symbol H+,a cation (positively charged ion) of hydrogen.
Elemental hydrogen, called protium to distinguish it from its isotopes, has just one proton and one electron—no neutrons. Therefore, the hydrogen cation, which has to lose its sole electron to gain a positive charge, is essentially nothing but a proton. It is thus at once an atom, an ion, and a proton, but the ionization of hydrogen constitutes the only case in which this is possible.
Thus when the term “proton donor” or “proton acceptor” is used, it does not mean that a proton is splitting off from an atom or joining another, as in a nuclear reaction. Rather, when an acid behaves as a proton donor, this means that the hydrogen proton/ion/atom is separating from an acidic compound; conversely, when a base acts as a proton acceptor, the positively charged hydrogen ion is bonding with the basic compound.

Reactions in br0nsted-lowry acid-base theory

In representing Br0nsted-Lowry acids and bases, the symbols HA and A-, respectively, are used. These appear in the equation representing the most fundamental type of Br0nsted-Lowry acid-base reaction: HA(aq) + H2O(Z) 4H3O+(aq) + A-(aq). The symbols (aq), (Z), and 4are explained in the Chemical Reactions essay. In plain English, this equation states that when an acid in an aqueous solution reacts with liquid water, the result is the creation of H3O+, known as the hydronium ion, along with a base. Both products of the reaction are dissolved in an aqueous solution.
Because water molecules are polar, the negative charges tend to congregate on one end of the molecule with the oxygen atom, while the positive charges remain on the other end with the hydrogen atoms. The Br0nsted-Lowry model emphasizes the role played by water, which pulls the proton from the acid, resulting in the creation of the hydronium ion.
The hydronium ion, in this equation, is an example of a conjugate acid, an acid formed when a base accepts a proton. At the same time, the acid has lost its proton, becoming A-,a conjugate base—that is, the base formed when an acid releases a proton. These two products of the reaction are called a conjugate acid-base pair, a term that refers to two substances related to one another by the donating of a proton.
Br0nsted and Lowry’s definition includes all Arrhenius acids and bases, as well as other chemical species not encompassed in Arrhenius theory. As mentioned earlier, ammonia is a base, yet it does not produce OH- ions; however, it does accept a proton from a water molecule. Water can serve either as an acid or base; in this instance, it is an acid, and in reaction with ammonia, it produces the conjugate acid-base pair of NH4+ (an ammonium ion) and OH-. Ammonia did not produce the hydroxide ion here; rather, OH- is the conjugate base that resulted when the water molecule lost its H+ atom (i.e., a proton.)

The Lewis Acid-Base Theory

The Br0nsted-Lowry model still had its limitations, in that it only described compounds containing hydrogen. American chemist Gilbert N. Lewis (1875-1946), however, developed a theory of acids and bases that makes no reference to the presence of hydrogen. Instead, it relates to something much more fundamental: the fact that chemical bonding always involves pairs of electrons.
Lewis acid-base theory defines an acid as the reactant that accepts an electron pair from another reactant in a chemical reaction, while a base is the reactant that donates an electron pair to another reactant. Note that, as with the Br0nsted-Lowry definition, the Lewis definition is reaction-dependant. Instead of defining a compound as an acid or base in its own right, it identifies these in terms of how the compound reacts with another.
The Lewis definition encompasses all the situations covered by the others, as well as many other reactions not described in the theories of either Arrhenius or Br0nsted-Lowry. In particular, Lewis theory can be used to differentiate the acid and base in chemical reactions where ions are not produced, something that takes it far beyond the scope of Arrhenius theory. Also, Lewis theory addresses situations in which there is no proton donor or acceptor, thus offering an improvement over Br0nsted-Lowry.
When boron trifluoride (BF3) and ammonia (NH3), both in the gas phases, react to produce boron trifluoride ammonia complex (F3BNH3), boron trifluoride accepts an electron pair. Therefore, it is a Lewis acid, while ammonia—which donates the electron pair—can be defined as a Lewis base. This particular reaction involves hydrogen, but since the operative factor in Lewis theory relates to electron pairs and not hydrogen, the theory can be used to address reactions in which that element is not present.

REAL-LIFE APPLICATIONS

Dissociation

Dissociation is the separation of a molecule into ions, and it is a key factor for evaluating the “strength” of acids and bases. The more a substance is prone to dissociation, the better it can
conduct an electric current, because the separation of charges provides a “pathway” for the current’s flow. A substance that dissociates completely, or almost completely, is called a strong electrolyte, whereas one that dissociates only slightly (or not at all) is designated as a weak electrolyte.
The terms “weak” and “strong” are also applied to acids and bases. For instance, vinegar is a weak acid, because it dissociates only slightly, and therefore conducts little electric current. By contrast, hydrochloric acid (HCl) is a strong acid, because it dissociates almost completely into positively charged hydrogen ions and negatively charged chlorine ones. Represented symbolically, this is: HCl 4H+ + Cl-.

A reaction involving a strong acid

It may seem a bit backward that a strong acid or base is one that “falls apart,” while the weak one stays together. To understand the difference better, let us return to the reaction described earlier, in which an acid in aqueous solution reacts with water to produce a base in aqueous solution, along with hydronium: HA(aq) + H2O(l) 4H3O+(aq) + A-(aq). Instead of using the generic symbols HA and A-,however, let us substitute hydrochloric acid (HCl) and chloride (Cl-) respectively.
The reaction HCl(aq) + H2O(l) 4H3O+(aq) + Cl-(aq) is a reversible one, and for that reason, the symbol for chemical equilibrium (!) can be inserted in place of the arrow pointing to the right. In other words, the substances on the right can just as easily react, producing the substances on the left. In this reverse reaction, the reactants of the forward reaction would become products, and the products of the forward reaction serve as the reactants.
However, the reaction described here is not perfectly reversible, and in fact the most proper chemical symbolism would show a longer arrow pointing to the right, with a shorter arrow pointing to the left. Due to the presence of a strong electrolyte, there is more forward “thrust” to this reaction.
Because it is a strong acid, the hydrogen chloride in solution is not a set of molecules, but a collection of H+ and Cl- ions. In the reaction, the weak Cl- ions to the right side of the equilibrium symbol exert very little attraction for the H+ ions. Instead of bonding with the chloride, these hydrogen ions join the water (a stronger base) to form hydronium.
The chloride, incidentally, is the conjugate base of the hydrochloric acid, and this illustrates another principal regarding the “strength” of electrolytes: a strong acid produces a relatively weak conjugate base. Likewise, a strong base produces a relatively weak conjugate acid.

The strong acids and bases

There are only a few strong acids and bases, which are listed below:
Strong Acids
• Hydrobromic acid (HBr)
• Hydrochloric acid (HCl)
• Hydroiodic acid (HI)
• Nitric acid (HNO3)
• Perchloric acid (HClO4)
• Sulfuric acid (H2SO4)
Strong Bases
• Barium hydroxide (Ba[OH]2)
• Calcium hydroxide (Ca[OH]2)
• Lithium hydroxide (LiOH)
• Potassium hydroxide (KOH)
• Sodium hydroxide (NaOH)
• Strontium hydroxide (Sr[OH]2) Virtually all others are weak acids or bases,meaning that only a small percentage of molecules in these substances ionize by dissociation. The concentrations of the chemical species involved in the dissociation of weak acids and bases are mathematically governed by the equilibrium constant K..

Neutralization

Neutralization is the process whereby an acid and base react with one another to form a salt and water. The simplest example of this occurs in the reaction discussed earlier, in which hydrochloric acid or HCl(aq) reacts with the base sodium hydroxide, designated as NaOH(aq), in an aqueous solution. The result is sodium chloride, or common table salt (NaCl[aq]) and H2O. This equation is written thus: HCl(aq) + NaOH(aq) 4NaCl(aq) + H2O.
The human stomach produces hydrochloric acid, commonly known as “stomach acid.” It is generated in the digestion process, but when a person eats something requiring the stomach to work overtime in digesting it—say, a pizza—the stomach may generate excess hydrochloric acid,and the result is “heartburn.” When this happens, people often take antacids, which contain a base such as aluminum hydroxide (Al[OH]3) or magnesium hydroxide (Mg[OH]2).
When a person takes an antacid, the reaction leads to the creation of a salt, but not the salt with which most people are familiar—NaCl. As shown above, that particular salt is the product of a reaction between hydrochloric acid and sodium hydroxide, but a person who ingested sodium hydroxide (a substance used to unclog drains and clean ovens) would have much worse heartburn than before! In any case, the antacid reacts with the stomach acid to produce a salt, as well as water, and thus the acid is neutralized.
When land formerly used for mining is reclaimed, the acidic water in the area must be neutralized, and the use of calcium oxide (CaO) as a base is one means of doing so. Acidic soil, too, can be neutralized by the introduction of calcium carbonate (CaCO3) or limestone, along with magnesium carbonate (MgCO3). If soil is too basic, as for instance in areas where there has been too little precipitation, acid-like substances such as calcium sulfate or gypsum (SaSO4) can be used. In either case, neutralization promotes plant growth.

Titration and ph

One of the most important applications of neutralization is in titration, the use of a chemical reaction to determine the amount of a chemical substance in a sample of unknown purity. In a typical form of neutralization titration, a measured amount of an acid is added to a solution containing an unknown amount of a base. Once enough of the acid has been added to neutralize the base, it is possible to determine how much base exists in the solution.
Titration can also be used to measure pH (“power of hydrogen”) level by using an acid-base indicator. The pH scale assigns values ranging from 0 (a virtually pure acid) to 14 (a virtually pure base), with 7 indicating a neutral substance. An acid-base indicator such as litmus paper changes color when it neutralizes the solution.
The transition interval (the pH at which the color of an indicator changes) is different for different types of indicators, and thus various indicators are used to measure substances in specific pH ranges. For instance, methyl red changes from red to yellow across a pH range of 4.4 to 6.2, so it is most useful for testing a substance suspected of being moderately acidic.

Buffered solutions

A buffered solution is one that resists a change in pH even when a strong acid or base is added to it. This buffering results from the presence of a weak acid and a strong conjugate base, and it can be very important to organisms whose cells can endure changes only within a limited range of pH values. Human blood, for instance, contains buffering systems, because it needs to be at pH levels between 7.35 and 7.45.
The carbonic acid-bicarbonate buffer system is used to control the pH of blood. The most important chemical equilibria (that is, reactions involving chemical equilibrium) for this system are: H+ + HCO3- ! H2CO3 ! H20 + CO2.In other words, the hydrogen ion (H+) reacts with the hydrogen carbonate ion (HCO3- ) to produce carbonic acid (H2CO3). The latter is in equilibrium with the first set of reactants, as well as with water and carbon dioxide in the forward reaction.
The controls the pH level by changing the concentration of carbon dioxide by exhalation. In accordance with Le Chatelier’s principle, this shifts the equilibrium to the right, consuming H+ ions. In normal blood plasma, the concentration of HCO3- is about 20 times as great as that of H2CO3, and this large concentration of hydrogen carbonate gives the buffer a high capacity to neutralize additional acid. The buffer has a much lower capacity to neutralize bases because of the much smaller concentration of carbonic acid.

Water: Both Acid and Base

Water is an amphoteric substance; in other words, it can serve either as an acid or a base. When water experiences ionization, one water molecule serves as a Br0nsted-Lowry acid, donating a proton to another water molecule— the Br0nsted-Lowry base. This results in the production of a hydroxide ion and a hydronium ion: H2O(Z) + H2O(Z) ! H3O+(aq) + OH-(aq).
This equilibrium equation is actually one in which the tendency toward the reverse reaction is much greater; therefore the equilibrium symbol, if rendered in its most proper form, would show a much shorter arrow pointing toward the right. In water purified by distillation, the concentrations of hydronium (H3O+) and hydroxide (OH-) ions are equal. When multiplied by one another, these yield the constant figure 1.0 • 10-14, which is the equilibrium constant for water. In fact, this constant—denoted as Kw—is called the ion-product constant for water.

KEY TERMS

Acid: A substance that in its edible form is sour to the taste, and in non-edible forms is often capable of dissolving metals. Acids and bases react to form salts and water. These are all phenomenological definitions, however, in contrast to the three structural definitions of acids and bases— the Arrhenius, Br0nsted-Lowry, and Lewis acid-base theories.
Alkali: A term referring to the soluble hydroxides of the alkali and alkaline earth metals. Once “alkali” was used for the class of substances that react with acids to form salts; today, however, the more general term base is preferred.
Alkalinity: An adjectival term used to identify the degree to which a substance displays the properties of a base.
Amphoteric: A term describing a substance that can serve either as an acid or a base. Water is the most significant amphoteric substance.
Aqueous solution: A substance in which water constitutes the solvent. A large number of chemical reactions take place in an aqueous solution.
Arrhenius acid-base theory:The first of three structural definitions of acids and bases. Formulated by Swedish chemist Svante Arrhenius (1859-1927), the Arrhenius theory defines acids and bases according to the ions they produce in an aqueous solution an acid produces hydrogen ions (H+), and a base hydroxide ions (OH-).
Base: A substance that in its edible form is bitter to the taste. Bases tend to be slippery to the touch, and in reaction withacids they produce salts and water. Bases and acids are most properly defined, however, not in these phenomenological terms, but by the three structural definitions of acids and bases—the Arrhenius, Br0nsted-Lowry, and Lewis acid-base theories.
Basic: In the context of acids and bases, the word is the counterpart to “acidic,” identifying the base-like quality of a substance.
Br0nsted-lowry acid-base theory: The second of three structural definitions of acids and bases. Formulated by English chemist Thomas Lowry (18741936) and Danish chemist J. N. Bronsted (1879-1947), Bronsted-Lowry theory defines an acid as a proton (H+) donor, and a base as a proton acceptor.
Chemical species: A generic term used for any substance studied in chemistry—whether it be an element, compound, mixture, atom, molecule, ion, and so forth.
Conjugate acid: An acid formed when a base accepts a proton (H+).
Conjugate acid-base pair:The acid and base produced when an acid donates a single proton to a base. In the reaction that produces this pair, the acid and base switch identities. By donating a proton, the acid becomes a conjugate base, and by receiving the proton, the base becomes a conjugate acid.
Conjugate base: A base formed when an acid releases a proton.
Dissociation: The separation of molecules into ions.
Ion: An atom or atoms that has lost or gained one or more electrons, and thus has a net electric charge. There are two types of ions: anions and cations.
Ionic bonding: A form of chemical bonding that results from attractions between ions with opposite electric charges.
Ionic compound: A compound in which ions are present. Ionic compounds contain at least one metal and nonmetal joined by an ionic bond.
Lewis acid-base theory: The third of three structural definitions of acids and bases. Formulated by American chemist Gilbert N. Lewis (1875-1946), Lewis theory defines an acid as the reactant that accepts an electron pair from another reactant in a chemical reaction, and a base as the reactant that donates an electron pair to another reactant.
Neutralization: The process whereby an acid and base react with one another to form a salt and water.
ph scale: A logarithmic scale for determining the acidity or alkalinity of a substance, from 0 (virtually pure acid) to 7 (neutral) to 14 (virtually pure base).
Phenomenological: A term describing scientific definitions based purely on experimental phenomena. These only convey part of the picture, however— primarily, the part a chemist can perceive either through measurement or through the senses, such as sight. A structural definition is therefore usually preferable to a phenomenological one.
Reactant: A substance that interacts with another substance in a chemical reaction, resulting in the creation of a product.
Salts: Ionic compounds formed by the reaction between an acid and a base. In this reaction, one or more of the hydrogen ions of an acid is replaced with another positive ion. In addition to producing salts, acid-base reactions produce water.
Solution: A homogeneous mixture in which one or more substances (the solute) is dissolved in another substance (the solvent)—for example, sugar dissolved in water.
Solvent: A substance that dissolves another, called a solute, in a solution.
Strong electrolyte: A substance highly prone to dissociation. The terms “strong acid” or “strong base” refers to those acids or bases which readily dissociate.
Structural: A term describing scientific definitions based on aspects of molecular structure and behavior rather than purely phenomenological data.
Titration: The use of a chemical reaction to determine the amount of a chemical substance in a sample of unknown purity. Testing pH levels is an example of titration.
Transition interval: The pH level at which the color of an acid-base indicator changes.
weak electrolyte: A substance that experiences little or no dissociation. The terms “weak acid” or “weak base” refer to those acids or bases not prone to dissociation.
Because the product of these two concentrations is always the same, this means that if one of them goes up, the other one must go down in order to yield the same constant. This explains the fact, noted earlier, that water can serve either as an acid or base—or, if the concentrations of hydronium and hydroxide ions are equal—as a neutral substance. In situations where the concentration of hydronium is higher, and the hydroxide concentration automatically decreases, water serves as an acid. Conversely, when the hydroxide concentration is high, the hydronium concentration decreases correspondingly, and the water is a base.

Making Dilutions From Acid and Base Solutions

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Making Dilutions From Acid and Base Solutions

Objective

The aim of the first part of the experiment was to make up dilutions
from acid and base solutions and record the results after adding
various indicators. The aim of the second part of the experiment was
to determine the pH values of the dilutions using both pH paper and a
pH metre recording the results for both tests.

Procedure - Part 1

For the first part of the experiment the correct amount of solid NaOH
needed to prepare 250 cm3 of a 0.1M solution was calculated and then
weighed out in a weighing bottle using balanced scales. The weight of
the empty bottle was recorded, minus lid before adding approximately
1g of solid NaOH, re-weighing the bottle plus contents and the new
weight recorded.

The Solid NaOH was then transferred to a 200 cm3 beaker where
approximately 150 cm³ of H2O was added to it and the solution stirred
with a glass stirrer (a magnetic stirrer was not available as
specified in the laboratory manual) until the solid was completely
dissolved. The solution was then transferred, via funnel, into a 250
cm³ volumetric flask and H2O added to increase the volume to 250 cm³,
the solution was then thoroughly mixed by inversion, ensuring the
stopper was held firmly in place.

Next a 1/10 dilution was prepared from the solution using volumetric
flasks and pipettes, although because the original solution was a 0.1M
concentration it was actually already a 1/10 dilution of a 1M
concentration making the 1/10 dilution prepared a 1/100 dilution of a
1M concentration. Therefore the two dilutions of NaOH were 1/10 and
1/100 of a 1M concentration both clearly labelled.

Dilutions of 1/10 and 1/00 using labelled volumetric flasks and
pipettes were then made up from a pre-prepared 1M solution of both HCl
and CH3COOH (adding the acid to water, not the other way around).

Nine test tubes were set out and numerically labelled 1-9 and a
pipette used to transfer a); 5 cm³ of 1/100 HCl into test tubes 1-3

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b); 5 cm³ of 1/100 CH3OOH into test tubes 4-6 and c); 5 cm³ of 1/100
NaOH into test tubes 7-9. The test tubes were then made into “sets”
compiled of one of each of the original sets and 1 drop of three
indicators added to then as shown below, the colour changes were then
recorded.

1 drop of methyl orange to each test tube in set 1 (test tubes 1, 4,
7)

1 drop of methyl red to each test tube in set 2 (test tubes 2, 5, 8)

1 drop of phenolphthalein to each test tube in set 3 (test tubes 3, 6,
9)

Procedure – Part2

An aliquot ~100 cm³ of each of the three concentrations (1M, 1/10M and
1/100M) of both acids HCl and CH3COOH were transferred into labelled
250 cm³ reaction flasks and the two concentrations (1/10M and 1/100M)
of the base NaOH into two labelled reaction flasks using pipettes. All
eight solutions were then tested using both pH paper and a freshly
calibrated pH meter to determine their pH values and the results
recorded.

Results

Table of results for part 1 of the experiment, Dilutions of acids and
base their reaction with indicators.

Solution

Reaction with methyl orange

Reaction with methyl red

Reaction with phenolphthalein

Set 1 – (1, 4, 7)

1. 5 cm³ 1/100 HCl

Pale Pink

4. 5 cm³ 1/100 CH3COOH

Pale Orange

7. 5 cm³ 1/100 NaOH

Orange

Set 2 – (2, 5, 8)

2. 5 cm³ 1/100 HCl

Very Pale Pink

5. 5 cm³ 1/100 CH3COOH

Pink

8. 5 cm³ 1/100 NaOH

No colour

Set 3 – (3, 6, 9)

3. 5 cm³ 1/100 HCl

No Colour

6. 5 cm³ 1/100 CH3COOH

No Colour

9. 5 cm³ 1/100 NaOH

Bright pink

Table of results for the experiment part 2, pH values of acids and
base, using pH paper and pH meter.

Aliquot ~100 cm³ of Solution

pH using pH paper

pH using pH meter

1M HCl

pH 1

pH 0.57

1/10M HCl

pH 2

pH 0.91

1/100M HCl

pH 4.5

pH 1.61

1M CH3COOH

pH 4

pH 2.30

1/10M CH3COOH

pH 5

pH 2.85

1/100M CH3COOH

pH 6

pH 3.30

1/10M NaOH

pH 11

pH 12.72

1/100M NaOH

pH 11

pH 11.70

Calculations

In this section I have included all the calculations and data used
during the experiment.

1) Calculations and data used to calculate how much solid NaOH needed
to make 250 cm³ of a 0.1M of NaOH.

Mass = mols × RM

RM = 23.00+16.00+1.00 = 40g

0.1 × 40 = 4g per dm³

1 dm³ = 1000 cm³ so to find how much in 250 cm³ need to 1000 cm³/4g
= 1g

1g of solid NaOH in a 250 cm³ 0.1M solution.

2) Weight of empty weighing bottle minus lid = 11.7956g

Weight of solid NaOH = 1.0056g

Total Weight =
12.8012g

3) Example of calculations of dilutions

1/10 NaOH → 25 cm³ NaOH in 225 cm³ H2O = 1/10 (or in this case 1/100
as the solution

0.1M was already to 1/10)

Discussion

The results from this experiment show that some errors have occurred
as the recorded data is not correct, compared with what we know should
have happened.

The tables below show what colour the three indicators should have
turned when added to the dilutions according to their pH range and
compares that with the recorded results.

1 2 3 4 5 6 7 8 9
10 11 12 13 14

[IMAGE]

pH range of indicators with acids – bases.

(Ref. Lewis and Evans, (1997), Chemistry 2nd edition. pg 302. Palgrave
Macmillan, UK).

Indicator

Lower pH colour

pH Range

Higher pH colour

Methyl Orange

Deep/Bright Red

3.2 ― 4.4

Bright/Yellow

Methyl Red

Bright/Deep Red

4.8 ― 6.0

Bright/Yellow

Phenolphthalein

No Colour

8.2 ― 10.0

Bright Pink

Comparison between recorded results and proven scientific data using
the pH results

from the pH meter as it is more accurate than using pH paper.

Solution

Ph value

From Ph meter.

Reaction with methyl orange

Reaction with methyl red

Reaction with phenolphthalein

What the reaction should have been

Set 1 – (1, 4, 7)

1. 5 cm³ 1/100 HCl

pH 1.16

Pale Pink

Red

4. 5 cm³ 1/100 CH3COOH

pH 3.30

Pale Orange

Reddish orange

7. 5 cm³ 1/100 NaOH

pH 11.75

Orange

Yellow

Set 2 – (2, 5, 8)

2. 5 cm³ 1/100 HCl

pH 1.16

Very Pale Pink

Red

5. 5 cm³ 1/100 CH3COOH

pH 3.30

Pink

Red

8.5 cm³ 1/100 NaOH

pH 11.75

No colour

Yellow

Set 3 – (3, 6, 9)

3. 5 cm³ 1/100 HCl

pH 1.16

No Colour

No Colour

6. 5 cm³ 1/100 CH3COOH

pH 3.30

No Colour

No Colour

9. 5 cm³ 1/100 NaOH

pH 11.75

Bright pink

Bright Pink

A base solution produces a greater concentration of HO- ions than
concentration of H3O+ ions and an acid solution produces a greater
concentration of H3O+ ions than concentration of HO- ions.

The comparisons show that only set 3 results, using phenolphthalein,
were correct. Phenolphthalein as an indicator reacts with colour
change to bases but not with acids. Basic solutions turn pink which
intensifies has base strength increases and remains transparent in
acidic solutions, therefore showing NaOH to be a strong base the
others to be acids. The results for the methyl red and orange
indicators were incorrect, a freshly calibrated pH meter was used and
so this indicates that the error/s occurred with the amount of
indicator used, the calculating and measuring of the dilutions or
there was contamination due to non sterile equipment.

If the methyl indicators had reacted with the dilutions as expected
then it would have shown CH3COOH to a weak acid and HCl to be strong
acid.

Conclusion

The experiment was partially successful. More attention to detail and
double checking of equipment to avoid contamination could have been
applied to ensure more a accurate result e.g. If the experiment was
carried out by an individual as apposed to a group then measurements
may have been more precise resulting in a more accurate overall
outcome.



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