Chem 14BL W1-4

how close a set of results is to be true value

how close a set of results are to each other, indicating consistency

when values are close to each other and to the true value

results are consistent but mot close to true value

systemic errors are consistent and directional, often due to flaws in experiment setup

EX: reading volume too high on graduated cylinder every time

random errors are unpredictable and are due to uncontrolled variables

EX: different readings from digital thermometer on same object

fixed amount of error in the same unit and decial places as the measurement

the error as percentage of the measurement

Calc: % Relative Error = (Absolute Error / Measurement) × 100%

± 0.1 cm

ΔC = (( ±0.02/ 5.00) + (±0.0002/0.4000)) x (5.00 x 0.3000)

= 0.009

C= AxB —> 5.00x.4000 = 2.00

--> 2.00 ± 0.01

we CANNOT write = 2.00 ± 0.009 round up the value of 0.009 to 0.01 to match precision of 2.00 (2 decial places)

ΔC = (ΔA/A) + (ΔB/ B) x C

C= A x B

ΔC/ C = (ΔA/A + ΔB/B)

ΔA ΔB are the absolute errors in A and B

ΔC/C = (( ±0.02/ 5.00) + (±0.0002/0.4000))

= 4.5 x10^-3

= 5x10^-3 (rounded)

= ± 0.5% (1 sig fig based on 5)

ΔC/ C = (ΔA/C + ΔB/C)

ΔC = (ΔA/C + ΔB/C) x C

C = A-B

ΔC/ C = ((±0.02/4.65) + (±0.02/4.65))

= 9x10^-3

= ± 0.9% (1 sig fig based on mathematics)

ΔC = ((±0.02/4.65) + (±0.02/4.65) x (5.05-0.40)

= 4x10^-2

= 4.65 ± 0.04

It does NOT affects %RAD, as % RAD measures variability, not consistency in one direction

it increases %RAD by adding variability to the measurement

Random error due to unpredictable fluctuations in the light source affecting each measurement differently

Systemic error, as it cosistently affects the volume delivered in one direction (too much liquid)

±0.01

absolute error found to be 0.3mL and absolute error is the fixed amount of error with the same amount of units and decimal places, the measurement should reported as 15.0mL with 3 sig figs to satisfy the 2% relative error requirement.

the absolute error cannot be in percentage form. if calculating for the relative error, must be done seperately.

5.00 mL (3 sig figs)

10.00 (4 sigs figs)

100.00mL (5 sig figs)

to the hundredth decimal place in mL

EX: the volume of a 100-mL volumetric flask is equivalent to 100.00mL OR 0.10000L (both values have the same number of significant figures!)

yes, deciamal places change due to unit conversion, but sig figs remains the same

Semi-quantitative provides approximate, rather than exact results, not perfectly precise, with a range of ±5 to 10%

rapid, allows for quick measurements

it’s easy to establigh quality control measurements, making results mroe reliable?

can be used for on-line monitoring to continuously track changes

the absorption profile of the chemical of interest must be distinct from other solution constituents to avoid interference

I = I₀ X 10^-ϵCL

I₀ is the intensity of the incident light beam

I is the intensity of the transmitted light beam

ϵ is the molar absorptivity constant (units: L/(mol·cm) or M⁻¹cm⁻¹).

ϵ depends on the identity of the chemical species and the wavelength of light selected.

C is the concentration of the chemical species in the solution (in mol/L).

L is the depth of the solution that the light beam traverses (in cm).

Transmittance (T) is the ratio of transmitted light to incident light, calculated as T=I/I₀=10^−ϵCL, where I₀ is the intensity of incident light and I is the intensity of transmitted

Absorbance (A) is calculated as A=−log⁡(T)

A=ϵCL, which is Beer’s Law.

Wavelength (λ) is fixed during the analysis.

The sample has a low concentration.

Low concentration is ambiguous, but it generally refers to concentrations where Beer’s Law holds without deviations.

Absorbance (A) is directly proportional to Concentration (C).

Absorbance (A) should fall between 0.1 and 1.0 for accurate results.

The limiting factor with the number of decimal places

ex answer: 40.9 —> 1 sig fig

the limiting factor

ex answer: 193 —> no sig figs

can’t add or subtract if the values have different exponents in since it affects decimal place, change exponents so they match

ex: 5.85×10² + 121×10² = 1.27 x10⁴

find limiting factor and use sig figs not decimal place

ex answer: limiting factor 2.9 w/ 2 sig figs, answer —> 18

49.8767 g - 49.2140 g =0.663 g 0.6627

2 SF

Relative Average Deviation

When one result deviates significantly from others, potentially causing large error in the average, especially if repeating the experiment isn’t possible

identifiable systemic error in the measurement or if the statistical deviation indicates low probabilty of the results fitting with the rest of the data

statistical method determines if a single deviant result in a small data set should be rejected based on how much it deviates from the others

R_q=∣x_q−x_n∣ / ∣x_q−x_f∣​

x_q = questionable result

x_n = result nearest to x_q

x_f = result farthest from x_q

Rq is compared to a critical ration (found in chart IV 7) specific to the number or measurements (n<7), if Rq greater, result likely error

graph must start at 0,0 bc of the blank reading

slope = ϵL

or = y_f-y_i / x_f-x_i

if curved or non-linear means Beer’s Law failed and concentration may be too high

at high concentrations, solution absorbs to much light that absorbance and concentration (beer’s Law) no longer accurate. Dilute the solution until absorbance falls between 0.1 to 1.0

ϵ = slope/ L

it "takes in" that color and doesn't reflect it back. Instead, we see the opposite (complementary) color.

the opposite color to the one absorbed. For example, if yellow is absorbed, we see purple.

to purify an impure product by forming crystals of the pure compound

soluability increases with temp, so compound dissolves better in hot solvent and forms crystals as it cools

impurities can have simililar soluability to the desired product, makaing them hard to seperate

1. dissole impure product in minimum hot solvent to make saturated solution

2. filter to remove insoluable impurities

2. Cool solution to form crystals

4. Filter crystals to get purified product

helps make the solution more saturated so impurities stay dissolved, and more pure crystal forms

Dissolve a lot of the product at high temperature

Not dissolve the product at low temperature

Dissolve impurities at all temperatures

Not react with the product

Two solvents can balance solubility better: one with high solubility and one with low solubility for the product, making crystallization more effective.

The two solvents must:

Be miscible (mix well together)

High solubility solvent dissolves the product easily

Low solubility solvent dissolves the product less, helping crystals form

Dissolve the product in the high-solubility solvent at high temperature.

Add low-solubility solvent slowly until the solution turns cloudy (crystals start forming).

Reheat to dissolve any crystals if needed; add a few drops of the high-solubility solvent if necessary.

NaOH reacts with aspirin in a 1:1 ratio to neutralize it. An acid-base indicator helps track this reaction by changing color at the end point.

It’s the point where the indicator changes color to show that the reaction is complete. It should be very close to the equivalence point.

point where moles of acidic protons (H+) equal moles of hydroxide (OH-). In the equation, this is expressed as MₐVₐ = MᵦVᵦ (where M = molarity, V = volume).

The indicator’s color change (end point) should occur close to the equivalence point pH to accurately signal the reaction’s completion.

Use a buret with NaOH of known concentration to measure volumes accurately, recording to 2 decimal places. The end point tells us moles of NaOH = moles of aspirin.

the attractive forces between molecules that hold crystalline structures together. They require energy (heat) to break.

the temperature at which a solid becomes a liquid. It can help determine the identity and purity of a substance.

Melting point alone doesn’t confirm identity because many organic compounds may have similar melting points.

A sharp melting point range indicates a pure compound, while a wider range suggests impurities. Pure substances typically melt within a range of ≤ 2°C.

The melting point range for a pure substance is usually very narrow, and in lab settings, it’s recorded as two temperatures:

Ti: First sign of melting

Tf: Complete meltingThe range is calculated as ΔT=Tf−TiΔT=Tf−Ti.

The strength of intermolecular forces influences both melting points and boiling points; stronger forces generally result in higher melting and boiling points.

Weight/volume percent (w/v%) is defined as the weight (in grams) of a substance per 100 mL of solution.

It means there are 5.0 grams of NaCl dissolved in 100 mL of solution.

concentration = 5g / 275 mL =0.0182 g/mL 0.018 g/mL

this is equivalent to 1.8g in 100mL, or 1.8% (w/v%)

Parts per million (ppm) is defined as milligrams of solute per liter of solution (1 mg = 10⁻³ g).

Parts per billion (ppb) is defined as micrograms of solute per liter of solution (1 μg = 10⁻⁶ g).

1. convert molarity (mol/L) to g/L using molar mass

2. divide g/L result by 1000 to get g/mL

3. multiply g/mL by 100 to obtain w/v% (g/100mL)

** no need to worry about sig figs

convert molarity (mol/L) to g/L using molar mass

ultiple g/L by 1000 to get ppm (mg/L)

theoretical yield = mols of limiting reactant x molar mass of product

(actual yield/ theoretical) 100%

actual yiel = theoretical x % yield

% yield = (actual yield/ theoretical ) 100

V = m/s

V is the volume of solvent required (in mL),

mm is the mass of the solute to be dissolved (in g),

S is the solubility of the solute in the solvent at a given temperature (in g/mL).

weak dispersion forces that occur due to temporary dipoles in non-polar molecules. they increase with size of the molecule

occur between molecules that have permanent dipoles (ex: polar molecules).

EX: bond in carbonyl (C=O) where they are partial charges ((δ+ and δ−).

strong type of of dipole-diple interactoin occurs when H is bonded to highly electronegative atoms O,N,F