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Dr Robert Craig PhD

This document provides instructions for standardizing a sodium hydroxide solution by titrating it with a primary standard of potassium hydrogen phthalate. It describes cleaning and preparing the buret, titrating the potassium hydrogen phthalate sample, and calculating the molarity of the sodium hydroxide solution based on the titration results. It also discusses using potassium hydrogen phthalate as a primary standard due to its stability and ease of accurate weighing, and explains the titration process and calculating concentration from titration data.

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PURPOSE:
• To standardize a solution of sodium hydroxide
by titration with a primary standard,
potassium hydrogen phthalate (KHC8H4O4)

Use ph electrode

ANALYTICAL CHEMISTRY
• Determination of iron in a razor blade
• Determination of acetic acid in vinegar
• Determination of Alcohol in Blood

Must do three things before You start
• 1. Clean the Buret
• 2. Delivering Solution from the Buret
• 3. Filing the Buret for a titration
• 4. Preparing the KHP Sample
• 5. Titrating the KHP
• * Must use Your Lab book today!
• *** This is one of the MOST complicated labs

KHP

Why use KHP?
• The hydrogen is slightly acidic, and it is often
used as a primary standard for acid-base
titrations because it is solid and air-stable,
making it easy to weigh accurately. It is not
hygroscopic. It is also used as a primary
standard for calibrating pH meters

as a buffering agent
• As a weak acid hydrogen phthalate reacts
reversibly with water to give hydronium (H3O+)
and phthalate ions.
HP- + H2O <->P2- + H3O+
• KHP can be used as a buffering agent(in
combination with hydrochloric acid (HCl) or
sodium hydroxide (NaOH) depending on which
side of pH 4.0 the buffer is to be)

What I would like . . .
• In this experiment you will determine the
amount of acid needed by titration with the
strong base NaOH.

What is a standardization???
• Since it is hard to prepare a NaOH solution of
accurately known concentration directly
from the solid, you will need to standardize
your NaOH solution against a precisely
weighed amount of standard acid.
• The acid used is the weak monoprotic acid,
potassium hydrogen.

Excellent to use!!!
Can actually obtain 1 mole of H+
by using a balance

A titration
• Knowing the volume of titrant added allows
the determination of the concentration of the
unknown. Often, an indicator is used to
usually signal the end of the reaction, the
endpoint

acid-base indicator
• An acid-base indicator is itself a weak acid (or
its conjugate base).
• An acid-base indicator is a weak acid having a
different colour in aqueous solution from its
conjugate base.

phenolphthalein

phenolphthalein
• Phenolphthalein
• Type: HIn + H2O <->In- + H3O+
pK: 9.5
Approximate pH range for color change:
• 8.0-9.8
Color of acid form: clear
Color of base form: red-violet

phenolphthalein

Commonly used equipment
• Buret
• Erlenmeyer Flask
• Magnetic Stir Plate
• Electronic Balance

accuracy
• The accuracy of the results of your
titration will be a reflection of the
care you took while performing it.
When done carefully, titrations give
very accurate, precise results.
•

accuracy
• SOMETIMES!!!!!!!
• Titrations of unknown solutions are done in
two steps: a scout titration used to determine
the approximate amount of titrant needed
followed by the actual titration that you will
use to make your calculations.

What You will do . . .
• An 0.8167 gram sample of primary standard
KHP (assay = 99.95%) required 38.25 mL of
NaOH to neutralize.
Calculate the molarity of the NaOH solution.

Goto pdf. . . .
• Path:
XXXX g sample -> g KHP -> mol KHP ->
mol NaOH -> M NaOH

path
0.8167 g sample x (99.95g KHP)/(100 gram)
x (1molKHP)/(204.22g KHP)
x (1mol NaOH)/(1 mol KHP)
x (1 )/(0.03825 L)=
0.1040 M NaOH

Using the pH Probe

Using the pH Probe
• Using the pH Probe
• Preparation of the pH Sensor
• A pH sensor connected to a computer will be
used to measure pH. Plug the pH probe into
channel 1 of the interface box. Plug the power
cord on the interface box into an electrical
outlet.

Using the pH Probe

Best site
• http://www.chemistry.mtu.edu/~djchesne/cla
sses/ch2212/Tutorials/Vernier_pH_tutorial.ht
m

Arrhenius acids.

The net ionic equation
H + (aq) + H2O(l) <-> H3O + (aq)
• Pure water already contains small quantities
of the H + ion (proton), H3O + (hydronium ion)

On Exam
• Example: Calculate the weight of primary
standard potassium hydrogen phthalate
• (assay = 99.95%) that would be required to
standardize a 0.1 N NaOH solution,
• assuming a 40 mL titration.
• PATH: L NaOH mol NaOH mol KHP g KHP 
g sample
• Note that 1 equiv = 1 mol for both KHP and NaOH
so N = M and molecular weight = equivalent
weight.

• Example: An 0.8167 gram sample of primary
standard KHP (assay = 99.95%) required
• 38.25 mL of NaOH to neutralize. Calculate the
molarity of the NaOH solution.
• PATH: g sample g KHP mol KHP mol
NaOH M NaOH

• Example: A 1.7734 gram sample of KHP
required 40.11 mL of 0.1036 N for titration.
• Calculate the assay of the KHP and report with
a relative error of 1 part per 1000.

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In Job's method, the variation in concentration of the reactants is performed which can reveal the empirical formula of a complex. The method is employed to find the formula of the compound formed by reacting two or more chemical species. The absorption is recorded against different wavelengths and wavelength having maximum absorption is selected. The intensity of solutions with different stoichiometric ratios of the reactants is measured. The highest observed intensity reveals the maximum amount of compound formed. In this experiment, the Fe3+ and salicylic acid are reacted and the wavelength at which the complex absorbs strongly is selected. The absorbance of the different stoichiometric ration of Fe3+ and salicylic acid at a selected wavelength is determined. The maximum absorbance of the solution of the stoichiometric ratio reveals the empirical formula of the complex as the maximum amount of coloured complex is formed.

Metalloporphyrins with special reference to Iron porphyrins ( Haemoglobin and...

Metalloporphyrins with special reference to Iron porphyrins ( Haemoglobin and...

Metalloporphyrins with special reference to Iron porphyrins ( Haemoglobin and...

Metalloporphyrins with special reference to Iron porphyrins ( Haemoglobin and Myoglobin ) Porphyrins are one of the most important groups of bioinorganic compounds in which a metal ion is surrounded by the four nitrogens of porphin ring. ❑ Porphines are made of four pyrrole rings linked together through methene bridges. ❑ Therefore, porphines have macrocylic pyrrole system with conjugated double bonds as shown here: ❑ These porphines act as tetradentate ligands with four nitrogen donor sites. Two of these are tertiary nitrogen donor positions which can form coordinate bonds by donating a pair of electrons each to the metal ion. ❑ The other two are secondary nitrogen donor positions. each of which lose a proton in forming a coordinate bond with a metal ion. ❑ Thus, a porphin ring acts as a tetradentate dinegative ligand (or dianion). ❑ Dipositive cations such as Mg2+ Fe2+ or Ni2+ are capable of forming neutral complexes with porphine as shown here: ❑ Four pyrrole rings of porphin carrying substituents other than hydrogen are called porphyrins. The complexes in which a metal ion is held in the porphyrin ring system are called metalloporphyrins. ❑ Such complexes play a vital role in biological systems.

Trans effect

1. The trans effect refers to the observation that certain ligands increase the rate of ligand substitution when positioned trans to the departing ligand. 2. This effect was first discovered in 1926 when studying platinum complexes, where it was found that ammonia preferentially substituted the chloride ligand cis rather than trans to the nitrite ligand in Pt(NO2)Cl3 complexes. 3. Two main theories have been proposed to explain the trans effect - the polarization theory involving electrostatic weakening of the trans metal-ligand bond, and the pi-bonding theory involving back-donation of electron density from the metal into the pi* orbitals of ligands like NO2 weakening the trans bond.

Preparation of tetraaminecopper(ii) sulphate complex

Preparation of tetraaminecopper(ii) sulphate complex

Preparation of tetraaminecopper(ii) sulphate complex

Mithil Fal Desai

This document provides instructions for preparing tetraaminecopper(II) sulfate complex in 3 steps: 1) Dissolving copper sulfate in water and adding ammonium hydroxide to form an intense blue solution. 2) Adding ethanol and heating the solution to obtain a clear deep blue solution. 3) Allowing the hot solution to crystallize for an hour, filtering and drying the blue-purple tetraaminecopper(II) sulfate crystals. Yield of the product is then determined.

BORAZINE- structure, preparation and properties

BORAZINE- structure, preparation and properties

BORAZINE- structure, preparation and properties

1. Borazine is called inorganic benzene because it has a planar hexagonal structure like benzene with alternating B and N atoms. 2. Borazine undergoes addition reactions readily unlike benzene due to the polarity in its B-N bonds from the difference in electronegativity between B and N atoms. 3. Borazine can be synthesized in the laboratory by heating a mixture of LiBH4 and NH4Cl in vacuum at 230°C, which yields 30% borazine.

Argentometric titration

Argentometric titration

Argentometric titration

This document describes different methods of argentometric titration: 1. Mohr's method uses silver nitrate solution and potassium chromate indicator to determine the content of chloride and bromide ions through a precipitation reaction. 2. Volhard's method uses thiocyanate ion as the titrant and ferric ion indicator to titrate excess silver ions from the initial reaction of silver nitrate with halide ions. 3. Fajans' method uses a pH indicator absorption method to determine the type of anion present based on the indicator used.

Alcohols,phenols, ethers

Alcohols,phenols, ethers

Alcohols,phenols, ethers

This document discusses alcohols, phenols, and ethers. It defines these compounds and describes their structures. Alcohols contain a hydroxyl group bonded to carbon, while phenols have a hydroxyl group bonded to an aromatic carbon. Ethers have an alkoxy or aryloxy group in place of a hydrogen. The document classifies these compounds based on the number and position of functional groups. It also discusses their naming conventions, physical properties, bonding, and methods of synthesis.

Molecular orbital theory

Molecular orbital theory

Molecular orbital theory

Pallavi Kumbhar

Molecular orbital theory (MOT) is an alternative model to valence bond theory that explains how atomic orbitals from different atoms combine to form molecular orbitals. The linear combination of atomic orbitals (LCAO) method considers the probability of finding electrons in atomic orbitals from different atoms. According to the LCAO method, molecular orbitals are formed from constructive and destructive interference of atomic orbitals. MOT can be used to explain bonding in homonuclear diatomic molecules like N2 and O2, heteronuclear diatomic molecules like CO and NO, and polyatomic molecules like CO2 and SF6. It can also describe bonding in octahedral transition metal complexes like hexaaquoferrate(II) ion

Colorimetry

Saravanan subramaniyam

This document discusses the principles and components of colorimetry. Colorimetry is a technique used to determine the concentration of colored compounds in a sample solution using visible light between 400-800 nm. The intensity of color is directly proportional to the concentration being measured. The key components of a colorimeter include a light source, filters to select the wavelength, a sample holder, and detectors to convert light energy to electrical signals. Common light sources are tungsten lamps and carbon arc lamps. Absorption and interference filters are used to select the desired wavelength. Barrier layer cells and photomultiplier tubes are common types of detectors.

TLC for chlorinated pesticide determination Bapi Mondal

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TLC for chlorinated pesticide determination Bapi Mondal

FISCHER AND NEWMAN PROJECTIONS

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FISCHER AND NEWMAN PROJECTIONS

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Transport number

Transport number

Transport number

M.sc. laboratory manual organic chemistry binary mixture seprations

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48695528 the-sulphur-system

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Metal Ligand Bonding in Transition metal complexes

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Metal Ligand Bonding in Transition metal complexes

P h metric titration

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Topicity

Pauli's Exclusion Principle.pptx

Pauli's Exclusion Principle.pptx

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Preparation of 2,4,6 tribromoaniline

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Preparation of 2,4,6 tribromoaniline

Determine the composition of the fe3+(jobs method)

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Determine the composition of the fe3+(jobs method)

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Preparation of tetraaminecopper(ii) sulphate complex

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TLC for chlorinated pesticide determination Bapi Mondal

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TLC for chlorinated pesticide determination Bapi Mondal

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Dr Robert Craig PhD

This document provides instructions for standardizing a sodium hydroxide solution by titrating it with potassium hydrogen phthalate (KHP) as the primary standard. It describes preparing the KHP sample, titrating the KHP with NaOH using a pH electrode and phenolphthalein indicator, and calculating the molarity of the NaOH solution based on the titration results. The document emphasizes that KHP is an ideal primary standard for acid-base titrations due to its stability and ease of accurate weighing.

PH, BUFFERS, PH METER.pptx

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PH and buffers can be summarized in 3 sentences: PH is a measure of acidity or alkalinity in solutions, determined by the concentration of hydrogen ions, with lower PH indicating more acidic solutions. Buffers resist changes in PH when acids or bases are added by maintaining an equilibrium between the acid/base forms. Common buffer systems used in biochemistry include acetate, borate, and bicarbonate buffers which are important for maintaining stable PH levels in processes like electrophoresis.

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This document discusses limitations of Lewis acid-base reactions, properties of buffer solutions, buffer capacity, the Henderson-Hasselbalch equation, and applications of buffers in pharmacy. It also covers general principles for adjusting solutions to isotonicity and its importance, types of impurities found in pharmaceutical substances, sources of impurities, effects of impurities, and limit tests for chloride, sulfate, iron, heavy metals, lead, and arsenic. Limit tests are used to identify and control small quantities of impurities that may be present in substances.

non aqueous titration, pharmaceutical and cosmetic analysis,Pca seminar

non aqueous titration, pharmaceutical and cosmetic analysis,Pca seminar

non aqueous titration, pharmaceutical and cosmetic analysis,Pca seminar

Non-aqueous titration involves titrating substances dissolved in solvents other than water. It has several advantages over aqueous titration, such as allowing the titration of very weak acids and bases. It provides a solvent in which organic compounds are soluble. Common non-aqueous titrations include titrating organic bases with perchloric acid in acetic acid. The document discusses the classification of solvents, reactions that take place, and examples of titrating substances like pyridine and benzoic acid in solvents like acetic acid and DMF. Practical applications involving titrating drugs and calculating percentage purity are also presented.

Buffers_Acidic and Basic buffer solutions

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Buffers_Acidic and Basic buffer solutions

ChimwemweGladysBanda

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Parthasarathy Ravichandran

The document describes several experiments related to protein analysis: 1. It provides procedures for preparing phosphate and acetate buffers, estimating protein concentration using Lowry's method, determining the absorption maximum of riboflavin, performing paper chromatography to separate amino acids, and titrating acetic acid to find its pKa value. 2. It also outlines protocols for purifying protein from a sample via ammonium sulfate precipitation and estimating the precipitated protein, separating amino acids from a mixture using ion exchange chromatography, and performing SDS-PAGE to separate proteins by molecular weight. 3. Reagents, materials, and step-by-step procedures are described for each experiment to analyze different properties of proteins and separate mixtures of

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P h meter use and calibration

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P h meter use and calibration

A pH meter measures the concentration of hydrogen ions (H+) in a solution using an ion-sensitive electrode. It has a glass electrode and reference electrode in one tube. When measuring pH, a salt bridge forms allowing sodium ions (Na+), not H+, to cross the glass membrane and allow the pH meter to measure the H+ concentration. To calibrate the pH meter, it is placed in high, neutral, and low pH buffer solutions. Litmus paper can also indicate pH but a pH meter provides a more precise measurement. The pH scale ranges from 0-14 with lower numbers indicating more acidity and higher numbers indicating more basicity.

Non aq. acid base titration

Non aq. acid base titration

Non aq. acid base titration

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Different techniques of analysis

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Poonam Aher Patil

Mrs. Poonam Sunil Aher discusses different analytical techniques including quantitative methods like solubility, melting point, and boiling point as well as qualitative methods like color, odor, and identification tests. She focuses on titrimetric techniques like acid-base titration using indicators, redox titration, iodimetry titration using starch, and precipitation titration. She also covers gravimetric analysis techniques like volatilization and precipitation, and discusses aqueous versus non-aqueous titration.

P h meter use and calibration

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A pH meter measures the concentration of hydrogen ions (H+) in a solution by using a glass electrode and reference electrode in one tube. It works by forming a salt bridge between the electrodes when placed in the solution, allowing hydrogen ions to cross the glass membrane and cause a change in free energy measured as pH. The meter must be calibrated using buffer solutions of different pH levels before taking measurements. Proper storage and cleaning of the electrodes is important for accurate readings. Litmus paper can also indicate pH but is not as precise as a pH meter.

18 pH Meter.pptx

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18 pH Meter.pptx

A pH meter measures the pH of a liquid by determining the concentration of hydrogen ions using a glass electrode and reference electrode. The pH meter must be calibrated using buffer solutions of known pH values in order to provide accurate readings. A pH of 7 indicates a neutral solution, while values below 7 are acidic and above 7 are alkaline. The pH scale is logarithmic, so each unit of change in pH represents a tenfold change in hydrogen ion concentration.

Maam queen 2

Kym Anne Surmion II

This document discusses pH and hydronium ion concentration in aqueous systems. It covers how pH is involved in hospital pharmacy parenteral admixture programs. pH values outside of certain limits can cause physical or chemical incompatibilities and adverse effects in patients. pH is measured using a standardized pH meter, pH indicator paper, or calculation. The document also discusses acid-base theories like Bronsted-Lowry and Arrhenius, and principles of pH, protolysis, autoprotolysis, amphoteric substances, and the Henderson-Hasselbalch equation. Examples are given for calculating pH using the Hendersen-Hasselbalch equation and determining buffer capacity.

Body buffers and acid base balance.pptx

Body buffers and acid base balance.pptx

Body buffers and acid base balance.pptx

The document discusses body buffers and acid-base balance. It covers several key points in 3 sentences: The human body maintains acid-base balance through various buffer systems, including bicarbonate, phosphate, and proteins, which help resist changes in pH. Respiratory mechanisms like breathing and renal mechanisms like kidney function also help regulate pH levels over different time periods. The Henderson-Hasselbalch equation relates the pH of a solution containing a weak acid and its conjugate base to their concentration ratio and defines acid-base properties crucial for cellular functions.

pH - MEASUREMENT .ppt

pH - MEASUREMENT .ppt

pH - MEASUREMENT .ppt

RosalindaGironPrinci

This document discusses pH measurement and provides details on: - The definition and scale of pH as a measure of acidity or alkalinity. - Why pH is measured in various industries and applications. - The principles of pH measurement using a glass electrode and pH meter. - Factors that affect pH measurement accuracy such as temperature, ionic strength, and electrode calibration. - The process of calibrating pH electrodes using buffer solutions and adjusting for the Nernstian slope.

calculations for preparation of solutions

calculations for preparation of solutions

calculations for preparation of solutions

MansurkhanTadakod1

1. This document provides guidelines for preparing chemical solutions accurately for laboratory use. It describes the types of solutions that may be required like percent, molar, and normal solutions and how they are calculated and prepared. 2. Quality control procedures are important to maintain precision and accuracy of results. Standards should be prepared and tested with each batch to check if instrument conditions and test performance are satisfactory. 3. The document emphasizes careful preparation of reagents, use of analytical grade chemicals, standardization of methods, and maintaining quality control through analysis of standards to ensure reproducibility of results.

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Introduction to Volumetric Analysis for Basic Chemistry

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HusseinHanibah1

HPLC method development by DrA.S Charan

HPLC method development by DrA.S Charan

HPLC method development by DrA.S Charan

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NCHE 211 UNIT 2 VOLUMETRY and Complexometry titration.pdf

NCHE 211 UNIT 2 VOLUMETRY and Complexometry titration.pdf

This module consists of six study units that are further divided into study sections. The study units include basic concepts, volumetry, gravimetry, surface characterization, atomic spectrometry, and separation methods. The document then provides an overview of analytical chemistry and its branches, classification of quantitative analysis methods, and learning outcomes related to volumetric analysis and titrations. [/SUMMARY]

Buffer presentation

Buffer presentation

Buffer presentation

Ashok Mourya @@@

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This document contains an astronomy homework assignment with multiple choice questions about the phases of the Moon and the scale of planetary orbits. It includes diagrams of the Moon at different positions in its orbit around Earth and asks the student to rank the Moon's appearance in terms of the illuminated area visible from Earth. The homework aims to test the student's understanding of the relative positions of Earth, the Sun and Moon and how this determines what lunar phase we see from Earth.

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1) The document describes a ranking task that orders major events in the history of the universe from longest ago to most recent. It then provides context about the "cosmic calendar" that compresses the 14 billion year history of the universe into a single calendar year. 2) On the cosmic calendar, the Big Bang occurred at the start of the year on January 1st, approximately 14 billion years ago. 3) Earth formed in early September on the calendar, around 4.5 billion years ago.

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This document provides an overview of astronomy and the scientific method. It discusses: 1) Astronomy as the study of objects beyond Earth and how they interact, with the goal of organizing our understanding of the universe's history. 2) The scientific method as a process of making observations, developing hypotheses, and testing them through experiments or further observations. Hypotheses must be falsifiable to be scientific. 3) Scientific laws as consistent rules that describe natural phenomena, allowing our understanding to be applied universally throughout the universe. Laws are subject to revision with new evidence.

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Khpacidbaseneutralization1slideonechm240xxxxxxxxxxxxxxxxxxdontletgo2 140110101250-phpapp02

1.

2. PURPOSE: • To standardize a solution of sodium hydroxide by titration with a primary standard, potassium hydrogen phthalate (KHC8H4O4)

3. Use ph electrode

4. ANALYTICAL CHEMISTRY • Determination of iron in a razor blade • Determination of acetic acid in vinegar • Determination of Alcohol in Blood

5. Must do three things before You start • 1. Clean the Buret • 2. Delivering Solution from the Buret • 3. Filing the Buret for a titration • 4. Preparing the KHP Sample • 5. Titrating the KHP • * Must use Your Lab book today! • *** This is one of the MOST complicated labs

7. Why use KHP? • The hydrogen is slightly acidic, and it is often used as a primary standard for acid-base titrations because it is solid and air-stable, making it easy to weigh accurately. It is not hygroscopic. It is also used as a primary standard for calibrating pH meters

8. as a buffering agent • As a weak acid hydrogen phthalate reacts reversibly with water to give hydronium (H3O+) and phthalate ions. HP- + H2O <->P2- + H3O+ • KHP can be used as a buffering agent(in combination with hydrochloric acid (HCl) or sodium hydroxide (NaOH) depending on which side of pH 4.0 the buffer is to be)

9. What I would like . . . • In this experiment you will determine the amount of acid needed by titration with the strong base NaOH.

10. What is a standardization??? • Since it is hard to prepare a NaOH solution of accurately known concentration directly from the solid, you will need to standardize your NaOH solution against a precisely weighed amount of standard acid. • The acid used is the weak monoprotic acid, potassium hydrogen.

11. Excellent to use!!! Can actually obtain 1 mole of H+ by using a balance

12. A titration • Knowing the volume of titrant added allows the determination of the concentration of the unknown. Often, an indicator is used to usually signal the end of the reaction, the endpoint

13. acid-base indicator • An acid-base indicator is itself a weak acid (or its conjugate base). • An acid-base indicator is a weak acid having a different colour in aqueous solution from its conjugate base.

14. phenolphthalein

15. phenolphthalein • Phenolphthalein • Type: HIn + H2O <->In- + H3O+ pK: 9.5 Approximate pH range for color change: • 8.0-9.8 Color of acid form: clear Color of base form: red-violet

16. phenolphthalein

17. Commonly used equipment • Buret • Erlenmeyer Flask • Magnetic Stir Plate • Electronic Balance

18. accuracy • The accuracy of the results of your titration will be a reflection of the care you took while performing it. When done carefully, titrations give very accurate, precise results. •

19. accuracy • SOMETIMES!!!!!!! • Titrations of unknown solutions are done in two steps: a scout titration used to determine the approximate amount of titrant needed followed by the actual titration that you will use to make your calculations.

20. What You will do . . . • An 0.8167 gram sample of primary standard KHP (assay = 99.95%) required 38.25 mL of NaOH to neutralize. Calculate the molarity of the NaOH solution.

21. Goto pdf. . . . • Path: XXXX g sample -> g KHP -> mol KHP -> mol NaOH -> M NaOH

22. path 0.8167 g sample x (99.95g KHP)/(100 gram) x (1molKHP)/(204.22g KHP) x (1mol NaOH)/(1 mol KHP) x (1 )/(0.03825 L)= 0.1040 M NaOH

23. Using the pH Probe

24. Using the pH Probe • Using the pH Probe • Preparation of the pH Sensor • A pH sensor connected to a computer will be used to measure pH. Plug the pH probe into channel 1 of the interface box. Plug the power cord on the interface box into an electrical outlet.

25. Using the pH Probe

26. Best site • http://www.chemistry.mtu.edu/~djchesne/cla sses/ch2212/Tutorials/Vernier_pH_tutorial.ht m

27.

28.

29. Arrhenius acids.

30. The net ionic equation H + (aq) + H2O(l) <-> H3O + (aq) • Pure water already contains small quantities of the H + ion (proton), H3O + (hydronium ion)

31. On Exam • Example: Calculate the weight of primary standard potassium hydrogen phthalate • (assay = 99.95%) that would be required to standardize a 0.1 N NaOH solution, • assuming a 40 mL titration. • PATH: L NaOH mol NaOH mol KHP g KHP  g sample • Note that 1 equiv = 1 mol for both KHP and NaOH so N = M and molecular weight = equivalent weight.

32. • Example: An 0.8167 gram sample of primary standard KHP (assay = 99.95%) required • 38.25 mL of NaOH to neutralize. Calculate the molarity of the NaOH solution. • PATH: g sample g KHP mol KHP mol NaOH M NaOH

33. • Example: A 1.7734 gram sample of KHP required 40.11 mL of 0.1036 N for titration. • Calculate the assay of the KHP and report with a relative error of 1 part per 1000.