Unit 1 - Error, 1.1, 1.4, 2.1

Causes of error

Causes of error

• imperfections in apparatus

• imperfections in methodology

• judgements made by humans

Random Uncertainty

• cannot be avoided

• degree of error

• sort of solved by doing many trials and averaging data

Analog uncertainly

• by eye measures

• half the smallest division available

Digital Uncertainty

• measured by equipment

• uncertainty is 1 of the smallest digit

Absolute uncertainty

the uncertainty associated with a given measurement

Relative uncertainty

absolute uncertainty / volume or substance

Adding and Subtracting Uncertainties

the absolute uncertainties are added

Multiplying and dividing

relative uncertainties added

Percent error

(experimental-theoretical)/(theoretical)

Systematic errors

consistent bias that always changes the outcome

Mole

1 mole =6.02×10²³

energy

the ability to preform work or produce heat

Atomic theory

All matter is comprised of atoms and matter can neither be created nor destroyed

Elementary substances

contain atoms of a single element

Chemical Compounds

contain atoms of two or more elements bound together

Filtration

mixture is pored through a paper filter or other porous material

Dissolution

mixture is added to water or an organic solvent

crystallization

mixture dissolved in hot water, solution then left to cool down, crystals formed

Sublimation

when a solid substance turns directly into a gas (skips melting phase)

Deposition

When gases turn directly into solids (no liquid phase)

Endothermic process

system absorbs heat from suroundings

exothermic process

System releases heat into the surroundings

Temperature

the measure of the average kinetic energy

0 degrees C

273.15 Kelvin

Relative Molecular mass

the ratio of the mass of a molecule to 1/12 of a Carbon-12 molecule

hydrates

compounds in which water molecules form coordinated bonds with ions

Molar Mass (M)

M= mass/moles

Molecular formula

shows the actual number of atoms of each element in the molecule of that substance

Empirical formula

shows the simplest ratio (integers) of atoms of the different elements

Percent composition

elemental composition in percent by mass

Solutions

homogenous mixtures of two or more substances

Solvent

what the liquid/dissolver is

Solute

what is being dissolved (salt)

Molar Concentration (AKA Molarity)

the ratio of the amount of solute to the volume of the solution

M= grams/volume in dm³

Stock solutions

Standard solutions

• can be stored and dilluted

Avagadro’s Law

the volume of two reacting gases under the same conditions is proportional to the amounts

• v1/n1 = V2/n2

Assumptions of the ideal gas model

1. all gas molecules are in constant random motion

2. collisions between molecules are perfectly elastic (no energy is lost)

3. the volume occupied by gas molecules is negligible compared to the volume of the container they occupy

4. there are not Intermolecular Forces (wont condense into a liquid)

5. the kinetic energy of the molecules is directly proportional to the temperature in kelvin

Boyle’s Law

pressure and volume are inversely related

P1V1=P2V2

KE decr. and IMF’s form

Effect of low temp. on the ideal gas model

relation between pressure and volume no longer inverse

Effect of high pressure on the ideal gas model

Combined gas law

Ideal gas law

PV=nRT

ideal gas constant

8.31 JK mol

Percent yield

experimental/theoretical *100

Common practices of Green Chemistry

1. aqueous or solvent-free

2. renewable starting materials

3. mid reaction conditions

4. efficient catalysts

5. utilization of any byproducts formed

Atom economy

the measure of the efficiency of a reaction

= (molar mass of desired)/(total mass of reaction) * 100%